Periodic Table & Periodic Trends

Chemistry • Secondary School

1. Introduction

The periodic table is a systematic arrangement of chemical elements. It allows chemists to organise elements according to their atomic numbers and to identify patterns in their properties.

Periodic table: A table in which elements are arranged in order of increasing atomic number so that elements with similar properties occur in the same groups.

Learning objectives

Explain how the modern periodic table is organised.
Identify groups and periods.
Relate electron arrangement to position in the table.
Describe broad periodic trends.
Compare properties of elements in the same group.
Explain trends using nuclear attraction and electron shells.
Predict some properties from an element's position.

2. Groups and Periods

Group: A vertical column in the periodic table.

Period: A horizontal row in the periodic table.

Elements in the same main group generally have the same number of outer-shell electrons and therefore often show similar chemical properties.

For the main-group elements, the period number is related to the number of occupied electron shells.

Electron arrangement	Period	Typical outer electrons
2	1	2
2,1	2	1
2,8,1	3	1
2,8,7	3	7

Exam skill: Given an electron arrangement, identify the period from the number of occupied shells and use the outer electrons to help identify the main-group position.

3. Simplified View of the Periodic Table

H
1

He
2

Li
3

Be
4

B
5

C
6

N
7

O
8

F
9

Ne
10

Na
11

Mg
12

Al
13

Si
14

P
15

S
16

Cl
17

Ar
18

K
19

Ca
20

The diagram shows the first 20 elements. The complete modern periodic table contains many more elements, including transition elements and the inner transition series.

4. Important Groups

Group 1 — Alkali metals

Have one outer electron.
Form \(+1\) ions.
Are generally soft metals.
Are highly reactive.
Reactivity generally increases down the group.

Examples include lithium, sodium and potassium.

\[ 2\mathrm{Na}+2\mathrm{H_2O} \rightarrow 2\mathrm{NaOH}+\mathrm{H_2} \]

Group 2 — Alkaline earth metals

Have two outer electrons.
Commonly form \(+2\) ions.
Are metals.
Reactivity generally increases down the group.

Group 17 — Halogens

Have seven outer electrons.
Are non-metals.
Usually form \(−1\) ions.
Exist as diatomic molecules such as \(F_2\), \(Cl_2\) and \(Br_2\).
Reactivity generally decreases down the group.

Group 18 — Noble gases

Have very stable outer electron arrangements.
Are generally very unreactive.
Exist as individual atoms rather than diatomic molecules.
Have low boiling points compared with many other elements.

5. Metals, Non-metals and Metalloids

The periodic table broadly separates metals from non-metals. A staircase-like boundary is commonly used to show the approximate transition between these categories.

Metals	Non-metals
Usually good conductors of heat and electricity	Usually poor conductors
Usually lustrous	Many have dull surfaces
Usually malleable and ductile	Solid non-metals are often brittle
Usually form positive ions	Often form negative ions or share electrons

Metalloid: An element showing a mixture of metallic and non-metallic properties.

Examples often discussed include silicon and boron.

6. Atomic Radius

Atomic radius: A measure of the size of an atom.

Across a period, atomic radius generally decreases from left to right. Down a group, atomic radius generally increases.

Why? Across a period, the nucleus gains protons while electrons are added to the same main shell. The increased nuclear attraction pulls the electron cloud closer to the nucleus. Down a group, additional electron shells are added, making the atom larger.

Across a period: atomic radius generally decreases.
Down a group: atomic radius generally increases.

7. First Ionisation Energy

First ionisation energy: The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous \(1+\) ions.

\[ X(g)\rightarrow X^{+}(g)+e^{-} \]

First ionisation energy generally:

increases across a period;
decreases down a group.

Across a period, nuclear charge generally increases and the outer electron is attracted more strongly. Down a group, the outer electron is farther from the nucleus and more shielded by inner shells.

Important: The overall periodic trend has some small exceptions. In examination questions, explain the general trend first, then consider electron subshell effects when a question specifically requires them.

8. Electronegativity

Electronegativity: The ability of an atom in a molecule to attract the bonding pair of electrons towards itself.

Electronegativity generally increases across a period and decreases down a group.

Fluorine is the most electronegative element on the commonly used Pauling scale.

Remember: Electronegativity describes attraction for bonding electrons. It is different from ionisation energy, which concerns removing an electron from a gaseous atom.

9. Reactivity Trends

Group 1

Group 1 metals become more reactive down the group because the outer electron is farther from the nucleus and is more easily lost.

\[ \mathrm{Li}<\mathrm{Na}<\mathrm{K} \]

in general order of increasing reactivity.

