Laws Of Chemical Combination

Antoine L. Lavoisier, considered the father of modern chemistry, along with Joseph L. Proust, established two fundamental laws based on experimental observations of chemical reactions. These laws govern how elements combine to form compounds.

---

Law Of Conservation Of Mass

This law addresses whether the total mass changes during a chemical reaction. Through careful experiments, such as reacting solutions of substances like copper sulphate and sodium carbonate in a sealed flask, it was observed that the total mass of the reactants before the reaction is equal to the total mass of the products after the reaction. The cork on the flask is crucial to ensure that no substances or gases involved in the reaction escape or enter the system, maintaining a closed environment for accurate mass measurement.

The **Law of Conservation of Mass** states that **mass can neither be created nor destroyed in a chemical reaction**. In other words, the total mass of the substances undergoing a chemical transformation remains constant.

---

Law Of Constant Proportions

Scientists like Lavoisier and Proust studied the composition of many compounds and consistently found that a pure chemical compound always contains the same elements combined in the same proportion by mass, regardless of its source or method of preparation. This led to the formulation of the **Law of Constant Proportions**, also known as the **Law of Definite Proportions**.

The law states: **In a chemical substance, the elements are always present in definite proportions by mass**.

For example, water ($\text{H}_2\text{O}$) is always formed by the chemical combination of hydrogen and oxygen in a mass ratio of 1:8. If 9 grams of water are decomposed, 1 gram of hydrogen and 8 grams of oxygen are always obtained, regardless of where the water came from. Similarly, ammonia ($\text{NH}_3$) always contains nitrogen and hydrogen in a fixed mass ratio of 14:3.

---

Dalton's Atomic Theory

The scientific community needed an explanation for these observed laws of chemical combination. John Dalton, a British chemist, provided the first comprehensive theory about the nature of matter, building upon the ancient philosophical idea of atoms. Dalton's **atomic theory**, proposed in 1808, was based on the laws of chemical combination and successfully explained them.

According to Dalton's atomic theory, all matter is made up of very small particles called atoms. The main postulates of his theory are:

1. All matter is composed of extremely tiny particles called **atoms**, which are the participants in chemical reactions.
2. Atoms are **indivisible particles** and cannot be created or destroyed during a chemical reaction (explains the Law of Conservation of Mass).
3. Atoms of a given element are **identical** in 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 (helps explain the Law of Constant Proportions).
6. The **relative number and kinds of atoms** are constant in a given compound (directly supports the Law of Constant Proportions).

Although later discoveries showed that atoms are indeed divisible and made of subatomic particles (electrons, protons, neutrons), Dalton's theory provided a fundamental framework for understanding chemical reactions and the composition of matter.

Answer:

Answer:

What Is An Atom?

Just as a building is constructed from bricks, or an ant-hill from grains of sand, all matter is ultimately built from fundamental particles called atoms. Atoms are the basic building blocks of all matter, whether it's an element, a compound, or a mixture.

---

How Big Are Atoms?

Atoms are incredibly tiny. Their size is so small that they are not visible to the naked eye. We cannot even see them with powerful light microscopes. Even stacking millions of atoms would only form a layer as thin as a sheet of paper.

The size of an atom is typically measured by its atomic radius, often expressed in nanometres (nm).

• 1 nanometre (nm) = $10^{-9}$ metre (m)
• 1 metre (m) = $10^9$ nanometres (nm)

To appreciate the scale, consider these relative sizes:

Size (in m)	Example
$10^{-10}$	Atom of hydrogen
$10^{-9}$	Molecule of water
$10^{-8}$	Molecule of haemoglobin
$10^{-4}$	Grain of sand
$10^{-3}$	Ant
$10^{-1}$	Apple

Despite their minuscule size, atoms are the fundamental units that make up everything around us and continuously influence our world. Modern scientific instruments, like scanning tunnelling microscopes, allow us to produce magnified images of surfaces, showing the arrangement of individual atoms.

---

What Are The Modern Day Symbols Of Atoms Of Different Elements?

To represent elements concisely, scientists use symbols. John Dalton was one of the first to use specific symbols for elements, with each symbol representing one atom of that element.

