Atoms and Molecules Notes

Everything around us — the air we breathe, the water we drink, the food we eat, and even our own bodies — is made up of matter. But what is matter made of? Scientists have discovered that matter is made up of tiny particles called atoms. These are the basic building blocks of all substances. When atoms combine, they form molecules, which make up different types of materials. This chapter helps us understand:

1. What are atoms and molecules ?
2. How they combined?
3. How we represent them using chemical symbols and formulas,
4. And how to calculate the mass of molecules.
5. Laws of Chemical Combination:

Laws of Chemical Combination

Scientists observed that chemical reactions always occur in a fixed, predictable manner. In a chemical reaction, two or more molecules interact to produce new compounds; they are called reactants, whereas the newly formed compounds are called products Based on these observations, two basic laws were formulated. In chemical industries, these laws help in calculating the amount of reactants required and products formed. They also help in writing balanced chemical equations.

The 2 laws of chemical combination are:

1. The law of conservation of mass
2. The law of constant proportions.

1. Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction. This means that the total mass of reactants (substances that react) is always equal to the total mass of products (substances formed) in a chemical reaction. For example: Hydrogen reacts with oxygen to form water: 2H₂+O₂→2H₂O. In this reaction, let’s say: 4 g of hydrogen reacts with 32 g of oxygen The total mass of the reactants = 4 g + 32 g = 36 g. The mass of the water formed = 36 g (same as the total reactants). Therefore there is no loss or gain of mass of atoms during the reaction. Antoine Lavoisier proved this law by conducting experiments in a closed system, where he burned phosphorus and sulfur and found the total mass remained the same before and after the reaction.

2. Law of Constant Proportions (also called Law of Definite Proportions) states that a chemical compound always contains the same elements in the same proportion by mass, regardless of the method of preparation or source No matter where a compound comes from or how it is made, it will always have the same proportion of elements. For example:
Water (H2O) always contains: Hydrogen and oxygen in the mass ratio 1:8. If you take:.

• 9 g of water → it will contain 1 g of hydrogen and 8 g of oxygen
• 90 g of water → it will contain 10 g of hydrogen and 80 g of oxygen

Joseph Proust studied different samples of copper carbonate from different sources. All samples had the same proportion of copper, carbon, and oxygen.

Concept of atom

An atom is the smallest unit of matter that takes part in a chemical reaction and retains the chemical properties of an element. Everything around us is made up of matter. This matter is made up of tiny particles called atoms. Atoms are extremely small — you cannot see them with a regular microscope. They combine to form molecules, which make up different substances. Maharishi Kanad (Indian philosopher) called the smallest particles as ‘parmanu’. Greek philosopher Democritus called it ‘atomos’, meaning indivisible. In the modern Concept (John Dalton), in 1808 proposed the Atomic Theory. He said that atoms are indivisible and indestructible, and combine in fixed ratios.

Characteristics of Atoms:

• Atoms of an element are identical in mass and properties.
• Atoms of different elements are different in mass and properties.
• Atoms can combine to form molecules and compounds.
• Atoms are neutral overall (equal protons and electrons).

Atomic Radius

The atomic radius is the distance from the nucleus of an atom to the outermost shell (valence shell) of electrons. It is measured in picometers (pm) or angstroms (Å). Atoms are extremely tiny and don’t have sharply defined boundaries. So, scientists estimate atomic radius using half the distance between two bonded atoms (in molecules or crystals). 1 pm=10−¹²m

 Examples of Atomic Radii (approximate)

Element	Atomic Number	Atomic Radius (pm)
H	1	37
Li	3	152
C	6	77
O	8	66
Na	11	186
Cl	17	99

Dalton’s Atomic Theory

According to Dalton’s Atomic Theory, atoms, which are indestructible and indivisible building blocks, make up all substances. Unlike other elements, which have atoms of different sizes and weights, an element’s atoms have all the same size and mass. Dalton proposed that the concept of atoms could be used to explain the laws of conservation of mass and definite proportions. He proposed that atoms, which he described as “solid, massy, hard, impenetrable, moving particle(s)”, are the smallest, indivisible units of matter.

Postulates of Dalton`s theory

1. The matter is made up of indivisible particles known as atoms.
2. The properties of all the atoms of a given element are the same, including mass. This can also be stated as – all the atoms of an element have identical mass and chemical properties; atoms of different elements have different masses and chemical properties.
3. Atoms of different elements combine in fixed ratios to form compounds.
4. Atoms are neither created nor destroyed. The formation of new products (compounds) results from the rearrangement of existing atoms (reactants) in a chemical reaction.
5. The relative number and kinds of atoms are constant in a given compound.

Drawbacks of Dalton's Atomic Theory

Dalton said that the atom is a smallest entity which cannot be divided further however, modern scientists proved that atom has divisible units called subatomic particles (protons, neutrons, electrons). Dalton’s theory doesn't explain isotopes, atoms of the same element with different masses. The table below briefs his limitations.

