CORE Chemistry (9222) | The periodic table

Welcome to The Periodic Table: The Chemist's Ultimate Cheat Sheet

Hi there! Don't worry if Chemistry sometimes feels like a mountain of facts. The great news is that the Periodic Table is the ultimate organization tool—a kind of library catalog for all the elements in the universe!

In this chapter, we'll learn how this table is structured, why elements are placed where they are, and how their position predicts their behavior (their chemistry!). Understanding the table is key to success in this whole section of the course. Let's dive in!

1. The Structure of the Periodic Table

The periodic table is arranged logically. Elements are ordered based on the number of protons they have, which is their Atomic Number (or Proton Number).

1.1 Groups (The Columns)

• What they are: The vertical columns on the table.
• Why they matter: Elements in the same Group have very similar chemical properties. Think of them as chemical "families."
• Analogy: Imagine an apartment building. Everyone in the same vertical column (Group) belongs to the same family, sharing characteristics.
• The Magic Link: The Group number (e.g., Group 1, Group 7) tells you the number of outer shell electrons (also called valency electrons) an atom has.

1.2 Periods (The Rows)

• What they are: The horizontal rows on the table.
• Why they matter: Properties change gradually as you move across a Period (from left to right).
• The Magic Link: The Period number (e.g., Period 2, Period 3) tells you the number of electron shells the atom has.

2. General Properties: Metals and Non-Metals

The table is broadly divided into two main categories: Metals and Non-metals.

2.1 Location

• Metals: Found mainly on the left side and centre of the table (e.g., Iron, Sodium).
• Non-metals: Found on the right side of the table (e.g., Oxygen, Chlorine).

2.2 Key Property Differences

When you look at their physical and chemical behaviors, the differences are clear:

Property	Metals	Non-metals
Appearance	Shiny (lustrous), solids (except mercury)	Dull, can be solids, liquids, or gases
Conductivity	Good conductors of heat and electricity	Poor conductors (Insulators)
Strength	Malleable (can be hammered into sheets) and Ductile (can be drawn into wires)	Brittle (shatter easily)
Ions Formed	Tend to lose electrons to form positive ions	Tend to gain electrons to form negative ions

3. The Reactive Families: Group 1 and Group 7

These two groups are the most reactive on the table because they are only one electron away from having a full, stable outer shell.

3.1 Group 1: The Alkali Metals

These are the first column, including Lithium (Li), Sodium (Na), and Potassium (K).

• Electron Structure: All have 1 outer electron. They desperately want to lose this electron to become stable positive ions (\(M^{+}\)).
• Physical Properties: They are soft (you can cut them with a knife!), have low melting points, and low density.
• Reactivity: Very reactive, especially with water and air. They must be stored under oil to prevent reactions.

The Trend in Reactivity for Group 1

As you move down Group 1, the elements become MORE reactive.

Step-by-step Explanation (Why they get stronger down the group):

1. Going down the group, atoms get bigger because they have more electron shells.
2. The outer electron is further away from the positively charged nucleus.
3. The attraction between the nucleus and the outer electron becomes weaker.
4. Therefore, it takes less energy to lose (or "ditch") that single outer electron, making the element more reactive.

Did you know? When Alkali metals react with water, they produce hydrogen gas and an alkaline solution (a hydroxide). The reaction for Potassium is so vigorous that the heat produced causes the hydrogen gas to ignite!

3.2 Group 7: The Halogens

These are non-metals, including Fluorine (F), Chlorine (Cl), Bromine (Br), and Iodine (I).

• Electron Structure: All have 7 outer electrons. They desperately want to gain 1 electron to become stable negative ions (\(X^{-}\)).
• Diatomic Molecules: Halogens exist as pairs of atoms (e.g., \(Cl_2\), \(Br_2\)).
• Physical Properties: Their melting and boiling points increase down the group.
  • Fluorine (F) and Chlorine (Cl) are gases.
  • Bromine (Br) is a liquid.
  • Iodine (I) is a solid.

The Trend in Reactivity for Group 7

As you move down Group 7, the elements become LESS reactive.

Step-by-step Explanation (Why they get weaker down the group):

1. Going down the group, atoms get bigger (more electron shells).
2. Halogens react by attracting an extra electron.
3. When the outer shell is further away from the nucleus, the positive nucleus has a harder time pulling in that new, incoming electron.
4. Therefore, the ability to gain an electron decreases, making the element less reactive.

4. The Unreactive Family: Group 0 (The Noble Gases)

These elements (Helium, Neon, Argon, etc.) sit on the far right of the table. They are extremely important because they represent the goal state for all other elements!

• Electron Structure: All have a full outer shell (usually 8 electrons, except Helium which only needs 2).
• Reactivity: They are unreactive or inert. They are already stable and do not need to lose, gain, or share electrons.
• State: They all exist as single atoms (monatomic) and are colorless gases at room temperature.
• Uses: Because they are inert, they are useful when you need an element that won't react. Example: Argon is used inside light bulbs to stop the hot metal filament from reacting with oxygen.