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New Simplified Chemistry Class 9 ICSE Solutions Study of Gas Laws

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Exercise

Question 1.(1988)
“When stating the volume of a gas the pressure and temperature should also be given”. Why ?
Answer:
Volume of a gas under goes significant change if its pressure or temperature is slightly changed.

Question 1.(1989)
Define or state : Boyle’s Law
Answer:
Boyle’s Law : “Temperature remaining constant the volume of a given mass of dry gas is inversely proportional to its pressure.”

Question 2.(1989)
Express Kelvin Zero in °C
Answer:
Kelvin zero or absolute zero = — 273°C.

Question 1.(1992)
A fixed volume of a gas occupies 760cm³ at 27° C and 70cm of Hg. What will be its vol. at s.t.p. [637 cm³]
Answer:

Question 1.(1993)
State : Boyle’s Law
Answer:
Boyle’s Law : “Temperature remaining constant the volume of a given mass of dry gas is inversely proportional to its pressure.”

Question 1.(1995)
At 0°C and 760 mm Hg pressure, a gas occupies a volume of 10Q cm³. The Kelvin temperature (Absolute temperature) of the gas is increased by one-fifth while the pressure is increased one and a half times. Calculate the final volume of the gas. [80 cc.]
Answer:

Question 1.(1996)
The pressure on one mole of gas at s.t.p. is doubled and the temperature is raised to 546 K. What is the final volume of the gas ? [one mole of a gas occupies a volume of 22.4 litres at stp.] [22.4 ltrs.]
Answer:

Question 1.(1997)
Is it possible to change the temperature and pressure of a fixed mass of gas without changing its volume. Explain your answer.
Answer:
No, it is not possible as change of any one of the parameters (pressure or temperature) has significant effect on volume.

Additional Questions

Question 1.
What volume will a gas occupy at 740 mm pressure which at 1480 nun occupies 500 cc ? [Temperature being constant] [1000 cc]
Answer:

Question 2.
The volume of a given mass of a gas at 27°C is 100 cc. To what temperature should it be heated at the same pressure so that it will occupy a volume of 150 cc ?[177°C]
Answer:

Question 3.
A fixed mass of a gas has a volume of 750 cc at—23°C and 800 mm pressure. Calculate the pressure for which its volume will be 720 cc. The temperature being —3°C. [900mm]
Answer:

Question 4.
What temperature would be necessary to double the volume of a gas initially at s.t.p. if the pressure is decreased by 50% ? [0°C]
Answer:

Question 5.
A gas cylinder having a capacity of 20 litres contains a gas at 100 atmos. How many flasks of 200 cm³ capacity can be filled from it at 1 atmos. pressure if the temperature remains constant ? [10,000]
Answer:

Question 6.
A certain mass of gas occupied 850 ml at a pressure of 760 mm of Hg. On increasing the pressure it was found that the volume of the gas was 75% of its initial value. Assuming constant temperature, find the final pressure of the gas? [1013.33 mm of Hg]
Answer:

Question 7.
It is required to reduce the volume of a gas by 20% by compressing it at a constant pressure. To do so, the gas has to be cooled. If the gas attains a final temperature of 157°C, find the initial temperature of the gas ? [264.5°C]
Answer:

Question 8.
At a given temperature the pressure of a gas reduces to 75% of its initial value and the volume increases by 40% of its initial value. Find this temperature if the initial temperature was —10°C. [3.15°C]
Answer:

Study Of Gas Laws – Unit Test Paper 7

Q.1. Name or state the following :

Question 1.
The law which states that pressure remaining constant the volume of a given mass of dry gas is directly proportional to its absolute [Kelvin] temperature.
Answer:
Charle’s Law.

Question 2.
The law which studies the relationship betweeir pressure of a gas and the volume occupied by it at constant temperature.
Answer:
Boyle’s Law.

Question 3.
An equation used in chemical calculations which gives a simultaneous effect of changes of temperature and pressure on the volume of a given mass of dry gas
Answer:
Gas equation.

Question 4.
The standard pressure of a gas in cm. of mercury corresponding to one atmospheric pressure.
Answer:
76 cm.

Question 5.
The absolute temperature value corresponding to 35°C.
Answer:
35 + 273 = 308K

Q.2. Give reasons for the following :

Question 1.
Gases unlike solids and liquids exert pressure in all directions.
Answer:
Impact of gas molecules with high velocity causes pressure to be exerted on the walls.

Question 2.
Gases have lower densities compared to solids or liquids.
Answer:
Gases have low densities as the inter-molecular distance between the molecules of gases is very large.

Question 3.
Temperature remaining constant the product of the vol. & the press, of a given mass of dry gas is a constant.
Answer:

Question 4.
All temperatures on the Kelvin scale are in positive figures.
Answer:
All temperatures on the Kelvin scale are in positive figures.
The temperature -273°C = OK [Absolute zero or zero Kelvin]
Hence it may be negative or positive on Celsius scale, it is always positive on Kelvin as 0°C = 0 + 273 = 273K
30°C = 30 + 273 = 303 K
– 70°C = – 70+ 273;= 203 K
– 273°C = – 273 + 273 = 0K

Question 5.
Volumes of gases are converted into s.t.p. conditions and then compared.
Answer:
Volumes of gases are converted into s.t.p. conditions and then compared as
volumes of gases change with temperature and pressure – hence a standard value of temperature and pressure is chosen to which gas volumes are referred.

Q.3. Calculate the following :

Question 1.
Calculate the temperature to which a gas must be heated, so that the volume triples without any change in pressure. The gas is originally at 57”C and having a volume 150 cc.
Answer:

Question 2.
A gas ‘X’ at -33°C is heated to 127° C at constant pressure. Calculate the percentage increase in the volume of the gas.
Answer:

Question 3.
Calculate the volume of a gas ‘A’ at s.t.p., if at 37°C and 775 mm of mercury pressure, it occupies a volume of 9 1/2 litres.
Answer:

Question 4.
Calculate the temperature at which a gas ‘A’ at 20°C having a volume, of 500 cc. will occupy a volume of 250 cc.
Answer:

Question 5.
A gas ‘X’ is collected over water at 17°C and 750 mm. pressure. If the volume of the gas collected is 50 cc., calculate the volume of the dry gas at s.t.p. [at 17°C the vapour pressure is 14 mm.]
Answer:

Q.4. Assuming temperature remaining constant calculate the pressure of the gas in each of the following :

Question 1.
The pressure of a gas having volume 1000 cc. originally occupying 1500 cc. at 720 mm. pressure.
Answer:

Question 2.
The pressure of a gas having volume 100 lits. originally occupying 75 dm³ at 700 mm. pressure.
Answer:

Question 3.
The pressure of a gas having volume 380 lits. originally occupying 800 cm³ at 76 cm. pressure.
Answer:

Question 4.
The pressure of a gas having volume 1800 ml. originally occupying 300 ml. at 6 atms. pressure.
Answer:

Question 5.
The pressure of a gas having volume 1500 cm³ originally occupying 750 cc. at 5 ats. pressure.
Answer:

Q.5. Calculate the following :

Question 1.
The temp, at which 500 cc. of a gas ‘X’ at temp. 293K occupies half it’s original volume [pressure constant].
Answer:

Question 2.
The volume at s.t.p. occupied by a gas “Y” originally occupying 760 cc. at 300K and 70 cm. press, of Hg.
Answer:

Question 3.
The volume at s.t.p. occupied by a gas ‘Z’ originally occupying 1.57 dm³ at 310.5K and 75 cm. press. of Hg.
Answer:

Question 4.
The volume at s.t.p. occupied by a gas ‘Q’ originally occupying 153.7 cm³ at 287K and 750 mm. pressure [vapour pressure of gas ‘Q’ at 287K is 12 mm of Hg.]
Answer:

Question 5.
The temperature to which a gas ‘P’ has to be heated to triple it’s volume, if the gas originally occupied 150 cm³ at 330K [pressure remaining constant].
Answer:

Q.6. Fill in the blanks with the correct word, from the words in bracket :

Question 1.
If the temperature of a fixed mass of a gas is kept constant and the pressure is increased, the volume correspondingly ____ [increases / decreases]
Answer:
If the temperature of a fixed mass of a gas is kept constant and the pressure is increased, the volume correspondingly decreases.

Question 2.
If the pressure of a fixed mass of a gas is kept constant and the temperature is increased, the volume correspondingly ____ [increases / decreases]
Answer:
If the pressure of a fixed mass of a gas is kept constant and the temperature is increased, the volume correspondingly increases.

Question 3.
1 dm3 of a gas is equal to ____ [1 litre / 100ml. / 100 cc.]
Answer:
1 dm3 of a gas is equal to 1 litre.

Question 4.
All the temperature on the kelvin scale are in ____ figures [negative positive]
Answer:
All the temperature on the kelvin scale are in positive figures.

Question 5.
At -273°C the volume of a gas is theoretically ____ [272 cc. / 0 cc. / 274 cc.]
Answer:
At -273°C the volume of a gas is theoretically 0 cc.

New Simplified Chemistry Class 9 ICSE Solutions The Periodic Table

ICSE SolutionsSelina ICSE SolutionsML Aggarwal Solutions

Viraf J Dalal Chemistry Class 9 Solutions and Answers

Simplified ChemistryPhysicsChemistryBiologyMathsGeographyHistoryCivics

Exercise

Question 1.(2000)
State the number of elements in Period 1, Period 2, and Period 3 of the Periodic Table.
Answer:
Period 1 is the shortest period has 2 elements Period 2 and 3 are short periods have 8, 8 elements.

Question 2.(2000)
Name the elements in Period 1.
Answer:
Names of elements are (H) Hydrogen and Helium (He).

Question 3.(2000)
What is the common feature of the electronic configuration of the elements at the end of Period 2 and Period 3 ?
Answer:
The elements at the end of period have 8 electrons in valence shell.
Ne = 2, 8
Ar = 2, 8, 8

Question 4.(2000)
If an element is in Group 17 [or Group 7A] is it likely to be metallic or non- metallic in character ?
Answer:
Element of group 17 or 7A has 7 electrons in valence shell hence non-metal.

Question 5.(2000)
Supply the missing word from those in brackets: If an element has one electron in its outermost energy level [shell] then it is likely to be [metallic / non- metallic].
Answer:
Supply the missing word from those in brackets: If an element has one electron in its outermost energy level [shell] then it is likely to be metallic.

Question 1.(2001)
Copy and complete the following sentences choosing the correct word or words from those given in bracket at the end of each sentence :
The similarities in the properties of a group of elements is because they have the same ____ (electronic configuration, number of outer electrons, atomic numbers.)
Answer:
The similarities in the properties of a group of elements is because they have the same electronic configuration.

Question 1.(2002)
What is meant by a Group in the Periodic Table.
Answer:
Groups : “The verticle columns from top to bottom are called groups.”

Question 2.(2002)
How many elements are there in Period 2.
Answer:
There are 8 elements in period 2.

Additional Questions

Question 1.
What are elements. Give a reason why elements need to be properly classified.
Answer:
Elements : “Are pure substances made up of one kind of atoms and cannot be broken down into simpler substances.” 118 elements are there.
Reason of classification of elements :

1. It makes the study of elements in an organized manner.
2. To help and correlate the properties of elements with the fundamental properties of different states of matter.
3. To help to define the relationship of one element with others.

Question 2.
How did the early chemists classify elements.
Why was this basis of classification discarded for future classification of elements.
Answer:
Early chemists classified the elements on the bases on their properties, nature, character, valency and whether they were metals or non-metals.
Basis of rejection :

1. Some characteristics being considered, varied under different conditions.
2. Certain elements resembled metallic as well as non-metallic both.
3. It did not serve the purpose of classification.

Question 3.
What was Dobereiner’s basis,of classification of elements.
State Dobereiner’s Law of Triads with suitable examples.
Answer:
Basis of classification was that “there was a close relationship between the properties of an element and its atomic mass”.
Dobereiner’s Law of Triads : “When elements are arranged in order of increasing atomic masses, groups of three elements (known as triads), having similar chemical properties are obtained. The atomic mass of the middle element of the triad being equal to the arithmetic mean of the atomic masses of the other two elements.”
Examples :

1. All of them are metals. Average wt. of sodium 7+39/2 = 46/2 = 23
2. All have valency 1.
3. They react with water form alkalis and hydrogen gas.

Halogen Group :

1. All of them are non-metals
2. All have valency one (1).
3. Each reacts with water to form acid.

Question 4.
Explain why Dobereiner’s method of classilication of elements did not hold much weightage for future classification.
Answer:
Reasons for discarding the law of triads :

• Only three triads could be formed and was not holding for other elements.
• The law was not applicable even in the same family.

Question 5.
What was the basis of classification proposed by Newland.
State Newland’s Law of Octaves with a suitable example.
Answer:

• Newland – Basis of classification : If elements are arranged in ascending order of their atomic masses every eighth element had properties similar to the properties of first element, just as eighth note of a musical octave is same as its first note.
• Newland’s Law of Octaves : “When elements are arranged in order of increasing atomic masses, the properties of every eighth element (starting from any element) are repetition of the properties of that starting element.

Examples :
Li (first element)
Na (eighth element)
Ca
Properties of sodium are similar to lithium and that of Ca are similar to sodium.

Question 6.
In which way was Newland’s Law of Octaves comparable to a musical note.
Answer:
As in above Question 5.

Question 7.
How did Mendeleeff arrange the elements in the periodic table.
What was the basis of his classification.
Answer:
Mendeleeff arranged the elements in the increasing order of their atomic weights. Basis of classification was that properties of elements are periodic functions of their atomic weights.
i.e. “Elements with similar properties appear at regular intervals.”

Question 8.
State Mendeleeff’s Periodic Law.
How did Mendeleeff’s arrangement of elements correlate with periodicity of properties of elements.
Answer:
Mendeleeff’s Periodic Law : “Physical and chemical properties of elements are periodic function of their atomic masses.”
Every ninth element resembles in properties with first element and tenth element with second and so on. That is after regular interval. This means elements in the same vertical column the properties resemble.

Question 9.
State the contributions made by Mendeleeff towards the periodic table.
Answer:
Contribution made by Mendeleeff :

1. Elements were arranged in increasing order of atomic weights in horrizontal rows called periods in vertical columns called groups.
2. Elements which are similar w.r.t. their chemical properties are grouped together and have atomic weights of nearly the same value.
3. Elements in the same group had same valency and similar chemical properties.
4. The properties of undiscovered elements left in the vacant gaps was predicted.
5. In correct atomic weights of some of the arranged elements were corrected with the knowledge of the atomic weights of the adjacent elements.
6. Number of gaps were left in the table for element not known at that time. Even predicted their properties by studying the properties of neighbouring elements. These elements were later on discovered and their properties were found to be corrects.

Question 10.
What were the defects and anomalies in Mendeleeff’s Periodic Table and how are they resolved by Moseley.
Answer:
Anomalies and Defects in Meneleeff’s Periodic Table :

1. Wrong order of atomic masses of some elements could not be explained on the bases of chemical properties cobalt At. mass 58.9 (higher mass) comes first and Nickel (Ni) with lower mass (58.7) comes latter. Argon (Ar) with At. mass (39.9) precedes K (39.1).
2. Position of isotopes i.e. atoms of the same element with similar chemical properties but different At. masses should have been given different places (According to Mendeleev), but were not placed separately.
3. Correct position could not be assigned to Hydrogen in the periodic table. Hydrogen resembles with alkali metals of group I. Also Hydrogen resembles Halogen of group 17.
4. Elements which are chemically similar such as gold and platinum have been placed in separate groups..
5. It does not explain the electronic arrangement of elements.
6. Elements such as copper and silver have no resemblance with alkali metals [Na, Li] but they have been placed together in first group.
   Moseley arranged the elements in increasing order of atomic number and most of the defects in Mendeleev’s periodic table dissappeared.

Question 11.
How were elements arranged in the long form of the periodic table or the Modern Periodic Table. State the Modern Periodic Law and compare it with Mendeleeff’s Periodic Law.
Answer:
Long form of periodic table elements were arranged in increasing order of atomic number.
Modern Period Law : “Properties of elements are periodic function of their atomic number”.
Comparison :
Modern Periodic Law :

1. Basis is “properties of elements depend on electronic configuration.” Elements are arranged in increasing atomic number.

