Atoms and Molecules

Laws of Chemical Combination

Ancient Philosophers’ Ideas

• Indian Philosophers:
  • Maharishi Kanad (500 BC): Proposed that dividing matter (padarth) repeatedly leads to the smallest particles called Parmanu.
  • Pakudha Katyayama: Expanded on Kanad’s idea, stating that these particles combine to form different types of matter.
• Greek Philosophers:
  • Democritus and Leucippus: Suggested that dividing matter repeatedly leads to indivisible particles called atoms (meaning indivisible).

Laws of Chemical Combination

• Antoine L. Lavoisier and Joseph L. Proust discovered two important laws:

1. Law of Conservation of Mass

• Experiment to Demonstrate:
  • Prepare a 5% solution of two chemicals, one from set X and one from set Y.
  • Mix them in a flask, ensuring they don’t mix initially, then weigh the flask.
  • After mixing, weigh again.
  • Observation: Mass remains unchanged.
• Law: Mass is neither created nor destroyed in a chemical reaction.

2. Law of Constant Proportions

• Observation:
  • Compounds are made of elements in fixed ratios by mass.
  • Examples:
    • Water (H₂O): Always has hydrogen and oxygen in a 1:8 mass ratio.
    • Ammonia (NH₃): Always has nitrogen and hydrogen in a 14:3 mass ratio.
• Law: Elements in a chemical substance are always in definite proportions by mass.

Dalton’s Atomic Theory

• John Dalton: Provided a theory to explain the above laws.
  • Key Points:
    1. Matter is made of tiny particles called atoms.
    2. Atoms are indivisible and indestructible.
    3. Atoms of the same element are identical.
    4. Atoms of different elements are different.
    5. Atoms combine in simple whole number ratios to form compounds.
    6. The number and type of atoms in a compound are constant.

Fun Fact

• John Dalton:
  • Born in 1766 in England to a poor family.
  • Became a teacher at 12, then a school principal.
  • Taught math, physics, and chemistry in Manchester.
  • Presented his atomic theory in 1808, which changed how we understand matter.

What is an Atom?

Building Blocks of Matter

• Think of how a mason builds walls and rooms to form a building.
• Just like that, atoms are the building blocks of all matter.

Size of Atoms

• Atoms are extremely small, smaller than we can imagine.
• Millions of atoms stacked together would form a layer as thin as a sheet of paper.
• Despite their small size, atoms are crucial because everything around us is made up of atoms.
• Modern techniques can magnify images of atoms on surfaces of elements.

Atomic Radius

• Measured in nanometers (nm).
  • 1 nm = 1/10⁹ meters.

Relative Sizes

• Examples of radii in meters:
  • Hydrogen atom: 10⁻¹⁰ m
  • Water molecule: 10⁻⁹ m
  • Hemoglobin molecule: 10⁻⁸ m
  • Grain of sand: 10⁻⁴ m
  • Ant: 10⁻³ m
  • Apple: 10⁻¹ m

Modern Symbols of Atoms

Historical Context

• Dalton first used specific symbols for elements representing one atom.
• Berzilius suggested using one or two letters from the element’s name.

Naming Elements

• Early names were based on places or colors (e.g., copper from Cyprus, gold meaning yellow).
• Today, IUPAC (International Union of Pure and Applied Chemistry) approves names and symbols.

Symbol Rules

• Most symbols are the first one or two letters of the element’s English name.
  • First letter: uppercase
  • Second letter: lowercase
• Examples:
  • Hydrogen: H
  • Aluminium: Al (not AL)
  • Cobalt: Co (not CO)

Symbols from Other Languages

• Some symbols come from Latin, German, or Greek names.
  • Iron: Fe (from Latin ferrum)
  • Sodium: Na (from Latin natrium)
  • Potassium: K (from Latin kalium)

Element	Symbol	Element	Symbol
Aluminium	Al	Copper	Cu
Argon	Ar	Fluorine	F
Barium	Ba	Gold	Au
Boron	B	Hydrogen	H
Bromine	Br	Iodine	I
Calcium	Ca	Iron	Fe
Carbon	C	Lead	Pb
Chlorine	Cl	Magnesium	Mg
Cobalt	Co	Neon	Ne
Nitrogen	N	Oxygen	O
Potassium	K	Silicon	Si
Silver	Ag	Sodium	Na
Sulphur	S	Uranium	U
Zinc	Zn

These are the symbols for some common elements. You’ll learn more as you keep studying!