As more elements were discovered, a more standardized system was needed. J.J. Berzelius suggested using one or two letters from the element's name as its symbol. The modern system, approved by the **International Union of Pure and Applied Chemistry (IUPAC)**, uses this approach.

Rules for writing element symbols:

• Symbols are derived from the English name, Latin name, German name, or Greek name of the element.
• If the symbol is a single letter, it is always written in **uppercase** (e.g., H for Hydrogen, C for Carbon, N for Nitrogen).
• If the symbol consists of two letters, the **first letter is uppercase** and the **second letter is lowercase** (e.g., Al for Aluminium, Co for Cobalt, Zn for Zinc, Cl for Chlorine). Writing 'AL' or 'CO' is incorrect as it could be misinterpreted as two different atoms.
• Some symbols are based on older names, often Latin. For example, the symbol for iron is Fe (from ferrum), sodium is Na (from natrium), and potassium is K (from kalium).

Each element has a unique name and a unique chemical symbol.

Here are symbols for some common elements:

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

(Note: You don't need to memorise all of these immediately. Familiarity will grow with practice.)

---

Atomic Mass

Dalton's atomic theory highlighted that each element has a characteristic atomic mass. Since individual atoms are extremely light, determining their absolute mass is very difficult. Scientists therefore decided to determine the relative atomic masses of elements by comparing them to the mass of a standard reference atom.

Initially, 1/16th the mass of a naturally occurring oxygen atom was used as the reference unit because oxygen forms compounds with many elements and this unit often gave atomic masses close to whole numbers.

However, in 1961, the **carbon-12 isotope** was chosen as the universally accepted standard reference. One **atomic mass unit (u)**, also known as a unified mass unit, is defined as exactly one-twelfth ($1/12$th) the mass of one atom of carbon-12. The relative atomic mass of an atom of an element is the average mass of the atom compared to 1/12th the mass of a carbon-12 atom.

This is analogous to a fruit seller using a watermelon divided into 12 equal pieces as a standard unit ("fruit mass unit" or fmu) to weigh other fruits relatively.

Atomic masses of elements are typically listed in atomic mass units (u). For example, the atomic mass of Hydrogen is approximately 1 u, Carbon is 12 u, and Oxygen is 16 u.

Atomic masses of a few common elements:

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 are highly reactive and for most elements, they cannot exist independently as single atoms under normal conditions. Instead, atoms combine with each other to form larger entities called molecules and ions. These molecules or ions then aggregate in vast numbers to form the matter that we can observe and interact with (solids, liquids, gases).

Answer:

Answer:

What Is A Molecule?

A molecule is generally defined as a group of two or more atoms that are held together by chemical bonds (attractive forces). A molecule is the smallest particle of an element or a compound that is capable of independent existence under ordinary conditions and exhibits all the properties of that substance.

Molecules can be formed by atoms of the same element or by atoms of different elements combining together.

---

Molecules Of Elements

The molecules of an element are made up of only one type of atom. The number of atoms that constitute a molecule of an element is called its **atomicity**.

• Many elements, particularly noble gases like Argon (Ar) and Helium (He), exist as single, independent atoms. Their molecules consist of just one atom, so their atomicity is 1. They are called **monoatomic** molecules.
• Most non-metals exist as molecules containing more than one atom. For instance, a molecule of oxygen consists of two oxygen atoms chemically bonded ($\text{O}_2$), so its atomicity is 2 (diatomic). If three oxygen atoms bond, they form an ozone molecule ($\text{O}_3$), which has an atomicity of 3 (triatomic).
• Phosphorus molecules contain four atoms ($\text{P}_4$), making them **tetra-atomic**.
• Sulphur molecules commonly exist as $\text{S}_8$ rings, so sulphur is considered **poly-atomic** (containing many atoms).

Elements that are metals (like iron, copper, gold) and some non-metals like carbon do not exist as simple, discrete molecules in the solid state. Instead, they form large structures where a vast and indefinite number of atoms are bonded together.

