September 26, 2022

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# Atom and Molecules Notes

Chapter -3

Atom and Molecules

Topics in the Chapter

• Introduction
• Laws of Chemical Combination

→ Law of Conservation of Mass

→ Law of Constant Proportions

• Dalton’s Atomic Theory
• Atoms

→ Modern day symbols of Elements

→ Atomic Mass

→ Atom existence

• Molecules
• Atomicity
• Chemical Formulae

→ Characteristics of chemical formulae

→ Rules for writing chemical formulae

• Molecular Mass

→ Formula Unit Mass

• Ions
• Mole Concept

→ Molar Mass

→ Important Formulae

Introduction

Maharishi Kanad – Around 500 B.C., Indian philosopher Maharishi Kanad, postulated the theory if we go on dividing matter (padarth), we will obtain smallest particle beyond which further division can’t be possible which is known as ‘parmanu’.

Pakudha Katyayama – He postulated that there are various forms of matter because the particles of matter exist together in combinations.

Democritus and Leucippus – Greek philosophers Democritus and Leucippus suggested that when we keep on dividing the matter there comes a time when no more division of particles can take place. Such particles are called atoms which means being invisible.

But all these ideas were not backed up by many experimental pieces of evidence until Antoine L. Lavoisier and Joseph L. Proust provided two laws of chemical combination.

Laws of chemical combination.

• Law of Conservation of Mass
• During a chemical reaction, the total mass of reactants will be equal to the total mass of the products.
• Mass can neither be created nor destroyed in a chemical reaction.
• Example: A (reactant) + B (reactant) → AB (product)
• Mass of A + Mass of B = Mass of AB
• Law of Constant Proportions
• In a chemical substance the elements are always present in definite proportions by mass.
• For example:
• 18 gm of H2O = 2 gm of hydrogen + 16 gm of oxygen

⇒ mass of hydrogen/mass of oxygen = 2/16 = 1/8

• 36 gm of H2O = 4 gm of hydrogen + 32 gm of oxygen

⇒ mass of hydrogen/mass of oxygen = 4/32 = 1/8

• 9 gm of H2O = 1 gm of hydrogen + 8 gm of oxygen

⇒ mass of hydrogen/mass of oxygen = 1/8

• This verifies law of constant proportions as the ratio of mass of hydrogen to oxygen is always same

Dalton’s Atomic Theory

• According to Dalton’s atomic theory, all matter, whether an element, a compound or a mixture is composed of small particles called atoms.

• Six Postulates of Dalton’s atomic theory:
1. All matter is made of very tiny particles called atoms.
2. Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.
3. All atoms of an element are same in size, mass and chemical properties.
4. Atoms of different elements have different masses and chemical properties.
5. Atoms combine in the ratio of small whole numbers to form compounds.
6. The relative number and kinds of atoms are constant in a given compound

Atoms

• Atoms are building blocks of all matter.
• An atom is the smallest particle of an element which takes part in chemical reaction.
• Atoms are very small and which can’t be seen even through very powerful microscope.
• Atomic radius is measured in nanometres. 1 Nanometres = 10-9 m
 Relative Sizes Radius (in meter) Example 10-10 Atom of hydrogen 10-9 Molecule of water 10-8 Molecule of haemoglobin 10-4 Grain of Sand 10-2 Ant 10-1 Watermelon

Symbols of Atoms or Elements

• Dalton was the first scientist to use the symbols for elements.

• Berzilius suggested that the symbols of elements should be made from one or two letters of the name of the element.
• Now-a-days, IUPAC (International Union of Pure and Applied Chemistry) approves names of elements
• The first letter of a symbol is always written as a capital letter (upper-case) and the second letter as a small letter (lower-case).
• Some other symbols have been taken from the names of elements in Latin, German or Greek.
• Symbols of some common elements:
 S.N. Name of the element Latin name Symbol 1 Hydrogen H 2 Helium He 3 Carbon C 4 Copper Cuprum Cu 5 Cobalt Co 6 Chlorine Cl 7 Cadmium Cd 8 Boron B 9 Barium Ba 10 Bromine Br 11 Bismuth Bi 12 Sodium Natrium Na 13 Potassium Kalium K 14 Iron Ferrum Fe 15 Gold Aurum Au 16 Silver Argentum Ag 17 Mercury Hydragyrum Hg

• Atomic Mass
• Dalton’s atomic theory proposed the idea of atomic mass which explained the law of constant proportions so well.
• The mass of an atom of an element is called its atomic mass.
• In 1961, IUPAC have accepted ‘atomic mass unit’ (u) to express atomic and molecular mass of elements and compounds.
• The atomic mass unit is defined as the quantity of mass equal to 1/12 of mass of an atom of carbon-12.
• 1 amu or u = 1/12 × Mass of an atom of C12
1 u = 1.66 × 10-27kg
 Element Atomic Mass Hydrogen 1 µ Carbon 12 µ Nitrogen 14 µ Oxygen 16 µ Sodium 23 µ Magnesium 24 µ Sulphur 32 µ Chlorine 35.5 µ Calcium 40 µ

• Atom existence
• Atoms of most of the elements are very reactive and does not exist in free state.
• Only the atoms of noble gases (such as He, Ne, Ar, Kr, Xe and Rn) are chemically unreactive and can exist in the free state as single atom.
• Atoms of all other elements combine together to form molecules or ions.

