13.2 Chemical Bonding and Compounds
Key Takeaways
- Ionic bonds form by electron transfer between metals and nonmetals; covalent bonds form by electron sharing between nonmetals.
- Valence electrons and the octet rule predict ion charges and simple molecular formulas at SHS depth.
- Electronegativity differences help classify bonds: large difference → ionic character; small difference → covalent.
- Ionic compounds form crystal lattices with high melting points; many covalent molecular substances are gases, liquids, or low-melting solids.
- Empirical formulas show simplest whole-number ratios; molecular formulas show actual atom counts in a molecule.
13.2 Chemical Bonding and Compounds
Why This Matters for the USTET
After atomic structure, USTET chemistry moves to how atoms stick together. Expect items that ask whether a bond is ionic or covalent, what charge a main-group ion should carry, which formula is correct for a binary compound, or which property matches salt-like solids versus molecular substances. These questions reward the same valence-electron counting you practiced with the periodic table.
Why Atoms Bond
Atoms bond to reach a more stable electron arrangement — often a noble-gas configuration. At SHS depth, that usually means an octet of valence electrons (duet for hydrogen). Bonding happens in two primary exam-relevant ways:
- Transfer of electrons → ionic bonding
- Sharing of electrons → covalent bonding
A third category, metallic bonding, describes metal lattices with delocalized electrons ("sea of electrons"), explaining conductivity and malleability. USTET items mention metallic bonding less often than ionic/covalent contrasts, but you should recognize that pure metals are not held by ionic salts or discrete covalent molecules.
Ionic Bonding
Ionic bonds form when electrons transfer from a metal (low ionization energy) to a nonmetal (high electronegativity). The metal becomes a cation; the nonmetal becomes an anion. Electrostatic attraction between opposite charges holds the ions in a crystal lattice.
Classic example — sodium chloride:
- Na (2-8-1) loses 1 e⁻ → Na⁺
- Cl (2-8-7) gains 1 e⁻ → Cl⁻
- Formula unit: NaCl (1:1 ratio so charges cancel)
Charge balance is the formula rule: total positive charge equals total negative charge.
| Cation | Anion | Compound formula | Reasoning |
|---|---|---|---|
| Na⁺ | Cl⁻ | NaCl | 1(+1) with 1(−1) |
| Mg²⁺ | O²⁻ | MgO | 1(+2) with 1(−2) |
| Mg²⁺ | Cl⁻ | MgCl₂ | Need two Cl⁻ to cancel +2 |
| Al³⁺ | O²⁻ | Al₂O₃ | Lowest whole numbers: 2(+3) and 3(−2) |
| Ca²⁺ | NO₃⁻ | Ca(NO₃)₂ | Polyatomic ion kept in parentheses when subscript > 1 |
Common main-group ion charges (memorize for speed):
| Group | Typical ion | Examples |
|---|---|---|
| 1 | +1 | Li⁺, Na⁺, K⁺ |
| 2 | +2 | Mg²⁺, Ca²⁺, Ba²⁺ |
| 13 (Al) | +3 | Al³⁺ |
| 16 | −2 | O²⁻, S²⁻ |
| 17 | −1 | F⁻, Cl⁻, Br⁻, I⁻ |
Polyatomic ions appear often enough that a short list pays off: NO₃⁻ (nitrate), SO₄²⁻ (sulfate), CO₃²⁻ (carbonate), OH⁻ (hydroxide), NH₄⁺ (ammonium), PO₄³⁻ (phosphate). Ammonium is the common polyatomic cation on entrance exams.
Properties of ionic compounds (exam checklist):
- High melting and boiling points (strong lattice attractions)
- Hard, brittle crystalline solids at room temperature for many salts
- Conduct electricity when molten or dissolved in water (mobile ions), but not as dry solids (ions locked in place)
- Often soluble in water (with important exceptions — know the idea, not every solubility rule)
Covalent Bonding
Covalent bonds form when nonmetal atoms share electron pairs. Shared pairs count toward each atom's octet. Single, double, and triple bonds share one, two, or three pairs.
| Molecule | Shared pairs (between central atoms) | Sketch idea |
|---|---|---|
| H₂ | 1 (single) | H–H |
| O₂ | 2 (double) | O=O |
| N₂ | 3 (triple) | N≡N |
| H₂O | O shares with two H atoms | Bent molecule |
| CO₂ | C double-bonded to each O | Linear O=C=O |
| CH₄ | C single-bonded to four H | Tetrahedral |
Lewis structures (SHS level): count valence electrons, place a skeleton with single bonds, then add lone pairs / multiple bonds to satisfy octets. Hydrogen is terminal and wants two electrons total. Carbon typically forms four bonds; nitrogen three; oxygen two; halogens one in many simple molecules.