Atomic Mass

Dalton’s Atomic Theory:

• Dalton proposed that each element has a unique atomic mass.
• This helped explain why elements combine in constant proportions.

Relative Atomic Mass:

• Measuring the mass of a single atom is tough.
• Scientists used chemical reactions to find relative atomic masses.

Example – Carbon Monoxide (CO):

• 3 g of carbon combines with 4 g of oxygen to make CO.
• This means carbon combines with 4/3 times its mass of oxygen.

Atomic Mass Units:

• Initially, scientists used 1/16 of an oxygen atom’s mass as the unit.
• Oxygen was chosen because it forms compounds easily.
• This method gave whole numbers for most elements.

Modern Standard – Carbon-12:

• In 1961, carbon-12 was chosen as the new standard.
• 1 atomic mass unit (u) = 1/12th the mass of a carbon-12 atom.
• Relative atomic mass is compared to carbon-12.

Example – Fruit Seller:

• Imagine a fruit seller using a watermelon as a standard.
• He divides it into 12 pieces.
• Other fruits are weighed relative to these pieces.

Summary:

• Relative atomic mass = average mass of an atom compared to 1/12th of a carbon-12 atom.

Certainly, here’s the table with four columns where entries are divided equally:

Element	Atomic Mass (u)	Element	Atomic Mass (u)
Hydrogen	1	Carbon	12
Nitrogen	14	Oxygen	16
Sodium	23	Magnesium	24
Sulphur	32	Chlorine	35.5
Calcium	40

How Do Atoms Exist?

Atoms and Existence:

• Most atoms cannot exist alone.
• They form molecules and ions.
• These combine in large numbers to form matter we can see and touch.

Key Points to Remember:

• Dalton’s theory introduced atomic mass.
• Relative atomic masses are easier to measure than actual atomic masses.
• Carbon-12 is the current standard for atomic mass units.
• Atoms usually combine to form molecules or ions.

What is a Molecule?

Definition:

• A molecule is a group of two or more atoms chemically bonded together.
• It is the smallest particle of an element or compound that can exist independently and show the properties of that substance.

Molecules of Elements

Same Type of Atoms:

• Molecules of elements are made up of the same type of atoms.
• Some elements have molecules with only one atom (monoatomic), like Argon (Ar) and Helium (He).
• Most non-metals have molecules with two atoms (diatomic), like Oxygen (O2) and Nitrogen (N2).

Examples of Atomicity:

Element	Type of Element	Atomicity
Argon	non-metal	Monoatomic (Ar)
Helium	non-metal	Monoatomic (He)
Oxygen	non-metal	Diatomic (O2)
Hydrogen	non-metal	Diatomic (H2)
Nitrogen	non-metal	Diatomic (N2)
Chlorine	non-metal	Diatomic (Cl2)
Phosphorus	non-metal	Tetra-atomic (P4)
Sulphur	non-metal	Poly-atomic (S8)

Complex Structures:

• Metals and some elements like Carbon do not have simple structures but are made of many atoms bonded together.

Molecules of Compounds

Different Elements Together:

• Atoms of different elements join in specific ratios to form molecules of compounds.
• Examples include:
  • Water (H2O): Hydrogen and Oxygen in a 1:8 mass ratio.
  • Ammonia (NH3): Nitrogen and Hydrogen in a 14:3 mass ratio.
  • Carbon Dioxide (CO2): Carbon and Oxygen in a 3:8 mass ratio.

What is an Ion?

Charged Particles:

• Compounds made of metals and non-metals have ions, which are charged particles.
• Ions can be single atoms or groups of atoms with a net charge.

Types of Ions:

• Anion: Negatively charged ion.
• Cation: Positively charged ion.

Examples:

• Sodium Chloride (NaCl): Sodium ions (Na+) and Chloride ions (Cl–).
• Polyatomic Ions: Groups of atoms with a charge.

