The periodic table - Chemistry (9202) - Oxford AQA IGCSE

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Welcome to the Ultimate Chemistry Filing Cabinet! The Periodic Table

Hello future chemist! Get ready to explore one of the most powerful tools in science: The Periodic Table. Don't worry if it looks intimidating—it's actually the most organised filing system you will ever see!

In this chapter, we will learn how this table is structured, why elements are placed where they are, and how that placement predicts their chemical behaviour. Understanding the periodic table is key to understanding almost all of chemistry. Let's dive in!

1. Understanding the Structure of the Periodic Table

The Periodic Table organises elements by increasing atomic number (the number of protons). But the real magic lies in how it organises them by chemical properties using rows and columns.

1.1 Groups (Vertical Columns)

Imagine groups like families living in the same apartment column.

Definition: Vertical columns are called Groups.
Key Property: All elements in the same group have the same number of electrons in their outermost shell (valence shell).
Result: Elements in the same group have very similar chemical properties because their reactions are determined by their outer electrons.
Example: Group 1 elements (Lithium, Sodium, Potassium) all have 1 electron in their outer shell, making them all highly reactive metals.

Memory Aid for Groups:

GReen (GRoup) stands up Great (Vertical).

1.2 Periods (Horizontal Rows)

Imagine periods like the different floors of the building.

Definition: Horizontal rows are called Periods.
Key Property: All elements in the same period have the same number of occupied electron shells.
Result: As you move across a period (from left to right), the properties change gradually from metallic to non-metallic.
Example: Elements in Period 3 (Sodium to Argon) all have three electron shells.

Quick Review: Key Takeaway

Groups tell us: How many outer electrons (chemical properties).

Periods tell us: How many electron shells (size of the atom).

2. Metals, Non-metals, and Transition Elements

The periodic table is naturally split into metals and non-metals, which are separated by a "staircase" line.

2.1 Location and Properties

Metals: Found on the left and centre of the table (e.g., Na, Fe, Al).
- Properties: Usually solid (except Mercury), shiny, high melting points, good conductors of heat and electricity, ductile (can be drawn into wires), malleable (can be hammered into shape).
- Chemical behaviour: Tend to lose electrons to form positive ions ($+$).
Non-metals: Found on the right of the table (e.g., O, S, Cl).
- Properties: Usually gases or brittle solids, dull, low melting points, poor conductors (insulators).
- Chemical behaviour: Tend to gain electrons to form negative ions ($-$).
Metalloids (or Semi-metals): These sit along the staircase boundary (e.g., Silicon, Boron). They have properties of both metals and non-metals.

2.2 The Transition Elements

The large block of elements in the centre of the table (between Group 2 and Group 3) are the Transition Elements (e.g., Iron, Copper, Gold).

These metals are typically:

Harder and stronger than Group 1 metals.
Have very high melting points (Mercury is an exception).
Often used as catalysts (substances that speed up reactions).

3. Exploring Specific Groups: Group 1 (Alkali Metals)

Group 1 elements (Lithium, Sodium, Potassium, Rubidium, Caesium, Francium) are extremely important!

3.1 General Properties

They are soft metals (can be cut with a knife).
They have surprisingly low densities and low melting points.
They are very reactive because they only need to lose 1 electron to achieve a stable full shell.
They are always stored under oil to prevent reaction with oxygen or water vapour in the air.

3.2 Reactions with Water (Crucial Concept)

When an alkali metal reacts with water, it produces a metal hydroxide (which is an alkali solution) and hydrogen gas.

Metal + Water $\to$ Metal Hydroxide + Hydrogen

Example (Sodium): $$2Na_{(s)} + 2H_2O_{(l)} \to 2NaOH_{(aq)} + H_{2(g)}$$

Observation (Lithium): Floats, fizzes slowly.
Observation (Sodium): Floats, fizzes vigorously, melts into a ball, often moves quickly across the surface.
Observation (Potassium): Floats, fizzes very vigorously, produces a lilac flame (it's hot enough to ignite the hydrogen gas!).

3.3 Trends Down Group 1: Reactivity Increases

As you go down Group 1 (from Li to K):

Atomic Size Increases: Each element has one more electron shell than the one above it.
Outer Electron is Further Away: The single valence electron is further from the positively charged nucleus.
Shielding Increases: The inner shells "shield" the outer electron from the pull of the nucleus.
Conclusion: It takes less energy to remove this single electron. Therefore, the elements get more reactive as you go down the group.

Analogy: Imagine holding a dog on a very long leash (Caesium). It's much easier for the dog to run away (react) than if it was on a very short leash (Lithium).

4. Exploring Specific Groups: Group 7 (The Halogens)

Group 7 elements (Fluorine, Chlorine, Bromine, Iodine, Astatine) are known as the Halogens (meaning "salt formers").

4.1 General Properties and Physical Trends

They are highly reactive non-metals, needing only 1 electron to fill their shell.
They exist as diatomic molecules (two atoms bonded together, e.g., $\text{Cl}_2$, $\text{Br}_2$).
They are toxic and corrosive.

