Atomic Structure and Properties
Periodic Trends — AP CHEMISTRY Revision Notes & Topical Questions
5 min read 10 parts Lesson 5 of 16 Unit quiz
Lesson summary
Periodic Trends — AP CHEMISTRY notes: Explain periodic trends in ionisation energy, radius, electron affinity and electronegativity, and predict ion charges.The table arranges elements by increasing atomic number, 1 to 118.Periods (rows) represent principal energy shells.Groups (…
- Configuration from position
- Periodicity
- Atomic radius
- Effective nuclear charge
- Valence electrons
- Groups (columns) share similar chemical properties because of equal valence electron numbers
Section 1 of 10Structure of the periodic table
Structure of the periodic table
The table arranges elements by increasing atomic number, 1 to 118.
Periods (rows) represent principal energy shells.
Groups (columns) share similar chemical properties because of equal valence electron numbers.
Blocks
- Blocks s-block has only s electrons in the valence shell, p-block has at least one p electron, d-block has d electrons, and f-block has f electrons.
Configuration from position
- Configuration from position germanium is in Group 14, Period 4, p-block position 2, so its configuration ends 4s2 3d10 4p2.
Section 2 of 10Explaining periodicity
Periodicity
Periodicity recurring, predictable trends in properties.
Trends are explained using Coulomb's law, the shell model, and shielding or effective nuclear charge.
Across a period nuclear charge rises while shielding stays about the same, so valence electrons are pulled closer.
Down a group extra inner shells increase shielding and distance, which weakens the attraction.
Section 3 of 10Ionisation energy
Ionisation energy
Ionisation energy the energy needed to remove one mole of electrons from one mole of gaseous atoms, for example Na(g) → Na+(g) + e−.
It generally increases across a period because the effective nuclear charge rises.
It decreases down a group because shielding and distance increase.
Section 4 of 10Atomic radius
Atomic radius
Atomic radius the distance from the nucleus to the outermost electron shell.
It decreases across a period as the effective nuclear charge rises.
It increases down a group as more shells are added.
It follows the opposite trend to ionisation energy.
Section 5 of 10Ionic radius
Ionic radius
Down a group ionic radius increases as shells are added.
Positive ions are smaller than their atoms because electrons are lost and repulsion falls; higher positive charge gives a smaller ion.
Negative ions are larger than their atoms because added electrons increase repulsion with the nuclear charge unchanged; higher negative charge gives a larger ion.
Section 6 of 10Electron affinity
Electron affinity
Electron affinity the energy change when one mole of electrons is added to one mole of gaseous atoms, for example Cl(g) + e− → Cl−(g).
The process is exothermic, so the value is negative.
Group 17 elements have the most negative values because they become stable 1− ions by completing the p subshell.
It generally becomes less negative down a group as distance and shielding increase, with fluorine an exception because of repulsion in its small 2p subshell.
Section 7 of 10Electronegativity
Electronegativity
Electronegativity the ability of an atom to attract the shared pair in a covalent bond.
It increases across a period and decreases down a group.
It does not apply to noble gases, which do not form covalent bonds.
Section 8 of 10Valence electrons and ion charges
1
Valence electrons are in the outermost shell and determine reactivity and bonding.
2
Across a period, ignoring transition metals, the count rises by one.
3
Atoms form ions to reach the configuration of the nearest noble gas.
4
Ionic bonds form between a metal and a nonmetal, and covalent bonds between two nonmetals.
| Group | Typical ion | Example |
|---|---|---|
| 1 | 1+ | Na+ |
| 2 | 2+ | Mg2+ |
| 13 | 3+ | Al3+ |
| 15 | 3− | N3− |
| 16 | 2− | O2− |
| 17 | 1− | F− |
- Group 14 elements tend not to form ions.
Section 9 of 10Key terms
| Term | Meaning |
|---|---|
| Periodicity | Recurring, predictable trends in properties across the periodic table. |
| Effective nuclear charge | Net attraction felt by outer electrons after shielding. |
| Ionisation energy | Energy to remove one mole of electrons from gaseous atoms. |
| Electron affinity | Energy change when one mole of electrons is added to gaseous atoms. |
| Electronegativity | Ability of an atom to attract bonding electrons. |
| Valence electrons | Electrons in the outermost shell. |
Section 10 of 10Exam technique
1
Saying ionisation energy decreases across a period — it generally increases
2
Thinking cations are larger than their atoms — losing electrons makes them smaller
3
Forgetting shielding when explaining group trends
4
Applying electronegativity to noble gases
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