Atomicity of some elements:

Type of Element	Name	Atomicity
Non-Metal	Argon	Monoatomic (1)
Non-Metal	Helium	Monoatomic (1)
Non-Metal	Oxygen	Diatomic (2)
Non-Metal	Hydrogen	Diatomic (2)
Non-Metal	Nitrogen	Diatomic (2)
Non-Metal	Chlorine	Diatomic (2)
Non-Metal	Phosphorus	Tetra-atomic (4)
Non-Metal	Sulphur	Poly-atomic (>4, commonly 8)

---

Molecules Of Compounds

Molecules of compounds are formed when atoms of **different elements** combine chemically in **fixed proportions**. These molecules represent the smallest unit of the compound that retains its chemical properties.

Examples of molecules of compounds:

Compound	Combining Elements	Ratio by Mass
Water ($\text{H}_2\text{O}$)	Hydrogen, Oxygen	1:8
Ammonia ($\text{NH}_3$)	Nitrogen, Hydrogen	14:3
Carbon dioxide ($\text{CO}_2$)	Carbon, Oxygen	3:8

We can determine the ratio by the *number* of atoms of each element in a compound molecule using the mass ratio and the atomic masses of the elements (Activity 3.2 from the text). For water ($\text{H}_2\text{O}$):

Given mass ratio H:O = 1:8

Atomic mass of H = 1 u, Atomic mass of O = 16 u

Ratio by number of atoms = $\frac{\text{Mass ratio of H}}{\text{Atomic mass of H}} : \frac{\text{Mass ratio of O}}{\text{Atomic mass of O}}$

Ratio by number of atoms = $\frac{1}{1} : \frac{8}{16}$

Ratio by number of atoms = $1 : \frac{1}{2}$

To get a whole number ratio, multiply by 2:

Ratio by number of atoms H:O = $1 \times 2 : \frac{1}{2} \times 2 = 2 : 1$

This confirms the formula $\text{H}_2\text{O}$, indicating two hydrogen atoms and one oxygen atom per molecule.

---

What Is An Ion?

Some compounds, particularly those formed between metals and non-metals, consist of **charged species** called ions. Unlike molecules which are neutral, ions are atoms or groups of atoms that have gained or lost electrons, resulting in a net electrical charge.

• A positively charged ion is called a **cation** (formed when an atom loses one or more electrons). Examples: Sodium ion ($\text{Na}^+$), Magnesium ion ($\text{Mg}^{2+}$).
• A negatively charged ion is called an **anion** (formed when an atom gains one or more electrons). Examples: Chloride ion ($\text{Cl}^-$), Oxide ion ($\text{O}^{2-}$).

Ions can be formed from a single atom (like $\text{Na}^+$ or $\text{Cl}^-$) or from a group of atoms bonded together that collectively carry a charge. A group of atoms carrying a net charge is called a **polyatomic ion**.

Examples of polyatomic ions: Ammonium ion ($\text{NH}_4^+$), Hydroxide ion ($\text{OH}^-$), Carbonate ion ($\text{CO}_3^{2-}$), Sulphate ion ($\text{SO}_4^{2-}$), Phosphate ion ($\text{PO}_4^{3-}$).

Ionic compounds, like sodium chloride (NaCl), are made up of a collection of positively charged ions ($\text{Na}^+$) and negatively charged ions ($\text{Cl}^-$) held together by strong electrostatic forces. The total positive charge balances the total negative charge, making the overall compound electrically neutral.

Names and symbols of some common ions:

Valency	Name of ion	Symbol	Non-metallic element	Symbol	Polyatomic ions	Symbol
1	Sodium	Na⁺	Hydrogen	H⁺ (Hydride H⁻)	Ammonium	NH₄⁺
	Potassium	K⁺	Chloride	Cl⁻	Hydroxide	OH⁻
	Silver	Ag⁺	Bromide	Br⁻	Nitrate	NO₃⁻
	Copper (I)*	Cu⁺	Iodide	I⁻	Hydrogen carbonate	HCO₃⁻
2	Magnesium	Mg²⁺	Oxide	O²⁻	Carbonate	CO₃²⁻
	Calcium	Ca²⁺	Sulphide	S²⁻	Sulphite	SO₃²⁻
	Zinc	Zn²⁺			Sulphate	SO₄²⁻
	Iron (II)*	Fe²⁺
	Copper (II)*	Cu²⁺
3	Aluminium	Al³⁺	Nitride	N³⁻	Phosphate	PO₄³⁻
	Iron (III)*	Fe³⁺

*Some elements like copper and iron can show more than one valency. The Roman numeral in parentheses indicates their specific valency in that context.

