Chemistry (0620) | Atomic structure and the Periodic Table

Welcome to Atomic Structure and the Periodic Table!

Hello future chemists! This chapter is absolutely crucial because it unlocks the secrets of what everything is made of and how it’s organized. Understanding the atom is like learning the alphabet of Chemistry—once you master it, you can start building words (molecules) and sentences (reactions)!

Don't worry if some of the terminology seems strange at first; we will break everything down using simple steps and helpful analogies. Let’s dive in and explore the tiny world that makes up our universe!

1. The Structure of the Atom (Syllabus 2.2)

Everything around us is made of tiny particles called atoms. While you can't see them, we know they have a specific structure.

1.1 Subatomic Particles

An atom has two main regions: the central core (the nucleus) and the outer region (the electron shells). Inside, there are three types of particles:

• Protons (\(\mathrm{p}\))
• Neutrons (\(\mathrm{n}\))
• Electrons (\(\mathrm{e}\))

The nucleus contains the protons and neutrons. The electrons orbit the nucleus in shells.

Quick Review: Relative Mass and Charge

It's essential to remember the properties of these particles:

Particle	Location	Relative Mass	Relative Electrical Charge
Proton	Nucleus	1	+1 (Positive)
Neutron	Nucleus	1	0 (Neutral)
Electron	Shells/Orbits	Negligible (almost 0)	-1 (Negative)

Analogy: Think of the atom as a stadium. The nucleus (protons and neutrons) is the heavy, dense center pitch, and the electrons are tiny fans running around the stands (shells).

1.2 Defining Elements using Numbers

What makes one element (like Carbon) different from another (like Oxygen)? It’s all about the numbers in the nucleus!

• Proton Number (Atomic Number), \(\mathrm{Z}\): This is the number of protons in the nucleus of an atom.
  • Importance: The proton number defines the element. If the proton number changes, the element changes!
• Mass Number (Nucleon Number), \(\mathrm{A}\): This is the total number of protons and neutrons in the nucleus.
  • Calculation: Mass Number = Proton Number + Number of Neutrons.

In a neutral atom:

Number of Protons = Number of Electrons. This is because the positive charges (\(\mathrm{p}\)) must equal the negative charges (\(\mathrm{e}\)) so the atom has no overall charge.

Interpreting Atomic Symbols (Syllabus 2.3.2)

Elements are often represented using the notation: \(^{A}_{Z}\mathrm{X}\)

Where:

• \(\mathrm{X}\) is the chemical symbol (e.g., C for Carbon).
• \(\mathrm{A}\) is the Mass Number (top).
• \(\mathrm{Z}\) is the Proton Number (bottom).

Example: A sodium atom is shown as \(^{23}_{11}\mathrm{Na}\).

• Protons (\(\mathrm{Z}\)) = 11
• Electrons (for a neutral atom) = 11
• Neutrons (\(\mathrm{A} - \mathrm{Z}\)) = 23 – 11 = 12

2. Electronic Configuration (Syllabus 2.2.5, 2.2.6)

Electrons orbit the nucleus in specific energy levels called shells. The way these electrons are arranged is the electronic configuration.

2.1 Rules for Filling Shells

For the first 20 elements (which is all you need for IGCSE):

1. The first shell (closest to the nucleus) can hold a maximum of 2 electrons.
2. The second shell can hold a maximum of 8 electrons.
3. The third shell can hold a maximum of 8 electrons (when forming stable compounds, though it can technically hold more later).

We write the configuration by listing the number of electrons in each occupied shell, separated by commas (e.g., 2, 8, 3).

Step-by-Step Example: Aluminium (Proton Number 13)

1. Total electrons = 13.
2. Fill the 1st shell: 2 electrons used. (Remaining: 11)
3. Fill the 2nd shell: 8 electrons used. (Remaining: 3)
4. The last 3 electrons go into the 3rd shell.
5. Configuration is 2, 8, 3.

2.2 Electrons in Ions

Atoms gain or lose electrons to achieve a stable electronic structure, usually a full outer shell, which mimics the structure of the nearest Noble Gas. When they do this, they become ions.

