AP® Chemistry review sheet from Aim for Five (aimforfive.com/chem/units/1/1-8)
Unit 1 · Topic 1.8
1.8 Valence Electrons and Ionic Compounds
Valence electrons decide how an element bonds and which ions it forms. Because elements in the same group have the same number of valence electrons, they form ions with the same charge and compounds with matching formulas.
Key terms
- valence electrons
- ion
- cation
- anion
- ionic compound
Valence electrons control bonding
Whether two elements bond, and how, depends on how their valence electrons and nuclei interact. A metal with a low ionization energy loses electrons easily. A nonmetal with a strong pull on electrons gains them easily. Put the two together and electrons transfer, forming ions that attract: an ionic compound.
Two nonmetals both hold their electrons tightly, so neither gives one up. Instead they share electrons in covalent bonds, which you'll study in Unit 2.
Predicting ion charges
Main-group atoms tend to gain or lose electrons until they have the same configuration as the nearest noble gas. Removing more electrons than that would mean pulling out core electrons, which takes far more energy (remember the jump in successive ionization energies).
You can see this in the data. The jump in successive ionization energies in topic 1.7 comes right after an atom loses its valence electrons, which is why Na stops at 1+ and Mg at 2+. Nonmetals on the right gain electrons until their outer s and p subshells are full: chlorine gains one, oxygen two, nitrogen three.
Transition metals often form more than one ion, such as Fe²⁺ and Fe³⁺, so you can't predict their charges from the group alone. The compound's name or formula will tell you.
| Group | Valence electrons | Typical ion | Examples |
|---|---|---|---|
| 1 | 1 | 1+ | Li⁺, Na⁺, K⁺ |
| 2 | 2 | 2+ | Mg²⁺, Ca²⁺ |
| 13 | 3 | 3+ (Al) | Al³⁺ |
| 15 | 5 | 3− | N³⁻, P³⁻ |
| 16 | 6 | 2− | O²⁻, S²⁻ |
| 17 | 7 | 1− | F⁻, Cl⁻, Br⁻ |
Writing ionic formulas
An ionic compound has zero net charge. Choose the smallest whole numbers of each ion that make the charges cancel. Calcium (2+) and nitrogen (3−) need 3 Ca²⁺ (+6) and 2 N³⁻ (−6), giving Ca₃N₂. Aluminum and oxygen give Al₂O₃.
Same group, similar compounds
Elements in the same group form analogous compounds. If you know sodium forms Na₂O, you can predict K₂O and Rb₂O. If MgCl₂ exists, so do CaCl₂ and SrCl₂. Exam questions use this to ask about unfamiliar elements: given the formula for one compound, predict the formula for a compound of a related element.
Reactivity trends
Group 1 metals get more reactive going down the group. Potassium reacts with water more violently than sodium, because its valence electron is farther from the nucleus and easier to remove.
Halogens get less reactive going down. Fluorine is the most reactive nonmetal: it's small, so an incoming electron can get very close to its nucleus and is strongly attracted.
Worked examples
Try each one yourself first, then open the solution.
- Example 1
Predicting a formula from group position
Gallium (Ga) is in group 13, below aluminum. Predict the formula of the compound formed between gallium and sulfur.
Show the solutionHide the solution
- Step 1: Ga is in group 13 like Al, so it forms Ga³⁺. Sulfur is in group 16, so it forms S²⁻.
- Step 2: The least common multiple of 3 and 2 is 6: two Ga³⁺ give +6 and three S²⁻ give −6.
- Step 3: Check by analogy: aluminum and sulfur form Al₂S₃, so gallium should match.
Answer: Ga₂S₃
- Example 2
Why Na⁺ and not Na²⁺? (classic trap)
Sodium's first two ionization energies are 496 and 4562 kJ/mol. Use them to explain why sodium forms Na⁺ in compounds and never Na²⁺.
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- Step 1: Removing the first electron (the 3s valence electron) takes 496 kJ/mol.
- Step 2: The second electron comes from the 2p subshell, a core shell much closer to the nucleus and much less shielded. It takes 4562 kJ/mol, about 9 times as much.
- Step 3: The energy released when ionic compounds form isn't nearly enough to pay for that second ionization, so sodium stops at 1+.
- Step 4: Saying 'Na⁺ is stable because it has an octet' isn't an explanation; the energy cost of reaching the core is.
Answer: Removing a second electron would mean pulling a core 2p electron, which costs about 9 times as much energy, so sodium forms only Na⁺.
Common mistakes
- Using the group number as the charge for nonmetals. Chlorine is in group 17 but forms Cl⁻, not 7+.
- Putting the numbers on the wrong atoms. Calcium nitride is Ca₃N₂: three Ca²⁺ (+6) balance two N³⁻ (−6). Ca₂N₃ would have a charge of +4 − 9 = −5.
- Assuming transition metals have one fixed charge. Many have more than one.
- Saying reactivity increases down every group. It increases for metals (group 1 and 2) but decreases for halogens.
On the exam
- Expect questions that give you a formula for one compound and ask for the formula of an analogous compound of another element in the same group.
- When asked why an element forms a particular ion, use ionization energies or Coulomb's law, not the octet rule alone.
Connected topics
Videos
Check yourself
4 questions on 1.8 Valence Electrons and Ionic Compounds. Pick an answer to see if you got it, and why.
Calcium reacts with phosphorus to form an ionic compound. What is the most likely formula of this compound?
Sodium and rubidium both react with chlorine to form white crystalline solids with the formula MCl. Which of the following best explains why the two compounds have the same formula?
Selenium is in the same group as oxygen. Which of the following is the most likely formula of the ionic compound formed between aluminum and selenium?
Which of the following is the most likely ion formed by barium, and why?
0 of 4 answered