The Periodic Table and Trends
Understand the organisation of the periodic table and explore key trends including atomic radius, ionisation energy, electronegativity and reactivity across periods and down groups.
Structure of the Periodic Table
The periodic table organises all known elements by increasing atomic number. Elements in the same group (vertical column) have the same number of valence electrons and similar chemical properties. Elements in the same period (horizontal row) have the same number of electron shells.
Key Regions of the Periodic Table
Group 1
Alkali Metals
1 valence electron
Group 2
Alkaline Earth Metals
2 valence electrons
Group 17
Halogens
7 valence electrons
Group 18
Noble Gases
Full outer shell
Metals
Left and centre. Shiny, conduct electricity, malleable, tend to lose electrons to form cations.
Non-metals
Upper right. Poor conductors, brittle (if solid), tend to gain electrons to form anions.
Metalloids
Along the staircase line. Have properties between metals and non-metals. E.g. silicon, germanium.
Periodic Trends
Several properties vary predictably across periods (left to right) and down groups. These trends are caused by changes in nuclear charge, atomic radius and shielding (inner electron repulsion).
Summary of Trends
| Property | Across a Period → | Down a Group ↓ |
|---|---|---|
| Atomic Radius | Decreases | Increases |
| Ionisation Energy | Increases | Decreases |
| Electronegativity | Increases | Decreases |
| Metallic Character | Decreases | Increases |
Why these trends occur: Across a period, the nuclear charge increases (more protons) while shielding stays roughly constant, so electrons are pulled closer. Down a group, new electron shells are added, increasing atomic size and reducing the pull on outer electrons.
Reactivity Trends
Reactivity depends on how easily an atom can achieve a stable electron configuration (full outer shell). Metals react by losing electrons; non-metals react by gaining electrons.
Metals (Groups 1 & 2)
Reactivity increases down the group.
As atomic radius increases, the outer electron is further from the nucleus and easier to remove. Francium is the most reactive metal.
Non-metals (Group 17)
Reactivity decreases down the group.
Smaller atoms attract electrons more strongly. Fluorine is the most reactive non-metal and the most electronegative element.
Remember: Metals and non-metals show opposite reactivity trends down a group because metals lose electrons (easier when far from nucleus) while non-metals gain electrons (easier when close to nucleus).
Key Vocabulary
Ionisation Energy
The energy required to remove the outermost electron from an atom in the gaseous state. Higher ionisation energy means the electron is harder to remove.
Electronegativity
The ability of an atom to attract shared electrons in a covalent bond toward itself. Fluorine has the highest electronegativity (4.0 on the Pauling scale).
Atomic Radius
Half the distance between the nuclei of two bonded atoms of the same element. A measure of how large an atom is.
Shielding Effect
The reduction in the nuclear attraction on outer electrons caused by the repulsion of inner-shell electrons. More shells means more shielding.
Worked Examples
Explain why sodium has a larger atomic radius than chlorine.
Both are in Period 3, so they have the same number of electron shells (3).
Chlorine (Z = 17) has more protons than sodium (Z = 11), but the shielding effect is similar.
The greater nuclear charge in Cl pulls electrons closer, resulting in a smaller atomic radius. Na has fewer protons pulling on the same number of shells, so it is larger.
Why does potassium react more vigorously with water than sodium?
Both are Group 1 metals that react by losing one electron.
Potassium has one more electron shell than sodium, so its outer electron is further from the nucleus and more shielded.
This means less energy is required to remove the outer electron, so potassium is more reactive.
Arrange these elements in order of increasing first ionisation energy: Na, Mg, Al, Si.
These are all Period 3 elements. Ionisation energy generally increases across a period.
However, Al has a slight dip because its outer electron is in a 3p orbital, which is slightly easier to remove than Mg's 3s electron.
Order: Na < Al < Mg < Si (noting the Al anomaly)
Knowledge Check
Select the correct answer for each question. Click "Check Answer" to see if you are right.
Question 1
Elements in the same group of the periodic table have the same number of:
Question 2
What happens to atomic radius as you move across a period from left to right?
Question 3
Which element has the highest electronegativity?
Question 4
In Group 1 (alkali metals), which element is the MOST reactive?
Question 5
First ionisation energy generally __________ across a period.
Key Concepts Summary
- ●Elements are arranged by atomic number. Groups share valence electrons; periods share electron shells.
- ●Atomic radius decreases across a period and increases down a group.
- ●Ionisation energy and electronegativity increase across a period and decrease down a group.
- ●Metal reactivity increases down a group; non-metal reactivity decreases down a group.
- ●These trends are explained by changes in nuclear charge, shielding and atomic size.