• Positive Ions (Cations): Formed when metals lose electrons. They have fewer electrons than protons. Example: Sodium (2, 8, 1) loses 1 electron to become \(\mathrm{Na}^{+}\) (2, 8).
• Negative Ions (Anions): Formed when non-metals gain electrons. They have more electrons than protons. Example: Chlorine (2, 8, 7) gains 1 electron to become \(\mathrm{Cl}^{-}\) (2, 8, 8).

Example of an Ion Symbol: Chloride ion is \(^{35}_{17}\mathrm{Cl}^{-}\).

• Protons: 17
• Electrons: 17 + 1 (gained one electron) = 18
• Configuration: 2, 8, 8

3. Isotopes (Syllabus 2.3)

3.1 Definition and Properties

Isotopes are different atoms of the same element that have the same number of protons but a different number of neutrons.

Example: Carbon-12 (\(^{12}_{6}\mathrm{C}\)) and Carbon-14 (\(^{14}_{6}\mathrm{C}\)).

• Both have 6 protons (so they are both Carbon).
• Carbon-12 has 6 neutrons (12-6=6).
• Carbon-14 has 8 neutrons (14-6=8).

Since isotopes belong to the same element, they have virtually identical chemical properties.

Why are the chemical properties the same? (Supplement)

Chemical reactions involve the outer electrons. Since isotopes have the same number of protons, they must also have the same number of electrons and thus the same electronic configuration. Therefore, they react in the exact same way! The only difference is their mass (due to the different number of neutrons).

3.2 Calculating Relative Atomic Mass, \(A_r\) (Supplement)

The Relative Atomic Mass (\(A_r\)) listed on the Periodic Table is rarely a whole number. This is because it is the average mass of all the naturally occurring isotopes of that element, compared to 1/12th the mass of a Carbon-12 atom.

To calculate the \(A_r\), you need the mass and abundance of each isotope.

Formula:

$$A_r = \frac{(\text{mass}_{1} \times \text{abundance}_{1}) + (\text{mass}_{2} \times \text{abundance}_{2})}{\text{Total Abundance (usually 100)}}$$

Step-by-Step Example Calculation

Chlorine has two common isotopes: Chlorine-35 (abundance 75%) and Chlorine-37 (abundance 25%).

1. Multiply mass by abundance for Isotope 1: \(35 \times 75 = 2625\)
2. Multiply mass by abundance for Isotope 2: \(37 \times 25 = 925\)
3. Add the results: \(2625 + 925 = 3550\)
4. Divide by total abundance (100): \(3550 / 100 = 35.5\)

The relative atomic mass of Chlorine is 35.5.

4. The Periodic Table: Arrangement and General Trends (Syllabus 8.1)

The Periodic Table organizes all known elements based on their atomic structure and properties.

4.1 Arrangement of Elements

The table is arranged in order of increasing Proton Number (Atomic Number).

• Periods (Horizontal Rows): The period number tells you the number of occupied electron shells in the atoms of that element.
• Groups (Vertical Columns, I to VIII): The group number tells you the number of outer shell electrons (valence electrons) for Groups I to VII (and VIII/0). These outer electrons determine the element's chemical properties.

Example: Sodium (\(\mathrm{Na}\)) is in Period 3 and Group I. This means it has 3 electron shells and 1 electron in its outer shell (2, 8, 1).

4.2 General Trends Across a Period

As you move from left to right across a period:

• The number of outer electrons increases (1 to 8).
• The elements change from metallic character to non-metallic character. (E.g., Period 3 starts with Sodium (metal) and ends with Argon (non-metal/Noble Gas)).
• Elements in Group I, II, and III form positive ions (\(\mathrm{Na}^{+}\), \(\mathrm{Mg}^{2+}\), \(\mathrm{Al}^{3+}\)) by losing electrons.
• Elements in Group VI and VII form negative ions (\(\mathrm{O}^{2-}\), \(\mathrm{Cl}^{-}\)) by gaining electrons.
• Elements in Group IV and V tend to share electrons (covalent bonding).

5. Exploring Specific Groups

5.1 Group I: The Alkali Metals (Syllabus 8.2)

Group I elements (Lithium, Sodium, Potassium, etc.) are known as Alkali Metals. They are highly reactive metals.

Structure and Reactivity:

• They all have 1 outer electron.
• They easily lose this one electron to form a stable positive ion (\(1^{+}\)).

