Principles of Chemistry: The Foundations
Welcome to the most fundamental chapter in Chemistry! Don't worry if this seems like a lot of definitions at first. Chemistry is like building with LEGOs; you first need to understand the individual bricks (atoms) before you can build a masterpiece (reactions).
This chapter is crucial because it teaches you the language of Chemistry—how matter is structured, why elements behave the way they do, and how they stick together. Let's get started!
1. The Structure of the Atom
Everything in the universe is made up of tiny particles called atoms. While you can't see them, understanding their structure is key. Think of the atom as a mini solar system.
1.1 Subatomic Particles
An atom is made of three main particles, called subatomic particles:
- Protons (p): Found in the centre (the nucleus). They carry a positive (+) electrical charge.
- Neutrons (n): Also found in the centre (the nucleus). They carry no charge (they are neutral).
- Electrons (e): Orbit the nucleus in shells. They carry a negative (-) electrical charge.
Key Fact: In a neutral (uncharged) atom, the number of Protons equals the number of Electrons. This keeps the atom electrically balanced.
1.2 Atomic Number and Mass Number
We use two main numbers to identify and describe an atom:
-
Atomic Number (Proton Number, Z):
This is the number of protons in the nucleus.
Why is it important? The Atomic Number is like the element’s fingerprint—it determines which element the atom is. Every atom of Carbon has 6 protons, no exceptions! -
Mass Number (Nucleon Number, A):
This is the total number of particles in the nucleus: Protons + Neutrons.
Quick Tip to Find Neutrons:
Neutron Number = Mass Number – Atomic Number
Did you know?
Elements that have the same number of protons but different numbers of neutrons are called isotopes. They are the same element, just slightly heavier or lighter versions (e.g., Carbon-12 and Carbon-14).
Quick Review: Atomic Structure
1. Protons give the element its identity (Atomic Number).
2. Electrons determine the charge and bonding behaviour.
3. Mass Number is the weight found in the nucleus (P + N).
2. The Periodic Table (The Chemist’s Map)
The Periodic Table organises all the known elements in a useful way. If you know how to read it, it tells you everything about an element’s behaviour!
2.1 Arrangement and Key Terms
- Periods (Rows): These run horizontally (left to right). The period number tells you how many electron shells the atoms of that element have.
-
Groups (Columns): These run vertically (up and down). The group number (usually 1 to 0/8) tells you how many electrons are in the outer shell (valence electrons).
Analogy: All the people in the same apartment building (Period) have the same number of floors, but all the people living in the same line (Group) have the same personality (chemical behaviour).
2.2 Important Groups to Remember
Elements in the same group have similar chemical properties because they have the same number of outer shell electrons.
Group 1 (Alkali Metals):
These are highly reactive metals. They have 1 outer electron, which they are desperate to lose to become stable.
Group 7 (Halogens):
These are highly reactive non-metals. They have 7 outer electrons, meaning they need just 1 more electron to fill their shell.
Group 0 / 8 (Noble Gases):
These are unreactive (inert). They have a full outer shell, meaning they are chemically stable and rarely form bonds.
Key Takeaway: Atoms want a full outer shell (usually 8 electrons, except for the first shell, which holds 2). This is the driving force behind all chemical bonding!
3. Chemical Bonding: Why Atoms Join Together
Atoms bond to achieve stability—they want that full outer shell just like the Noble Gases. There are two main ways they do this: by transferring electrons (Ionic) or by sharing electrons (Covalent).
3.1 Ionic Bonding (Giving and Taking)
Ionic bonding occurs between a Metal and a Non-Metal.
Step-by-Step Process:
- The Metal (Group 1, 2, 3) has few outer electrons and loses them entirely. Losing negative electrons makes the metal atom positively charged. This positive particle is called a positive ion (or cation).
- The Non-Metal (Group 6, 7) has almost a full shell and gains these electrons. Gaining negative electrons makes the non-metal atom negatively charged. This negative particle is called a negative ion (or anion).
- The strong electrical force of attraction between the oppositely charged ions (positive and negative) is the ionic bond.
Example: Sodium Chloride (NaCl)
Sodium (Na, Group 1) loses 1 electron to become \(\text{Na}^+\).
Chlorine (Cl, Group 7) gains 1 electron to become \(\text{Cl}^-\).
The strong attraction between \(\text{Na}^+\) and \(\text{Cl}^-\) forms the ionic compound.
Common Mistake: Students often forget that ionic bonding requires the complete transfer of electrons, resulting in charged ions.