Group 17

Halogens become less reactive down the group because it becomes more difficult for the atom to attract and gain an additional electron.

\[ \mathrm{F_2}>\mathrm{Cl_2}>\mathrm{Br_2}>\mathrm{I_2} \]

in general order of decreasing ability to gain an electron.

10. Periodic Trends Across Period 3

The elements sodium to argon provide an important example of periodic trends.

Property	General trend Na → Ar	Main reason
Atomic radius	Decreases	Increasing nuclear attraction
First ionisation energy	Generally increases	Increasing nuclear attraction
Electronegativity	Generally increases	Increasing attraction for bonding electrons
Metallic character	Decreases	Atoms hold electrons more strongly

Example: Which is larger, Na or Cl?

Answer: Na is larger because atomic radius generally decreases from left to right across a period.

11. Position and Electronic Configuration

Example: An element has electron arrangement \(2,8,2\).

It has three occupied shells, so it is in Period 3. It has two outer electrons, so it belongs to the main-group family associated with Group 2.

The element is magnesium, Mg.

Example: An element has electron arrangement \(2,8,7\).

It has three occupied shells and seven outer electrons. It is therefore chlorine, Cl, in Period 3 and Group 17.

12. Explaining Periodic Trends

Good examination answers should use the following ideas where appropriate:

Nuclear charge.
Number of occupied electron shells.
Distance of the outer electron from the nucleus.
Shielding by inner electrons.
Attraction between the nucleus and electrons.

Strong explanation pattern: State the trend → identify the change in nuclear charge or shells → explain the effect on attraction → state the resulting property.

13. Worked Questions

Question 1: Explain why atomic radius increases down Group 1.

Solution:

Each element down the group has an additional occupied electron shell. The outer electron is therefore farther from the nucleus and is more shielded by inner electrons. The attraction between the nucleus and outer electron is weaker, so the atomic radius increases.

Question 2: Which has the greater first ionisation energy: Na or Mg?

Solution:

Both are in Period 3, but Mg has a greater nuclear charge while the outer electrons are in the same main shell. Therefore Mg generally holds its outer electron more strongly and has the greater first ionisation energy.

Question 3: Why does the reactivity of Group 17 elements decrease down the group?

Solution:

Halogens react by gaining an electron. Down the group, additional electron shells increase the distance and shielding between the nucleus and the incoming electron. The attraction becomes weaker, so the tendency to gain an electron decreases.

14. MANEB Examination Focus

Identify groups and periods from electron arrangements.
Describe properties of Groups 1, 2, 17 and 18.
Explain periodic trends.
Compare atomic sizes.
Compare first ionisation energies.
Explain reactivity trends.
Distinguish electronegativity from ionisation energy.
Use nuclear charge, shells and shielding in explanations.
Relate position in the periodic table to electronic structure.

15. Practice Exercises

1. Define a group in the periodic table.

2. Define a period.

3. An element has electron arrangement \(2,8,1\). Identify its period and main group.

4. State the general trend in atomic radius across a period.

5. State the general trend in atomic radius down a group.

6. Explain why atomic radius increases down a group.

7. State the general trend in first ionisation energy across a period.

8. Which is more reactive: Na or K? Explain.

9. Which halogen is generally more reactive, Cl₂ or Br₂? Explain.

10. Explain why electronegativity generally increases across a period.

Answers

A vertical column of elements.
A horizontal row of elements.
Period 3; Group 1.
It generally decreases from left to right.
It generally increases down the group.
Additional electron shells increase atomic size and shielding.
It generally increases from left to right.
K is more reactive because its outer electron is farther from the nucleus and easier to remove.
Cl₂ is generally more reactive because it is higher in Group 17 and attracts an incoming electron more strongly.
Nuclear charge increases while electrons are added to the same main shell, increasing attraction for bonding electrons.

16. Common Mistakes

Confusing groups with periods.
Assuming all periodic trends are absolute without considering exceptions.
Confusing atomic radius with ionisation energy.
Confusing electronegativity with electron affinity or ionisation energy.
Explaining trends without mentioning nuclear attraction or electron shells.
Forgetting that Group 1 and Group 17 reactivity trends move in opposite directions.

17. Examination Tips

Draw or refer to the periodic table when it helps your explanation.
When asked “explain why”, give the trend and the reason.
Use the words nuclear charge, shielding and distance appropriately.
For electronic configurations, count shells carefully.
Learn the characteristic properties of Groups 1, 2, 17 and 18.

18. Topic Summary

The periodic table arranges elements according to increasing atomic number. Elements in the same group have related outer-electron structures and often similar chemical properties. Periods represent horizontal rows and are related to occupied electron shells.

Across a period, atomic radius generally decreases while ionisation energy and electronegativity generally increase. Down a group, atomic radius generally increases. These trends can be explained using nuclear charge, electron shells, shielding and distance from the nucleus.