Examples of Ionic Compounds

Compound	Ionic Constituting Elements	Ratio by Mass
Calcium oxide	Calcium and oxygen	5:2
Magnesium sulphide	Magnesium and sulphur	3:4
Sodium chloride	Sodium and chlorine	23:35.5

Writing Chemical Formulae

Chemical Formula:

• A symbolic way to represent the composition of a compound.
• To write chemical formulae, know the symbols and valencies of elements.

Valency:

• Valency is the combining capacity of an element.
• Think of valency like hands: humans have 2, an octopus has 8.

Rules for Writing Chemical Formulae:

1. Valencies or charges must balance.
2. For compounds with a metal and a non-metal, write the metal first.

• Examples:
  • Calcium oxide (CaO)
  • Sodium chloride (NaCl)

1. For polyatomic ions, use brackets if there’s more than one ion.

• Example: Mg(OH)₂

Formulae of Simple Compounds

Binary Compounds:

• Made of two different elements.
• Use valencies to write formulae by criss-crossing the valencies of the combining atoms.

Examples:

1. Hydrogen chloride: HCl
2. Hydrogen sulphide: H₂S
3. Carbon tetrachloride: CCl₄
4. Magnesium chloride: MgCl₂

• Mg²⁺ and Cl⁻ criss-cross to give MgCl₂.

More Examples:

• Aluminium oxide: Al₂O₃
• Calcium oxide: CaO (simplified from Ca₂O₂)
• Sodium nitrate: NaNO₃
• Calcium hydroxide: Ca(OH)₂ (brackets show two OH groups)
• Sodium carbonate: Na₂CO₃
• Ammonium sulphate: (NH₄)₂SO₄

Key Points:

• Balance charges to write correct formulae.
• Use brackets for multiple polyatomic ions.
• Metals are written first in compound names.

Molecular Mass

Molecular Mass:

• Definition: The molecular mass of a substance is the sum of the atomic masses of all atoms in a molecule.
• Unit: Expressed in atomic mass units (u).

Examples:

1. Water (H₂O):
   • Atomic mass of hydrogen = 1 u
   • Atomic mass of oxygen = 16 u
   • Molecular mass of water = (2 × 1) + (1 × 16) = 18 u
2. Nitric Acid (HNO₃):
   • Atomic mass of hydrogen = 1 u
   • Atomic mass of nitrogen = 14 u
   • Atomic mass of oxygen = 16 u
   • Molecular mass of HNO₃ = 1 + 14 + (3 × 16) = 63 u

Formula Unit Mass:

• Definition: The formula unit mass of a compound is the sum of the atomic masses of all atoms in a formula unit of an ionic compound.
• Difference: Used for substances with ions.

Examples:

1. Sodium Chloride (NaCl):
   • Atomic mass of sodium (Na) = 23 u
   • Atomic mass of chlorine (Cl) = 35.5 u
   • Formula unit mass of NaCl = (1 × 23) + (1 × 35.5) = 58.5 u
2. Calcium Chloride (CaCl₂):
   • Atomic mass of calcium (Ca) = 40 u
   • Atomic mass of chlorine (Cl) = 35.5 u
   • Formula unit mass of CaCl₂ = 40 + (2 × 35.5) = 111 u

Key Points:

• Molecular mass is for molecules (covalent compounds).
• Formula unit mass is for ionic compounds.
• Both are calculated by adding atomic masses of all atoms in the molecule or formula unit.

Chapter Summary:

• During a chemical reaction, the sum of the masses of the reactants and products remains unchanged. This is known as the Law of Conservation of Mass.
• In a pure chemical compound, elements are always present in a definite proportion by mass. This is known as the Law of Definite Proportions.
• An atom is the smallest particle of the element that cannot usually exist independently and retain all its chemical properties.
• A molecule is the smallest particle of an element or a compound capable of independent existence under ordinary conditions. It shows all the properties of the substance.
• A chemical formula of a compound shows its constituent elements and the number of atoms of each combining element.
• Clusters of atoms that act as an ion are called polyatomic ions. They carry a fixed charge on them.
• The chemical formula of a molecular compound is determined by the valency of each element.
• In ionic compounds, the charge on each ion is used to determine the chemical formula of the compound.