Writing Chemical Formulae

The **chemical formula** of a compound is a symbolic representation that shows the types of elements present and the relative number of atoms of each element in one molecule or formula unit of the compound. Writing correct chemical formulae requires knowledge of the symbols of elements and their combining capacity.

The **valency** of an element is its combining power or capacity – it indicates how many atoms of another element it can combine with. Valency can be thought of as the 'arms' of an atom that connect to the 'arms' of other atoms. For example, if an octopus has 8 arms (valency 8) and humans have 2 arms (valency 2), one octopus could theoretically hold 4 humans if all arms are engaged, leading to a combination with a formula like $\text{OH}_4$ (where O represents octopus and H represents human). The subscript number shows the quantity of the second entity.

---

Formulae Of Simple Compounds

Compounds made up of two different elements are called **binary compounds**. We can write their formulae using the concept of valency and a method often called the 'criss-cross' method.

General steps for writing chemical formulae:

1. Write the symbols of the constituent elements. For ionic compounds (metal + non-metal), write the metal symbol (cation) first, followed by the non-metal symbol (anion).
2. Write the valency (or charge) of each element/ion above or below its symbol.
3. Cross-over the valencies (ignore the sign for ions, just use the magnitude). This means the valency of the first element becomes the subscript for the second, and the valency of the second becomes the subscript for the first.
4. Simplify the subscripts to the smallest whole-number ratio if possible.

Examples of writing formulae:

1. **Formula of Hydrogen Chloride**

Elements: Hydrogen ($\text{H}$), Chlorine ($\text{Cl}$)

Valency: H = 1, Cl = 1

Criss-cross: $\text{H}_1\text{Cl}_1$

Simplified: HCl

Formula: HCl

2. **Formula of Hydrogen Sulphide**

Elements: Hydrogen ($\text{H}$), Sulphur ($\text{S}$)

Valency: H = 1, S = 2

Criss-cross: $\text{H}_2\text{S}_1$

Simplified: $\text{H}_2\text{S}$

Formula: $\text{H}_2\text{S}$

3. **Formula of Carbon Tetrachloride**

Elements: Carbon ($\text{C}$), Chlorine ($\text{Cl}$)

Valency: C = 4, Cl = 1

Criss-cross: $\text{C}_1\text{Cl}_4$

Simplified: $\text{CCl}_4$

Formula: $\text{CCl}_4$

4. **Formula of Magnesium Chloride** (Ionic Compound)

Ions: Magnesium ion ($\text{Mg}^{2+}$), Chloride ion ($\text{Cl}^-$)

Charge magnitude: Mg = 2, Cl = 1

Criss-cross: $\text{Mg}_1\text{Cl}_2$

Simplified: $\text{MgCl}_2$

Formula: $\text{MgCl}_2$

This means one magnesium ion combines with two chloride ions to form a neutral compound. The charges must balance: $(+2) + 2 \times (-1) = 0$.

When dealing with **polyatomic ions**, special rules apply:

• If the compound contains a polyatomic ion, and more than one unit of that ion is needed to balance the valencies/charges, the formula of the polyatomic ion is enclosed in **brackets** before writing the subscript number.
• If only one unit of the polyatomic ion is needed, brackets are not used.

Examples with polyatomic ions:

(a) **Formula for Aluminium Oxide** (Ionic)

Ions: Aluminium ion ($\text{Al}^{3+}$), Oxide ion ($\text{O}^{2-}$)

Charge magnitude: Al = 3, O = 2

Criss-cross: $\text{Al}_2\text{O}_3$

Formula: $\text{Al}_2\text{O}_3$

Charges balance: $2 \times (+3) + 3 \times (-2) = +6 - 6 = 0$.

