Unit 2: Chemistry — Atomic Structure and the Periodic Table
Welcome to one of the most exciting topics in chemistry! Everything around you — the screen you are reading on, the air you are breathing, and even your own body — is made of tiny building blocks called atoms. Think of atoms like nature's ultimate Lego bricks.
Don't worry if science has felt tricky in the past. In this chapter, we will break everything down into bite-sized, easy-to-understand steps. By the end of this guide, you will be able to look at the Periodic Table and decode the secrets of any element!
1. Inside the Atom: Subatomic Particles
For a long time, scientists thought atoms were solid spheres that could not be split. Today, we know that an atom is made up of three even smaller parts, called subatomic particles:
• Protons: Heavy particles found in the centre of the atom.
• Neutrons: Heavy particles also found in the centre of the atom.
• Electrons: Tiny, almost weightless particles that whizz around the outside in shells (or energy levels).
The centre of the atom is called the nucleus. It contains all the protons and neutrons packed tightly together. Almost all the mass of the atom is squashed into this tiny nucleus.
Properties of Subatomic Particles
You need to know the relative charge, relative mass, and location of each particle:
• Proton: Relative Mass = \(1\), Relative Charge = \(+1\) (Positive), Location = Nucleus
• Neutron: Relative Mass = \(1\), Relative Charge = \(0\) (Neutral / No charge), Location = Nucleus
• Electron: Relative Mass = \(\frac{1}{1840}\) (or negligible / nearly \(0\)), Relative Charge = \(-1\) (Negative), Location = Shells / Orbits
Memory Aid Tricks:
• Proton = Positive (\(+1\))
• Neutron = Neutral (\(0\))
• Electron = nEgative (\(-1\))
Why Are Atoms Neutral Overall?
An atom has no overall electrical charge. Why? Because the number of positive protons is always equal to the number of negative electrons. The positive and negative charges cancel each other out completely!
Example: If an atom has \(6\) protons (\(+6\)) and \(6\) electrons (\(-6\)), the total charge is \(+6 + (-6) = 0\).
Key Takeaway: The nucleus contains protons (\(+1\)) and neutrons (\(0\)). Electrons (\(-1\)) orbit in shells. In any neutral atom, \(\text{number of protons} = \text{number of electrons}\).
2. Atomic Number and Mass Number
When you look at an element on the Periodic Table, you will see two numbers beside its chemical symbol. Let's look at Sodium (\(\text{Na}\)) as an example:
\(^{23}_{11}\text{Na}\)
1. The Atomic Number (The Smaller Number)
• The atomic number is the number of protons in the nucleus of an atom.
• Since atoms are neutral, it also tells you the number of electrons.
• Every element has its own unique atomic number. For example, only sodium has an atomic number of \(11\).
2. The Mass Number (The Larger Number)
• The mass number is the total number of protons + neutrons in the nucleus.
• Electrons are so light that they do not count towards the mass number.
How to Calculate the Number of Each Particle
You can find the number of all three subatomic particles using simple subtraction:
1. \(\text{Number of Protons} = \text{Atomic Number}\)
2. \(\text{Number of Electrons} = \text{Atomic Number}\)
3. \(\text{Number of Neutrons} = \text{Mass Number} - \text{Atomic Number}\)
Worked Example: Sodium (\(^{23}_{11}\text{Na}\))
• Protons: \(11\) (the smaller number)
• Electrons: \(11\) (same as protons)
• Neutrons: \(23 - 11 = 12\) (mass number minus atomic number)
Common Mistake to Avoid:
Never just guess which number is which! The mass number is always the larger number (except for standard hydrogen, where both are \(1\)).
Key Takeaway: Protons = Atomic Number, Electrons = Atomic Number, Neutrons = Mass Number \(-\) Atomic Number.
3. Electronic Configuration (Arranging Electrons)
Electrons do not just fly around randomly; they sit in fixed paths called shells. Filling electron shells follows strict rules:
• 1st Shell (closest to nucleus): holds a maximum of \(2\) electrons.
• 2nd Shell: holds a maximum of \(8\) electrons.
• 3rd Shell: holds a maximum of \(8\) electrons.
• 4th Shell: for GCSE Single Award, you only need to know up to Calcium (atomic number \(20\)), which puts up to \(2\) electrons in the 4th shell.
Step-by-Step: Writing Electronic Configurations
Let's find the electron structure for a few elements:
Example 1: Carbon (Atomic Number = \(6\))
• We have \(6\) electrons to place.
• 1st shell takes \(2\). That leaves \(6 - 2 = 4\) electrons.
• 2nd shell takes the remaining \(4\).
• Electronic configuration = \(2, 4\)
Example 2: Magnesium (Atomic Number = \(12\))
• We have \(12\) electrons to place.
• 1st shell takes \(2\) (leaving \(10\)).
• 2nd shell takes \(8\) (leaving \(2\)).
• 3rd shell takes the remaining \(2\).
• Electronic configuration = \(2, 8, 2\)
Example 3: Chlorine (Atomic Number = \(17\))
• We have \(17\) electrons to place.
• 1st shell = \(2\), 2nd shell = \(8\), 3rd shell = \(7\).
• Electronic configuration = \(2, 8, 7\)
Key Takeaway: Fill the shells from the inside out: \(2, 8, 8, 2\). The outer electrons determine how the element reacts chemically.
4. Understanding the Periodic Table
The modern Periodic Table arranges all known elements in order of increasing atomic number.
Groups and Periods
• Groups (Vertical Columns \(\downarrow\)): The group number tells you how many electrons are in the outer shell of the atom. Elements in the same group share similar chemical properties because they have the same number of outer-shell electrons.
