AP® Chemistry review sheet from Aim for Five (aimforfive.com/chem/units/1/1-7)
Unit 1 · Topic 1.7
1.7 Periodic Trends
Atomic radius, ionization energy, electron affinity and electronegativity all follow patterns across the periodic table. Two ideas explain every pattern: effective nuclear charge rises across a period, and outer electrons sit in higher, more shielded shells as you go down a group.
Key terms
- atomic radius
- ionization energy
- electronegativity
- electron affinity
- effective nuclear charge
- shielding
Why the periodic table repeats
Elements in a group have the same valence configuration: every alkali metal ends in s¹ and every halogen ends in s² p⁵. That's why their properties repeat. Full shells and subshells, as in the noble gases, are especially hard to add to or remove from.
Explain every trend with two factors from Coulomb's law: how much charge the valence electrons feel (effective nuclear charge) and how far they are from the nucleus (which shell they're in).
The four trends
Across a period, protons are added but electrons go into the same shell. Electrons in the same shell shield each other poorly, so effective nuclear charge rises. The valence electrons are pulled closer and held more tightly: radius shrinks, and ionization energy and electronegativity rise.
Down a group, each element adds a new shell. Valence electrons are farther out and shielded by more core electrons, so they're held less tightly: radius grows and ionization energy falls.
Electronegativity is how strongly an atom attracts shared electrons in a bond. Fluorine has the highest value. Electron affinity is the energy change when a neutral atom gains an electron; halogens release the most energy. Electron affinity has more exceptions than the other trends (chlorine releases more energy than fluorine, for example), so exam questions about it stay qualitative.
| Property | Across a period (left to right) | Down a group |
|---|---|---|
| Atomic radius | decreases | increases |
| First ionization energy | generally increases | decreases |
| Electronegativity | increases | decreases |
| Energy released on gaining an electron (electron affinity) | generally increases, with exceptions | generally decreases, with exceptions |
Two ionization energy exceptions
- Group 2 to group 13 (Mg 738 → Al 578 kJ/mol): aluminum's outer electron is in a 3p subshell, which is higher in energy and slightly more shielded than magnesium's 3s, so it's easier to remove.
- Group 15 to group 16 (P 1012 → S 1000 kJ/mol): sulfur's fourth 3p electron has to share an orbital with another electron. The repulsion between the paired electrons makes it a little easier to remove.
Ionic radii and successive ionization energies
A cation is smaller than its atom: it often loses its whole outer shell, and the remaining electrons feel more pull. An anion is larger than its atom: the added electrons repel each other while the nuclear charge stays the same.
In an isoelectronic series (ions with the same number of electrons), more protons means a smaller ion: O²⁻ > F⁻ > Na⁺ > Mg²⁺, all with 10 electrons.
Successive ionization energies rise steadily, then jump sharply when you start removing core electrons. The jump tells you how many valence electrons the atom has.
Predicting values
Trends let you estimate a missing value. If you know the atomic radii of the elements above and below an element in its group, its radius will fall between them.
Worked examples
Try each one yourself first, then open the solution.
- Example 1
Successive ionization energies
An element in period 3 has these ionization energies (kJ/mol): IE₁ = 738, IE₂ = 1451, IE₃ = 7733, IE₄ = 10,543. Identify the element and explain.
Show the solutionHide the solution
- Step 1: Look for the big jump. IE₂ → IE₃ goes from 1451 to 7733, more than five times larger.
- Step 2: The jump means the third electron comes from a core shell, which is much closer to the nucleus and far less shielded.
- Step 3: So the atom has 2 valence electrons and is in group 2. In period 3, that's magnesium.
Answer: Magnesium. The large jump after IE₂ shows two valence electrons; the third electron is a core (n = 2) electron.
- Example 2
Explaining an exception (classic trap)
Aluminum has a lower first ionization energy than magnesium even though Al has more protons. Explain.
Show the solutionHide the solution
- Step 1: Configurations: Mg is [Ne] 3s²; Al is [Ne] 3s² 3p¹.
- Step 2: The electron removed from Al is in the 3p subshell, which is higher in energy than 3s and partly shielded by the 3s electrons.
- Step 3: The extra proton in Al doesn't make up for this, so less energy is needed to remove Al's 3p electron than Mg's 3s electron.
- Step 4: An answer that just says 'ionization energy increases across a period' predicts the wrong order.
Answer: Al's outermost electron is in a higher-energy 3p subshell, shielded by the 3s electrons, so it is removed more easily than a 3s electron from Mg.
- Example 3
Ranking ion sizes
Rank Mg²⁺, Na⁺, F⁻ and O²⁻ from largest to smallest radius.
Show the solutionHide the solution
- Step 1: All four have 10 electrons (1s² 2s² 2p⁶), so they're isoelectronic and the shielding is about the same.
- Step 2: Protons: O has 8, F 9, Na 11, Mg 12. More protons pull the same 10 electrons in closer.
- Step 3: So the order from largest to smallest follows fewest to most protons.
Answer: O²⁻ > F⁻ > Na⁺ > Mg²⁺
Common mistakes
- Saying atoms get smaller across a period because they 'have more electrons'. The cause is more protons with electrons in the same shell, so effective nuclear charge rises.
- Explaining with 'stability' or 'wanting an octet'. Use Coulomb's law: charge and distance.
- Forgetting the IE exceptions at groups 2→13 and 15→16.
- Confusing electronegativity (a bonded atom's pull on shared electrons) with electron affinity (energy change when a lone atom gains an electron).
On the exam
- Free-response questions often say 'explain in terms of atomic structure'. A full answer names both atoms' relevant features (protons, shells or shielding) and states the effect on attraction. Comparing just one atom loses points.
- Expect data tables of ionization energies or radii where you identify an element or predict a missing value.
Connected topics
Videos
Check yourself
4 questions on 1.7 Periodic Trends. Pick an answer to see if you got it, and why.
Mg²⁺, Na⁺, Ne and F⁻ all have 10 electrons. Which of these species has the largest radius?
The first ionization energy of potassium is less than that of sodium. Which of the following best accounts for this difference?
Fluorine is more electronegative than chlorine. Which of the following best explains this?
Which of the following lists the elements in order of increasing first ionization energy?
0 of 4 answered