Physical Trends Down Group 7 (Cl to I):

Colour: Gets darker.
Melting/Boiling Point: Increases (due to stronger intermolecular forces).
State at Room Temperature: Changes from gas $\to$ liquid $\to$ solid.
- Fluorine ($\text{F}_2$): Pale yellow gas
- Chlorine ($\text{Cl}_2$): Green-yellow gas
- Bromine ($\text{Br}_2$): Red-brown liquid
- Iodine ($\text{I}_2$): Grey solid (forms purple vapour when heated)

4.2 Chemical Trend Down Group 7: Reactivity Decreases

Halogens react by gaining an electron to form a negative ion ($1-$).

As you go down Group 7 (from F to I):

Atomic Size Increases: More electron shells.
Attraction Decreases: The incoming electron is attracted by the positive nucleus, but this attraction is weaker because the electron is further away and heavily shielded by the increasing inner shells.
Conclusion: It gets harder to attract the extra electron needed for stability. Therefore, the elements get less reactive as you go down the group.

Encouragement: Don't worry if the trend seems tricky! Just remember: Group 1 activity increases down, Group 7 activity decreases down.

4.3 Halogen Displacement Reactions

A crucial reaction for Group 7 involves one halogen displacing another from a salt solution.

The rule is simple: A more reactive halogen will displace a less reactive halogen from its compound.

Since reactivity decreases down the group, Chlorine ($\text{Cl}_2$) is more reactive than Bromine ($\text{Br}_2$) or Iodine ($\text{I}_2$).
Bromine ($\text{Br}_2$) is more reactive than Iodine ($\text{I}_2$).

Step-by-Step Example: Chlorine displacing Bromine

If you add chlorine water ($\text{Cl}_2$) to a solution of potassium bromide ($\text{KBr}$):

Chlorine is more reactive than bromine.
Chlorine displaces the bromine, forming potassium chloride and elemental bromine.

$$Cl_{2(aq)} + 2KBr_{(aq)} \to 2KCl_{(aq)} + Br_{2(aq)}$$

Observation: The solution changes colour as the colourless bromide ion is converted into brown bromine liquid.

Common Mistake to Avoid: A less reactive halogen can never displace a more reactive one. Chlorine water will react with potassium iodide, but Iodine water will NOT react with potassium chloride.

5. Exploring Specific Groups: Group 0 (The Noble Gases)

Group 0 (or Group 8) contains Helium, Neon, Argon, Krypton, Xenon, and Radon. These are the cool, laid-back elements of the table!

5.1 Properties and Stability

Outer Shell: They all have a full outer electron shell (2 for Helium, 8 for the others).
Reactivity: Because their shells are full, they are chemically inert (unreactive). They do not easily gain or lose electrons.
State: They are all colourless, odourless gases at room temperature.
Structure: They exist as monatomic gases (single atoms, e.g., $\text{He}$, not $\text{He}_2$).

5.2 Practical Uses

Their inertness makes them extremely useful where a non-flammable, non-reactive atmosphere is needed:

Argon ($\text{Ar}$): Used in filament light bulbs to stop the hot metal filament from reacting with oxygen. Also used as an inert shield during welding.
Helium ($\text{He}$): Used to fill balloons and airships because it is very light and non-flammable (unlike hydrogen).
Neon ($\text{Ne}$): Used in "neon" advertising signs because it glows brightly when electricity passes through it.

Did you know?

Helium's name comes from the Greek word for the sun, helios. It was detected in the sun’s spectrum before it was discovered on Earth!

6. General Trends Across a Period

As we move left to right across any period (e.g., Period 3: Na $\to$ Mg $\to$ Al $\to$ Si $\to$ P $\to$ S $\to$ Cl $\to$ Ar):

6.1 Nuclear Charge and Atomic Size

Number of Shells: Stays the same (e.g., all have 3 shells in Period 3).
Number of Protons (Nuclear Charge): Increases by one with each step.
Atomic Size Trend: The increased positive charge in the nucleus pulls the electron shells closer. Therefore, atomic size decreases across a period.

6.2 Metallic to Non-Metallic Character

As we move from left to right:

Elements start as strong metals (Group 1 and 2, tend to lose electrons).
They transition through metalloids (e.g., Silicon).
They end as strong non-metals (Group 7, tend to gain electrons).

6.3 Reactivity Trend

Reactivity follows a "valley" shape across a period:

High reactivity on the far left (metals that want to lose 1 or 2 electrons).
Low reactivity in the middle (metalloids and elements requiring complex sharing).
High reactivity on the far right (halogens that want to gain 1 electron).
Zero reactivity in Group 0 (Noble Gases).

Final Key Takeaway Summary

Group 1 (Metals): Reactivity increases down the group (easier to lose electron).
Group 7 (Non-metals): Reactivity decreases down the group (harder to gain electron).
Across a Period: Atomic size decreases, and elements change from metallic to non-metallic.

Great job! You now understand the basic blueprint of all matter. Keep practising those group trends—they are the key to predicting chemical reactions!