(b) **Formula for Calcium Oxide** (Ionic)

Ions: Calcium ion ($\text{Ca}^{2+}$), Oxide ion ($\text{O}^{2-}$)

Charge magnitude: Ca = 2, O = 2

Criss-cross: $\text{Ca}_2\text{O}_2$

Simplified (divide subscripts by 2): $\text{Ca}_1\text{O}_1$

Formula: CaO

(c) **Formula of Sodium Nitrate** (Ionic, with polyatomic ion)

Ions: Sodium ion ($\text{Na}^{+}$), Nitrate ion ($\text{NO}_3^{-}$)

Charge magnitude: Na = 1, $\text{NO}_3$ = 1

Criss-cross: $\text{Na}_1(\text{NO}_3)_1$

Simplified (no brackets needed as subscript is 1): $\text{NaNO}_3$

Formula: $\text{NaNO}_3$

(d) **Formula of Calcium Hydroxide** (Ionic, with polyatomic ion)

Ions: Calcium ion ($\text{Ca}^{2+}$), Hydroxide ion ($\text{OH}^{-}$)

Charge magnitude: Ca = 2, OH = 1

Criss-cross: $\text{Ca}_1(\text{OH})_2$

Simplified (brackets needed as subscript is 2): $\text{Ca(OH)}_2$

Formula: $\text{Ca(OH)}_2$. This indicates one calcium ion for every two hydroxide ions. The bracket around OH shows that the subscript 2 applies to both oxygen and hydrogen atoms within the hydroxide group.

(e) **Formula of Sodium Carbonate** (Ionic, with polyatomic ion)

Ions: Sodium ion ($\text{Na}^{+}$), Carbonate ion ($\text{CO}_3^{2-}$)

Charge magnitude: Na = 1, $\text{CO}_3$ = 2

Criss-cross: $\text{Na}_2(\text{CO}_3)_1$

Simplified (no brackets needed as subscript is 1): $\text{Na}_2\text{CO}_3$

Formula: $\text{Na}_2\text{CO}_3$

(f) **Formula of Ammonium Sulphate** (Ionic, with polyatomic ions)

Ions: Ammonium ion ($\text{NH}_4^{+}$), Sulphate ion ($\text{SO}_4^{2-}$)

Charge magnitude: $\text{NH}_4$ = 1, $\text{SO}_4$ = 2

Criss-cross: $(\text{NH}_4)_2(\text{SO}_4)_1$

Simplified (brackets needed for $\text{NH}_4$ as subscript is 2; no brackets for $\text{SO}_4$ as subscript is 1): $\text{(NH}_4\text{)}_2\text{SO}_4$

Formula: $\text{(NH}_4\text{)}_2\text{SO}_4$

Molecular Mass

The concept of atomic mass can be extended to determine the mass of molecules and ionic compounds.

---

Molecular Mass

The **molecular mass** of a substance is the sum of the atomic masses of all the atoms present in one molecule of that substance. It represents the relative mass of a molecule compared to the carbon-12 standard, and is expressed in atomic mass units (u). To calculate the molecular mass, multiply the atomic mass of each element by the number of atoms of that element in the molecule and then add up these values for all elements in the molecule.

Answer:

---

Formula Unit Mass

For ionic compounds, which exist as a collection of ions rather than discrete molecules, the term **formula unit mass** is used instead of molecular mass. The formula unit mass is the sum of the atomic masses of all the atoms (which form the ions) in one **formula unit** of the compound. The calculation is done in the same way as molecular mass calculation, but the term 'formula unit mass' is specifically applied to ionic compounds.

For example, the formula unit of sodium chloride is NaCl. Its formula unit mass is calculated by adding the atomic mass of sodium and the atomic mass of chlorine.

Atomic mass of Na = 23 u

Atomic mass of Cl = 35.5 u

Formula unit mass of NaCl = (1 $\times$ Atomic mass of Na) + (1 $\times$ Atomic mass of Cl)

Formula unit mass of NaCl = $(1 \times 23 \text{ u}) + (1 \times 35.5 \text{ u}) = 23 \text{ u} + 35.5 \text{ u} = 58.5 \text{ u}$

Answer:

Page No. 27 - 28

Answer:

Answer:

Answer:

Answer:

---

Page No. 30

Answer:

Answer:

---

Page No. 34

Answer:

Answer:

Answer:

Answer:

---

Page No. 35

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer: