AP Chemistry 1.7: Periodic trends

Khan Academy's current periodic-trends section covers Coulomb's law, atomic and ionic radii, ionization energy, successive ionization energies, electron affinity, and electronegativity.

The most important thing in AP Chemistry is explaining why trends happen.

The attraction between charged particles can be understood with Coulomb's law.

A simplified relationship is:

• F ∝ (q₁q₂)/r²

• where:

• q₁ and q₂ represent charges

• r represents distance between charges.

For atoms:

• positive nucleus attracts negative electrons.

Stronger attraction occurs when:

• nuclear charge is greater

• electron is closer to the nucleus

Weaker attraction occurs when:

• electron is farther away

• shielding reduces the attraction

Khan Academy directly connects periodic trends to Coulombic attraction.

More protons mean a larger positive nuclear charge.

• Example:

C has 6 protons.

N has 7.

O has 8.

• Moving left → right across a period increases the number of protons.

Inner electrons partially reduce the attraction between outer electrons and the nucleus.

• This effect is called shielding .

Core electrons lie between:

• the nucleus

• and

valence electrons.

• They partially "shield" the valence electrons from the full nuclear charge.

An electron experiences a net positive attraction called effective nuclear charge , often written:

• Zeff

A simplified idea is:

• Zeff ≈ nuclear charge − shielding

• This is not normally used as an exact AP Chemistry calculation. It is mainly a reasoning tool.

Across a period:

• number of protons increases

• but shielding does not increase as dramatically because electrons are entering the same general shell.

• So effective nuclear attraction generally increases.

Atomic radius describes atomic size.

Across a period:

• radius generally decreases left → right

• Why?

More protons are added.

Electrons are generally added to the same principal energy level.

The greater effective nuclear attraction pulls the electron cloud closer.

• Therefore:

• smaller radius

Moving downward:

• atomic radius increases

• Why?

Each new period adds another electron shell.

• Example:

• Li has outer electrons in n = 2.

• Na has outer electrons in n = 3.

• K has outer electrons in n = 4.

Additional shells mean:

valence electrons farther from nucleus

• increased shielding

• Therefore the atoms become larger.

Overall:

• largest → bottom-left

• smallest → top-right

Think:

Radius grows:

• ← and ↓

Forming ions changes atomic size.

Cations

Cations are generally smaller than their neutral atoms.

• Example:

• Na⁺ < Na

• Why?

Na loses its outer 3s electron.

Na:

• [Ne]3s¹

Na⁺:

• [Ne]

The entire third shell is no longer occupied.

• Therefore, radius decreases dramatically.

Anions are generally larger than their neutral atoms.

• Example:

• Cl⁻ > Cl

• Why?

Adding an electron increases:

• electron-electron repulsion

• while the number of protons stays the same.

The electron cloud spreads out more.

Isoelectronic species have the same number of electrons.

• Example:

• O²⁻

• F⁻

• Ne

• Na⁺

• Mg²⁺

Each contains 10 electrons .

If electron count is identical, compare protons.

• O²⁻: 8 protons

• F⁻: 9

• Ne: 10

• Na⁺: 11

• Mg²⁺: 12

More protons pull the same number of electrons more strongly.

• Therefore size decreases as proton number increases:

O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺

• This is extremely useful on AP questions.

Ionization energy is the energy required to remove an electron from a gaseous atom or ion.

• Example:

• Na(g) → Na⁺(g) + e⁻

Removing an electron requires overcoming the attraction between the electron and nucleus.

Generally:

• increases left → right

• Why?

Effective nuclear attraction increases.

Electrons are held more tightly.

• Therefore more energy is needed to remove one.

Generally:

• decreases going down

• Why?

Outer electrons are:

• farther from nucleus

• more shielded

• Therefore they are easier to remove.

Overall:

• higher toward:

• top-right

• lower toward:

• bottom-left

• So it moves opposite to atomic radius.

Radius:

• ← ↓

Ionization energy:

• → ↑

You should not treat periodic trends as perfect rules.

Some exceptions arise because:

• s and p subshells have different energies

• paired electrons repel each other

• especially stable configurations can occur

AP questions often care more about whether you can explain experimental values using electron structure than whether you memorized a simple arrow.

Ionization can happen repeatedly.

IE₁:

• removing first electron

IE₂:

• removing second

IE₃:

• removing third

• and so on.

Generally:

IE₁ < IE₂ < IE₃ < IE₄...

Each successive electron becomes harder to remove.

This is one of the most important Unit 1 skills.

• Imagine:

• IE₁ = 738

• IE₂ = 1451

• IE₃ = 7733

• IE₄ = 10543 kJ/mol

There is a huge jump:

• between IE₂ and IE₃

• This means:

• the first two electrons are relatively easy to remove.

After those two valence electrons are gone, the third electron would have to be removed from a much more stable inner shell.

• Therefore:

• the element has 2 valence electrons .

Khan Academy uses this exact style of reasoning in Unit 1.

Huge jump after IE₁:

• → 1 valence electron

Huge jump after IE₂:

• → 2 valence electrons

Huge jump after IE₃:

• → 3 valence electrons

For main-group elements, this can help identify the periodic-table group.

Electron affinity describes the energy change associated with adding an electron to a gaseous atom.

General idea:

• X(g) + e⁻ → X⁻(g)

Electron affinity trends are more complicated than radius or electronegativity.

In general, many atoms toward the right side of a period have a greater tendency to accept electrons, particularly halogens.

• However:

• do not memorize electron affinity as a perfectly smooth trend.

There are several exceptions.

Electronegativity measures how strongly an atom attracts shared electrons in a chemical bond.

Trend:

• left → right:

• increases

down a group:

• decreases

• Therefore:

• highest toward top-right

Fluorine is the most electronegative element.

Ionization energy

How difficult is it to remove an electron ?

Electron affinity

What is the energy change when an isolated gaseous atom gains an electron ?

Electronegativity

How strongly does a bonded atom pull shared electrons toward itself ?

• These are related but not interchangeable.