Science (Double Award) | Principles of chemistry

Welcome to Principles of Chemistry!

Hello future chemist! This chapter is the absolute foundation of everything we will learn in chemistry. Think of it as learning the alphabet before writing a novel. If you understand the basic rules about atoms and how they interact, the rest of the course will be much easier!

Don't worry if some concepts seem a bit abstract at first. We will break them down step-by-step, use real-life examples, and ensure you master these chemical building blocks!

Section 1: The Structure of the Atom

1.1 The Atom: The Smallest Building Block

Everything in the universe is made of incredibly tiny particles called atoms. Atoms themselves are made up of even smaller subatomic particles.

Imagine the atom like a miniature solar system:

• The Nucleus is in the center (like the sun). This is where most of the mass is concentrated.
• Electrons orbit the nucleus (like planets).

1.2 Subatomic Particles

The three key particles are Protons, Neutrons, and Electrons. You must know their relative mass and charge:

Key Takeaway: Atoms are electrically neutral because the number of positive protons equals the number of negative electrons.

1.3 Defining Atoms: Atomic Number and Mass Number

Every element is defined by the number of protons it has.

• Atomic Number (Z): This is the number of protons in the nucleus. It is the atom's unique ID card.
  Mnemonic: P = E = Z (Protons = Electrons = Atomic Number) in a neutral atom.
• Mass Number (A): This is the total number of particles in the nucleus (protons + neutrons).

How to find Neutrons:
\( \text{Neutrons} = \text{Mass Number (A)} - \text{Atomic Number (Z)} \)

1.4 Isotopes (Same Element, Different Mass)

Isotopes are atoms of the same element (meaning they have the same number of protons and electrons), but they have a different number of neutrons.

Example: Carbon-12 and Carbon-14. Both have 6 protons. Carbon-12 has 6 neutrons, while Carbon-14 has 8 neutrons. They behave chemically in almost the same way.

Section 2: Electron Arrangement and the Periodic Table

2.1 Electron Shells

Electrons don't orbit randomly; they occupy specific energy levels or shells around the nucleus. When drawing these arrangements, we follow specific rules:

1. First Shell (closest to nucleus) holds a maximum of 2 electrons.
2. Second Shell holds a maximum of 8 electrons.
3. Third Shell holds a maximum of 8 electrons (for the first 20 elements).

The electrons in the outermost shell are called valence electrons. These are the electrons involved in bonding and determining how the element reacts.

Example: Sodium (Na) has 11 electrons. Its arrangement is 2, 8, 1. It has 1 valence electron.

2.2 The Periodic Table: A Chemist's Map

The Periodic Table arranges elements in order of increasing Atomic Number (Z). It is designed to group elements with similar chemical properties together.

Groups (The Vertical Columns)

The vertical columns are called Groups (numbered 1 to 7, plus Group 0).
The Rule: Elements in the same group have the same number of valence electrons and therefore have similar chemical properties.

• Group 1 (Alkali Metals): All have 1 valence electron. Highly reactive metals.
• Group 7 (Halogens): All have 7 valence electrons. Highly reactive non-metals.
• Group 0 (Noble Gases): All have a full outer shell (usually 8, or 2 for Helium). They are unreactive or inert.

Periods (The Horizontal Rows)

The horizontal rows are called Periods.
The Rule: Elements in the same period have the same number of electron shells.
Example: Period 3 elements all have electrons in three shells (e.g., Sodium, Magnesium, Chlorine).

Metals vs. Non-metals:
The periodic table is divided by a "zigzag line."

• Metals are found mainly on the left and center. They tend to lose electrons.
• Non-metals are found on the right (and Hydrogen). They tend to gain or share electrons.

Section 3: Chemical Bonding

Atoms bond together to achieve a stable structure—usually a full outer shell (8 electrons, known as the octet rule).

3.1 Ionic Bonding (Transferring Electrons)

Ionic bonding occurs when electrons are transferred from one atom to another, usually between a metal and a non-metal.

Step 1: Forming Ions

When an atom loses or gains electrons, it becomes a charged particle called an ion.

• Metals: Tend to lose electrons to form positive ions called cations.
  Example: Na (2, 8, 1) loses 1 electron to become Na+ (2, 8).
  Mnemonic: CaTions are always posiTive!
• Non-metals: Tend to gain electrons to form negative ions called anions.
  Example: Cl (2, 8, 7) gains 1 electron to become Cl- (2, 8, 8).

Step 2: The Electrostatic Attraction

The strong electrostatic force of attraction between the oppositely charged ions (positive cation and negative anion) holds the compound together. This forms a Giant Ionic Lattice structure.

3.2 Covalent Bonding (Sharing Electrons)

Covalent bonding occurs when atoms share electrons to achieve a full outer shell, usually between two non-metals.

When atoms share electrons, they form molecules.

• A single covalent bond means 1 pair of electrons is shared.
• A double covalent bond means 2 pairs of electrons are shared.

Example: Hydrogen atoms (H) both need 1 electron to fill their shell (2 max). They share one pair of electrons, forming an H₂ molecule.

Did you know? Ionic compounds are always solids at room temperature and usually dissolve in water, while covalent compounds can be solids, liquids, or gases and often don't dissolve in water.

Section 4: Relative Mass and the Mole Concept

Chemistry relies on accurate measurement. We need ways to compare the masses of different atoms and compounds.

4.1 Relative Atomic Mass (\(A_r\))

The Relative Atomic Mass (\(A_r\)) is the average mass of one atom of an element compared to 1/12th the mass of a Carbon-12 atom.

In IGCSE chemistry, you generally find the \(A_r\) value by looking at the larger number given for the element on the Periodic Table (this is essentially the Mass Number, A, adjusted for isotopes).

4.2 Relative Formula Mass (\(M_r\))

The Relative Formula Mass (\(M_r\)) is the sum of the Relative Atomic Masses (\(A_r\)) of all the atoms shown in the chemical formula.

Step-by-Step Calculation of \(M_r\)

Let's calculate the \(M_r\) of Water, \(H_2O\). (Assume \(A_r\) for H=1 and O=16).

1. Identify the atoms present: Hydrogen (H) and Oxygen (O).
2. Count how many of each atom: 2 Hydrogen atoms, 1 Oxygen atom.
3. Multiply the number of atoms by their \(A_r\):
   • H: \( 2 \times 1 = 2 \)
   • O: \( 1 \times 16 = 16 \)
4. Add the totals together: \( M_r = 2 + 16 = 18 \)

The Relative Formula Mass (\(M_r\)) of water is 18. This value has no units.

4.3 The Mole Concept (Mass and Moles)

Since atoms are too small to count individually, we use a unit called the Mole.

A mole is simply a specific, very large number of particles (known as Avogadro's constant).

The mass of one mole of a substance (in grams) is numerically equal to its Relative Formula Mass (\(M_r\)). This is called the Molar Mass.

Example: If the \(M_r\) of water is 18, then 1 mole of water weighs 18 grams.

We can use the following formula to convert between mass and moles:

$$ \text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass } (M_r)} $$

Summary and Key Takeaways

You have now mastered the chemical foundations! Remember these core principles:

• The number of protons defines the element.
• Atoms seek stability by achieving a full outer shell (usually 8 electrons).
• Ionic bonding involves the transfer of electrons (metal + non-metal).
• Covalent bonding involves the sharing of electrons (non-metal + non-metal).
• The Relative Formula Mass (\(M_r\)) allows us to calculate the mass of substances accurately using the mole concept.

Keep practicing those \(M_r\) calculations and drawing electron shells. You've got this!