Postulate by Dalton	Why it’s a Drawback
Atoms are indivisible	Modern scientists disproved this statement by the discovery of electrons, protons
All atoms of an element are identical	Disproved by isotopes such as carbon isotopes.
Atoms of different elements are different	Disproved by the discovery of isobars
No concept of subatomic particles	Could not explain bonding or reactions
No concept of molecular structure	Could not explain compounds and valency

Elements and Their Symbols

John Dalton introduced a system of notations to represent elements in the early 1800s using circles and patterns. Later, Jöns Jakob Berzelius, a Swedish chemist, proposed the modern system of element symbols. Simple alphabetic representations still used today.

 Berzelius' Symbols of Elements

Berzelius suggested that each element should be represented by one or two letters of its English or Latin name. The first letter is always capitalized, and if there's a second letter, it is written in lowercase.

Types of Symbols

1. Single-Letter Symbols
   These are derived from the first letter of the element's name:
• C → Carbon
• B → Boron
• O → Oxygen
2. Two-Letter Symbols
   These use the first letter and one more letter from the name:
• Al → Aluminium
• Cl → Chlorine
• Ca → Calcium
3. Symbols from Latin Names
   Some elements use abbreviations of their Latin names:
• Fe → Iron (Ferrum)
• Na → Sodium (Natrium)
• K → Potassium (Kalium)
• Cu → Copper (Cuprum)

 Examples Table for symbols

Element	Symbol	Derived From
Carbon	C	English name
Aluminium	Al	English name
Iron	Fe	Latin (Ferrum)
Sodium	Na	Latin (Natrium)
Potassium	K	Latin (Kalium)
Copper	Cu	Latin (Cuprum)

Atomic Number

The atomic number of an element is the number of protons present in the nucleus of an atom of that element.

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Key Points:

• Symbol of atomic number = Z

• It is unique for every element.

• All atoms of the same element have the same atomic number.

• In a neutral atom,

  Number of protons=Number of electrons

Atomic No.	Element	Symbol	Common Use / Importance
1	Hydrogen	H	Found in water and fuels
2	Helium	He	Used in balloons and cooling systems
6	Carbon	C	Basis of all organic life and fuels
7	Nitrogen	N	Major component of air (78%)
8	Oxygen	O	Essential for respiration and combustion
9	Fluorine	F	Used in toothpaste and refrigerants
10	Neon	Ne	Used in neon signs and lighting
11	Sodium	Na	Found in table salt (NaCl)
12	Magnesium	Mg	Used in fireworks and medicines
13	Aluminium	Al	Used in utensils and aircraft
14	Silicon	Si	Used in electronics and glass
15	Phosphorus	P	Found in fertilizers and DNA
16	Sulphur	S	Used in gunpowder and medicines
17	Chlorine	Cl	Used in disinfectants and PVC
18	Argon	Ar	Used in light bulbs and welding
19	Potassium	K	Essential mineral for plants
20	Calcium	Ca	Important for bones and teeth
26	Iron	Fe	Used in construction and blood (hemoglobin)
29	Copper	Cu	Used in electrical wires and coins
79	Gold	Au	Used in jewellery and electronics

Atomic Mass and Atomic Mass Unit (AMU)

Atomic mass is the mass of a single atom of an element. Since atoms are extremely tiny, we cannot express their mass in grams directly because it would be an extremely small and impractical number. So, scientists use a relative mass scale based on a standard atom to compare the masses of different atoms. Atomic mass is the mass of one atom of an element, expressed relative to the mass of 1/12th of a carbon-12 atom. In simple terms, If the mass of a carbon-12 atom is taken as exactly 12 units, then the atomic mass of other elements is compared with it.

Why was Carbon-12 chosen as the standard?

Earlier, hydrogen (lightest element) was used as a standard. Later, carbon-12 was chosen because:

• It is more stable and common.
• It can form many organic and inorganic compounds.
• Its mass is more precise for comparisons.

Atomic Mass Unit (AMU or 'u')

Since atoms are incredibly small, their masses are also very small. So, scientists created a unit specifically for measuring atomic masses, called the atomic mass unit (u).

Definition of 1 Atomic Mass Unit:

Atomic mass unit (1 u) is defined as one-twelfth (1/12) the mass of one atom of the carbon-12 isotope. In simple terms, the mass of a single carbon-12 atom is taken as exactly 12 atomic mass units. Therefore, 1 u = 1/12 × mass of one carbon-12 atom. In terms of grams, 1 u=1.66×10⁻²⁴ grams 

Atomic Masses of Some Elements (Approximate):

Element	Symbol	Atomic Mass (in u)	Notes
Hydrogen	H	1.008 u	Lightest atom
Helium	He	4.0026 u	Noble gas
Carbon	C	12.01 u	Used as standard
Nitrogen	N	14.01 u	In proteins and air
Oxygen	O	16.00 u	Essential for respiration
Sodium	Na	22.99 u	Metal
Chlorine	Cl	35.5 u	Average due to isotopes

Atomicity

Atomicity is the number of atoms present in a single molecule of an element or compound. In simple words, atomicity tells us how many atoms are bonded together to form one molecule of a substance. Atomicity helps us understand the structure of molecules and it tells us how elements and compounds exist in nature — whether as single atoms or as molecules made of multiple atoms.

Types of Atomicity and Examples

Atomicity	No. of Atoms	Examples
Monatomic	1	Helium (He), Neon (Ne), Argon (Ar)
Diatomic	2	Hydrogen (H2), Chlorine (Cl2), Nitrogen (N2)
Triatomic	3	Ozone (O3), Water (H2O)
Tetratomic	4	Phosphorus (P4)
Polyatomic	More than 4	Sulfur (S8)

 

 

Molecular Mass

Molecular mass is the sum of the atomic masses of all the atoms present in a molecule of a substance. A molecule is made up of two or more atoms chemically bonded together. To calculate the molecular mass, we simply add the atomic masses of each atom in the molecule. It is expressed in atomic mass units (u) because atomic masses are measured in u.

How to Calculate Molecular Mass

Step 1: Write the chemical formula of the compound. The chemical formula tells you which atoms and how many of each are present in the molecule.

Step 2: List the atomic masses of all the elements involved by using periodic table or standard values. Example: Hydrogen (H) = 1 u, Oxygen (O) = 16 u, Carbon (C) = 12 u, etc.

Step 3: Multiply the atomic mass of each element by the number of atoms of that element in the molecule. Therefore, mass of each element=Number of atoms X Atomic mass.

Step 4: Add the total mass of all atoms in the molecule. So, Molecular Mass=Sum of all atomic masses in the molecule

Examples to calculate atomic masses in the molecule

 Example 1: Water (H2O)

1. Chemical formula = H2O
2. Atomic masses = H = 1 u, O = 16 u
3. Number of atoms:
   • 2 Hydrogen → 2×1=2
   • 1 Oxygen → 1×16=16
4. Total molecular mass =H₂0 . Therefore= 2+16=18 u

Example 2: Sodium Chloride (NaCl)

1. Formula: NaCl
2. Atomic masses: Na = 23 u, Cl = 35.5 u
3. Total molecular mass of NaCl= 23+35.5=58.5 u

Molecule concept

A molecule is the smallest particle of a substance that can exist independently and still retain the chemical properties of that substance. Molecules are made up of two or more atoms that are chemically bonded together. These atoms can be of the same element or of different elements. Molecules can be solid, liquid, or gas depending on the substance.

 Types of Molecules:

Type	Description	Example
Elemental Molecule	Made of atoms of the same element	O2 (Oxygen), N2 (Nitrogen), H2 (Hydrogen)
Compound Molecule	Made of atoms of different elements	H2O (Water), CO2 (Carbon dioxide), NH3 (Ammonia)

 Examples of Molecules:

Molecule	Composition	Type of Molecule
H2	2 Hydrogen atoms	Elemental molecule
O2	2 Oxygen atoms	Elemental molecule
N2	2 Nitrogen atoms	Elemental molecule
H2O	2 Hydrogen + 1 Oxygen	Compound molecule
CO2	1 Carbon + 2 Oxygen	Compound molecule
CH4	1 Carbon + 4 Hydrogen	Compound molecule
NH3	1 Nitrogen + 3 Hydrogen	Compound molecule
C6H12O6	6 Carbon, 12 H, 6 O	Compound molecule (Glucose)

Molecule vs Atom:

Atom	Molecule
Smallest unit of an element	Smallest unit of a compound or element
May or may not exist freely	Always exists independently
Example: H, O, Na, Cl	Example: H2, O2, H2O, CO2

 Real-Life Examples of molecules

Water (H2O) – Every drop is made of water molecules. Oxygen (O2) – The air we breathe contains oxygen molecules. Sugar (C12H22O11) – A sweet substance made of large compound molecules.

Example: Formation of Water

When2 hydrogen atoms (H) Combine with 1 oxygen atom (O)
→ They form 1 molecule of water (H2O). So, H2 + O → H2O

Table: Molecular Mass of Some Common Compounds

Compound	Formula	Molecular Mass (u)
Water	H2O	18
Carbon dioxide	CO2	44
Oxygen molecule	O2	32
Methane	CH4	16
Ammonia	NH3	17

Difference Between Molecule and Compound

Feature	Molecule	Compound
Definition	A molecule is formed when two or more atoms bond chemically.	A compound is a molecule made of two or more different elements.
Same/Different Elements	Can be made of same or different atoms.	Must contain different types of atoms.
Example	O2 (oxygen gas), H2 (hydrogen gas)	H2O (water), CO2 (carbon dioxide), NaCl (salt)
Type	All compounds are molecules. Not all molecules are compounds.	All compounds are molecules.
Bond Type	Can involve covalent or ionic bonds	Usually involve covalent or ionic bonds
Existence	Can exist freely as single units	Often form crystalline or structured arrangements

In Simple Words, A molecule = 2 or more atoms bonded (same or different) but Compound = Molecule made of different elements only