• Periods (Horizontal Rows \(\rightarrow\)): The period number tells you the number of occupied electron shells in the atom.
Memory Tip:
• Periods go sideways across the page, just like a Pentence ends with a full stop (period)!
• Groups go up and down, like tall Giraffes!
Metals vs Non-Metals
The Periodic Table is split into two main sections by a "stepped" line (staircase) on the right side:
• Metals: Found on the left and in the centre. They are shiny, conduct electricity, conduct heat, and are malleable (can be hammered into shape).
• Non-Metals: Found on the far right. They are dull, poor conductors (insulators), and brittle when solid.
Key Takeaway: Group number = outer electrons. Period number = number of electron shells. Metals are on the left, non-metals on the right.
5. Group 1: The Alkali Metals
Group 1 elements include Lithium (\(\text{Li}\)), Sodium (\(\text{Na}\)), and Potassium (\(\text{K}\)).
Physical Properties of Group 1 Metals:
• They are very soft (they can be easily cut with a butter knife).
• They have a shiny surface when freshly cut, but quickly go dull (tarnish) in air as they react with oxygen.
• They have low densities (Lithium, Sodium, and Potassium all float on water!).
• They have low melting points compared to other metals.
• They are stored under oil to prevent them reacting with oxygen and moisture in the air.
Reactions with Water
All Group 1 metals react vigorously with water to form an alkaline solution (a metal hydroxide) and hydrogen gas:
\(\text{Alkali Metal} + \text{Water} \rightarrow \text{Metal Hydroxide} + \text{Hydrogen}\)
Example:
\(\text{Sodium} + \text{Water} \rightarrow \text{Sodium Hydroxide} + \text{Hydrogen}\)
\(2\text{Na} + 2\text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{H}_2\)
Observations During the Reaction with Water:
• General observations for all: The metal floats, fizzes (effervescence as hydrogen gas is released), moves across the surface of the water, and gets smaller until it disappears. If Universal Indicator is added, it turns purple/blue (showing an alkali has formed).
• Lithium: Fizzes steadily, moves on the water.
• Sodium: Fizzes rapidly, melts into a shiny ball, dashes across the surface.
• Potassium: Reacts very violently, catches fire instantly, and burns with a characteristic lilac flame!
Trend in Reactivity
Reactivity increases as you go down Group 1:
\(\text{Lithium} < \text{Sodium} < \text{Potassium} < \text{Rubidium} < \text{Caesium}\)
Why? As you go down the group, atoms have more electron shells, so the single outer electron is further away from the positive nucleus. This makes it easier to lose.
Key Takeaway: Group 1 metals are soft, stored in oil, react with water to make hydrogen + metal hydroxide, and get more reactive down the group.
6. Group 7: The Halogens
Group 7 elements are known as the Halogens. They are reactive non-metals with \(7\) electrons in their outer shell. They exist as diatomic molecules (pairs of atoms joined together, like \(\text{F}_2\), \(\text{Cl}_2\), \(\text{Br}_2\), \(\text{I}_2\)).
Appearance and States at Room Temperature:
• Chlorine (\(\text{Cl}_2\)): Pale green gas (toxic/poisonous).
• Bromine (\(\text{Br}_2\)): Red-brown liquid that gives off dense orange/brown fumes.
• Iodine (\(\text{I}_2\)): Dark grey/black solid that turns into a purple gas when heated (sublimation).
Trend in Reactivity
Reactivity decreases as you go down Group 7:
\(\text{Fluorine} > \text{Chlorine} > \text{Bromine} > \text{Iodine}\)
Why? Group 7 elements need to gain \(1\) electron to get a full outer shell. As you go down the group, the outer shell is further from the nucleus, so it is harder to attract and pull in that extra electron.
Safety Note:
All experiments using halogens must be carried out in a fume cupboard because their vapours are toxic and irritating to the lungs.
Key Takeaway: Halogens are diatomic non-metals. Chlorine is a green gas, Bromine is a red-brown liquid, and Iodine is a dark grey solid. Reactivity decreases down the group.
7. Group 0: The Noble Gases
Group 0 (sometimes called Group 8) contains elements such as Helium (\(\text{He}\)), Neon (\(\text{Ne}\)), and Argon (\(\text{Ar}\)).
Properties of Noble Gases:
• They are all colourless, odourless gases at room temperature.
• They exist as single atoms (monatomic).
• They are completely unreactive (inert) and do not easily form chemical bonds.
Why Are Noble Gases Unreactive?
They have a full outer shell of electrons (Helium has \(2\); the rest have \(8\)). Because their outer shell is already full and stable, they have no need to gain, lose, or share electrons!
Everyday Uses:
• Helium: Used in party balloons and airships because it is much less dense than air and will not catch fire.
• Argon: Used inside traditional filament light bulbs because it is inert and prevents the hot metal filament from burning.
• Neon: Used in glowing advertising signs when an electric current passes through it.
Key Takeaway: Noble gases are unreactive because they have full outer electron shells.
Quick Chapter Summary Checklist
Before you move on, make sure you can answer "yes" to these core skills:
• I can state the mass, charge, and location of protons, neutrons, and electrons.
• I can calculate the number of protons, neutrons, and electrons from atomic and mass numbers.
• I can write the electronic configuration for elements up to atomic number \(20\).
• I know that Group number = outer electrons, and Period number = number of shells.
• I can describe how Group 1 metals react with water and state the trend in reactivity.
• I know the colours, physical states, and reactivity trend of the Group 7 halogens.
• I can explain why Group 0 noble gases are inert.