Atom (electrically neutral)

Ion (electrically charged)                                    Molecules (electrically neutral)

• Molecules
• A molecule is in general a group of two or more atoms that are chemically bonded together.
• A molecule is the smallest particle of matter (except element) which is capable of an independent existence and show all properties of that substance.
• Examples: ‘H2O’ is the smallest particle of water which shows all the properties of water.
• A molecule may have atom of same or different elements, depending upon this, molecule can be categorized into two categories:
(i) Homoatomic molecules(containing atom of same element)
Examples: H2, O2, O3, S8, P4

(ii) Heteroatomic molecules or compounds (containing atoms of different elements)
Examples:  H2O, CO2, NaCl, CaCO3 etc.

• Atomicity
• The number of atoms present in one molecule of an element is called its atomicity.
 S.N. Name of the element Atomicity Molecules formula 1 Helium Monoatomic He 2 Sodium Monoatomic Na 3 Hydrogen Di-atomic H2 4 Chlorine Di-atomic Cl2 5 Phosphorus Polyatomic (Tetra) P4 6 Sulphur Polyatomic (Octa) S8

• Ions
• When atoms, groups of atoms or molecules lose or gain an electron (s) they become charged. These elements include charged particle are known as ions.
• Anion– Negatively charged ion.       g., Cl, Br, S–- etc.
• Cation– Positively charged ion.        g., Na+, K+, Al+++ etc.
• A group of atoms carrying charge a single entity is known as a polyatomic ion. e.g., NO3, SO4
• We can classify ions in two types:
(i) Simple ions
• Mg2+ (Magnesium ion)
• Na+(Sodium ion)
• Cl (Chloride ion)
• Al3+(Aluminium ion)

(ii) Compound ions

• NH4+(Ammonium ion)
• CO32-(Carbonate ion)
• SO42-(Sulphate ion)
• OH(Hydroxide ion)
• NO3(Nitrate ion)
• Valency
• The capacity of an atom to lose, gain or share valence electrons in order to complete its octet determines the valency of the atom.
• Every atom wants to become stable, to do so it may lose, gain or share electrons.
• If an atom consists of 1, 2 or 3 electrons in its valence shell then its valency is 1, 2 or 3 respectively,
• If an atom consists of 5, 6 or 7 electrons in the outermost shell, then it will gain 3, 2 or 1 electron respectively and its valency will be 3, 2 or 1 respectively.
• If an atom has 4 electrons in the outermost shell than it will share this electron and hence its valency will be 4. If an atom has 8 electrons in the outermost electron and hence its valency will be 0.
• Chemical Formulae
• It is the symbolic representation of the composition of a compound.
• The valencies or charges on ion must balance.
• When a compound is formed of metal and non-metal, symbol of metal comes first. E.g., CaO, NaCl, CuO.
• When polyatomic ions are used, the ions are enclosed in brackets before writing the number to show the ratio. E.g., Ca(OH)2, (NH4)2SO4
• Rules for writing chemical formulae

(i) We first write symbols of elements which form compound.

(ii) Below the symbol of each element, we should write their valency.

(iii) Now cross over the valencies of combining atoms.

Examples:

• Molecule
• A molecule is a collection of various atoms that combine chemically with each other.
• Atoms of the same elements or different elements can bind together to form molecules.
• Therefore, a molecule is the smallest particle of a substance that can exist independently and shows all the properties of that substance.
• Molecular Mass
• It is the sum of the atomic masses of all the atoms in a molecule of the substance. It is expressed in atomic mass units (amu).
• For example, the molecular mass of HNO3 can be calculated as:

Atomic mass of H =1u

Atomic mass of N =14u

Atomic mass of O =16u

Molecular mass of HNO3 = 1 + 14 + (16*3) = 63u

• Formula Unit Mass
• It is the sum of the atomic masses of all atoms in a formula unit of a compound. The constituent particles are ions.
• For example, formula unit mass of Sodium Chloride (NaCl) can be calculated as:

(1*23) + (1*35.5) = 58.5u

• Mole Concept
• Wilhelm Ostwald Introduce the word “mole” in 1896. It is derived from Latin word moles meaning a ‘heap’ or ‘pile’
• 1 mole of any substance = 6.022 X 1023 number of particles (atoms, ions or molecules)
• This is called the Avogadro number or Avogadro Constant which is represented as N0
• The mass of 1 mole of a substance is the same as that its atomic mass or molecular mass expressed in grams.

Gram atomic mass of a substance – the atomic mass of a substance when expressed in grams is known as its gram atomic mass.

Gram molecular mass of a substance – the molecular mass of a substance when expressed in grams is known as its gram molecular mass.

• For example, the atomic mass of Sulphur is 32u. Gram atomic mass of Sulphur is 32g.
• Also, 32u of Sulphur has 1 atom of Sulphur. 32g of Sulphur has 1 mole atoms, that is, 6.022 X 1023 atoms of Sulphur.

Consider these formulae