 Examples to Understand

Substance	Chemical Formula	Molecule or Compound?	Why?
Hydrogen gas	H2	Molecule only	2 same hydrogen atoms
Oxygen gas	O2	Molecule only	Same element (oxygen)
Water	H2O	Molecule and Compound	2 hydrogen + 1 oxygen
Carbon dioxide	CO2	Molecule and Compound	1 carbon + 2 oxygen
Sodium chloride	NaCl	Molecule and Compound	Sodium + Chlorine

 Molecules of Compounds

Compounds are of two main types, based on the type of bond holding the atoms together:

1. Molecular Compounds

• Held together by covalent bonds (shared electrons).
• Made of non-metal + non-metal.
• Exist as discrete molecules.
• Examples: H2O (water), CO2 (carbon dioxide), NH3 (ammonia)

2. Salts / Ionic Compounds

• Held together by ionic bonds (electron transfer).
• Formed between metals and non-metals.
• Do not exist as single molecules but as crystalline solids or ionic lattices.
• Examples: NaCl (sodium chloride), MgCl2 (magnesium chloride), KBr (potassium bromide)

 How to Identify if It’s a Molecule or Compound

If it is...	Then it is...
Made of same atoms only	A molecule, not a compound (e.g., O2)
Made of different atoms	A compound and molecule (e.g., H2O)
Held together by covalent bonds	A molecular compound
Held together by ionic bonds	An ionic compound or salt

Mole Concept

The mole is a standard scientific unit for measuring large quantities of very small entities like atoms, molecules, or ions. One mole of any substance contains exactly 6.022 × 10²³ particles. These particles can be atoms, molecules, ions, or electrons, depending on the context. This number is known as Avogadro’s Number.

Avogadro’s Number

Avogadro's Number= 6.022 × 10²³ particles/mole.It is named after the Italian scientist Amedeo Avogadro. It helps relate the number of microscopic particles to a macroscopic mass of a substance. Example:1 mole of hydrogen atoms (H) = 6.022 × 10²³ atoms of hydrogen. Similarly, 1 mole of H2O molecules = 6.022 × 10²³ water molecules.

Molar Mass

The molar mass of a substance is the mass of one mole of that substance. It is expressed in grams per mole (g/mol). Molar mass = Sum of atomic masses of all atoms in the molecule. Some examples are: Molar mass of H2O = (2 × 1.008) + (1 × 16.00) = 18.02 g/mol. Molar mass of CO2 = (1 × 12.01) + (2 × 16.00) = 44.01 g/mol

Atomic Valency

Valency is the combining capacity of an atom. It tells us how many electrons an atom can gain, lose, or share to become stable.

Examples:

Element	Atomic Number	Valency
Hydrogen	1	1
Oxygen	8	2
Nitrogen	7	3
Carbon	6	4

 Molecules and Atomicity

A molecule is the smallest unit of a compound or element that retains its chemical properties. Molecules are made of two or more atoms bonded together. Atomicity refers to the number of atoms present in one molecule of an element.

Atomicity of Some Elements

Element	Atomicity	Example Formula
Hydrogen	2	H2
Oxygen	2	O2
Nitrogen	2	N2
Ozone	3	O3
Phosphorus	4	P4
Sulfur	8	S8

 Structure of an Atom

An atom is the basic unit of matter. It consists of three fundamental subatomic particles:

Particle	Charge	Location
Proton	+1 (positive)	In the nucleus
Neutron	0 (neutral)	In the nucleus
Electron	-1 (negative)	Outside the nucleus, in orbits or shells

Important Points:

• The nucleus is at the center of the atom and contains protons and neutrons.
• Electrons revolve around the nucleus in specific energy levels called shells or orbitals.
• The total mass of an atom is concentrated in the nucleus.

Example:

• An atom of carbon (C) has:
  • 6 protons
  • 6 neutrons
  • 6 electrons

Valency

Valence electrons are those electrons which are present in the outermost orbit of the atom.

• The capacity of an atom to lose, gain or share valence electrons in order to complete its octet determines the valency of the atom.

Writing Chemical Formulae

Compounds

• When two or more elements chemically combine in a fixed ratio by mass, the obtained product is known as a compound.
• Compounds are substances consisting of two or more different types of elements in a fixed ratio of its atoms.

Ions

• An ion is defined as an atom or molecule which has gained or lost one or more of its valence electrons, giving it a net positive or negative charge.
• A negatively charged particle is called an anion, and a positively charged particle is called a cation.

Ionic Compounds: Chemical Formula

Each constituent element in a chemical formula is identified by its chemical symbol, along with the relative number of atoms that make up each element. These ratios are used in empirical equations to start with a key element and then assign atom counts for the remaining elements in the compound in relation to the key element.

• Ionic compounds are chemical compounds in which ions are held together by specialised bonds called ionic bonds.
• An Ionic compound always contains an equal amount of positive and negative charge.
• For example: In calcium chloride, the ionic bond is formed by oppositely charged calcium and chloride ions.
  The calcium atom loses 2 electrons and attains the electronic configuration of the nearest noble gas (Ar). By doing so, it gains a net charge of +2.
  The two chlorine atoms take one electron each, thus gaining a charge of -1 (each) and attain the electronic configuration of the nearest noble gas (Ar).

TEXTBOOK QUESTION ANSWERS

Question 1. In a reaction, 5.3 g of sodium carbonate reacted with 6 g of ethanoic acid. The products were 2.2 g of carbon dioxide, 0.9 g water and 8.2 g of sodium ethanoate. Show that these observations are in agreement with the law of conservation of mass. Sodium carbonate + ethanoic acid ? sodium ethanoate + carbon dioxide + water
Answer: Sodium carbonate + ethanoic acid ? sodium ethanoate + carbon dioxide + water
5.3 + 6 ? 8.2 + 2.2 + 0.9
= 11.3 g = 11.3 g
Weight of reactants is equal to weight of products. This observation is in agreement with the law of conservation of mass.

Question 2. Hydrogen and Oxygen combine in the ratio or 1 : 8 by mass to form water. What mass of oxygen gas would be required to react completely with 3 g of hydrogen gas?
Answer: 24 g of oxygen gas would be required to react completely with 3 g of hydrogen gas. Because in water the ratio of the mass of hydrogen to the mass of oxygen is always 1 : 8.

Question 3. Which postulate of Dalton’s atomic theory is the result of the law of conservation of mass?
Answer: The relative number and kinds of atoms are constant in a given compound. This postulate is the result of the law of conservation of mass.

Question 4. Which postulate of Dalton’s atomic theory can explain the law of definite proportions?
Answer: Atoms combine in the ratio of small whole numbers to form compounds. This postulate explains the law of definite proportions.

Question 1. Define the atomic mass unit.
Answer: One atomic mass unit is a mass unit equal to exactly one-twelfth (1/12^th) the mass of one atom of carbon-12.

Question 2. Why is it not possible to see an atom with naked eyes?
Answer: Atoms are very small, they are smaller than anything that we can imagine or compare with. Therefore it is not possible to see an atom with naked eyes.

Text book questions-2

Question 1. Write down the formulae of:

1. sodium oxide
2. alluminium chloride
3. sodium sulphide
4. magnesium hydroxide

Answer:

1. sodium oxide: Na₂O
2. aluminium chloride: Al₂Cl₃
3. sodium sulphide : NaS
4. magnesium hydroxide: Mg(OH)₂

Question 2. Write down the names of compounds represented by the following formulae:

1. Al₂(SO₄)₃
2. CaCl₂
3. K₂SO₄
4. KNO₃
5. CaCO₃

Answer:

1. Al₂(SO₄)₃: Aluminium sulphate
2. CaCl₂: Calcium chloride
3. K₂SO₄: potassium sulphate
4. KNO₃: potassium nitrate
5. CaCO₃: calcium carbonate

Question 3. what is meant by the term chemical formula?
Answer: The chemical formula of a compound is a symbolic representation of its composition.

Question 4. How many atoms are present in a
i) H₂S molecule and
ii) PO₄^3- ion?

Species	Calculation	Total Atoms
H₂S	H: 2, S: 1 → 2 + 1	3
PO₄³⁻	P: 1, O: 4 → 1 + 4	5

Question 1. Calculate the molecular masses of H₂, O₂, Cl₂, CO₂, CH₄, C₂H₆, C₂H₄, NH₃, CH₃OH.

Molecule	Calculation	Molecular Mass (g/mol)
H₂	1 × 2	2
O₂	16 × 2	32
Cl₂	35.5 × 2	71
CO₂	12 + (16 × 2)	44
CH₄	12 + (1 × 4)	16
C₂H₆	(12 × 2) + (1 × 6)	30
C₂H₄	(12 × 2) + (1 × 4)	28
NH₃	14 + (1 × 3)	17
CH₃OH	12 + (1 × 4) + 16	32

Question 2. Calculate the formula unit masses of ZnO, Na₂O, K₂CO₃ 

Compound	Calculation	Formula Unit Mass (u)
ZnO	Zn: 65 + O: 16 → 65 + 16	81
Na₂O	2 × Na: 2 × 23 + O: 16 → 46 + 16	62
K₂CO₃	2 × K: 2 × 39 + C: 12 + 3 × O: 3 × 16 → 78 + 12 + 48	138

Question 1. If one mole of carbon atoms weighs 12 gms, what is the mass (in gms) of 1 atom of carbon?
Answer:
Number of moles = n
Given mass = m
molar mass = M
Given the number of particles = N
Avogadro number of particles = N₀
i) Atomic mass of carbon = 12u.
Atomic mass of one mole of carbon=12g

Question 2. Which has more number of atoms, 100 g, 100 gms of sodium or 100 gms of iron (given, the atomic mass of Na=23u, Fe = 56u).
Answer:
Atomic mass of sodium = 23u (data)
It means the gram atomic mass of sodium = 23 gm
Now atoms present in 23 gm sodium = 6.022 × 10²³

It means Number of atoms in 100 gm sodium = 1.6753 × 10²⁴
∴ 100 gm sodium has more number of atoms rather than 100 gms of iron.