Mendeleeff’s Periodic Law :

1. Basis is “properties of elements are periodic function of atomic mass.”

Question 12.
What are periods in a periodic table.
What is meant by ‘period number’. What does it signify.
Answer:
Periods : “Horizontal rows from left to right in a periodic table are called periods.”
Period number : “Period numbers 1, 2, 3, 4, 5, 6, 7 signifies – no of electron shells of an element.”
i.e. period no. 2 signifies second shell from the nucleus,

Question 13.
Name the period which is the shortest period and state the number of elements present in it. State the number of elements in the second and the third periods of the periodic table.
Answer:
Shortest period is period 1.
Period 1 has 2 elements
Second period has 8 elements.
Third period has 8 elements.

Question 14.
(i) Name the elements in the correct order of their increasing atomic numbers present in the first, second and third short periods of the periodic table.
State which of the elements are – (a) metallic (b) non-metallic (c) noble gases in each of the periods 2 and 3.
(ii) State the property trends in general of elements on moving from left to right in a period of the periodic table.
Answer:

(a) Metals : Li, Be, Na, Mg, Al
(b) Non-metals : C, N, O, F, P, S, Cl
(c) Noble – gases : Ne, Ar

(ii) Property trends of elements from left to right in a period

(a) Number of electron shell remains the same.
(b) Valence electrons increases by one.
Transition from metallic to non-metallic character increases.
(c) Size of atom decreases.

Question 15.
What are groups in a periodic table.
State the property trends in general of elements on moving down in a group of the periodic table.
State the characteristics which remain similar on moving down in a group of the periodic table.
Answer:
Groups : “Vertical columns in a periodic table are called groups.”
Trends in a group :

1. Number of shells and valence electrons : As we move down in a group the number of shells increase and valence electrons in a group remains same.
2. Valency : Valency of all elements is same.
3. Properties of elements : (a) Metallic character increases, (b) Non-metallic character decreases
4. Atomic size increases.
   

Character which remains similar is physical and chemical properties, i.e. The alkali metal (Group 1-A) are very reactive Li < Na < K, Rb < Cs reactivity increases. The Halogen atoms also very reactive.

Question 16.
Write short notes on the following types of elements –

(a) alkali metals
(b) alkaline earth metals
(c) halogens
(d) noble gases
(e) transition and inner transition elements
(f) normal elements.

Answer:

(a) Alkali metals : Elements of I-A group like Li, Na, K, Rb, Cs, fr are called alkali metals because these metals react with water to form alkalis and hydrogen.
(b) Alkaline earth metals : Elements of group II-A Be, Mg, Ca, Sr, Ba, Ra are called alkali earth metals as oxides of all these metals are alkaline in nature.
(c) Halogens : Elements of groups 17 (F, Cl, Br, I, At.) are called Halogens as they react with water to form acids.
(d) Noble gases : Element of 18th or O group (He, Ne, Ar, Kr, Xe, Rn) are called noble gases as they are inert gases and have valency 6. All are mono-atomic.
(e) Transition and inner transition elements

• Transition elements : d-block elements (40 elements) which are included in groups 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12 [IB to VII-B and VIII] are called transition elements as most of them exhibit variable valency and are used as catalysts.
• Inner transition elements : f-block elements i.e. Lathanides and actinides are called inner transition elements.

(f) Normal elements OR Representative elements : “Elements of groups 1, 2, 13, 14, 15, 16, 17 which have all inner shells completely filled with electrons. Only outermost shells are incomplete are called normal elements.”

Question 17.
‘The periodic table contains elements methodically grouped together’.
State the main helpful features of the long form of the periodic table.
Answer:
The periodic table contains elements methodically grouped together like [I-A] alkali group [II-A] alkali earth group, halogens [17 OR 7-B] group, normal elements. Transition elements, noble gases [O-group] etc. If we know the properties of one member of a family, we know the properties of other members of the same family will be similar.
Moreover long form of the periodic table has 7 periods horizontal rows and 18 groups vertical columns and properties show regular gradations in periods and also in groups, i.e.

Question 18.
Name or state following with reference to the elements of the first three periods of the periodic table.

(a) The noble gas having duplet arrangement of electrons.
(b) The noble gas having an electronic configuration 2, 8, 8.
(c) A metalloid in period 2 and in period 3.
(d) The number of electron shells in elements of period 1, period 2 and period 3.
(e) The valency of elements in group 1 [IA].
(f) The group whose elements have zero valency.
(g) An alkaline earth metal in period 3.
(h) The non-metallic element present in period 3 other than sulphur and chlorine.
(i) A non-metal in period 2 having electronic configuration 2, 6.
(j) An electrovalent compound formed between an alkali metal and a halogen.
(k) A covalent compound formed between an element in period 1 and a halogen.
(l) An alkali metal in period 3 which dissolves in water giving a strong alkali.
(m) A metal in period 3 having valency 3.
(n) The bridge elements of period 3 of group 1 [IA], 2 [IIA] and 13 [IIIA].
(o) The periods which contain the inner transition elements.
(p) The formula of the hydroxide of the element having electronic configuration 2, 8, 2.
(q) The valency of the element in period 3 having atomic number 17.
(r) A non-metal in period 2 which is tetravalent.

Answer:

(a) Element is Helium
(b) Argon is a noble gas
(c) Metalloid in period 2 is B (Boron) and in period 3 is Si (Silicon)
(d) In period 1 number of electron shells is [1] In period 2 is [2], in period 3 is [3]
(e) Valency of elements in group [I-A] is +1
(f) Group is 18th or O group
(g) An alkaline earth metal in period 3 is Mg [Magnesium]
(h) is P [Phosphorus]
(i) is O [Oxygen]
(j) [NaCl] Sodium chloride
(k) HCl [Hydrogen chloride]
(l) Alkali metal is Na dissolves in water gives strong alkali NaOH
(m) Metal in period 3 is A1 [Aluminium]
(n) Elements of period 3 are Mg, A1 and Si which form bridge relationship i.e. diagonal relationship.

(o) Inner transition elements in 6th period and 7th period.
(p) Hydroxide of element [2, 8, 2] is Mg [OH], where Mg is the element
(q) Valency is 1 negative [-1]
(r) Tetravalent non-metal is carbon [C]

Periodic Table – Unit Test Paper 5

Q.1. Select the correct answer from the words in bracket.

1. He arranged elements in increasing order of atomic numbers. [Dobereiner / Moseley / Mendeleeffj
2. Is a metal in period 2 having electronic configuration 2, 1. [Beryllium / Lithium / Sodium]
3. Is a period having elements from atomic no. 11 to 18. [period – 1 / period – 2 / period – 3]
4. The most reactive halogen from group 17. (chlorine / fluorine / bromine)
5. Is the group number of the element whose atomic number is 4. [group -1 / group – 2 / group – 18]

Q.2. Fill in the blanks from the words A to F given below’.
A: Decreases
B: Increases
C: Remains same
D: Increases by one
E: Electropositive
F: Electronegative
Answer:

1. Across a period from left to right in the Modern Periodic Table.
   No. of electron shells remains same; No. of valence electrons increases by one : Electronegativity increases Character of elements changes from electropositive to electronegative.
2. Down a group in the Modern Periodic Table.
   No. of electron shells increases by one; No. of valence electrons remains same; Electronegativity decreases Character of elements changes from electronegative to electropositive.

Q.3. Give reasons for the following.
1. Mendeleeff’s contributions to the periodic table, laid the foundation for the Modern Periodic Table.
2. Properties of elements are periodic functions of their atomic numbers and not atomic weights.
3. A transition from metallic to non-metallic elements is seen on moving from left to right in a period of the periodic table.
4. Noble gases do not form compounds readily.
5. Group 1 [IA] elements are called alkali metals.
Answer:
1. Mendeleeff found that there was close relation between atomic weights and properties of elements. He also found that when elements are arranged in increasing order of their atomic weights, the properties repeat at regular interval.
He even left some gaps for the elements not know at that time whose properties resembled the neighbouring elements. Predicted properties of eKa Aluminium (before discovery) are exactly the same as actual properties of gallium element (after discovery).
Moreover he arranged all the 63 elements known at that time. Still there were many draw-backs in this periodic table, but Mendeleeff laid the foundation of modern periodic table.
2. Atomic number : Number of electrons in an atom, it helps in arranging the elements according to their electronic configuration so the bases of atomic number is better than atomic mass.
3. In a period valence electrons increase from 1 to 7. But shell remains the same. Element having valence electron 1, 2, 3 lose electrons and are metals on the left. And elements having 5, 6, 7 valence electrons gain electrons and become non – metals are on the right of periodic table.
Hence transition from metallic to non-metallic character is seen from left to right a period.
4. Noble gases do not form compounds readily as noble gases have their outer most orbit completely filled and have stable configuration.
5. Elements of I-A group are metals, react with water to form alkalis.

Q.4. State the following.

1. The grotip to which the -element with electronic configuration of 2, 8, 2 belongs.
2. The group from the groups 1 [IA], 2 [IIA], 16 [VIA] and 17 [VIIA] whose elements are most electronegative.
3. The group which contains highly electropositive’ metals including sodium.
4. The group whose elements are unreactive or inert.
5. The group which contains highly reactive electronegative non-metals including chlorine.

Answer:

1. 1. Element [2, 8, 2] belongs to group 2 or [II-A]
   2. Most electronegative 17 [VII-A].

1. 1 [I-A] group contain highly electropositive metals.
2. Inert elements group 18 [0 group].
3. Highly reactive group 17 [VII-A],

Q.5. Match the elements of List-I with their type from List-II.

Answer:

Q.6. Complete the table pertaining to the following elements given in Column 1.

Answer:

New Simplified Chemistry Class 9 ICSE Solutions Study of The First Element Hydrogen

ICSE SolutionsSelina ICSE SolutionsML Aggarwal Solutions

Viraf J Dalal Chemistry Class 9 Solutions and Answers

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Equation Worksheet

Complete and balance the equations
Hydrogen



Answer:

Exercise

Question 1.(1984)
Name an element whsich reacts violently with water at room temperature.
Answer:
Element reacts with water violently at room temperature is potassium.

Question 2.(1984)
What do the following symbols [or formula] denote : 2H ; H₂ ; H⁺. [two atoms, molecule, ion]
Answer:
2H donotes — 2 atoms of hydrogen
H₂ — a molecule of hydrogen
H⁺ — an ion

Question 3.(1984)
Write correctly balanced equation for the following “word equation” :
calcium + water → calcium hydroxide + hydrogen
Answer:
Ca + 2H₂O → Ca[OH]₂ + H₂

Question 4.(1984)
When steam is passed over red-hot iron, magnetic oxide of iron and hydrogen are obtained. “The reaction between steam and red-hot iron is a Reversible Reaction.” What is meant by this statement.
Answer:
Statement means magnetic oxide of iron and hydrogen recombine to form iron and steam.

Question 5.(1984)
How can you obtain hydrogen from sodium hydroxide [not by electrolysis].
Answer:
When powdered zinc reacts with sodium hydroxide it produces hydrogen.

Question 1.(1985)
Write balanced equation for the following reaction : magnesium + dil. hydrochloric acid →
Answer:
Mg + 2HCl (dil.) → MgCl₂ + H₂

Question 1.(1986)
Name a gas which burns in air or oxygen forming water.
Answer:
The gas that burns in oxygen to form water is hydrogen.

Question 2.(1986)
Write correctly balanced equation for the following : When steam is passed over red hot iron.
Answer:

Question 3.(1986)
Explain the following : Two jars of H₂ are collected – “one burns quietly and the other does not”.
Answer:
‘One bums quietly’ is pure hydrogen and ‘which does not bum quietly’ is mixed hydrogen with air it explodes on burning.

Question 1.(1987)
Write correctly the balanced equation for the following : ‘When zinc filings are added to a concentrated solution of sodium hydroxide’.
Answer:

Question 2.(1987)
Describe one chemical test applied to the following gases, which would enable you to distinguish between them :‘carbon monoxide and hydrogen’.
Answer:

Question 1.(1988)
Write down the “word equation” for the following reaction : sodium hydroxide solution + zinc →
Answer:
‘Word equation’
Sodium hydroxide + zinc → Sodium zincate + hydrogen

Question 2.(1988)
Explain briefly how hydrogen is manufactured on a large scale, from steam.
Answer:
Manufacture of H₂ from steam :
From natural gas : Natural gas is obtained from petroleum wells and consists mainly CH₄ (methane). It is mixed with steam at 30 atm. and passed over heated nickel 800°C, when water gas is formed.

Separation of CO : The gaseous mixture of [CO and H₂] is passed through ammonical cuprous chloride solution in order to dissolve uncombined CO.

Thus hydrogen gas is left over.

Question 1.(1989)
State the products of the reaction “when steam is passed over red-hot iron”.
Answer:

When steam is passed over red hot iron, teriferric tetroxide and hydrogen are the ; products formed.

Question 1.(1990)
How can you obtain hydrogen from a mixture of hydrogen and carbon monoxide.
Answer:
See Q. 2. 1988. By dissolving the mixture in ammonical cuprous chloride. CO dissolves and H₂ is left over.

Question 2.(1990)
What do you observe when a piece of sodium is dropped into cold water ?
Answer:
When a piece of sodium metal dropped in cold water we observe : Sodium floats on water surface melts forming a silvery globule which darts about the surface of water catches fire and burns with golden yellow flame.
Bubbles of hydrogen evolve and solution is soapy, slightly warm (alkaline) hence turns red litmus blue.

Question 3.(1990)
Give reasons for the following : ‘Though hydrogen is lighter than air, it is not collected by the downward displacement of air’.
Answer:
Though hydrogen is lighter than air it is not collected by displacement of air as it forms explosive mixture with air.

Question 4.(1990)
Complete the following word equations :

1. Sodium hydroxide + zinc → hydrogen + ____
2. Calcium + water → calcium hydroxide + ____

Answer:

1. Sodium hydroxide + zinc → hydrogen + sodium zincate
2. Calcium + water → calcium hydroxide + hydrogen

Question 1.(1991)
How would you obtain ‘hydrogen from sodium hydroxide’ solution other than by electrolysis ?
Answer:
By adding aluminium powder in cone, sodium hydroxide.

Question 1.(1992)
Complete and balance the following equations :
Al + NaOH + ____ → ____ + ____
Answer:

Question 2.(1992)
What do the following symbols represent : 2H and H₂.
Answer:
2H → represents → 2 atoms of hydrogen
H₂ → represents → 1 molecule of hydrogen

Question 1.(1993)
Write balanced equation of the reaction in the preparation of : hydrogen from a solution of potassium hydroxide [other than by electrolysis].
Answer:
Preparation of H₂ from potassium hydroxide

Question 2.(1993)
Describe briefly, with equations, the Bosch Process for the large scale production of hydrogen.
Answer:
Bosch process for large production of H₂ :

Question 3.(1993)
Account for the following facts :

1. Though lead is above hydrogen in the activity series, it does not react with dilute hydrochloric acid or dilute sulphuric acid.
2. Potassium and sodium are not used to react with dilute ‘hydrochloric acid or dilute sulphuric acid in the laboratory preparation of hydrogen.

Answer:

1. Pb (lead) is above hydrogen in reactivity series. With dil. HCl, Pb forms PbCl₂ and with dil. H₂SO₄, Pb forms PbSO₄ both PbCl₂ and PbSO₄ are insoluble and forms respective coating to stop further reaction.
2. Sodium and potassium with dil. HCl or dil. H₂SO₄ but reaction is highly explosive and practically not-feasible.

Question 1.(1994)
Place the metals calcium, iron, magnesium and sodium in order of their activity with water, placing the most active first. Write the equation for each of the above metals which react with Water.
Answer:

Question 2.(1994)
Why is copper not used to prepare hydrogen by the action of dilute hydrochloric acid or dilute sulphuric acid on the metal, [copper [Cu] below hydrogen – no reaction]
Answer:
Cu is below hydrogen in reactivity series and cannot displace H₂ from acid and no-reaction takes place.