Trends Down Group I (Li \(\rightarrow\) Na \(\rightarrow\) K)

As you go down the group:

1. Melting Point Decreases: They get softer and melt more easily.
2. Density Increases: They get heavier for the same volume.
3. Reactivity Increases: They lose their outer electron more easily because the electron is further from the nucleus (more shielding by inner shells), meaning the force of attraction is weaker.

Mnemonic: Think of Group I getting "Bigger and Angrier" (more reactive) as you go down.

Properties: They are relatively soft, shiny metals (though they tarnish quickly in air), and have low densities compared to transition metals. They react violently with water to produce hydrogen gas and a metal hydroxide (an alkali).

5.2 Group VII: The Halogens (Syllabus 8.3)

Group VII elements (Fluorine, Chlorine, Bromine, Iodine) are known as Halogens. They are highly reactive non-metals.

Structure and Reactivity:

• They all have 7 outer electrons.
• They easily gain one electron to form a stable negative ion (\(1^{-}\)).
• They exist as diatomic molecules (\(\mathrm{Cl}_{2}\), \(\mathrm{Br}_{2}\), \(\mathrm{I}_{2}\)).

Physical Appearance at Room Temperature (r.t.p.)

• Chlorine (\(\mathrm{Cl}_{2}\)): Pale yellow-green gas.
• Bromine (\(\mathrm{Br}_{2}\)): Red-brown liquid.
• Iodine (\(\mathrm{I}_{2}\)): Grey-black solid.

Trends Down Group VII (\(\mathrm{Cl} \rightarrow \mathrm{Br} \rightarrow \mathrm{I}\))

As you go down the group:

1. Density Increases: They become heavier.
2. Reactivity Decreases: It becomes harder to attract and gain an extra electron because the outer shell is further away from the nucleus.

Displacement Reactions

A more reactive halogen will displace a less reactive halide from its salt solution.

Reactivity Order (Most to Least): \(\mathrm{Cl}_{2} > \mathrm{Br}_{2} > \mathrm{I}_{2}\)

Example: Chlorine water added to Potassium Bromide solution:

$$\mathrm{Cl}_{2}(\mathrm{aq}) + 2\mathrm{KBr}(\mathrm{aq}) \rightarrow 2\mathrm{KCl}(\mathrm{aq}) + \mathrm{Br}_{2}(\mathrm{aq})$$

(Chlorine is more reactive than bromine, so it successfully displaces it. The solution changes colour as bromine is formed.)

5.3 Group VIII (or 0): The Noble Gases (Syllabus 8.5)

Group VIII elements (Neon, Argon, Helium, etc.) are known as Noble Gases.

Properties:

• They are unreactive (inert).
• They are monatomic gases (exist as single atoms, e.g., \(\mathrm{Ne}\), not \(\mathrm{Ne}_{2}\)).
• They have a full outer electron shell (2 for Helium, 8 for the rest).

This full shell means they are already extremely stable. They neither need to gain nor lose electrons, hence they do not react easily.

6. Transition Elements (Syllabus 8.4)

The Transition Elements are the large block of metals located in the middle of the Periodic Table (Groups 3 to 12). Think of familiar metals like Iron, Copper, and Gold.

6.1 Characteristic Properties (Core)

Transition metals exhibit distinct properties that set them apart from the highly reactive Group I and II metals:

• High Densities: They are typically very dense (heavy).
• High Melting Points: They require high temperatures to melt (Mercury is the notable exception).
• Form Coloured Compounds: Their compounds are often brightly coloured. Example: Copper(II) compounds are blue, Iron(III) compounds are brown/yellow.
• Often Act as Catalysts: Both as elements and in compounds, they speed up chemical reactions without being used up themselves. Example: Iron is used in the Haber process.

6.2 Variable Oxidation States (Supplement)

Unlike Group I elements which only form +1 ions, transition metals often have ions with different charges (known as variable oxidation numbers or states).

Example: Iron can form two stable ions:

• Iron(II), \(\mathrm{Fe}^{2+}\) (often green in solution).
• Iron(III), \(\mathrm{Fe}^{3+}\) (often yellow or brown in solution).

This ability to change oxidation state is part of why they make such good catalysts.