Textbook Exercises

Question 1. A 0.24 g sample of a compound of oxygen and boron was found by analysis to contain 0.096g of boron and 0.144 g of oxygen. Calculate the percentage composition of the compound by weight.
Answer:
Mass of Boron = 0.096 g (Data)
Mass of Oxygen = 0.144g (Data)
Given mass = 0.24 g (Data)
∴ The percentage composition of Boron

∴ The percentage of O₂ = 0.144/0.24 × 100
= 60%.

Question 2. When 3.00 g of carbon is burnt in 8.009 of oxygen, 11.00 g of carbon dioxide is produced. What mass of carbon dioxide will be formed when 3.00 g of carbon is burnt in 50,000 g of oxygen? Which law of Chemical combination will govern your answer?
Answer:
3.0 g of carbon is burnt in 8.00 g of oxygen, 11.00 g of carbon dioxide is produced.
But when 3 g of carbon is burnt in 50 g of oxygen, only 3 g of carbon reacts with 8 g of oxygen.
Remaining 42 gm of oxygen will not react.
11 gm of carbon dioxide is produced.
∴ Our answer obeys law of constant proportion.

Question 3. What are polyatomic ions? Give examples.
Answer: A group of atoms carrying a charge is known as a polyatomic ion.

Question 4. Write the chemical formulae of the following.
(a) Magnesium chloride
(b) Calcium oxide
(c) Copper nitrate
(d) Aluminium chloride
(e) Calcium carbonate
Answer:
(a) Magnesium chloride : MgCl₂
(b) Calcium oxide : CaO
(c) Copper nitrate : Cu(NO₃)₂
(d) Aluminium chloride : AlCl₃
(e) Calcium carbonate : CaCO₃.

Question 5. Give the names of the elements present in the following compounds.
(a) Quick lime
(b) Hydrogen bromide
(c) Baking powder
(d) Potassium sulphate
>Answer:
(a) Quick lime: Calcium, carbon, oxygen
(b) Hydrogen bromide: Hydrogen, Bromine
(c) Baking powder: Sodium, Bicarbonate
(d) Potassium sulphate: Potassium, Sulphur, Oxygen.

Question 6.Calculate the molar mass of the following substances.
(a) Ethyne, C₂H₂
(b) Sulphur molecule, S₈
(c) Phosphorus Molecule, P₄ (Atomic mass of phosphorus = 31)
(d) Hydrochloric acid, HCl
(e) Nitric acid HNO₃.

Answer: (a) Ethyne C₂H₂
Molar Mass = 2 × 12 + 2 × 1
= 24 + 2
= 26 g.

(b) Molar mass of Sulphur molecule
= 8 × 32
= 256 g.

(c) Molar Mass of Phosphorus molecule = 4 × 31
(Atomic mass of Phosphorus) = 124 g.

(d) Molar mass of Hydrochloric acid = HCl
= 1 + 35.5
= 36.5 g.

(e) Molar mass of HNO₃
= 1 + 14 + 3 × 16
= 15 + 48
= 63 gm.

Question 7.What is the mass of –
(a) 1 mole of nitrogen atoms?
(b) 4 moles of aluminium atoms (Atomic mass of Aluminium = 27)
(c) 10 moles of Sodium Sulphite (Na₂SO₃)
g>Answer:
(a) Mass of 1 mole of nitrogen= 14g.

(b) Mass of 4 moles of Aluminium
= 4 × 27
= 108 g.

(c) Mass of 10 moles of Sodium sulphite
= 10 × [2 × 23 + 32 + 3 × 16]
= 10 × 126
= 1260 gm.

Question 8.Convert into mole:
(a) 12 g of oxygen gas
(b) 20 g of water
(c) 22 g of carbon dioxide

Answer: (a) 32gm of oxygen means 1 mole
12gm oxygen means 12/32 mole
∴ 12g of oxygen gas = 0.375 mole
(b) 18 gm of water means = 1 mole
20 gm of water means 20/18 mole = 1.11 mole
(c) 22 g of carbon dioxide means
22/44 = 0.5 mole.

Question 9. What is the mass of:
(a) 0.2 mole of oxygen atoms?
(b) 0.5 mole of water molecules?
Answer:(a) Mass of 1 mole of oxygen = 16 g
Mass of 0.2 mole of oxygen = 0.2 × 16
= 3.2 g.
(b) Mass of 1 molecule of water= 18 gm.
Mass of 0.5 mole of water = 0.5 × 18
= 9 gm.

Question 10. Calculate the number of molecules of sulphur (S₈) present in 16 g of solid sulphur.
Answer: One mole of Sulphur (S) = 8 × 32
= 256 gm.
256 g of solid sulphur = 6.022 × 10²³ molecules.