Additional Questions

Question 1.
State the electronic configuration of hydrogen [at. no. 1].e
Give a reason why hydrogen can be placed in group 1 [1A] and group 17 [VIIA] of the periodic table.
Answer:
Electronic configuration of hydrogen is 1 i.e. it has le, IP in nucleus. Hydrogen can be placed in group 1[1A] as it forms a positive ion as in HC1 like alkali metals H

Question 2.
Give the general group characteristics applied to hydrogen with respect to similarity in properties of hydrogen with –

(a) alkali metals of group 1 [IA]
(b) halogens of group 17 [VIIA].

with special reference to valency electrons & ion formation.
Answer:
General group characteristics of
(a) Alkali metals of group 1 [IA] : Alkali metals lose electron to become electro positive ion
Na – le^– → Na¹⁺
H – le^– → H¹⁺
(b) Hydrogen gains 1 electron like halogens to be come electronegative ion
Cl + le^– → Cl^1-
H + le^– → H^1-

Question 3.
How did the name ‘hydrogen’ originate. How does hydrogen occur in the combined state.
Answer:
Hydrogen means in Greek – water former. Hydrogen initially called inflammable gass that it bums in air. It produces water. Lavoisier in 1783 established its name ‘hydrogen’ meaning water producer hydrogen in combined state : in a animal and plant tissues. As a constituent of : proteins, carbohydrates, fats, acids, alkalis, petrolium products and organic substances.

Question 4.
Give balanced equation for obtaining hydrogen from cold water using –

(a) A monovalent active metal
(b) A divalent active metal

Answer:
(a) Monovalent active metal Na :
2Na + 2H₂O → 2NaOH + H₂
(b) Divalent active metal Ca :
Ca + 2H₂O → Ca[OH]₂ + H₂

Question 5.
Give balanced equations for obtaining hydrogen from ?

(a) Boiling water using a divalent metal
(b) Steam using a trivalent metal
(c) Steam using a metal – and the reaction is reversible.

Answer:
To obtain H₂ using boiling water and
(a) A divalent metal Mg
Mg + H₂O → MgO + H₂
(b) Steam and trivalent metal Al
2Al + 3H₂O → Al₂O₃ + 3H₂
(c) Steam and a metal – a reaction is reversible

Question 6.
State why hydrogen is not prepared in the laboratory by the action of –

(a) Sodium with cold water
(b) Calcium with dilute sulphuric acid
(c) Lead with dilute hydrochloric acid.

Answer:

(a) As sodium is very reactive and reacts with water violently, darts on the surface of water and collection of hydrogen becomes difficult.
(b) The reaction between calcium and dil. H₂SO is highly explosive and practically not feasible.
(c) With lead dil. HCl forms PbCl₂ which is in soluble and forms coating and stops the further reaction.

Question 7.
Give balanced equations for the following conversions.

(a) Sodium zincate from zinc
(b) Sodium plumbite from lead
(c) Sodium aluminate from aluminium.

Answer:

Question 8.
In the laboratory preparation of hydrogen from zinc and dil. acid. Give reasons for the following :

(a) The complete apparatus is air-tight.
(b) Dilute nitric acid is not preferred as the reactant acid.
(c) The lower end of the thistle funnel should dip below the level of the acid in the flask.
(d) Hydrogen is not collected over air.

Answer:

(a) The gas is highly inflammable, any leakage can cause explosion.
(b) Hydrogen produced is oxidised to water as nitric acid is powerful oxidizing agent.
(c) Otherwise gas produced will escape through thistle funnel.
(d) Air forms an explosive mixture with H₂.

Question 9.
‘Magnesium reacts with very dilute nitric acid at low temperatures liberating hydrogen.’ Give reasons.
Answer:
Magnesium reacts with dil. HNO₃ at low temperature liberating H₂, since oxidising action of dil. HNO₃ is much reduced due to dilution.

Question 10.
State the conditions and give balanced equations for the conversion of – (a) coke to water gas, (b) water gas to hydrogen – in the Bosch process.
Answer:
Conditions and balance equations for conversion of

Question 11.
How are the unreacted gases separated out in ‘Bosch process’ in the manufacture of hydrogen.
Answer:
See Q. 2 Step (iii) and (iv) 1993
CO₂ is removed by dissolving mixture in water under pressure or in caustic potash solution to dissolve CO₂. CO is removed by dissolving mixture in ammonical CuCl solution. CuCl + CO + 2H₂O → CuCl. C0.2H₂O.

Question 12.
Compare the combustibility of –

(a) pure hydrogen
(b) hydrogen-air mixture.

Answer:
Combustibility of :
(a) Pure hydrogen : Bums quietly in air with a pale blue flame forming water
2H₂ + O₂ → 2H₂O
(b) Hydrogen air mixture : It explodes on burning.

Question 13.
State the reactant added to hydrogen to obtain the respective product in each case.

(a) Ammonia
(b) Hydrogen chloride
(c) Water
(d) Hydrogen sulphide

Answer:
Reactant added to obtain :

(a) Ammonia : Nitrogen and hydrogen.
(b) Hydrogen chloride : Hydrogen and chlorine.
(c) Water : Hydrogen and oxygen.
(d) Hydrogen sulphide : Hydrogen and sulphur.

Question 14.
State the use of hydrogen –

(a) As a fuel
(b) In hydrogenation of oil & coal
(c) In extraction of metals

Answer:
Use of hydrogen :
(a) As a fuel in the form of coal gas, water gas, liquid hydrogen.
(b) In hydrogenation of oil & coal :

1. Vegetable oil [palm oil] turns to [vegetable ghee] semisolid fats by hydrogenation in presence of nickel catalyst at 170°C.
2. Passage of H₂ under high pressure in presence of catalyst and at suitable temperature produces a product similar to petroleum.

(c) In extraction of metals : by reducing oxides of metals.

Question 15.
Explain the terms – oxidation and reduction in terms of addition and removal of oxygen/hydrogen with suitable examples.
Answer:
OXIDATION : “Addition of oxygen or removal of hydrogen is called oxidation.

Here H₂S is oxidised to sulphur as removal of H₂ takes place.
Reduction : Addition of hydrogen or removal of oxygen is called reduction.

chlorine is reduced, as addition of hydrogen.

Question 16.
Explain the term – redox reaction with an example involving – the reaction of hydrogen sulphide with chlorine.
Answer:
Redox Reaction : “The reaction in which oxidation and reduction takes place simultaneously is called Redox Reaction.”

Removal of hydrogen is oxidation. Here H₂S is oxidised to sulphur whereas addition of hydrogen is reduction.
Here chlorine is being reduced to hydrogen chloride

Question 17.
State what are – oxidising and reducing agents. Give examples of oxidising and reducing agents in the gaseous, liquid and solid form. Give two tests each generally answered by oxidising and reducing agents respectively.
Answer:

Test For An Oxidising Agent :

1. On heating strongly, oxidising agent liberates oxygen and to test oxygen : O₂ gas rekindles a glowing splinter.
2. On warming oxidising agent with cone. HCl, it liberates chlorine, that bleaches moist litmus paper.

Test For A Reducing Agent :

1. When reducing agent is warmed with HNO₃, BROWN FUMES of N02 are given out.
2. Reducing agent, decolourises the pink colour of (KMnO₄) dil. potassium permanganate solution.

Hydrogen – Unit Test Paper 6

Q.1. Select from A to G the reactant added, to give the products 1 to 5, in the preparation of hydrogen gas.
A : dilute acid
B : dilute alkali
C : cold water
D : cone, alkali
E : boiling water
F : cone, acid
G : steam

Answer:

Q.2. Give balanced equations for the following conversions, 1 to 5.

Answer:

Q.3. Give reasons for the following.

1. Nitric acid in the dilute form is not used in the laboratory preparation of hydrogen from metals.
2. Granulated zinc is preferred to metallic zinc in the preparation of hydrogen using dilute acid.
3. Hydrogen and alkali metals of group 1 [IA] react with copper [II] oxide to give copper.
4. Hydrogen is collected by the downward displacement of water and not air even though – it is lighter than air.
5. A mixture of hydrogen and chlorine can be separated by passage through a porous pot.

Answer:

1. Nitric acid is a strong oxidising agent and nascent oxygen formed oxidises hydrogen produced to water.
2. Zinc granules are preferred rather than pure zinc as impurity copper present in it has a catalysing effect and speeds up the rate of the reaction.
3. Hydrogen and alkali metals act as reducing agent and reduce CuO to Cu.
   CuO + H₂ → Cu + H₂O
4. H₂ is insoluble in water. Moreover H, forms explosive mixture with air.
5. H₂ and Cl₂ differ in densities H₂ = 0.9 g/lit
   Cl₂ = 35.5 g/lit
   H₂ diffuses faster than Cl₂ and is separated.

Q.4. Name the following.

1. A metal below iron but above copper in the activity series of metals which has no reaction with water.
2. A metal which cannot be used for the preparation of hydrogen using dilute acids.
3. The salt formed when aluminium reacts with potassium hydroxide, during the preparation of hydrogen from alkalis.
4. A gaseous reducing agent which is basic in nature.
5. A compound formed between hydrogen and an element from group 17 [VIIA] – period 3.

Answer:

Q.5. Select the correct answer from the symbols in bracket.

1. The element placed below hydrogen in group 1 [IA]. [Na, Li, K, F],
2. The element other than hydrogen, which forms a molecule containing a single covalent bond. [Cl, N, O]
3. The element, which like hydrogen has one valence electron. [He, Na, F, O]
4. The element, which like hydrogen is a strong reducing agent. [Pb, Na, S, Cl]
5. The element which forms a diatomic molecule. [C, Br, S, P]

Answer:

1. Li
2. Cl
3. Na
4. Na
5. Br

Q.6. The diagram represents the preparation & collection of hydrogen by a standard laboratory method.

1. State what is added through the thistle funnel ‘Y’
2. State what difference will be seen if pure zinc is added in the distillation flask ‘X’ instead of granulated zinc.
3. Name a solution which absorbs the impurity – H₂S.
4. State why hydrogen is collected after all the air in the apparatus is allowed to escape.
5. Name a gas other than hydrogen collected by the same method.

Answer:

1. Dil. HCl
2. Rate of production of H₂ will be slow.
3. Lead nitrate solution absorbs H₂S.
4. Air with H₂ forms explosive mixture.
5. Oxygen gas.

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Exercise

Question 1.(1984)
Define the terms : atomic number, mass number and electron.
Answer:
Atomic number [Z] : “Is the number of protons in the nucleus of an atom.”
OR “Is the number of positive charges in the nucleus of its atom.”
OR “Is the number of electrons in the complete atom.”
Z = P = e
Mass number [A] : “Is the total number of protons and neutrons in the atom of an element.”
A = P + n
Electron [e] : “A sub-atomic particle having a unit negative charge and mass equal to 1/1837 of that of hydrogen atom, revolve round the nucleus.” [_-1e⁰]

Question 1.(1985)
From the symbol He for the element helium, write down the mass number and the atomic number of the element.
Answer:

• Mass number — 4
• Atomic number — 2

Question 1.(1986)
Sulphur has an atomic number of 16 and a mass number of 32. State the number of protons and neutrons present in the nucleus of sulphur.
Answer:
Protons → Atomic number = 16 = P
Number of Neutrons = A – P = 32 – 16 = 16

Question 2.(1986)
Give a simple diagram to show the arrangement of the electrons in an atom of sulphur.
Answer:
In sulphur atom.
Number of electrons = P = e = 16

Question 1.(1987)
24₁₂Mg and 26₁₂Mg are symbols of two isotopes of magnesium. Compare the ‘ atoms of these isotopes with respect to :

1. the composition of their nuclei.
2. their electronic configurations.
3. give the reason why the two isotopes of magnesium have different mass numbers.

Answer:

Question 2.(1987)
Chlorine is an element of atomic number 17. It is a mixture of two isotopes having mass number of 35 and 37.

1. What is meant by “atomic number of an element” ? What do you understand by an ‘atom’
2. Write down the electronic configuration of the chlorine atom.
3. State the number of protons, electrons and neutrons in the following isotopes: ³⁵Cl₁₇, ³⁷Cl₁₇
4. Explain why the two atoms in (iii) above have the same chemical reactions.
5. If molten magnesium chloride is electrolysed suggest a suitable electrode [anode].

Answer:

1. Atomic number is the number of protons in an atom.
   Atom : The basic unit of matter is the smallest particle of an element which can take part in a chemical reaction.
2. Number electrons in chlorine atom is 17 [2, 8, 7] and their distribution in various shells in K = 2, L = 8, M = 7.
3. Two isotopes of chlorine are
   ³⁵₁₇Cl number of protons = 17 and number neutrons = 35 – 17 = 18
   Number of electrons = 17
   ³⁷₁₇ Cl number of protons = 17
   electrons = 17
   neutrons = 20 (37 – 17)
4. Two isotopes of chlorine have same chemical reactions as their atomic number i. e. number of electrons i.e. electronic configuration is same (17).
5. When molten magnesium chloride is electrolysed suitable anode used is carbon.

Question 1.(1988)
Five atoms are labelled V to Z

1. Which one of these atoms (1) contains 7 protons ; (2) has an electronic configuration 2, 7 ?
2. Write down the formula of the compound formed by atoms X and Y.

Answer:
(i)
1. Contains 7 protons
Z – has atomic number of Z = 7
2. Has atomic configuration (2, 7)
W – atomic number = of electrons = 9 (2, 7)
(ii) X-atomic number 3
= number of electron = (2, 1)
∴ X has valency [1+]
Y-atomic number = 8
= number of electrons = (2, 6)

Question 1.(1989)
Elements X, Y, Z have atomic numbers 6, 9 and 12 respectively. Which one :

1. forms anion – negative ion ;
2. forms cation – positive ion ;
3. has 4 electrons in the outermost orbit. [Y, Z, X]

Answer:
Atomic number of
X = 6 = number of electrons (2, 4)
∴ X has 4 electrons in outermost orbit.
Y = 9 = number of electrons (2, 7)
∴ Y form anion = negative ion
Z = 12 = [2, 8, 2]
∴ Z forms cation – positive ion.

Question 2.(1989)
Ordinary chlorine gas has two isotopes : ³⁵₁₇Cl and ³⁷₁₇Cl in the ratio of 3 : 1. Calculate the relative atomic mass [atomic weight] of chlorine.
Answer:

Question 1.(1990)
The atom of aluminium is represented by ²⁷Al₁₃. Write down the number of

(a) electrons
(b) protons
(c) neutrons
(d) the arrangement of electrons

in the different orbits or shells in one atom of aluminium.
Answer:

Question 2.(1990)
Name the clement which does not contain any neutrons in its nucleus.
Answer:
Element is [H] Hydrogen.

Question 3.(1990)
Elements A, B and C have atomic numbers 9, 20 and 10 respectively.

1. State which one is (1) a non-metal, (2) a metal, (3) chemically inert._[A,B,C]
2. Write down the formula of the compound formed by two of the above elements._[BA2]

Answer:

Question 4.(1990)
Define : Proton, Electron, Neutron.
Answer:
Proton is subatomic particle found in nucleus of atom having unit positive charge, mass = 1 hydrogen atom.
Found in nucleus of atom and number of protons in an atom = number of electrons.
Electron = e = is sub atomic particle revolving around the nucleus of atom having
unit – ve charge, mass = 1/1840 of 1 hydrogen atom and number of electrons = number of protons.
Neutrons = n = is sub-atomic particle found in nucleus of atom, having no charge, mass = mass of 1 proton.

Question 1.(1991)
Write down the electronic configuration of the following (i) ²⁷₁₃X, (ii) ³⁵₁₇Y.
Write down the number of electrons in X and neutrons in Y and the formula of the compound formed by X and Y._[XY3]
Answer:

Question 1.(1992)
According to the Dalton’s Atomic Theory, atoms of the same element are identical in all respects. But according to the Modern Atomic Theory, this postulate is proved wrong. Explain.
Answer:
According to modern atomic theory atoms of the same element may not be alike in all respect.
e.g. isotopes of chlorine ³⁵₁₇Cl and ³⁷₁₇ Cl have different atomic masses 35 and 37. This proves Dalton’s Atomic Theory wrong.

Question 1.(1993)
What are isotopes.
Answer:
Isotopes : “Atoms of the same element having same atomic number, different mass number and similar chemical properties.”

Question 2.(1993)
Write down :

1. the mass number of the atom having 20 neutrons and 15 protons.
2. the number of neutrons in the nucleus of an atom having atomic number 17 and mass number 37.

Answer:

1. Mass number (A) = Neutrons (n) + Protons (p)
   Mass number = 20 + 15 = 35
2. A = n + p (At. No.) A is mass number
   37 = n + 17
   Number of neutrons = n = 37 – 17 = 20

Question 1.(1994)
What is a proton ? What is the significance of the number of protons found in the atoms of each of the different elements.
Answer:
Proton : “Is positively charged sub-particle of atom present in nucleus.”
Proton helps in determining and understanding the structure and behaviour of an atom. Elements are arranged on the bases of number of protons in their respective atoms.

Question 2.(1994)
What is the relation between the number of protons and the number of electrons in an atom.
Answer:
The number of protons in an atom is equal to the number of electrons.

Question 3.(1994)
What would be the reason for an element to have atoms with differing mass numbers.
Answer:
Atoms of same element have different mass number as these isotopes have different physical constants i.e. different physical properties. OR “Atoms have different number of neutrons.”

Question 4.(1994)
Copy and complete the following table relating to the atomic structure of some elements :

Answer:

Question 5.(1994)
The electronic structure [configuration] of fluorine can be written as 2, 7. In a similar way give the electronic configuration of :

1. aluminium
2. phosphorus

Answer:
Electronic configuration of :

1. Aluminium 13 is 2, 8, 3.
2. Phosphorus 15 is 2, 8, 5

Additional Questions

Question 1.
State the main postulates of Dalton’s atomic theory. How does the modern atomic theory contradict and correlate with Dalton’s atomic theory.
Answer:
Postulates of Dalton’s Atomic Theory :

1. Matter consist of small indivisible particles called atoms.
2. Atoms are neither created nor destroyed.
3. Atoms of the same element are alike in every respect i.e. atoms of oxygen are similar (have same At. No. Mass No. etc.) and atoms of hydrogen are similar.
4. Atoms of different elements are different in every respect i.e. atoms of hydrogen atoms are different from atoms of oxygen.
5. Atoms combine with other atoms in simple whole number ratio forming compound atom or molecules.

Contradiction by Modern Theory :

1. Atoms are no longer indivisible. Atoms are made up of subatomic particles protons, electrons and neutrons.
2. Atoms can be changed in energy.
3. Atoms of same elements may not be alike i.e. isotopes ¹²₆C ¹³₆ C ⁴₆C.
4. Atoms of different elements may be similar i.e. isobars.
5. Ratio in which atoms combine may not be simple whole number ratio i.e. ratio of atoms in C₁₂ H₂₂ O₁₁ (sugar) is not simple whole number.

Question 2.
Explain in brief the experimental proofs which led to the discovery of –

1. Electrons
2. Protons
3. atomic nucleus
4. neutrons.

Answer:
(i) Electrons : English physicist William Crookes in 1878 discovered the cathode rays

He found that when an electric discharge is passed through a tube containing a gas at low pressure (0.01 mm of Hg).
Blue rays were emitted from the negative plate [i.e. cathode]. These rays were called cathode rays and consist of negatively charged particles now called electrons. Electrons are essential part of each and every atom.
(ii) Protons : Since atoms are electrically neutral there must be equal positively charged particles in atom. This led to discovery of the protons.

Goldstein discovered protons. He used a perforated disc as cathode and found positively charged rays travelling in opposite direction to cathode rays.
These were protons i.e. positively charged particles, 1837 times heavier than an electron.
(iii) Discovery of Atomic Nucleus : Lord Rutherford in 1911 directed alpha particles (⁴₂He) towards gold foil. The deflection of the alpha particles was observed and concluded that there was nucleus at the centre of atom which was positively charged dense very small space (solid) called nucleus.

(iv) Discovery of Neutrons : As atom contains protons and electrons. Atomic mass of electron in negligible.
∴ An atom of helium which contains 2 protons should have a mass = (2 × 1) + 0 = 2 a.m.u. But the atomic mass of a helium atom was found to be approximately 4.0 a.m.u. It was therefore proved that, in the nucleus of an atom, there is an other particle called a neutron which has no electrical charge but is almost of an equal mass as the protons.

Question 3.
State in brief the drawbacks of Rutherfords atomic model correlating them with the postulates of Bohr’s atomic model.
Answer:
Drawbacks of Rutherfords’s Atomic Model :

1. The comparison of electrons with the planets in the solar system is the main drawback of Rutherford’s model. Because electrically charged particles in motion lose energy (radiate energy). As a result electrons revolving the nucleus should after losing energy, fall in spiral path and colLapse.
2. If it was so atom should be highly unstable, but we know that atom is structurally stable. Rutherford’s model could not explain this stability.
   

Postulates of Bohr’s Atomic Model :

1. Electrons revolve around the nucleus in fixed orbits or energy levels or shells.
2. While revolving around the nucleus in an orbit, an electron does not lose energy nor does it gain energy.
3. The integer ‘n’ represents the various energy levels 1, 2, 3 or k, l, m starting from the innermost.
4. An electron revolving in a particular orbit, on gaining a certain amount of energy, jumps to the next orbit and vice versa.
   These postulates explain the cause of stability of the atom.

Question 4.
What is meant by the terms :

1. subatomic particles,
2. nucleus,
3. orbits,
4. atomic number
5. mass number with reference to an atom.

Answer:

1. Subatomic particles : “Particles into which atom is divided are called subatomic particles.” Protons, electrons, neutrons.
2. Nucleus : “Subatomic part of atom which is very small as compared to the size of, atom +vely charged, massive and dense present at the centre of atom.”
3. Orbits : “Shells into which electrons revolve around the nucleus.”
4. Atomic number : “Number of protons or numbers of electrons present in an atom.” [Z]
5. Mass number [A] : “Total number of protons + number of neutrons” A = p + n.
   

Question 5.
Represent each of the following :

1. a proton ‘p’
2. an electron ‘e’
3. a neutron ‘n’ in terms of its symbols showing the subscript and superscript values.

Answer:

Question 6.
What arc ‘energy levels’ ? Explain the arrangement and distribution of electrons in the various shells with reference to an atom in general and to an atom of potassium ‘³⁹₁₉K’ with special referencence to the 2n² rule.
Answer:
Energy levels : ‘Electrons revolved the nucleus in definite orbits called shells or energy levels since while revolving in a shell an electron possesses a certain amount of energy.”
Arrangement of electrons : Direction of electrons in various shells is given by a formula 2n² where is the number
n = 1 In first shell there can be maximum of 2 × (1)² = 2 electrons.
n= 2 In second shell there can be maximum of 2 (2)² = 8 electrons.
n = 3 In third shell there can he maximum of 2 (3)² = 2 × 9 = 18 electrons and soon.
But there cannot be more than 8 electrons in the outermost shell and cannot be more than 18 electrons in penultimate orbit
Higher shell will start only when the lower shells are completely filled.
³⁹₁₉k has 19 electrons
In first shell or K shell will have 2 electrons.
2nd shell or L shell 8 electrons.
3rd shell or M shell 8 electrons
and 4th or N shell 1 electron.

Question 7.
An element ‘A’ has mass number 23 and atomic number 11. State the —

1. no. of neutrons in its shell,
2. electronic configuration of the element ‘A’.

Answer:
Mass number A = n + p (atomic number)
A = n + Z
23 = n + 11
∴ n = 23 – 11 = 12
(i) Number of neutrons = 12
(ii) e = p = 11
∴ Number of electrons e = 11
∴ Electron configuration of A = 2, 8, 1

Question 8.
The following elements U to Z are given ₃U, ₆V, ₉W, ₁₄X, ₁₈Y, ₂₀Z.
State the electronic configuration of each and state whether they arc metals, non-metals or inert gases.
Answer:

Question 9.
Draw the geometric atomic structure of each of the following atoms showing the number of electrons, protons and neutrons in each of them :

1. ¹²₆ C
2. ²³₁₁ Na
3. ³¹₁₅ P
4. ³⁹₁₉ K
5. ⁴⁰₂₀ Ca.

Answer:

Question 10.
Define an ‘isotope’. Give reasons why isotopes have same chemical but different physical properties.
Answer:
Isotopes : “Are the atoms of the same element having same atomic number but different mass number.”
Chemical properties are same because they have same atomic number hence same number of valence electrons physical properties are different because they have different mass number.

Question 11.
Draw the geometric atomic structure of the three isotopes of hydrogen and the two isotopes of chlorine.
Answer:
Geometric atomic structures of three isotopes of hydrogen.

Question 12.
Four elements A, B, C, D are given :
A shows the presence of 20 neutrons, 17 protons and 17 electrons.
B shows the presence of 18 neutrons, 17 protons and 17 electrons.
C shows the presence of 10 neutrons, 9 protons and 10 electrons.
D shows the presence of 4 neutrons, 3 protons and 2 electrons.
State which of the above is —

(a) an anion
(b) a cation
(c) a pair of isotopes.

Write the formula of the compound formed between D and C.
Answer:
For (a) and (b) and (c) we need to find the valence electrons.
Element A has 17 electrons [2, 8, 7]
As there are 7 electrons in outermost shell. It will gain 1 electron to attain stable octet structure of nearest gas — Neon [2, 8] ∴ A^1- is Anion.
Element B also has 17 electrons [2, 8, 7] is short of 1 electron in the valence shell will gain 1 electron and B^– is Anion.
Element C has 10 electrons and electronic configuration [2, 8] has 8 electrons in valence shell. Its octet is complete, but there are 9 protons.
∴ Element C must have 9 electrons
The 10th electron is gained to have stable octet structure and hence 8 electrons in valence shell.
∴ C^– is Anion.
Element D has protons 3 and electrons 2
It should have 3 electrons. This means it has lost 1 electron from valence shell to have a stable duplet structure of nearest gas (inert) [2] Helium.
∴ D+ is Cation.
(c) A and B are isotopes as both have same atomic number but different number of neutrons.
(d) D¹⁺ C^1-
∴ Compound formed is [DC]

Question 13.
What are noble gases. Give a reason why noble gases have stable electronic configuration.
Answer:
Noble gases are gases which have their outermost orbit completely filled i.e. have stable electronic configuration. They do not gain or lose or share electrons and noble gases have stable electronic configuration as they are in minimum state of energy.

Question 14.
Explain the reason for chemical activity of an atom with reference to its electronic configuration.
Answer:
Reason for chemical activity of atom is its unstable electronic configuration.

Question 15.
Differentiate between the terms —

(a) Stable and unstable electronic configuration
(b) Duplet and octet rule.

Answer:
(a) Stable electronic configuration : Atoms having octet (8 electrons) or duplet 2 electrons in outermost shell is stable electronic configuration 2 electrons is K shell the valence shell.
i.e. “Atoms having valence shell completely filled” shell.

• Unstable electronic configuration : “Atoms of element which do not have their valence shells completely filled.”

(b) Duplet Rule : For an atom to achieve stable electronic configuration it must have 2 electrons in the first shell [outermost if it is K^– shell] like nearest noble gas Helium.

• Octet Rule : For an atom to achieve stable electronic configuration it must have 8 electrons in outermost orbit like that of noble gas other than Helium.

Question 16.
Explain the octet rule for formation of —

(a) Sodium chloride from a sodium atom and a chlorine atom.
(b) Nitrogen molecule from two nitrogen atoms.

Answer:
(a) Formation of sodium chloride NaCl, Na atom has atomic number 11 [2, 8, 1], There is 1 electron in outermost M shell it is not possible to gain 7 electrons and to lose 1 electron is easier, by losing sodium atom because Na⁺ ion.
While Cl atomic number 17 [2, 8, 7] has 7 electrons in outermost M shell for Cl atom it is easier to gain 1 electron, becomes Cl^– ion.
By losing 1 electron and by gaining 1 electron both sodium and chlorine attain octet configuration and become stable and there is electrostatic force which pulls them and become sodium chloride.

Each atom of nitrogen has 5 electrons in outermost orbit. For both it is not possible to gain 3 electrons or lose 5 electrons to attain nearest inert gas electronic configuration. So they share 3 electrons each and attain octet configuration which is stable and form N₂ molecule.

Atomic Structure – Unit Test Paper 4

Q.1. Select the correct answer from the answers in brackets to complete each sentence.

Question 1.
An element has electronic configuration 2, 8, 1 and 12 neutrons. Its mass no. is ____ [11 / 23/ 12]
Answer:
An element has electronic configuration 2, 8, 1 and 12 neutrons. Its mass no. is 23.

Question 2.
The maximum number of electrons in M-shell is ____ [8 / 32 / 18]
Answer:
The maximum number of electrons in M-shell is 18.

Question 3.
Isotopes have same ____ [no. of neutrons / electronic configuration / atomic masses].
Answer:
Isotopes have same electronic configuration.

Question 4.
An ____ [atom / ion] is capable of independent existence in solution.
Answer:
An ion is capable of independent existence in solution.

Question 5.
An atom with electronic configuration 2, 7 and mass number 19 will have ____ neutrons. [8 / 10 / 12]
Answer:
An atom with electronic configuration 2, 7 and mass number 19 will have 10 neutrons.

Q.2. Give reasons for the following.

1. The physical properties of isotopes of the same element are not identical.
2. The mass number of an atom is slightly less than the actual atomic mass.
3. The shells surrounding the nucleus of an atom are also called ‘energy levels’.
4. Helium is chemical extremely unreactivc.
5. Mass number is slightly less than the actual atomic mass.

Answer:
1. Physical properties depend on mass number and isotopes have different mass numbers.
2. Atomic mass = number of protons + number of neutrons + number electrons, and mass number = protons + neutrons
∴ Atomic mass = mass number + mass of electrons
Though mass of electrons is very small 1/1837 of hydrogen atom, yet the mass of electrons is there.
∴ Mass number is slightly less than atomic mass.
3. Shells are called energy levels since while revolving in a shell an electron possesses a certain amount of energy.
4. Helium has its outermost shell completely filled i.e. 2 electrons in K shell.
5. Same as answer (2) of this question above.

Q.3. Differentiate between the following terms.

1. Electron and proton
2. Atomic number and mass number
3. Nucleus and nucleons
4. Valence shell and penultimate shell
5. Octet and duplet.

Answer:
Differentiate between :
(i) Electron :

1. e^– has unit -ve charge
2. has 1/1837 th of 1 hydrogen atom mass negligible.
3. Revolve around nucleus in different shells.

Proton :

1. ¹₁ P has unit +ve charge
2. has unit mass i.e. mass equal to 1 hydrogen atom
3. Are present in nucleus.

(ii) Atomic number :

1. Z-is the number of protons or number of electrons in an atom.
2. Present in nucleus = P
   Present in orbits = e

Mass number :

1. A is the sum of protons and neutrons.
2. Present in nucleus of atom.

(iii) Nucleus :

1. Nucleus is central space of atom. Is solid massive, positively charged, dense and very-very small as compared to size of atom.

Nucleons :

1. Is the sum of protons and neutrons, present in nucleus.

(iv) Valence shell :

1. Is the outermost shell containing free electrons.

Penutimate shell :

1. Is the shell before the valence shell.

(v) Duplet :

1. Having 2 electrons in the outermost shell -K shell (if outermost) is called duplet i.e. Helium.

Octet :

1. Having 8 electrons in the outermost shell except He is called octet i.e. Ne, Ar, Kr, Xe, Rn.

Q.4. Name or state the following.

1. The three isotopes of hydrogen.
2. Two elements having same number of protons and electrons but different number of neutrons.
3. The valency of an element whose electronic configuration is 2, 8, 3.
4. The shell closest to the nucleus of an atom.
5. An element having valency ‘zero’.

Answer:

Q.5. State the number of neutrons in each of the atoms A to E. Also state which of the atoms A to E is a metal.

Answer:
Number of neutrons = A – Z
= Mass number – Atomic number
A. Number of neutrons =1 – 1 = 0
B. Number of neutrons = 16 – 8 = 8
C. Number of neutrons = 14 – 7 = 7
D. Number of neutrons = 23 – 11 = 12 (2, 8, 1 metal)
E. Number of neutrons = 35 – 17 = 18

Q.6. Match the elements A to E in List 1 with their valencies in List 2 and with their nature in List 3.

Answer:

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Exercise

Question 1.(2005)
Compound X consists of molecules Bonding in X will be :
A : ionic
B : electrovalent
C : covalent
D : molecular.
Answer:
C : covalent

Question 1.(2006)
What is a lone pair of electrons. Draw an electron dot diagram of a hydronium ion (with lone pair).
Answer:

Question 1.(2007)
The electronic configuration of nitrogen is 2, 5. How many electrons in the outer shell of a nitrogen atom are not involved in the formation of a nitrogen molecule.
Answer:
2 electrons in the outer shell of each nitrogen atom are not involved in sharing during formation of nitrogen molecule.

Question 1.(2008)
What is the term defined in a bond formed by a shared pair of electrons each bonding atom contributing one electron to the pair.
Answer:
Single covalent bond.

Question 1.(2009)
Draw the atomic orbit structure of carbon tetrachloride and state the type of bond present in it
Answer:

Question 1.(2010)
Sodium chloride – covalent bond /ionic bond/covalent and coordinate bond.
Answer:
Ionic bond.

Question 2.(2010)
Carbon tetrachloride – covalent bond/ionic bond/covalent and coordinate bond.
Answer:
Covalent bond

Question 1.(2011)
(i) In covalent compounds, the bond is formed due to the ____ (sharing/ transfer) of electrons.
(ii) A molecule of ____ contains a triple bond, (hydrogen, ammonia,nitrogen)
Answer:
(i) In covalent compounds, the bond is formed due to the sharing of electrons.
(ii) A molecule of nitrogen contains a triple bond.

Question 1.(2014)
The molecule containing a triple covalent bond is :

(A) ammonia
(B) methane
(C) water
(D) nitrogen

Answer:
(D) nitrogen

Question 1.(2015)
Explain the bonding in methane molecule using electron dot structure.
Answer:
Formation of methane molecule – Non-polar covalent compound :

Additional Questions

Question 1.
Explain the term ‘chemical bond’ & ‘chemical bonding’?
Answer:

• Chemical bond : The linkage force which acts between two or more atoms to hold them together as a stable molecule is called a chemical bond.
• Chemical bonding : The concept of chemical bond is called chemical bonding.

Question 2.
State why noble gases have stable electronic configuration while atoms of other elements have unstable electronic configuration.
Answer:
Noble gases have stable electronic configuration because they have their outermost orbit completely filled with electrons, 2 incase of the helium and 8 electron in rest of the noble gases.
Atoms of other elements are unstable as they do not have electronic configuration of their nearest noble gas i.e. They have uncomplete valence shell and do not have their outer shell completely filled.

Question 3.
Explain in brief how atoms of other elements other than noble gases, attain stable electronic configuration of the nearest noble gas.
Answer:
Elements of atoms of other than noble gas attain stable electronic configuration by lose of electron/s incase they have 2, 3 valence electron or by gain of electron/ s incase they have 5, 6 or 7 valence electrons or by sharing of electrons.

Question 4.
Define the terms

1. Electrovalent or ionic bond
2. Electrovalent or ionic compound
3. covalent bond
4. covalent compound

with suitable examples wherever required.
Answer:

1. Electrovalent or ionic bond : The chemical bond formed between two atoms by transfer of one or more electrons from the atom.
2. Electrovalent or ionic compound : The chemical compound formed as a result of transfer of one or more electrons from atom.
3. Covalent bond : Sharing of electrons between two atoms form Covalent bond.
4. Covalent compound : The chemical compound formed due to mutual sharing of electrons between the given pairs of atoms forming a covalent bond is called covalent compound.

Question 5.
State what is meant by :

(a) duplet rule
(b) octet rule.

Answer:

(a) Duplet rule : For an atom to achieve stable electronic configuration it must have 2 electrons in its first shell if first shell is valence shell like noble gas Helium it is called duplet rule.
(b) Octet rule : For achieving stable electronic configuration an atom must have 8 electrons in it’s valence shell like noble gas other than helium.

Question 6.
Sodium chloride is formed os a result of — transfer of electrons from metallic sodium atom to non-metallic chlorine atom. With the help of atomic orbit structure diagram show the formation of sodium chloride. [Atomic numbers: Na = 11, Cl = 17]
Answer:
Sodium chlorine (NaCl)
Atomic orbit structure :

Question 7.
Atomic number of calcium is 20 and of oxygen is 8. State the number of electrons – calcium loses & oxygen gains to attain stable electronic configuration of the nearest noble gas – during formation of electrovalent molecule, calcium oxide.
Answer:
Atomic number of calcium is 20, electronic configuration 2, 8, 8, 2 has 2 valence electrons which Ca loses as it is not possible to gain 6 electron (to have octet arrangement) to have stable nearest noble gas K configuration.
Oxygen O 8=2,6 has 6 valence electrons and gains 2 electrons to attain stable nearest noble gas Ne configuration (octet)
∴ Calcium oxide CaO electrovalent molecule is formed.

Question 8.
Explain with the help of atomic orbit structure diagram the formation of calcium oxide.
Answer:
Calcium oxide (CaO)

Question 9.
Give a reason why – in the formation of electrovalent compound – magnesium chloride, one magnesium atom combines with two chlorine atoms to give magnesium chloride, (at. no.: Mg = 12, Cl = 17]
Answer:
At. no. of Mg = 12
electronic configuration 12 = 2, 8, 2
Nearest noble gas Ar = (2, 8)
Mg-atom loses 2 electrons from its valence shell and becomes stable configuration and cation Mg²⁺

But Cl at. no. 17 electron configuration 8, 7 has 7 valence electrons.
It accepts 1 electron to have stable nearest noble gas configuration of argon (2, 8, 8)
So to accept 2 electrons lost by Mg atom there must be 2 Cl-atom. Hence to form MgCl₂
One magnesium atom and two chlorine atoms are required.

Question 10.
‘Formation of hydrogen molecule takes place by sharing of electrons’. Give a reason why the molecule of hydrogen is not formed by — transfer of electrons.
Answer:
Atom of hydrogen has unstable electronic configuration and has 1 valence electron. To attain stable electronic configuration. To attain stable electronic configuration of nearest noble gas configuration i.e. 2 electrons in valence shell (Duplet rule) share one electron.
Both atoms of hydrogen in hydrogen molecule are short of 1 electron in valence shell and to have stable electronic configuration of nearest noble gas they share 1 electron each in valence shell and no transfer of electron takes place in that case H-atom will cease to exist.

Question 11.
Draw the atomic orbit structure diagram for formation of —

(a) hydrogen molecule
(b) chlorine molecule
(c) oxygen molecule
(d) nitrogen molecule [at.no.: H = 1, Cl = 17, O = 8, N = 7]

Answer:

Question 12.
Give reasons for formation of a single covalent bond between two chlorine atoms and a double covalent bond between two oxygen atoms – during formation of a covalent chlorine molecule and formation of-a covalent oxygen molecule, respectively.
Answer:
At. no. of Cl-atom is 17
17 = 2, 8, 7
each atom has 7 valence electrons which one atom cannot lose 7 electron to other nor it can accept one electron from other Cl atom as it will be left will 6 electrons. Hence both atoms of chlorine share 1 electron and both have 8 electrons in valence shell i.e. both have stable configuration and since both share 1 electron Cl – Cl forms single covalent bond.
O 8 = 2,6 each atom of oxygen atom has 6 valence electrons and short of 2 electron
to have stable octet. Both of the atoms cannot lose 6 electrons nor gain 2 electrons as explained above. Hence both atoms share 2 electrons each and each has 8 electrons revolving in valence shell and forms oxygen molecule with double covalent bond.

Question 13.
Draw the atomic orbit structure diagram for formation of a —

(a) water molecule
(b) ammonia molecule [at.no.:H = 1, O = 8, N = 7]

Answer:
Hydrogen atom has 1 electron in first valence shell and O – atom has 6 electrons in valence shell

Question 14.
State why water has two lone pairs of electrons in its covalent molecule while ammonia has one lone pair.
Answer:

In water there are two shared as there are 2H-atoms pair and two lone pairs.
In ammonia there are three shared pairs as there are 3H-atoms and one lone pair.

Question 15.
With the help of ail atomic orbit structure diagram – explain the formation of —

(a) carbon tetrachloride
(b) methane, [at. no.: C = 6, Cl = 17, II = 1]

Answer:
Atomic orbit structure diagram of
(a) CCl₄ (carbon tetrachloride)
C atomic number 6 = 2, 4

Question 16.
Give a reason why one atom of carbon — shares four electron pairs, one with each of the four atoms of chlorine during the formation of covalent molecule — carbon tetrachloride.
Answer:
C-atomic no. = 6 = 2, 4, has 4 valence electrons, it cannot lose 4 electrons also it cannot gain 4 electrons for energy considerations. Hence C atom shares valence electrons. Each atom of Cl has 7 electrons in outer most orbit and can gain 1 electron. Hence to attain octet i.e. 8 electrons in valence shell it needs 4 Cl-atoms to form covalent molecule of CCl₄.

Chemical Bonding – Unit Test Paper 4

Q.1. Name the following :

Question 1.
An electrovalent compound formed by transfer of one electron from a metallic atom to a non-metallic atom.
Answer:
The electrovalent compound is sodium chloride (NaCl)

Question 2.
A covalent hydrocarbon molecule, having four single covalent bonds.
Answer:
A covalent hydrocarbon molecule having four single covalent bonds is methane

Question 3.
A covalent molecule, having two lone pair of electrons.
Answer:
A covalent molecule having two lone pair of electrons is water

Question 4.
A covalent molecule having a triple covalent bond.
Answer:
A covalent molecule having triple covalent bond is nitrogen (N = N)

Question 5.
A covalent molecule having two shared pairs of electrons in its molecule.
Answer:
A covalent molecule having two shared pairs of electrons in its molecule is oxygen (O = O).

Q.2. Give reasons for the following :
To attain stable electronic configuration of the nearest noble gas —

Question 1.
Hydrogen atom – needs one electron.
Answer:
¹₁H needs 1 electron to attain nearest noble gas He configuration i.e. 2 electrons in first valence shell (Duplet-rule)

Question 2.
Oxygen atom ¹⁶₈O – needs two electrons.
Answer:
¹⁶₈O_=2,6 to achieve nearest noble gas Ne, electronic configuration (octet-rule) Oxygen atom has 2 electrons short of 8.

Question 3.
Nitrogen atom ¹⁴₇N – needs three electrons.
Answer:
¹⁴₇N_=2,5 Nitrogen atom has 3 electrons short of 8 electrons in valence shell to attain nearest noble gas Ne configuration, (octet rule).

Question 4.
Carbon atom 126C – needs four electrons.
Answer:
¹²₆C_=2,4 C – atom has 4 electrons short of 8 electrons in valence shell. It needs 4 electrons to attain nearest noble gas Ne electronic configuration.

Question 5.
Chlorine atom ³⁵₁₇Cl – needs one electron.
Answer:
³⁵₁₇Cl_=2.8,7 atom Has 1 electron short of nearest noble gas Ar configuration (octet rule).

Q.3. Complete the table given below, pertaining to formation of covalent compounds :

Answer:

(a) Single i.e. 1 covalent bond
(b) 3 covalent bonds
(c) 2 covalent bonds
(d) 1 covalent bond
(e) 1 covalent bond

Q.4. Differentiate between the following with a suitable example :

Question 1.
Lone pair & shared pair of electrons.
Answer:
Difference between lone pair and shared pair :

• Lone pair : A pair of electrons which is alone i.e. does not share with electron of other atom to complete valence shell of both at m.
• Shared pair : A pair of electrons, one froms each atom are share for both atom to complete their respective valence shell.

Example, In case of water

Shared pair makes complete valence shell (Duplet) of hydrogen and octet in oxygen.
Lone pair is used in making octet of oxygen only.

Question 2.
Duplet rule & octet rule.
Answer:

• Duplet rule : If valence shell is in first orbit (or K orbit), it has maximum of 2 electrons i.e. He.
• Octet rule : An atom can have maximum of 8 electrons in valence shell.

Question 3.
Stable electronic configuration & unstable configuration.
Answer:

• Stable electronic configuration : When an atom of an element has an valence shell of 2 electrons, only in case of noble gas He i.e. in 1st outer most shell or 8 valence electrons in valence shell like other noble gases, that is called stable electronic configuration.
• Unstable electronic configuration : Atoms of the elements other than noble gases have their valence shell incomplete i.e. have 1, 2 … 7 electrons and are reactive. They have unstable electronic configuration.

Question 4.
Electrovalent bond & covalent bond
Answer:

• Electrovalent bond : “A bond formed between metal and non-metal by transfer of electron in called electrovalent bond e.g. (NaCl)
• Covalent bond : “Bond formed between non-metal or molecule of gas by sharing equal number of electrons by each atom is covalent bond e.g. Cl – Cl, H – H,
  

Question 5.
Single covalent bond & triple covalent bond.
Answer:

• Single covalent bond : “Sharing one pair of electrons betwen two reacting atoms So that each atom attains stable electronic configuration of the nearest noble gas Helium results in the formation of single covalent bond.”
• Triple covalent bond : “A bond formed as a result of sharing three pairs of electrons between two reacting atoms so that each atom attain stable electronic configuration of nearest noble gas is called triple covalent bond.”

Q.5. Atomic numbers of the following elements are given below:
A = 8; B = 7; C = 17; D = 11; E = 20

Question 1.
State which of the above is a divalent metal.
Answer:

Question 2.
State which of the above is a non-metal.
Answer:

Question 3.
State the type of bonding between two atoms of ‘A’.
Answer:
Type of bonding between two atoms of ‘A’ is Double covalent bond

Question 4.
State the type of bonding between ‘D’ & ‘C’
Answer:
Type of bonding between ‘D’ and ‘C’ is Electrovalent bond

Question 5.
State the number of covalent bonds formed in a molecule of ‘B₂’
Answer:
Covalent bonds formed in molecule of B₂ is Triple covalent bond

Q.6. The representation below shows the outline formation of an electrovalent compound.

If the atomic number of element ‘X’ is 11 and of element ‘Y’ is 17 —

Question 1.
State why an electron is transferred from ‘X’ to ‘Y’ during the formation of ‘X Y’
Answer:
X— has 1 electron in valence shell which it can lose and become stable.
Y — has 7 electrons in valence shell and by gaining 1 electron attains stable octet.

Question 2.
Give a reason why electrons are not shared between ‘X’ & ‘Y’ — during the formation of ‘XY’
Answer:
Both X and Y can not share equal number of electrons to attain their stable configuration. Moreover X has 1 electron to lose and Y needs 1 electron to attain stable valence shell.

Question 3.
State the difference between ‘X’ and ‘X¹⁺’
Answer:
X is atom and has no charge X¹⁺ is ion and has unit +ve charge called cation.

Question 4.
Does ‘Y^1-’ have a stable or an unstable electronic configuration.
Answer:
Y^1- has stable electronic configuration.

Question 5.
If a compound is formed from atom A [at. no.19] and an atom Y [at. no. 17], would the compound ‘AY’ be an electrovalent or covalent compound. Give reasons.
Answer:

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Exercise

Question.1 (1985)
XCl₂ is the chloride of a metal X. State the formula of the sulphate and the hydroxide of the metal X.
Answer:

Question.1 (1987)
An element X is trivalent. Write the balanced equation for the combustion of X in oxygen.
Answer:
Combustion of X³⁺ in oxygen means oxide of X

Question.1 (1991)
The formula of the nitride of a metal X is XN, state the formula of :

1. its sulphate
2. its hydroxide.

Answer:
Formula of nitride of metal X is XN
Since valency of nitrogen is 3-
∴ Valency of X is 3+

Question.1 (1992)
What is the valency of nitrogen in :

1. NO
2. N₂O
3. NO₂

Answer:

Additional Questions

Question 1.
What is meant by the term ‘symbol’. Give the qualitative and quantitative meaning of the term ‘symbol’.
Answer:
Symbol : “Is the short form that stands for the atom of a specific element.”
Qualitative meaning : C is the symbol of atom of element carbon
S is the symbol of atom of element sulphur.
This means symbol stand for a specific element.
No two elements can have the same symbol.
Quantitative meaning : A symbol also represents quantity of the element i.e. atomic mass of element. Symbol C represents 12 g of carbon. In other words how many times that element is heavier than 1/12 th C¹²

Question 2.
Name three metals whose symbols are derived from :

(a) the first letter of the name of the element
(b) from their Latin names.

Answer:

Question 3.
Explain the meaning of the term ‘valency’.
State why the valency of the metal potassium is +1 and of the non-metal chlorine is -1.
Answer:
Valency : “Is the combining capacity of atom or of a radical.” OR
“Valency is the number of electrons, which an atom can lose/gain/share during a chemical reaction.”
Metal potassium (K) has 1 electron in outer most shell which it loses and becomes K+ has valency [+1] whereas
Non-metal chlorine [Cl] has 7 electrons in valence shell and gains 1 electron and becomes [Cl^–] has valency [-1]

Question 4.
What is meant by the term ‘variable valency’. Give a reason why silver exhibits a valency of +1 and +2.
Answer:
Variable valency : “Certain elements exhibit more than one valency and show variable valency.”
Reason : Why silver exhibits valency +1 and +2 Ag 47 [2, 8, 18, 18, 1] has 1 electron in the outermost shell when loses this electron shows [+1] valency but when penultimate shell has not attained stability and one more electron jumps to the outermost shell there by increasing valency electron and new configuration [2, 8, 18, 17, 2] loses two electrons and has valency [+2]
∴ Silver exhibits Ag¹⁺ [ous] and Ag²⁺ [ic]

Question 5.
Give examples of eight metals which shows variable valency. State the valency of sulphur in :

(a) SO₂
(b) SO₃

Answer:
Eight metals which show variable valency are

Question 6.
State the valency in each case and name the following elements or radicals given below

Answer:

Question 7.
State the variable valencies of the following elements and give their names.
(a) Cu, (b) Ag, (c) Hg, (d) Fe, (e) Pb, (f) Sn, (g) Mn, (h) Pt, (i) Au
Answer:

Question 8.
State which of the following elements or radicals are divalent –
(a) Lithium, (b) Nickel, (c) Ammonium, (d) Bromide, (e) Sulphite, (f) Nitride, (g) Carbide, (h) Chromium, (i) Bisulphite, (j) Dichromate, (k) Permanganate.
Answer:

(b) Nickel
(e) Sulphite
(j) Dichromate are divalent

Question 9.
Explain the meaning of the term ‘compound’ with a suitable example. State the main characteristics of a compound with special reference to the compound iron [II] sulphide.
Answer:
Compound : “Is a pure substance made up of two or more elements combined chemically by in a fixed proportion.”
Example : CO₂ carbon dioxide is made up of two elements carbon and oxygen C,
12 parts by weight and oxygen
2 × 16 = 32 parts by weight
i.e. in ratio C : O = 12 : 16 = 3 : 4
CHARACTERISTICS OF COMPOUND :

1. Components in definite proportion e.g. compound iron [II] sulphide FeS element Iron and sulphur are in definite ratio.
2. Compound is Homogeneous.
3. Particles in a compound are of J kind composition of iron (II) sulphide is uniform and components cannot be seen separately.
4. Compound has definite set of properties.
5. Component in FeS do not retain their original properties, i.e. iron cannot be attracted by a magnet, sulphur is insoluble in CS₂ and Fe does not giyes H₂ with dil. acid. This means compound formed has new properties.

Question 10.
Name the elements in the compound and give the formula – of the following compounds :
(a) Nitric acid, (b) Carbonic acid, (c) Phosphoric acid, (d) Acetic acid, (e) Blue vitriol, (f) Green vitriol, (g) Glauber’s salt, (h) Ethane, (i) Ethanol
Answer:

Question 11.
Explain the term ‘chemical formula’. State why the molecular formula of zinc carbonate is ZnCO₃
Answer:
Chemical Formula : “A molecule of a substance element or compound could be represented by symbols. Representation known as chemical formula i.e. molecular formula of oxygen gas is O₂ water H₂O, hydrochloric acid HCl etc.
Molecular formula of zinc carbonate : Zinc and carbonate both have valency

Question 12.
Write the formula of the following compounds :



Answer:

Question 13.
Write the names of the following compounds :

Answer:

Question 14.
Explain the term ‘chemical equation’. What is meant by ‘reactants’ and ‘products’ in a chemical equation.
Answer:
Chemical equation : “Is the symbolic representation of a chemical reaction using symbols and the formulae of the substances involved in the reaction.”
Sulphur bums in oxygen is a chemical reaction and sulphur dioxide is formed. This can be represented as sulphur + oxygen \(\underrightarrow { heating } \) sulphur dioxide is word equation.

Reactants : Substances taking part in a reaction separated by (+) sign on the left hand side of arrow are called reactants.
Products : Substances formed in the reaction are called products.

Question 15.
Give an example of a chemical equation in which two reactants form –

(a) one product
(b) two products
(c) three products
(d) four products

Answer:

Question 16.

(a) State what is a ‘balanced equation’.
(b) Give a reason why the above equation is balanced.
(c) State why the compound MnO₂ is written above the arrow.

Answer:

(a) Balanced equation : “Equation in which the total number of atoms of each element in the reactants, on the left side of the equation is equal to the number of atoms of each element in the products formed, on the right side of the equation.”
(b) As the total number of [K, Cl, 0] atoms on the L.H.S. is equal to the number of the given atoms on the R.H.S., the given equation is balanced.
(c) MnO₂ is a catalyst in the reaction which does not under go any change and simply increases the rate of reaction is written above the arrow.

Question 17.
What do the symbols —

Answer:

Question 18.

(a) State the information provided by the above chemical equation.
(b) State the information not conveyed by the above chemical equation.

Answer:
(a) Information provided by equation

1. Calcium carbonate reacts with [dil.] hydrochloric acid to produce calcium chloride, water and Carbondioxide.
2. One molecule of calcium carbonate reacts with two molecules of acid to produce one molecule of calcium chloride, one molecule of water and one molecule of carbondioxide.
3. About the chemical composition of respective, molecules like one molecule of calcium chloride contains one atom of calcium, one atom of carbon and three atoms of oxygen.
4. Molecular masses
   
   that 100 parts by weight of calcium carbonate reacts with 73 parts by weight of hydrochloric acid to produce 111 parts by weight of calcium chloride 18 parts by weight of water and 44 parts by weight of carbon dioxide.
5. That reaction is irreversible.
6. That about the state of substances present i.e. solid, liquid or gas.

(b) Information not conveyed are :

1. Time, reaction takes to complete.
2. About the concentrations of reactants and products.
3. Speed of reaction.
4. Changes in colour occuring – during the reaction.
5. Whether heat is given out or absorbed during the reaction.
6. The physical state of reactants and products.

Question 19.
Balance the following simple equation :


Answer:

Question 20.
Write balanced equations for the following word equations :

1. Potassium nitrate → Potassium nitrite + Oxygen
2. Calcium + Water → Calcium hydroxide + Hydrogen
3. Iron + Hydrochloric acid → Iron [II] chloride + Hydrogen
4. Nitrogen dioxide + Water + Oxygen → Nitric acid
5. Lead dioxide [lead (IV) oxide] → Lead monoxide + Oxygen
6. Aluminium + Oxygen → Aluminium oxide
7. Iron + Chlorine → Iron [III] chloride
8. Potassium bromide + Chlorine → Potassium chloride + Bromine
9. Potassium bicarbonate → Potassium carbonate + Water + Carbon dioxide
10. Calcium hydroxide + Ammonium chloride → Calcium chloride + Water + Ammonia

Answer:

Question 21.
Balance the following important equations :




Answer:

Question 22.
Give balanced equations for (1) & (2) by partial equation method, [steps are given below]
(1) Reaction of excess ammonia with chlorine – Ammonia as a reducing agent

(a) Ammonia first reacts with chlorine to give hydrogen chloride and nitrogen.
(b) Hydrogen chloride then further reacts with excess ammonia to give ammonium chloride.

(2) Oxidation of Lead [II] Sulphide by Ozone

(a) Ozone first decomposes to give molecular oxygen & nascent oxygen.
(b) Nascent oxygen then oxidises lead [II] sulphide to lead [II] sulphate.

Answer:
(1)

(a) Ammonia reacts with chlorine to give hydrogen chloride and nitrogen.
(b) Hydrogen chloride reacts with excess ammonia to give ammonium chloride.

(2)

(a) ozone decomposes to O₂ (molecular) and nascent oxygen.
(b) Nascent oxygen oxidises Pb [II] sulphide to Lead [II] sulphate and needs 4 atoms of oxygen to form sulphate. Hence 1st equation must be multiplied by 4. Cancelling 4 atom of oxygen from both sides and adding.

Question 23.
Define the terms – (a) Relative atomic mass (b) Relative molecular mass. State why indirect methods are utilised to determine the absolute mass of an atom. Explain in brief the indirect method used.
Answer:
(a) Relative atomic mass [RAM] of an element :
“is the number of times one atom of an element is heavier than 1/12 the mass of an atom of carbon [C¹²]”
or
“Mass of an atom of an element as compared with 1/12 mass of an atom of carbon [C¹²]”
(b) Relative molecular mass [RMM] of an element/compound : “Is the number of times one molecule of the substance is heavier than 1/12 the mass of an atom of carbon [C¹²].

or
“Mass of one molecule as compared with the 1/12 mass of an atom of carbon[C¹²]”

To determine the absolute mass of an atom indirect methods are utilised as Atom are extremely small and very light.
An isotope of carbon C¹² [carbon -12 atom] has been assigned atomic mass of exactly 12 atomic mass unit] is used.

Question 24.
1. Calculate relative molecular mass of

(a) ZnCO₃
(b) CaSO₄ [Zn = 65, S=32, O = 16, Ca = 40, C = 12]

2. Calculate the percentage composition of

(a) calcium chloride
(b) calcium nitrate [Ca = 40 , Cl = 35.5 , N = 14 , O = 16]

Answer:

The Language Of Chemistry Unit Test Paper 1

Q.1. Match the names of ions and radicals from 1 to 10 with their correct answer from A to Q.

Answer:

Q.2. State which of the following formulas of compounds A to J are incorrect. incorrect write the correct formula.

Answer:

Q.3. Fill in the blanks with the correct word from the words in brackets :

Question 1.
A symbol represents a short form of a / an ____ [atom / element / molecule]
Answer:
A symbol represents a short form of a / an element.

Question 2.
Compounds are always ____ (heterogeneous/homogeneous) in nature.
Answer:
Compounds are always homogeneous in nature.

Question 3.
Variable valency is exhibited, since electrons are lost from an element from the ____ [valence / penultimate] shell.
Answer:
Variable valency is exhibited, since electrons are lost from an element from the penultimate shell.

Question 4.
A chemical equation is a shorthand form for a ____ [physical / chemical] change.
Answer:
A chemical equation is a shorthand form for a chemical change.

Question 5.
Relative molecular mass of an element/compound is the number of times one ____ of the substance is heavier than -1/12th the mass of an atom of carbon [C₁₂]. (atom/ion/molecule)
Answer:
Relative molecular mass of an element/compound is the number of times one molecule of the substance is heavier than -1/12 th the mass of an atom of carbon
[C12].

Q.4. Underline the compound in each equation given below, which is incorrectly balanced and write the correct balancing for the same.

Answer:

Q.5. With reference to a chemical equation state which of the statements 1 to 5 pertain to A or B.
A : Information provided by a chemical equation. .
B : Limitations of a chemical equation

1. The nature of the individual elements.
2. The speed of the reaction.
3. The state of matter in which the substance is present.
4. The completion of the reaction.
5. The direction of the reaction.

Answer:
A : Information provided by a chemical equation.

1. The nature of the individual elements.
3. The state of matter in which the substance is present.
5. The direction of the reaction.

B : Limitations of a chemical equation

2. The speed of the reaction.
4. The completion of the reaction.

New Simplified Chemistry Class 9 ICSE Solutions Chemical Changes and Reactions

ICSE SolutionsSelina ICSE SolutionsML Aggarwal Solutions

Viraf J Dalal Chemistry Class 9 Solutions and Answers

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Exercise

Question 1.(1986)
Explain : Silver nitrate solution is kept in coloured reagent bottles in the laboratory.
Answer:

Silver nitrate gets decomposed by sunlight. Hence silver nitrate is stored in coloured reagent bottles.

Question 1.(1987)
Give an example of an endothermic reaction.
Answer:

Question 1.(1988)
State in each case if the reaction represents oxidation or reduction

Answer:

Question 1.(1989)
Reactions can be classified as : Direct combination, decomposition, simple displacement, double decomposition, Redox reactions. State which of the following types, takes place in the reactions given below

1. Cl₂ + 2KI → 2KCl + l₂
2. 2Mg + O₂ → 2MgO
3. SO₂ + 2H₂O + Cl₂ → 2HCl + H₂SO₄
4. AgNO₃ + HCl → AgCl + HNO₃
5. 4HNO₃ → 4NO₂ + 2H₂O + O₂

Answer:

1. Cl₂ + 2KI → 2KCl + I₂ — Simple displacement
2. 2Mg + O₂ → 2MgO — Direct combination
3. SO₂ + 2H₂O + Cl₂ → 2HCl + H₂SO₄ — Redox reaction
4. AgNO₃ + HCl → AgCl + HNO₃ — Double decomposition
5. 4HNO₃ → 4NO₂ + 2H₂O + O₂ — Decomposition reaction

Question 1.(1990)
Give one reason why magnetizing a piece of steel is a physical change.
Answer:
There is no change in composition.

Question 1.(1993)
State whether the following are oxidation or reduction

Answer:

Additional Questions

Question 1.
Explain the term chemical reaction with special reference to reactants and products.
Answer:
Chemical change : “Is reprsentation of a chemical change in substances taking part and are called reactants written on left side of →

When Zn react with dil HCl, new substance ZnCl and H₂ are produced as a result of chemical change reactants are seperated by + sign and products formed are also separated by + sign.

Question 2.
Give a suitable example with equation to show the representation of a chemical reaction.
Answer:
Iron and sulphuric acid dilute which react to produce ferrous sulphate and hydrogen gas can be represented by a chemical equation.
Iron + dil. sulphuric acid → ferruous sulphate and hydrogen

Question 3.
A chemical reaction is often accompanied by external indications or characteristics.
Give two examples where a chemical reaction is accompanied by a change in colour of the reactants & products on completion of the reaction.
Answer:
A chemical reaction is accompanied by change in colour of reactants and products

Question 4.
Give balanced equations for reactions involving evolution of a gas on addition of dilute acid to

(a) sodium sulphite
(b) calcium carbonate.

Answer:

Question 5.
Give a balanced equation for conversion of

(a) an ammonium salt to a basic gas
(b) a soluble lead salt to an insoluble lead salt – formed as a white precipitate.

Answer:
(a) Ammonium salt [NH₄Cl] on heating produces NH₃(g) which is a basic gas

(b) Lead nitrate when reacts with sodium chloride insoluble white ppt. of lead chloride is formed

Question 6.
Chemical reactions may proceed with evolution or absorption of heat. Give an example of each.
Answer:
Chemical reaction is characterised by evolution of heat :

Chemical reaction is characterised by absorption of heat :

Question 7.
Define the following types of chemical changes or reactions with a suitable example each.

(a) Direct combination reaction or synthesis
(b) Decomposition reaction
(c) Displacement reaction or substitution reaction
(d) Double decomposition reaction

Answer:
(a) Direct Combination Reaction Of Synthesis :
“Those reactions in which two or more substances [element or elements and compound or compounds] combine and form a new substance”
H₂ + O₂ → H₂O Hydrogen combines with oxygen to produce a new substance water.
(b) Decomposition Reaction :
“Those reactions in which a compound splits up into two or more simpler substances”

(c) Displacement Reaction Or Substitution Reaction :
“Those reactions in which one element takes place of another element in a compound, are known as substitution reactions.”

(d) Double-De-Composition Reaction :
“Those reactions in which two compounds react by an exchange of ions to form two new compounds are called double displacement reactions.”

Question 8.
Give a balanced equation for a direct combination reaction involving :

(a) Two elements — one of which is a neutral gas and the other a yellow non¬metal
(b) Two elements – one of which is a neutral gas and the other a monovalent metal
(c) Two compounds – resulting in formation of a weak acid

Answer:
(a) Direct Combination of two elements one of which is a neutral gas and the other a yellow non-metal

(b) Direct Combination of two elements one of which is a neutral and the other a monovalent metal

(c) Direct Combination of two compounds to form weak acid.

Question 9.
Give balanced equations for the following reactions of synthesis involving formation of :

(a) An acid – from sulphur dioxide gas
(b) An alkali – from a basic oxide – sodium oxide
(c) A salt – from a trivalent metal and a coloured gas.

Answer:

Question 10.
Convert — (a) nitrogen to ammonia (b) hydrogen to hydrogen chloride – by a direct combination reaction.
Answer:
By direct combination reaction :
(a) Nitrogen to ammonia

(b) Hydrogen to hydrogen chloride

Question 11.
Give balanced equations for thermal decomposition of :

(a) lead carbonate
(b) lead nitrate
(c) ammonium dichromate
(d) mercury [II] oxide
(e) calcium hydroxide

Answer:
Decomposition of :

Question 12.
Define – a thermal dissociation reaction with a suitable example. Give an example of a photochemical decomposition reaction. Name a metallic oxide which on thermal decomposition is reduced to a metal.
Answer:
Thermal Dissociation :
“A decomposition reaction — in which a substance dissociates — into two or more simpler substances on application of heat.”

Question 13.
Define a displacement reaction with a suitable example. State how it is represented. Give a reason why zinc displaces hydrogen from dilute sulphuric acid but copper docs not.
Answer:
Displacement reaction : “Is a chemical reaction in which an element [or radical] replaces another element in a compound.”

Zinc (more electropositive) and being above [H] in activity series displaces hydrogen from dilute sulphuric acid but copper is below [H] in electrochemical series cannot displace hydrogen from sulphuric acid.

Question 14.
Explain the term double decomposition — precipitation reaction. Give a balanced equation for the preparation of two different insoluble lead salts from their salt solutions by — double decomposition — precipitation.
Answer:
Double decomposition reaction is precipitation reaction.
i. e. in which ppt. is formed
“Reaction between two compounds in aqueous state, to give two new compounds one of which is precipitate (or insoluble).”

Question 15.
Explain with the help of balanced equations, how precipitation reactions are used for identifying the positive radicals in three different salts, each having a different cation [positive ion].
Answer:
Precipitation reactions are used to identify the positive ions from their colours.

Question 16.
Define the term – double decomposition – neutralization reaction with a suitable representation.
Convert :
(a) an insoluble base
(b) a soluble base to their respective soluble salts by neutralization reaction.
Answer:
When acid reacts with a base salt and water are formed and this is called Neutralisation.
Double-decomposition — Neutralisation reaction :
“Is the chemical reaction between two compounds (acid and base) to interchange radicals and produce salt and water.”
This is represented as :

Question 17.
Explain the term energy changes in a chemical change or reaction. Give an example with a balanced equation, for each of the following reactions:

(a) exothermic reaction
(b) endothermic reaction
(c) photochemical reaction
(d) electrochemical reaction.

Answer:
Energy Changes In A Chemical Reaction :
“Is the difference between the chemical energy of the REACTANTS and the PRODUCTS”.
Example of :

Question 18.
Supply of energy maybe required to initiate a reaction. State the different forms with a suitable example of reactions initiated by supply of energy.
Answer:
To Start A Reaction Energy Needed In The Form :

Chemical Changes & Reactions – Unit Test Paper 2

Q.1. Complete the statements by filling in the blank with the correct word/s :

Question 1.
Direct combination reaction of sulphur dioxide with water gives [H₂SO₄/H₂SO₃/H₂S₂O₇].
Answer:
Direct combination reaction of sulphur dioxide with water gives H₂SO₃.

Question 2.
Formation of hydrogen chloride from hydrogen and chlorine is an example of [photochemical reaction/electrochemical reaction].
Answer:
Formation of hydrogen chloride from hydrogen and chlorine is an example of photochemical reaction.

Question 3.
The reaction of hydrogen burning in oxygen to give a neutral liquid is an example of [exothermic/ endothermic] reaction.
Answer:
The reaction of hydrogen burning in oxygen to give a neutral liquid is an example of exothermic reaction.

Question 4.
The neutral gas evolved when lead nitrate undergoes thermal decomposition is [nitrogen dioxide/oxygen/nitrogen].
Answer:
The neutral gas evolved when lead nitrate undergoes thermal decomposition is nitrogen dioxide.

Question 5.
The reddish brown precipitate obtained during a double decomposition precipitation reaction between an iron salt and an alkali is [iron [III] hydroxide/iron [III] hydroxide]
Answer:
The reddish brown precipitate obtained during a double decomposition – precipitation reaction between an iron salt and an alkali is iron [III] hydroxide.

Q.2. Select the correct answer from A, B, C, D and E for each statement given below :
A : Ammonia
B : Hydrogen chloride
C : Hydrogen
D : Nitrogen dioxide
E : Nitric oxide
State the gaseous product formed, when

Question 1.
An active metal reacts with dilute sulphuric acid.
Answer:
C : Hydrogen

Question 2.
A metallic nitrate undergoes thermal decomposition giving a coloured gas.
Answer:
D : Nitrogen dioxide

Question 3.
Two gases one of them neutral, combines by absorption of light energy.
Answer:
B : Hydrogen chloride

Question 4.
An ammonium salt reacts with an alkali.
Answer:
A: Ammonia

Question 5.
An exothermic reaction takes place between ammonia and a neutral gas.
Answer:
E : Nitric oxide

Q.3. Give a balanced equation for each of the following types of reactions :

Question 1.
A direct combination reaction between phosphorus and a neutral gas.
Answer:
Phosphorus with neutral gas (O₂)

Question 2.
A soluble salt of lead formed from an insoluble base by double decomposition – neutralisation.
Answer:
Soluble salt of lead formed from insoluble base by double-decomposition by neutralization.

Question 3.
A thermal decomposition reaction of a salt – which results in the formation of nitrogen gas.
Answer:
Thermal decomposition of a salt with the formation of nitrogen gas.

Question 4.
A synthesis reaction between a metal & a non-metal resulting in formation of an insoluble salt of iron.
Answer:
Synthesis reaction between a metal and non-metal to form insoluble salt of iron [FeS]

Question 5.
A decomposition reaction of a salt which leaves behind a silvery metal.
Answer:
Decomposition reaction of a salt leaving behind a silvery metal.

Q.4. Differentiate between the following :

Question 1.
Synthesis reaction & a substitution reaction.
Answer:
Synthesis and substitution reaction :

• Synthesis reaction : “Is a chemical reaction in which two or more elements or an element and a compound or two compound combine to form a new compound.”
  
• Substitution reaction : “Is a chemical reaction takes place when an element or radical (more reactive) replaces another element in a compound.”
  

Question 2.
Electrolytic decomposition & photochemical decomposition.
Answer:
Electrolytic decomposition and photochemical decomposition :

• Electrolytic decomposition : “A chemical reaction which takes place with absorption of — electrical energy”
  
• Photochemical decomposition : “A chemical reaction which takes place with absorption of light energy :
  

Question 3.
Thermal decomposition & thermal dissociation.
Answer:
Thermal decomposition and thermal dissociation.

• Thermal decomposition : “A chemical reaction in which a compound decomposes into two new substance by heat.” It is not reversible reaction.
  CaCO₃ Cao + CO₂
• Thermal dissociation : “Is decomposition of a compound into two or more new substances by heat energy.” It is reversible reaction.
  

Question 4.
Decomposition reaction & a double decomposition reaction
Answer:
Decomposition reaction and double decomposition reaction :

• Decomposition reaction : A compound decomposes into two or more new substances.
  
• Double-decomposition : “Two substances exchange their ions and two new substances are produced one of which is insoluble i.e. ppt.”
  

Question 5.
Neutralization reaction & a precipitation reaction.
Answer:
Neutralization reaction and a precipitation reaction :
Neutralization reaction “is a double — decomposition reaction in which acid neutralizes a base and salt and water is formed.”

Q.5. Match the chemical reactions in List I with the appropriate answer in List II. List I


Answer:

Q.6. Name the solid residual product formed in each reaction and state its colour during – thermal decomposition of the following substances.

1. Copper nitrate.
2. Ammonium dichromate.
3. Zinc carbonate.
4. Lead nitrate.
5. Calcium hydroxide.

Answer:
Solid residual product of thermal decomposition of :

New Simplified Chemistry Class 9 ICSE Solutions – Water

ICSE SolutionsSelina ICSE SolutionsML Aggarwal Solutions

Viraf J Dalal Chemistry Class 9 Solutions and Answers

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Exercise

Question 1.(1984)
Name one substance which is ‘deliquescent’.
Answer:
Deliquescent substance is NaOH (sodium hydroxide), KOH, CaCl₂

Question 2.(1984)
How does an increase in temperature affect :

1. the solubility of NaCl
2. the solubility of CaSO₄ in water ?

Answer:
Increase in temperature :

1. Increases the solubility of NaCl only a little.
2. Solubility of CaSO₄ first increases and then decreases.

Question 1.(1985)
Give reasons for the following :

1. table salt becomes moist and sticky during the rainy season
2. a white power forms on the surface of washing soda crystals which are left exposed to the air.

Answer:

1. Table salt contains impurities of calcium and magnesium chloride which are Deliquescent and absorb moisture from air and common salt becomes Sticky and Wet During humid conditions [rainy season].
2. Washing Soda is Efflorescent substance when exposed to air loses its moisture [water of crystallisation] partly and becomes powder which forms white powder on the crystals.

Question 1.(1986)
Name a salt

(a) which contain of crystallization
(b) which does not contain.

Answer:

(a) Blue vitriol [copper sulphate] CuSO₄. 5H₂O
(b) Potassium chloride KCl or sodium chloride NaCl

Question 1.(1987)
Name a deliquescent substance’.
Answer:
Iron [III] chloride [FeCl₃]

Question 1.(1988)
Explain the following observations :

1. Washing-soda become coated with a white powder when left exposed to the atmosphere.
2. In the expression anhydrous copper sulphate, what is meant by “anhydrous”.
3. Why is fused calcium chloride or cone, sulphuric acid used in a desiccator.

Answer:

1. Washing Soda : is efflorescent substance and lose its water of crystallisation partly when exposed to atmosphere and change into white powder (amorphous state).
2. Copper sulphate loses water of crystallization when heated and becomes white powder of CuSO₄. This white powder (CuSO₄) without water is anhydrous.
3. Fused CaCl₂ or cone. H₂SO₄ is used in desiccator as drying agent to remove moisture from the substance (to dry it) but does not change its composition.

Question 2.(1988)
Complete the following : The solubility of a gas at constant pressure may be increased by decreasing the
Answer:
The solubility of a gas at constant pressure may be increased by decrease in temperature.

Question 1.(1991)
What is ‘water of crystallization’ ? Name a crystalline salt which does not contain water of crystallization.
Answer:
Water Of Crystallization : “The fixed number of water molecules which enter into a loose Chemical Combination with the substance when the substance is crystallised from its hot saturated solution is called water of crystallization.”
The substance is crystalline but does not contain water of crystallization is KCl

Question 2.(1991)
What would you observe, when the water of crystallization of a salt Is removed by heating it.
Answer:
It will turn Amorphous in nature lose its Geometric shape, become powder.

Question 3.(1991)
Define :

1. Hygroscopy
2. Efflorescence.

Answer:

1. Hygroscopy : “is the phenomenon of a subtance to absorb water (moisture) from the atmosphere but does not change its state.”
2. Efflorescence : “Is the phenomenon in which a substance loses it moisutre (water of crystallisation partly or completely to atomsphere and change to amorphous state, when exposed to atmosphere.”

Question 4.(1991)
What is the effect of temperature on the solubility of KNO, and calcium sulphate in water.
Answer:
Solubility of KNO₃ increases with temperature and solubility of calcium sulphate first increases upto 50°C and then decreases upto 100°C.

Question 1.(1992)
What test would you do to find out whether a given solution is saturated or unsaturated.
Answer:
If on adding more of solute in the given solution and stirring, the solute dissolves, it is unsaturated solution. If the solute settles down and does not dissolve, it is saturated solution.

Question 2.(1992)
How can you increase the solubility of a given volume of gas in water.
Answer:
Solubility of a given volume of gas can be increased by increasing the pressure on the surface of water.

Question 1.(2007)
Define ‘eutrophication’.
Answer:
Eutrophication : “Organic matter in sewage poured into water bodies generally results in excessive growth of algae – which deoxygenates water and produces deadening atmosphere there.”

Question 2.(2007)
What is meant by the term ‘oil spill’.
Answer:
Oil spill or Leaks : Oil is lighter and insoluble in water and oil layer floats on the surface of water and prevents oxygen transfer from atmosphere.
Oil leakage and petroleum products into sea water due to accidents of ships and oil tankers or leakage of pipe lines and storage tanks. They occur when oil is being produced from offshore well and may happen due to transportation by pipes and tanks, oil refineries, petrochemical plants etc.
Damage caused by oil spills :

1. Fish, birds, reptiles and amphibians living in such water cannot breathe and die. The oil penetrates the birds feathers and they cannot fly, and affect their insulation and damages reproductive system.
2. It interrupts the food chain which may cause extinction of species.

Question 1.(2008)
State any two sources of water pollution.
Answer:
Two sources of water pollution are :

(a) Industrial waste
(b) Sewage.

Question 2.(2008)
State the causes and consequences of ‘eutrophication’.
Answer:
Causes of eutrophication are :

(a) Increase in chemical nutrients in an ecosystem.
(b) Organic matter in sewage poured into water bodies.

Question 3.(2008)
What is meant by the term ‘offshore drilling’. State the main environmental effects of offshore drilling.
Answer:
Offshore drilling : Involves exploring for oil and gas beneath the ocean floor. Steps involved are location of wells, exploring wells to find out if there is oil below and if oil and gas ‘production well’ is drilled and an ‘oil rig’ is built to replace the exploratory drilling rig.
Environmental effects : It affects the health and reproduction of marine animals. It interrupts the food chain, which may cause extinction of species.

Question 4.(2008)
Explain why oil spills have an adverse effect on marine life.
Answer:
Oil is lighter and insoluble in water and oil layer floats on the surface of water and prevents oxygen transfer from atmosphere.
Oil leakage and petroleum products into sea water due to accidents of ships and oil tankers or leakage of pipe lines and storage tanks. They occur when oil is being produced from offshore well and may happen due to transportation by pipes and tanks, oil refineries, petrochemical plants etc.
Damage caused by oil spills :

1. Fish, birds, reptiles and amphibians living in cannot breathe and die. The oil penetrates the birds feathers and they cannot Tly, and affect their insulation and damages reproductive system.
2. It interrupts the food chain which may cause extinction of species.

Question 1.(2009)
Explain any two environmental impacts of an ‘oil spill’.
Answer:
See Question 2, 2007.

Question 1.(2010)
Explain the methods of controlling water pollution.
Answer:
Methods of controlling water pollution.

1. Control of water pollution can be done by properly handling of sewage waste.
2. Use drilling fluids which are biodegradable and have low aquatic toxicty.
3. Develop better pollution control measures which include removal of oil, dispersing oil, removing of oil clumps.
4. Oil and grease is removed from industrial waste by oil-water separator.
5. Acid and alkalies are neutralized.
6. Organic materials in waste water are removed by distillation adsorption.
7. Biodegradable organics are removed by (a) Activated sludge process (an aerobic biochemical process) (b) Biological trickling filter process.

Additional Questions

Question 1.
State the importance of water for all general uses.
Answer:
Importance of water : Water is vital for the growth of plants and animal life. Water is also essential for industrial processes, agricultural processes transportation and for power generation and cleanliness.

Question 2.
How does it occur in the free state and in the combined state.
Answer:
Water occurs in free state : In the form of ice, snow, frost
As – river water, lake water, sea water, spring water
As – water vapour, clouds, mist, fog
In combined state : In plants and animals, in hydrated salts i.e., MgCl₂.6H₂O and in certain minerals, milk, dry cereals and in green vegetables.

Question 3.
State a reason to prove that water is a compound and not a element.
Answer:
Three reasons are :

1. Water has new properties than its constituents [i.e. H₂ and O₂]. H₂ is combustible gas, O₂ gas is supporter of combustion but water is liquid.
2. In water hydrogen and oxygen elements combine in fixed ratio by weight 1 : 8.
3. Components of water i.e. H₂ and oxygen can be separated by chemical means e.g. by electrolysis of water.

Question 4.
State why ‘Water is considered a universal solvent’. Give the reason for the same.
Answer:
Water is universal solvent : Water dissolves many substances forming aqueous solutions. Not only solids, but gases and other liquids also dissolve in water. When water is put in the glass vessel an extremely small amount of glass dissolves in it. Organic compounds [carbohydrates, proteins] also dissolve in water. Hence water is called universal solvent.

Question 5.
Define the terms :

1. solute
2. solvent
3. solution

Answer:

• Solute : In a salt solution, salt is added to water then salt [NaCl] is solute i.e. substance dissolves.
• Solvent : Medium (liquid) in which solute is dissolved i.e. water is solvent to which salt dissolves to form salt solution.
• Solution : “A homogenous mixture of a solute in a solvent” is called solution.

Question 6.
State the characteristics of a true solution.
Answer:
Characteristics of a true solutions are :

1. Is homogenous in nature.
2. (a) Particles can pass through the filter paper, (b) Cannot be seen under microscope, (c) Do not settle down.
3. A true solution is a mixture and not a compound.