3.2 Covalent Bonding (Sharing is Caring)
Covalent bonding occurs between Non-Metals only.
Neither atom wants to lose electrons, so they decide to share one or more pairs of electrons to achieve a temporary full outer shell. This shared electron pair is the covalent bond.
- Covalently bonded atoms form molecules (like \(\text{H}_2\text{O}\), \(\text{O}_2\), \(\text{CH}_4\)).
- A single covalent bond involves sharing one pair of electrons (2 electrons).
- A double bond involves sharing two pairs (4 electrons), and a triple bond involves three pairs (6 electrons).
Example: Hydrogen Gas (\(\text{H}_2\))
Each Hydrogen atom has 1 electron. They share their single electrons so that both atoms feel like they have 2 electrons (a full first shell).
Key Takeaway: Ionic bonding makes strong structures (often solids with high melting points). Covalent bonding makes individual molecules (often gases or liquids with low melting points).
4. Basic Calculations: Relative Formula Mass (\(M_r\))
To understand how much substance we need for a reaction, we need to know the mass of its atoms. This is where Relative Atomic Mass (\(A_r\)) and Relative Formula Mass (\(M_r\)) come in.
4.1 Relative Atomic Mass (\(A_r\))
The Relative Atomic Mass (\(A_r\)) of an element is the number found below the element symbol on the Periodic Table (usually the larger number). It is the mass of one atom of that element compared to a standard (Carbon-12).
- For Hydrogen (H), \(A_r\) is approximately 1.
- For Carbon (C), \(A_r\) is approximately 12.
- For Oxygen (O), \(A_r\) is approximately 16.
4.2 Relative Formula Mass (\(M_r\))
The Relative Formula Mass (\(M_r\)) is the sum of the relative atomic masses of all the atoms shown in the chemical formula.
The Calculation Rule:
\( M_r = \text{Sum of all } A_r \text{ values in the formula} \)
Step-by-Step Example: Calculate the \(M_r\) of Water (\(\text{H}_2\text{O}\))
- Identify the atoms and their quantities:
- Hydrogen (H): 2 atoms. \(A_r = 1\)
- Oxygen (O): 1 atom. \(A_r = 16\)
- Calculate the total mass for each element:
- Total H mass: \(2 \times 1 = 2\)
- Total O mass: \(1 \times 16 = 16\)
- Add the masses together for the final \(M_r\):
\( M_r \text{ of } \text{H}_2\text{O} = 2 + 16 = 18 \)
Example 2: \(\text{CO}_2\)
C (1 atom, \(A_r=12\)) + O (2 atoms, \(A_r=16\))
\(M_r = 12 + (2 \times 16) = 12 + 32 = 44\)
Remember: Don't confuse the \(A_r\) (for single atoms) with \(M_r\) (for compounds). The \(M_r\) calculation is simple addition, so practice makes perfect!
5. States of Matter and the Kinetic Theory
Matter exists in three main states: Solid, Liquid, and Gas. The behaviour of these states is explained by the Kinetic Theory, which states that all particles are constantly moving.
5.1 Characteristics of the Three States
| State | Particle Arrangement | Particle Movement | Volume/Shape |
|---|---|---|---|
| Solid | Regular, fixed pattern (lattice) | Vibrate only in fixed positions | Fixed volume and fixed shape |
| Liquid | Random, close together | Move around and slide past each other | Fixed volume, takes shape of container |
| Gas | Random, far apart | Move quickly and randomly in all directions | No fixed volume or shape; fills container |
Why? The difference is the strength of the forces between the particles. These forces are strongest in solids, moderate in liquids, and almost non-existent in gases.
5.2 Changes of State (Heating and Cooling)
When you add energy (heat) to a substance, the particles gain kinetic energy and move faster, leading to a change of state.
- Solid \(\to\) Liquid: Melting (at the melting point)
- Liquid \(\to\) Gas: Boiling/Evaporating (at the boiling point)
- Gas \(\to\) Liquid: Condensing
- Liquid \(\to\) Solid: Freezing
- Solid \(\to\) Gas (skips liquid): Subliming
Did you know?
When a substance is melting or boiling, the temperature remains constant even though you are still adding heat. This energy is being used to break the bonds between particles, not to increase their speed (kinetic energy).
Key Takeaway: Adding heat increases movement and breaks forces; removing heat decreases movement and allows forces to re-form.
You have now built a strong foundation! Review these principles regularly—they are the starting point for every chemistry concept you will encounter.