= 3.76 × 10²²
(Approximate)

Question 11. Calculate the number of aluminium ions present in 0.0519 of aluminium oxide.
(Hint: The mass of an ion is the same as that of an atom of the same element. Atomic mass of Al = 27 u).
Answer: One mole of Aluminium oxide
= 2 × 27+ 3 × 16
= 54 + 48
= 102 g.
102 g of Al₂O₃ = 6.022 × 10²³ moles (aluminium oxide)

It means Aluminium present in 0.051 gm
= 3.011 × 10²⁰ Aluminium oxide molecules
Number of Al ions in one mole of Al₂O₃ = 2
∴ Number of Al Ions In 3.011 × 10²⁶ molecules
0.051 Al₂O₃
= 2 × 3.011 × Number of Al ions In 3.011 × 10²⁰
= 6.022 × 10

Additional Questions

Question 1. Write the symbols of the following:
a) Iron b) lead c) Zinc d) Oxygen e) Chlorine
Answer:
a) Iron = Fe b) lead = Pb c) Zinc = Zn d) Oxygen = O e) Chlorine = Cl

Question 2. What is a molecule?
Answer: The smallest particle of an element or a compound that is capable of independent existence and shows all the properties of that substance.

Question 3. What is an ion?
Answer: An ion is a charged species present in metals and non-metals.

Question 4. What is molecular mass?
Answer: The molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance.

Question 5. What is a mole?
Answer: One mole of any species (atoms, molecules, ions, or particles) is that quantity in number having a mass equal to its atomic or molecular mass in grams.

Question 6. An element Z forms an oxide with formula Z₂O₃. What is its valency?
Answer :
Valency is 3⁺

Question7. Mention the elements present in (1) quick line (2) sodium hydrogen carbonate.,
Answer :
(1) Quick line (calcium oxide) (CaO) element present are calcium and oxygen.
(2) Sodium hydrogen carbonate (NaHCO₃) elements are sodium, hydrogen, hydrogen, carbon and oxygen.

Question 8. Calculate the total number of ions in 0.585 g of sodium chloride.
Answer : Gram formula mass of NaCl = 23 + 35-5 = 58.5 g
58.5 g of NaCl have ions = 2 × N_A
58.5 g of NaCl have ions
=2×N_A×0.585/58⋅5=0⋅02×N_A
= 0.02 × 6.022 x 10²³
= 1.20 × 10²² ions

Question 9.> Which of the following represents a correct chemical formula? Name it.

(a) CaCl

(b) BiPO₄

(c) NaSO₄

(d) NaS

Answer: (b), BiPO₄ represents the correct formulae of bismuth phosphate.

Question 10. Write the molecular formulae for the following compounds

(a) Copper bromide

(b) Aluminium nitrate

(c) Calcium phosphate

(d) Iron sulphide

(e) Mercury chloride

(f) Magnesium acetate

Answer:

(a) The molecular formula of Copper bromide is CuBr.

(b) The molecular formula of Aluminium nitrate is Al(NO₃)₃.

(c) The molecular formula of Calcium phosphate is Ca₃(PO₄)₂.

(d) The molecular formula of Iron sulphide is Fe₂S₃.

(e) The molecular formula of Mercury chloride is HgCl₂.

(f) The molecular formula of Magnesium acetate is Mg(CH₃COO)₂.

Question 11.Write the molecular formulae of all the compounds that can be formed by the combination of the following ions. Cu²⁺, Na⁺, Fe³⁺, Cl^–, SO₄^2-, PO₄^3-.

Answer:The molecular formula of the compounds formed by the combination of Cu²⁺, Na⁺, Fe³⁺, Cl^–, SO₄^2-, and PO₄^3- are CuCl₂, CuSO₄, NaCl, Na₂SO₄, FeCl₃, and Fe₂(SO₄)₃.

Question 12. Write the cations and anions present (if any) in the following compounds

(a) CH₃COONa

(b) NaCI

(c) H₂

(d) NH₄NO₃

Answer:

(a) The cation and anion present in CH₃COONa are Na⁺ and CH₃COO^–.

(b) The cation and anion present in NaCI are Na⁺ and Cl^–.

(c) There is no cation and anion in H₂.

(d) The cation and anion present in NH₄NO₃ are NH₄⁺ and NH₃^–.

Question 13. Give the formulae of the compounds formed from the following sets of elements

(a) Calcium and fluorine

(b) Hydrogen and sulphur

(c) Nitrogen and hydrogen

(d) Carbon and chlorine

(e) Sodium and oxygen

(f) Carbon and oxygen

Answer:

(a) The formulae of the compound formed by Calcium and fluorine is CaFl₂.

(b) The formulae of the compound formed by Hydrogen and sulphur is H₂S.

(c) The formulae of the compound formed by Nitrogen and hydrogen is NH₃.

(d) The formulae of the compound formed by Carbon and chlorine is CCl₄.

(e) The formulae of the compound formed by Sodium and oxygen is Na₂O.

(f) The formulae of the compound formed by Carbon and oxygen is CO₂.

Question 14. Which of the following symbols of elements are incorrect? Give their correct symbols

S. No.	Element	Formula
1.	Cobalt	CO
2.	Carbon	c
3.	Aluminium	AL
4.	Helium	He
5.	Sodium	So

Answer: The formula for cobalt, carbon, aluminium, and sodium is incorrect, while the formula for helium is correct.

The correct formulas are enlisted below.

S. No.	Element	Correct Formula
1.	Cobalt	Co
2.	Carbon	C
3.	Aluminium	Al
4.	Helium	He
5.	Sodium	Na

Question 15 Give the chemical formulae for the following compounds and compute the ratio by mass of the combining elements in each one of them. (You may use appendix-III).

(a) Ammonia

(b) Carbon monoxide

(c) Hydrogen chloride

(d) Aluminium fluoride

(e) Magnesium sulphide

Answer:

(a) The chemical formula of ammonia is NH_3, and its mass ratio is Mass of N: Mass of H = 14: 3.

(b) The chemical formula of Carbon monoxide is CO_, and its mass ratio is Mass of C: Mass of O = 12: 16 = 3:4.

(c) The chemical formula of Hydrogen chloride is HCl_, and its mass ratio is Mass of H: Mass of Cl = 1: 35.5.

(d) The chemical formula of Aluminium fluoride is AlF_3, and its mass ratio is Mass of Al: Mass of F = 27: 19.

(e) The chemical formula of Magnesium sulphide is MgS_, and its mass ratio is Mass of Mg: Mass of S = 24: 32 = 3: 4.

Question 16 State the number of atoms present in each of the following chemical species

(a) CO₃^2-

(b) PO₃^3-

(c) P₂O₅

(d) CO

Answer:

(a) There are four atoms in CO₃^2-.

(b) There are four atoms in PO₃^3-.

(c) There are seven atoms in P₂O₅.

(d) There are two atoms in CO.

Question 17. What is the fraction of the mass of water due to neutrons?

Answer:

The mass of one neutron = 1 amu.

The mass of one water molecule = 18 amu.

The oxygen atom has eight neutrons, while the hydrogen atom has 0 neutrons.

So the mass of neutrons in one water molecule is eight amu.

The fraction of mass of water due to neutrons = 8 / 18 = 4 / 9.

Q18. Does the solubility of a substance change with temperature? Explain with the help of an example.

Answer:

Yes, the solubility of a substance changes with temperature. The solubility generally increases with an increase in temperature.

Example: You can dissolve more sugar in hot water than in cold water.

Q19. Classify each of the following based on their atomicity.

(a) F₂ (b) NO₂ (c) N₂O (d) C₂H₆  (e) P₄ (f) H₂O₂ (g) P₄O₁₀ (h) O₃ (i) HCI (i) CH₄

(k) He (l) Ag

Substance	Name	No. of Atoms in One Molecule	Atomicity	Type
(a) F₂	Fluorine	2	Diatomic	Molecule of element
(b) NO₂	Nitrogen dioxide	3 (1 N + 2 O)	Triatomic	Molecule of compound
(c) N₂O	Nitrous oxide	3 (2 N + 1 O)	Triatomic	Molecule of compound
(d) C₂H₆	Ethane	8 (2 C + 6 H)	Octatomic	Molecule of compound
(e) P₄	Phosphorus	4	Tetra-atomic	Molecule of element
(f) H₂O₂	Hydrogen peroxide	4 (2 H + 2 O)	Tetra-atomic	Molecule of compound
(g) P₄O₁₀	Phosphorus pentoxide	14 (4 P + 10 O)	Tetradecatomic	Molecule of compound
(h) O₃	Ozone	3	Triatomic	Molecule of element
(i) HCl	Hydrogen chloride	2 (1 H + 1 Cl)	Diatomic	Molecule of compound
(j) CH₄	Methane	5 (1 C + 4 H)	Pentatomic	Molecule of compound
(k) He	Helium	1	Monoatomic	Atom of element
(l) Ag	Silver	1	Monoatomic	Atom of element

Q12. You are provided with a fine white colored powder, either sugar or salt. How would you identify it without tasting it?

Answer: We can differentiate sugar and salt by

(a) Heating salts separately. Sugar will melt while salt will not.

(b) Dissolving them separately in water. The salt solution will conduct electricity due to Na⁺ ion and Cl^– while the sugar solution will not conduct electricity. So, we can immediately tell the difference by testing a drop of the solution with an ohmmeter.

Q13. Calculate the number of moles of magnesium present in a magnesium ribbon weighing 12 g. The molar atomic mass of magnesium is 24g mol^-1.

Answer:Given

Mass of magnesium ribbon = 12 g

Molar mass of magnesium = 24 g

Number of moles = Mass / Molar Mass

Number of moles = 12 / 24

Number of moles = 0.5 moles.

Hence, there are half moles of magnesium in a 12 g magnesium ribbon.