unit 3

⚛ Ions & Charges

An ion is an atom or group of atoms that has an electric charge.

• Electrons (−) = negative

• Protons (+) = positive

Atoms become ions when they gain or lose electrons:

• Lose electrons → positive ion (cation)

• Gain electrons → negative ion (anion)

⚛ Cation vs. Anion

• Cation = positive ion (+)

• Anion = negative ion (−)

⚛ Polyatomic Ions

A polyatomic ion is a group of atoms that act as one charged unit.

⚛ Transition Metals

Transition metals can have more than one possible charge and use Roman numerals in their names.

Examples:

• Fe²⁺ → Iron (II)

• Fe³⁺ → Iron (III)

• Au⁺ → Gold (I)

• Au³⁺ → Gold (III)

• Mn²⁺ → Manganese (II)

• Mn⁷⁺ → Manganese (VII)

Roman Numerals 1–10

I = 1

II = 2

III = 3

IV = 4

V = 5

VI = 6

VII = 7

VIII = 8

IX = 9

X = 10

⚛ Alkali & Alkaline Earth Metals

These metals have fixed charges:

• Lithium (Li⁺) → always +1

• Sodium (Na⁺) → always +1

• Calcium (Ca²⁺) → always +2

⚛ Chemical Bonds – Types

There are three main types of chemical bonds:

• Ionic → metal + nonmetal

• Covalent → nonmetal + nonmetal

• Metallic → metal + metal

⚛ Ionic Bonds

Ionic bonds happen when metals and nonmetals connect:

• Metals lose electrons → become cations

• Nonmetals gain electrons → become anions

• Opposite charges attract and bond

✅ Ionic = metal + nonmetal

⚛ Binary Ionic Formulas

Steps to write ionic formulas:

• Write element symbols with charges

• Balance charges (cross over numbers)

• Drop charges in final formula

Examples:

• Li⁺ + O²⁻ → Li₂O

• K⁺ + N³⁻ → K₃N

• Na⁺ + Cl⁻ → NaCl

• Al³⁺ + O²⁻ → Al₂O₃

⚛ Covalent Bonds

Covalent bonds form between nonmetals.

• Electrons are shared

• Usually between gases/nonmetals

Examples:

CO₂, H₂O, CH₄

⚛ Lewis Dot Structures

• Dots = valence electrons

• Only show outer electrons

Bond types:

• Single bond = 1 line (2 shared electrons)

• Double bond = 2 lines (4 shared electrons)

• Triple bond = 3 lines (6 shared electrons)

✏ Steps to Draw Lewis Structures

• Count total valence electrons

• Pick central atom (least electronegative, never H)

• Place single bonds

• Fill octets with lone pairs

• Make double/triple bonds if needed

⚛ SAN Method for Lewis Structures

• S = Sum valence electrons

• N = Needed electrons (8 per atom, H = 2)

• A = Available electrons

• S = Shared → (N − A) ÷ 2 = bonds

⚛ Naming Covalent Compounds

Uses prefixes:

PrefixNumbermono1di2tri3tetra4penta5hexa6hepta7octa8nona9deca10

Naming Rules:

• First element: prefix + name (no “mono” on first element)

• Second element: prefix + root + -ide

Examples:

• CO₂ → carbon dioxide

• N₂O₄ → dinitrogen tetroxide

• SF₆ → sulfur hexafluoride

⚛ Periodic Trends

Atomic Radius

• Across → decreases

• Down → increases

Ionization Energy

• Across → increases

• Down → decreases

Electronegativity

• Across → increases

• Down → decreases

• Highest = Fluorine (4.0)

Electron Affinity

• Across → increases

• Down → decreases

✅ Quick Trend Chart

TrendAcrossDownAtomic Radius⬇⬆Ionization Energy⬆⬇Electronegativity⬆⬇Electron Affinity⬆⬇

⚛ Lewis Structure Example

N₂

Total electrons = 10

:N≡N: