Valence Electrons and Ionic Compounds

Learning Objective 1.8.A Explain the relationship between trends in the reactivity of elements and periodicity.

Quick Notes

Valence (outermost) electrons control bonding and reactivity; periodic trends arise because elements in the same group have the same number of outer (valence) electrons.
Group number = number of valence electrons: Group 1 = 1, Group 2 = 2, etc.
Octet idea: Atoms tend to gain/lose electrons to reach a full outer shell of 8 electrons; transition metals are an exception.
Metals usually lose electrons forming positively charged ions, cations:
- Group 1: +1 (e.g. Na⁺)
- Group 2: +2 (e.g. Mg²⁺)
- Group 3: +3 (e.g. Al³⁺)
Non‑metals usually gain electrons forming negatively charged ions, anions:
- Group 5: −3 (e.g. N³⁻)
- Group 6: −2 (e.g. O²⁻)
- Group 7: −1 (e.g. Cl⁻)
Ionic compounds form when oppositely charged ions attract; formulas reflect charge balance:
- Na⁺ + Cl⁻ → NaCl (1:1)
- Mg²⁺ + O²⁻ → MgO (1:1)
- Ca²⁺ + 2F⁻ → CaF₂ (1:2)
Exam rule of thumb: To predict an ion’s charge, count valence electrons and think about how many must be lost or gained to reach a full shell.

Full Notes

Atoms form chemical bonds by gaining, losing, or sharing electrons. The type of bond and the ions formed depend on the valence electrons – the electrons in the outermost shell of an atom.

Valence Electrons and Reactivity

Valence (outermost) electrons are responsible for nearly all chemical behavior. The number of valence electrons an atom has can be predicted from its group number:

AP Chemistry periodic table labelled with groups and periods to show where valence electron counts come from

Group 1 (alkali metals): 1 valence electron
Group 2 (alkaline earth metals): 2 valence electrons
Group 17 (halogens): 7 valence electrons
Group 18 (noble gases): 8 valence electrons (except helium, which has 2)

The most stable configuration for most atoms is to have eight electrons in their outermost shell (octet rule). Note, transition metals are an exception to this (see Transition Metals). Atoms of elements can often lose or gain electrons to achieve a full outer shell of electrons. This forms charged ions with the charge of the ion based on the number of electrons lost or gained by the atom.

Metals (Groups 1, 2, 3) usually lose electrons to form positively charged ions (cations).
- Group 1 metals form ions with 1+ charge (e.g. Na⁺)
- Group 2 metals form ions with 2+ charge (e.g. Mg²⁺)
- Group 3 metals form ions with 3+ charge (e.g. Al³⁺)
Non-metals (Groups 5, 6, 7) usually gain electrons to form negatively charged ions (anions).
- Group 5 non-metals form ions with 3− charge (e.g. N³⁻)
- Group 6 non-metals form ions with 2− charge (e.g. O²⁻)
- Group 7 non-metals form ions with 1− charge (e.g. Cl⁻)

Examples of Ionic Compounds

Ions form in complementary pairs – when one atom loses electrons (forming a cation), another atom (or atoms) must gain those electrons (forming an anion). The resulting oppositely charged ions are then held together by strong electrostatic forces, forming an ionic compound.

a) Sodium chloride (NaCl)

AP Chemistry diagram showing sodium losing one electron to form Na plus, chlorine gaining one electron to form Cl minus, producing NaCl lattice

Sodium atom (Na) loses one electron → Na⁺
Chlorine atom (Cl) gains one electron → Cl⁻
Ions combine to form NaCl, held in a lattice.

b) Magnesium oxide (MgO)

AP Chemistry diagram showing magnesium losing two electrons to form Mg 2 plus and oxygen gaining two electrons to form O 2 minus, giving MgO

Mg loses two electrons → Mg²⁺
O gains two electrons → O²⁻
1:1 ratio gives the formula MgO.

c) Calcium fluoride (CaF₂)

AP Chemistry diagram showing calcium losing two electrons to form Ca 2 plus and two fluorine atoms each gaining one electron to form F minus, giving CaF2

Ca loses two electrons → Ca²⁺
Each F gains one electron → F⁻
Two F⁻ ions needed to balance one Ca²⁺ → formula is CaF₂.

Periodicity and Common Ion Charges

The position of an element on the periodic table helps predict the charge of the ion it will form:

AP Chemistry periodic table annotated with common ion charges by group for main-group elements

Group 14 (carbon group) often shares electrons instead of forming ions, though it can vary.

Transition metals are in the middle block of the periodic table and can form ions with multiple charges (e.g. Fe²⁺ and Fe³⁺) due to their electronic configurations.

Reactivity Trends and Periodicity

Metals (especially in groups 1 and 2) are highly reactive because they lose electrons easily.
Nonmetals (especially in group 17) are highly reactive because they gain electrons easily.
Group trends: elements in the same group tend to form similar compounds because they have the same number of valence electrons.

Examples:

Group 1 metals (Li, Na, K) all form +1 ions and react with halogens to form salts (e.g. NaCl, KBr).
Group 17 nonmetals (F, Cl, Br) all form −1 ions and react with metals to form similar halide compounds.

Worked Example

Why does calcium (Ca) tend to form a Ca²⁺ ion?

Answer: Calcium is in group 2, so it has 2 valence electrons. To achieve a stable configuration (like argon), it must lose both outer electrons, forming a Ca²⁺ ion.

Matt’s exam tip

Always relate ion formation back to electron configuration and periodic position. In multiple-choice questions, if you're unsure of an ion's charge, count the valence electrons and think about how many need to be lost or gained to reach a full shell.

Summary

The reactivity of elements and the formation of ions are directly related to their valence electrons and position on the periodic table. Periodicity allows us to predict common ion charges, the types of compounds elements will form, and how readily they will gain or lose electrons.
Key points to remember:
- Valence electrons determine how atoms bond
- Metals → cations; Nonmetals → anions
- Elements in the same group form similar compounds
- Periodic trends allow us to predict ionic charges and reactivity