Question 7.
Differentiate between unsaturated, saturated & supersaturated solutions.
Answer:

• Saturated : Solution cannot dissolve any more of solute at a given temperature.
• Unsaturated : solution can dissolve more solute at a given temperature.
• Supersaturated : Solution which has more solute than its saturated sol. at that temperature.

Question 8.
How would you convert a saturated solution to an unsaturated solution and vice versa.
Answer:

• A saturated solution can be made unsaturated by adding more of solvent.
• An unsaturated solution can be made saturated by adding more of solute.

Question 9.
Define solubility. Give the main steps with the calculations involved of the method to determine the solubility of a given salt ‘X’ in water.
Answer:
Solubility : “The maximum amount of a SOLUTE which can be dissolved in 100 grams of a solvent at a specified temperature is called Solubility.
To Determine The Solubility Of ‘X’ In Water :
Steps :
(i) Preparation of saturated solution of ‘X’

1. Take 100 ml of distilled water in a boiling test tube.
2. Add crystals of ‘X’ to distilled water and stir slowly.
3. Continue adding and ktirring till the crystals dissolve. Repeat the process till no more salt can dissolve.
4. Pour the saturated solution — in a clean dry test-tube.

Steps :
(ii) Calculations To Determine Solubility Of Solute :

1. Weigh a clean and dry vaporating dish = M₁g
2. Add above saturated solution to it and weigh = M₂g
   ∴ Weight of solution (solute + solvent water) = (M₂ – M₁) g
3. Heat the solution to dryness and weigh the dish with residue = M₃g
   Weight of solute = (M₃ – M₁) g
   weight of solvent (water) = M₂ – M₃ = solution – solute (M₂ – M₁) = (M₃ – M₁) = M₂ – M₁ – M₃ + M₁
4. Note the temperature of saturated solution = t° C

Question 10.
From the following list of salts : Na₂SO₄, 10H₂O, NaCl, KClO₃, NaNO₃, Ca(OH)₂, NH₄Cl, KCI, CaSO₄.
State the salts whose solubility (a) increases, (b) decreases, (c) is fairly independent or slightly increases – with rise in temperature of water.
Answer:

(a) Solubility Increase with rise of temperature KClO₃, NH₄Cl
(b) Decreases with rise in temperature CaSO₄, Ca(OH)₂ above 70° C
(c) Slightly increases with rise in temperature
NaCl, KCl, Ca(OH)₂ below 70°C

Question 11.
What is a solubility curve. State two applications and two benefits of the solubility curve.
Answer:
Solubility curve : “The effect of temperature on solubility of solute in a solvent shown by a curve in the graph (temperature-solubility) is called solubility curve.”
Two applications of the solubility curve :

1. MEDICAL : Enables a pharmacist to determine the amount of drugs that must be dissolved together in a given quantity of solvent at different temperatures to give a prescribed drug preparation.
2. Chemists And Research Workers : Helps them to determine the most suitable solvent to be used at various temperatures for Extraction Of Essential Chemicals from their natural sources.

Two benefits of the solubility curve :

1. To compare solubilities of different solutes in a solvent at a given temperature.
2. To determine solubility of a given solute at a specific temperature.

Question 12.
Give the influence of

1. pressure
2. temperature on the solubility of gases in water.

Answer:
Influence on solubility of gases :

1. Pressure : “is directly proportional to pressure” at a given temperature. Solubility of gases increases with increase in pressure and decreases with decrease in pressure.
2. Temperature : “is inversely proportional to temperature”, i.e. increase in temperature of water causes decrease in solubility of gases. At low, temperature the solubility of the gas is more compared to higher or ordinary temperature.

Question 13.
State the reasons why

1. boiled water tastes flat
2. a soda water bottle opens with a ‘fizz’.

Answer:

1. Boiled Water Tastes Flat : Soluble gases in water contribute to the taste of water.
   When water is boiled Solubility Of Gases decrease and gases conic out of water and water tastes flat.
2. Soda Water Bottle Opens With A ‘Fizze’.
   An unopened soda bottle is virtually bubble-free because the pressure inside the bottle keeps the carbon dioxide dissolved in the liquid. When you crack open the bottle, and allow the gas bubbles to wiggle the pressure is released and allow the gas bubbles to wiggle free from the liquid and rise to the surface.

Question 14.
What is meant by the terms :

(a) crystal,
(b) crystallization,
(c) seed crystal.

Explain with examples
Answer:

(a) Crystal : “On cooling a Hot Saturated Solution, homogenous solids arranged symmetrically are obtained, called Crystals.”
(b) Crystallization : “The process by which crystals are separated or deposited from a hot saturated solution of a substance on cooling gently is called crystallization.”
(c) Seed crystal : “A well-formed crystal from the cooled filtrate of saturated solution used as seed for the formation (growing) of large sized crystal is called SEED CRYSTAL.”
e.g. a large size crystal of potassium nitrate is prepared from saturated solution of KNO₃ saturated solution.

Question 15.
Define the term ‘water of crystallization’.
Answer:
Water Of Crystallization : “Is The Fixed Number Of Water Molecules Which Enter Into A Loose Chemical Combination With The Substance When The Substance Is Crystallized From Its Hot Saturated Solution, Is Called Water Of Crystallization.”

Question 16.
Differentiate between hydrated and anhydrous crystals with examples. State three defined changes which occur when hydrated copper sulphate is heated.
Answer:
Difference between hydrated and anhydrous crystals :

1. Colour changes from blue to colourless.
2. Changed to AMORPHOUS white powder
3. Geometric shape vanishes
4. On cooling and adding water colour restores but crystalline does not shape.

Question 17.
Washing soda and iron [III] chloride are separately exposed to the atmosphere. State

1. the observations seen
2. the phenomenon which occurs
3. the reason for the phenomenon occurring in each case.

Would a similar phenomenon occur in case of exposure of common salt. Explain giving reasons.
Answer:
Observations When Exposed To Atmosphere :

(i) Washing soda :

(a) Loses its moisture (water of crystallization)
(b) Becomes amorphous
(c) Coated with white powder

(ii) Iron (III) Chloride (FeCl₂)
Absorbs moisture from atmsophere dissolves in moisture changes to Liquid state
(iii) In case of washing soda — the phenomenon is Efflorescence.
In case of FeCl₃ it is Deliquescence
No, in case of common salt (NaCl) no such phenomenon can occur as, it is Anhydrous salt and does not contain water of crystallization.
When common salt contains impurities of calcium or magnesium chloride dilquescent it absorbs moisture from atmosphere and become sticky and wet.

Question 18.
Why is fused calcium chloride and not potassium chloride kept in a desiccator?
Answer:
Fused Calcium Chloride (CaCl₂) Is Deliquescent in nature, absorbs moisture and hence is used as Drying Agent in desiccator for drying other substances in the laboratory but KCl (potassium chloride) has no such property and cannot be used.

Question 19.
How does fused calcium chloride differ from iron [III] chloride when exposed to the atmosphere?
Answer:
Fused CaCl₂ is Hygroscopic substance, absorbs moisture from atmosphere like FeCl₃ but does not change its state. FeCl₃ on absorbing moisture changes to Liquid State.

Question 20.
Cone. H₂SO₄ acts as a ‘drying agent’ & a ‘dehydrating agent’. Explain and differentiate the words in italics.
Answer:
Cone. H₂SO₄ when acts as drying agent, it absorbs only moisture from the substance and makes it dry without changing its composition.
When cone. H₂SO₄ acts as dehydrating agent, it removes water molecule from the composition of substance and reacts chemically, produces a new substance with new properties.

Question 21.
Explain the meaning of the terms – hard water & soft water.
Answer:

• Hard Water : “Water that does not lather readily-with ordinary soap and hence wastes soap” is called hard water.
• Soft Water : “Water that latters readily with ordinary soap and — hence soap is not wasted” is called soft water.

Question 22.
State the causes of hardness in water.
Answer:
Presence of salts of calcium and magnesium i.e. Ca(HCO₃) or Mg(HCO₃) or calcium or magnesium sulphate or chloride in water are the Causes Of Hardness Of Water.

Question 23.
Give two natural sources of hard water.
Answer:
Two natural sources of hard water :

1. Water from springs
2. River water.

Question 24.
Differentiate between temporary hard water & permanent hard water.
Answer:
Differences between temporary hard water and permanent hard water :
Temporary hard water :

1. is hard due to the presence of Bicarbonates of Calcium or magnesium.
2. Hardness can be removed by boiling.

Permanent hard water :

1. is hard due to the presence of salts like chlorine or sulphate of calcium or magnesium.
2. Hardness cannot be removed by boiling.

Question 25.
State the cause of hardness in temporary & permanent hard water.
Answer:
Cause Of Hardness In :

• Temporary hard water — bicarbonate of calcium or bicarbonate of magnesium.
• Permanent hard water — CaCl₂ or CaSO₄ or MgCl₂ or MgSO₄

Question 26.
State the disadvantages of hardness in water.
Answer:
Disadvantages of hardness in water :

1. Soap is wasted and clothes do not get clean.
2. Unsafe for drinking.
3. Unfit for laundries.
4. Forms boiler scale or fur in boilers.
5. Not suitable for preparing solutions.

Question 27.
Temporary hardness in water can be removed by boiling. Give balanced equations to explain how, hardness in temporary hard water is removed by boiling.
Answer:
Temporary hardness in water which is due to the presence of HCO₃ of calcium or magnesium.
On boiling bicarbonate changes to carbonate which is insoluble in water and is filtered out, CO₂ gas escapes and water becomes soft.

Question 28.
Both temporary & permanent hardness in water can be removed by addition of washing soda. Give balanced equations for the same.
Answer:
Using washing soda, removal of :

Question 29.
A sample of water is given in a trough. State how would you prove experimentally whether the given sample is hard water or soft water.
Answer:
On adding ordinary soap in the sample and shaking if this water gives lather easily it is soft water otherwise it is hard water.

Question 30.
Two samples of water are placed in a beaker individually. State how you will determine experimentally which of the two samples contains permanent hard water.
Answer:
Boil each sample of water separately and then two each sample add ordinary soap, the sample that lathers is temporary hard water and the sample which does not lather is permanent hard water.

Question 31.
State what are synthetic detergents. Explain experimentally how you will determine the advantage of synthetic detergents over soap using a sample of hard water.
Answer:
Divide the sample of hard water into two parts.
To the first part of hard water add synthetic detergent and rub slowly with hand, it will lather easily.
Now add to the second part of hard water an ordinary piece of soap and rub with hand, it will not lather but scum is formed and soap is wasted.
Hence in case of hard water, it is advantageous to use synthetic detergent.

Water – Unit Test Paper 3

Q.1. Select the correct word from the words in brackets to complete each sentence :

Question 1.
If pressure on the surface of water increases its boiling point ____ and freezing point ____ [increases / decreases].
Answer:
If pressure on the surface of water increases its boiling point increases and freezing point decreases.

Question 2.
A saturated solution can be converted to an unsaturated solution by ____ [increasing / decreasing] the amount of the solvent.
Answer:
A saturated solution can be converted to an unsaturated solution by increasing the amount of the solvent.

Question 3.
Dissolved air in water contains a ____ [higher / lower] percentage of oxygen than ordinary air.
Answer:
Dissolved air in water contains a higher percentage of oxygen than ordinary air.

Question 4.
At low temperatures the solubility of a gas in water is ____ compared to that at ordinary temperatures.
Answer:
At low temperatures the solubility of a gas in water is more compared to that at ordinary temperatures.

Question 5.
Efflorescence occurs when the vapour pressure of the hydrated crystals is ____ (more/less) than the vapour pressure of the atmospheric humidity.
Answer:
Efflorescence occurs when the vapour pressure of the hydrated crystals is more than the vapour pressure of the atmospheric humidity.

Q.2. Select the correct answer from the choice given in the brackets.

Question 1.
An anhydrous crystal, (blue vitriol/epsom salt/lead chloride).
Answer:
An anhydrous crystal lead chloride.

Question 2.
A substance which causes hardness in water. (NH₄Cl/CaCl₂/NaCl)
Answer:
A substance which causes hardness in water CaCl₂.

Question 3.
A deliquescent salt of a divalent metal. (CuCl/CaCl₂/FeCl₂/PbCl₂)
Answer:
A deliquescent salt of a divalent metal CaCl₂.

Question 4.
An anhydrate of a heptahydrate salt. (Cu(N0₃)₂/Ca(NO₃)₂/FeSO₄/CaSO₄)
Answer:
An anhydrate of a heptahydrate salt FeSO₄

Question 5.
A drying agent, deliquescent in nature used in a dessicator. (cone. H₂SO₄/ fused CaCl₂FeCl₃)
Answer:
A drying agent, deliquescent in nature used in a dessicator fused CaCl₂.

Q.3. Give reasons for the following

Question 1.
Solubility curves find utility in separation and purification of solutes.
Answer:
Those fractions with very low solubilities will be the first to crystallise out from the solution and will be separated out.

Question 2.
Pressure and temperature influence the solubility of gases in water.
Answer:

• An Increase In Pressure On The Surface Of Water Causes Increase In Solubility Of Gas In Water.
• An Increase In Temperature Of Water Causes Decreas In Solubility Of Gas In Water.

Question 3.
Heating a hydrated copper sulphate cryttal is deemed a chemical change.
Answer:
Heating hydrated copper sulphate crystal is a chemical change as on adding water drops on white powder of CuSO₄ (amorphous) colour restores but geometric shape is not restored i.e. after heating new composition is found to give crystalline shape.

Question 4.
Efflorescence is minimum during humid conditions.
Answer:
Efflorescence Occurs : When The Vapour Pressure Of Hydrated Crystals Exceeds The Vapour Pressure Of The Atmospheric Humidity.
Hence efflorescence is minimum during humid conditions.

Question 5.
A crusty ‘boiler scale’ is formed in boilers, when hard water is used.
Answer:
A crusty ‘boiler scale’ is formed in boilers, when hard water is used because insoluble CaCO₃ or MgCO₃ is formed in large scale which gets deposited as crusty ‘boiler scale’. Its formed when bicarbonates of calcium or magnesium are converted into carbonates.

Q.4. Name or state the following.

1. An efflorescent decahydrate salt.
2. A deliquescent salt of a trivalent metal.
3. A liquid hydroscopic substance.
4. A salt whose solubility decreases with rise in temperature of the solvent water.
5. A substance added to remove both temporary and permanent hardness in water.

Answer:

1. An efflorescent decahydrate salt is washing soda (Na₂CO₃.10H₂O)
2. A deliquescent salt of a trivalent metal is iron (III) chloride (FeCl₃)
3. A liquid hygroscopic substance — concentrated sulphuric acid (cone. H₂SO₄)
4. CaSO₄
5. Washing soda (Na,CO₃)

Q.5. Differentiate between the following :

1. Natural water and treated water
2. Saturated solution and a supersaturated solution
3. Solubility and solubility curve
4. Deliquescent salt and hygroscopic salt
5. Solute and solvent – forming a solution.

Answer:
Difference between :
1. Natural water and treated water :

• • Natural water : Water obtained in nature, river water, lake water, sea water, snow, rain water.
  • Treated water : Water which has received some form of treatment – ice, mineral water, distilled water, water in swimming pool.

2. Saturated solution and a supersaturated solution :

• Saturated : Solution cannot dissolve any more of solute at a given temperature.
• Supersaturated : Solution which has more solute than its saturated sol. at that temperature.
  A saturated solution can be made unsaturated by adding more of solvent.

3. Solubility :

(a) Is the ability of solute to dissolve in a particular solvent
(b) It is used to find the amount of solute required.

Solubility curve :

(a) Is a line graph that plots the changes in the solubility of a solute in a solvent against changes in temperature
(b) It is used to compare the solubility of different solutes.

4. Deliquescent salt : “When exposed to atmosphere lose (moisture) water of crystallization partly or wholly and become amoi’phous.” e.g. washing soda (Na₂CO₃.10H₂O)
Hygroscopic salt : “absorb water or moisture from the atmosphere when exposed but donot change their state like deliquent salts.” e.g. Cone. H₂SO₄, quick lime CaO.
5. Solute :

(a) It is the substance which is dissolved to form solution
(b) It is small in quantity.

Solvent :

(a) It is medium in which solute is dissolved
(b) It is large in quantity.

Q.6. Match the terms in List I with the correct answers in List II.


Answer: