Prepare for AP Chemistry Topic 1.8 with this free practice quiz. Test your understanding of valence electrons, ionic charges, and chemical formulas to maximize your AP score.
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Valence Electrons
Valence electrons are electrons in the atom’s outermost occupied principal energy level. They are the electrons most directly involved in chemical bonding and largely determine an element’s chemical behavior.
Main-group elements
For AP Chemistry, the fastest way to determine the number of valence electrons is from the group number on the periodic table:
| Group | Typical valence electrons | Common behavior |
|---|---|---|
| 1 | 1 | Usually loses 1 e⁻ |
| 2 | 2 | Usually loses 2 e⁻ |
| 13 | 3 | Often loses 3 e⁻ |
| 14 | 4 | Often shares electrons |
| 15 | 5 | Often gains or shares 3 e⁻ |
| 16 | 6 | Usually gains 2 e⁻ |
| 17 | 7 | Usually gains 1 e⁻ |
| 18 | 8* | Generally unreactive |
*Helium has 2 valence electrons because its first energy level is full with 2 electrons.
Why valence electrons matter
Atoms tend to undergo chemical changes that produce more stable electron arrangements.
For many main-group atoms, stability is associated with a filled valence shell:
• First shell: maximum 2 electrons
• Second shell: maximum 8 electrons
• Third shell: commonly treated as having an octet for introductory bonding predictions
This leads to the octet rule:
Main-group atoms often gain, lose, or share electrons to obtain 8 valence electrons.
The octet rule is a useful model, not an absolute law. AP Chemistry later requires recognition of exceptions when working with Lewis structures and molecular bonding.
Electron configurations and valence electrons
Example: sodium
Na: 1s² 2s² 2p⁶ 3s¹
The highest occupied principal energy level is n = 3, containing one electron.
Therefore:
Na has 1 valence electron.
Example: chlorine
Cl: 1s² 2s² 2p⁶ 3s² 3p⁵
The n = 3 shell contains:
2 + 5 = 7 electrons
Therefore:
Cl has 7 valence electrons.
Valence Shells
The valence shell is the highest occupied principal energy level of an atom.
For a main-group atom, the valence electrons occupy this outer shell.
Principal energy levels
Principal energy levels are represented by n:
n = 1, 2, 3, 4, …
Example:
Mg: 1s² 2s² 2p⁶ 3s²
The highest occupied shell is n = 3.
Therefore:
• Valence shell = n = 3
• Valence electrons = 2
Relationship to periodic trends
Across a period:
• The number of occupied principal energy levels generally remains the same.
• The number of valence electrons increases.
• Effective nuclear charge generally increases.
• Atomic radius generally decreases.
Down a group:
• The number of occupied principal energy levels increases.
• The number of electron shells increases.
• Atomic radius generally increases.
• The number of valence electrons remains similar for main-group elements.
These patterns help explain why elements in the same group have similar chemical properties.
Stable valence-shell configurations
Noble gases have especially stable electron configurations.
Example:
Ne: 1s² 2s² 2p⁶
Neon has 8 valence electrons, so its outer shell is full.
This helps explain the common ion charges of main-group elements:
• Group 1 metals tend to lose 1 electron.
• Group 2 metals tend to lose 2 electrons.
• Group 17 nonmetals tend to gain 1 electron.
• Group 16 nonmetals tend to gain 2 electrons.
Ionic Compounds
An ionic compound consists of positively charged ions, called cations, and negatively charged ions, called anions. They are held together by electrostatic attractions.
A typical ionic compound forms when electrons are transferred from a metal to a nonmetal.
Example:
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
The resulting ions attract:
Na⁺ + Cl⁻ → NaCl
The electron is transferred from sodium to chlorine.
Cations
A cation is a positively charged ion.
Cations form when atoms lose electrons.
Atom → cation + electrons
Example:
Mg → Mg²⁺ + 2e⁻
Because electrons are negatively charged, losing electrons makes the species more positive.
Anions
An anion is a negatively charged ion.
Anions form when atoms gain electrons.
Atom + electrons → anion
Example:
O + 2e⁻ → O²⁻
Ionic compounds must be electrically neutral
The total positive charge must equal the total negative charge.
Therefore:
Total charge = 0
Example:
Mg²⁺ and O²⁻ combine in a 1:1 ratio:
Mg²⁺ + O²⁻ → MgO
For aluminum oxide:
Al³⁺ and O²⁻
The smallest combination with zero total charge is:
2 Al³⁺ + 3 O²⁻
Total charge:
2(+3) + 3(−2) = 0
Therefore:
Al₂O₃
Ionic Charges
The charge of a monatomic ion can often be predicted from the element’s group.
Common main-group ion charges
| Group | Typical ion | Reason |
|---|---|---|
| 1 | +1 | Loses 1 e⁻ |
| 2 | +2 | Loses 2 e⁻ |
| 13 | +3 | Often loses 3 e⁻ |
| 15 | −3 | Gains 3 e⁻ |
| 16 | −2 | Gains 2 e⁻ |
| 17 | −1 | Gains 1 e⁻ |
| 18 | 0 | Already has a stable valence shell |
Group 14 does not have one simple predictable ionic charge. Elements such as carbon and silicon commonly form covalent compounds rather than simple monatomic ions.
Transition-metal charges
Transition metals commonly have multiple possible ionic charges.
Examples:
Fe²⁺ and Fe³⁺
Cu⁺ and Cu²⁺
Therefore, their charges usually cannot be predicted simply from their periodic-table group.
The Roman numeral in a compound’s name specifies the charge of the metal.
Examples:
• iron(II) = Fe²⁺
• iron(III) = Fe³⁺
• copper(I) = Cu⁺
• copper(II) = Cu²⁺
Example: iron(III) oxide
Fe³⁺ and O²⁻ must combine in a ratio that produces zero total charge.
2 Fe³⁺ + 3 O²⁻ → Fe₂O₃
Therefore:
iron(III) oxide = Fe₂O₃
Common fixed-charge metals
Important ions to know:
• Group 1 metals → +1
• Group 2 metals → +2
• Al³⁺ → +3
• Zn²⁺ → +2
• Cd²⁺ → +2
• Ag⁺ → +1
These metals generally have predictable charges in ionic compounds.
Predicting Chemical Formulas
The AP Chemistry skill is to use ion charges and charge neutrality to construct the correct formula.
Step 1: Identify the ions
Determine the cation and anion and their charges.
Example: calcium chloride
Calcium:
Ca²⁺
Chloride:
Cl⁻
Step 2: Balance the charges
One Ca²⁺ requires two Cl⁻ ions.
Ca²⁺ + 2 Cl⁻ → CaCl₂
Total charge:
(+2) + 2(−1) = 0
Therefore:
CaCl₂
Step 3: Use the smallest whole-number ratio
The formula must contain the simplest whole-number ratio of ions.
Example:
Mg²⁺ and O²⁻
A 1:1 ratio gives:
MgO
Do not write Mg₂O₂ because the subscripts can be reduced.
Charge-balance method
For ions with charges:
Mᵃ⁺ and Xᵇ⁻
choose the smallest whole-number ratio that makes the total positive and negative charges equal.
Example:
Al³⁺ and O²⁻
The least common multiple of 3 and 2 is 6.
Two Al³⁺ ions give:
2(+3) = +6
Three O²⁻ ions give:
3(−2) = −6
Therefore:
Al₂O₃
Cross-over method
A quick method is to use the magnitudes of the charges as subscripts.
Example:
Al³⁺ and O²⁻
Cross the charge numbers:
Al₂O₃
Then check whether the subscripts can be reduced.
Example:
Ca²⁺ and O²⁻
Crossing would initially give Ca₂O₂, but the ratio reduces to:
CaO
The charge-balance method is conceptually safer because it emphasizes the fundamental requirement that the compound must be electrically neutral.
Ionic Compounds with Polyatomic Ions
A polyatomic ion is a group of covalently bonded atoms with an overall charge.
Important polyatomic ions for AP Chemistry include:
| Ion | Name |
|---|---|
| NH₄⁺ | ammonium |
| OH⁻ | hydroxide |
| NO₃⁻ | nitrate |
| NO₂⁻ | nitrite |
| SO₄²⁻ | sulfate |
| SO₃²⁻ | sulfite |
| CO₃²⁻ | carbonate |
| HCO₃⁻ | hydrogen carbonate / bicarbonate |
| PO₄³⁻ | phosphate |
| ClO₃⁻ | chlorate |
| ClO₄⁻ | perchlorate |
| CN⁻ | cyanide |
When balancing formulas, treat a polyatomic ion as one charged unit.
Calcium nitrate
Ions:
Ca²⁺
NO₃⁻
Two nitrate ions are required:
Ca²⁺ + 2 NO₃⁻ → Ca(NO₃)₂
The parentheses are necessary because the subscript 2 applies to the entire nitrate ion.
Therefore:
Ca(NO₃)₂
Aluminum sulfate
Ions:
Al³⁺
SO₄²⁻
Two Al³⁺ ions give +6.
Three SO₄²⁻ ions give −6.
Therefore:
Al₂(SO₄)₃
When parentheses are not needed
If only one polyatomic ion is present, parentheses are unnecessary.
Correct:
NaNO₃
Not:
Na(NO₃)
Predicting Formulas from Compound Names
Use this sequence:
Name → identify ions → determine charges → balance charges → write formula → check neutrality
Binary ionic compounds
Example: magnesium bromide
Mg²⁺ and Br⁻
Two Br⁻ ions are required to balance Mg²⁺.
Formula:
MgBr₂
Transition-metal ionic compounds
Use the Roman numeral to determine the metal charge.
Example: copper(II) chloride
Cu²⁺ and Cl⁻
Therefore:
CuCl₂
Example: iron(III) sulfide
Fe³⁺ and S²⁻
Two Fe³⁺ ions give +6.
Three S²⁻ ions give −6.
Therefore:
Fe₂S₃
Polyatomic ionic compounds
Identify the entire polyatomic ion before balancing.
Example: sodium phosphate
Na⁺ and PO₄³⁻
Three Na⁺ ions are required:
Na₃PO₄
Example: ammonium sulfate
NH₄⁺ and SO₄²⁻
Two NH₄⁺ ions are required:
(NH₄)₂SO₄
Formula Units vs. Molecules
Ionic compounds do not consist of discrete molecules in the same way that many covalent substances do.
Instead, ions form an extended ionic lattice.
Therefore, the term formula unit is used for ionic compounds.
For example:
NaCl
represents the simplest whole-number ratio of ions in the lattice:
1 Na⁺ : 1 Cl⁻
Similarly:
Al₂O₃
represents:
2 Al³⁺ : 3 O²⁻
The formula does not mean that an isolated Al₂O₃ molecule exists in the ionic solid.
High-Yield AP Chemistry Rules
Memorize the main charge patterns
Group 1 → +1
Group 2 → +2
Group 15 → −3
Group 16 → −2
Group 17 → −1
Also know these common fixed-charge ions:
Al³⁺
Zn²⁺
Ag⁺
Cd²⁺
Remember electron-transfer direction
Metals lose electrons → cations
Nonmetals gain electrons → anions
Loss of electrons → more positive
Gain of electrons → more negative
Always enforce charge neutrality
For every neutral ionic compound:
Total positive charge + total negative charge = 0
If the total charge is not zero, the formula is incorrect.
Do not write ionic charges in the final neutral formula
Correct:
CaCl₂
Incorrect:
Ca²⁺Cl₂⁻
The charges are used to determine the formula. They are not written in the formula of the neutral compound.
Use the simplest whole-number ratio
Correct:
NaCl
MgO
Al₂O₃
Incorrect:
Na₁Cl₁
Mg₂O₂
Al₄O₆
Distinguish ionic and covalent compounds
At the introductory AP Chemistry level:
• Metal + nonmetal → usually ionic
• Nonmetal + nonmetal → usually covalent/molecular
• Metal + polyatomic ion → ionic compound containing covalent bonds within the polyatomic ion
Examples:
NaCl → ionic
CO₂ → molecular/covalent
NaNO₃ → ionic overall; covalent bonds exist within NO₃⁻
AP Exam Quick Check
When asked to predict an ionic formula, use this exact process:
- Identify the cation.
- Identify the anion.
- Determine the charge of each ion.
- Balance total positive and negative charge.
- Use the smallest whole-number ratio.
- Use parentheses if more than one polyatomic ion is required.
- Check that the final compound has zero net charge.
Example: aluminum hydroxide
Al³⁺ and OH⁻
Three hydroxide ions are required:
Al³⁺ + 3 OH⁻ → Al(OH)₃
Therefore:
Al(OH)₃
Example: ammonium sulfate
NH₄⁺ and SO₄²⁻
Two ammonium ions are required:
2 NH₄⁺ + SO₄²⁻ → (NH₄)₂SO₄
Therefore:
(NH₄)₂SO₄
Core idea
The entire topic can be reduced to one logical chain:
Valence electrons → tendency to gain or lose electrons → ion charge → charge neutrality → simplest whole-number ratio → chemical formula.
For AP Chemistry, do not merely memorize formulas. Be able to explain why the formula has its particular subscripts using electron configuration, ion charge, and electrical neutrality.