Molecular vs empirical formulas:
- Molecular formula — actual number of atoms in one molecule (H₂O₂, C₆H₁₂O₆).
- Empirical formula — simplest whole-number ratio (HO for hydrogen peroxide; CH₂O for glucose).
Ionic compounds are written as empirical formula units (NaCl, MgCl₂) because there is no single "molecule" of NaCl in the lattice sense taught at this level.
Electronegativity and Bond Polarity
Electronegativity is an atom's pull on shared electrons. When two atoms share unequally, the bond is polar covalent; equal (or nearly equal) sharing gives a nonpolar covalent bond. A very large electronegativity difference usually means electron transfer and ionic character.
Useful Grade 11 comparisons (not a full Pauling table):
- Nonpolar covalent examples: H₂, Cl₂, O₂, N₂ (identical atoms); C–H is often treated as essentially nonpolar in intro courses.
- Polar covalent examples: H–Cl, H–O in water, C–O, N–H.
- Ionic examples: NaCl, MgO, CaF₂ (metal + nonmetal with large ΔEN).
Molecular polarity depends on both bond polarity and shape. CO₂ has polar C=O bonds but is a nonpolar molecule because the linear shape cancels dipoles. H₂O is a polar molecule because the bent shape does not cancel the O–H dipoles — which is why water is an excellent solvent for many ionic compounds.
Naming Binary Compounds (Quick Rules)
Ionic (metal + nonmetal): name the metal, then the nonmetal with an -ide ending: NaCl sodium chloride; MgO magnesium oxide; Al₂O₃ aluminum oxide. For metals with variable charge (Fe, Cu), stock notation uses Roman numerals: FeCl₂ iron(II) chloride; FeCl₃ iron(III) chloride.
Covalent (two nonmetals): use prefixes mono-, di-, tri-, tetra-, penta-, hexa- on the second element always, and on the first when the count is not one: CO carbon monoxide; CO₂ carbon dioxide; N₂O₅ dinitrogen pentoxide. The mono- prefix is usually omitted on the first element (CO, not monocarbon monoxide).
Ionic vs Covalent Property Contrast
| Property | Typical ionic | Typical covalent molecular |
|---|---|---|
| Particles | Ions in lattice | Discrete molecules |
| Melting point | High | Often low |
| Electrical conductivity | Molten/aqueous yes; solid no | Usually poor (unless acid solutions, etc.) |
| Examples | NaCl, CaCO₃, MgSO₄ | H₂O, CO₂, CH₄, C₁₂H₂₂O₁₁ |
| Formation partners | Metal + nonmetal | Nonmetal + nonmetal |
Network covalent solids (diamond, SiO₂ quartz) are an important exception: covalent bonding throughout a giant network → very high melting points. If a stem says "covalent but extremely hard with very high melting point," think network solid, not sugar or methane.
Worked Mini-Problems
- Predict the formula of the compound between Al and Br. Al³⁺ and Br⁻ → AlBr₃.
- Is the bond in F₂ ionic or covalent? Covalent (identical nonmetals sharing electrons); nonpolar.
- Why does solid NaCl not conduct, but salt water does? In the solid, ions are fixed; in solution, Na⁺ and Cl⁻ are free to move and carry current.
- Empirical formula of C₂H₆? Divide by 2 → CH₃.
Exam Strategy
- Classify partners first: metal+nonmetal → lean ionic; nonmetal+nonmetal → lean covalent.
- Write ions with charges, then crisscross magnitudes to get subscripts (then simplify if needed).
- For conductivity questions, ask whether charged particles can move.
- Watch for CO₂ polarity traps: polar bonds do not automatically mean a polar molecule.
Bonding is the bridge from atoms to the formulas and reactions that dominate the rest of USTET chemistry.
Which pair is most likely to form an ionic compound?
What is the correct formula for the ionic compound formed from Mg²⁺ and Cl⁻?
Why does an aqueous solution of NaCl conduct electricity while solid NaCl typically does not?
Carbon dioxide has polar carbon–oxygen bonds yet is often classified as a nonpolar molecule. Why?