Difficulty: Intermediate | Prerequisites: Electron configuration, valence electrons, periodic table layout
Once you know how electrons are arranged in atoms, the next question is: how do atoms stick together? This topic covers the types of chemical bonds, how to draw Lewis structures, how to predict molecular shape using VSEPR theory, and how to name compounds. It also covers molar mass calculations and the law of conservation of mass, which underpin all of stoichiometry. Bonding and geometry are central to the second third of CHM 11500 and form the basis for understanding intermolecular forces, polarity, and reactivity.
Ionic bonds transfer electrons; covalent bonds share them. Lewis structures and VSEPR theory predict molecular geometry. Nomenclature follows straightforward naming rules for ionic and covalent compounds. Molar mass is the sum of atomic masses in a formula.
Ionic bond
A bond formed by the complete transfer of one or more electrons from a metal atom to a nonmetal atom. The resulting cation and anion are held together by electrostatic attraction.
In simple terms, one atom gives electrons to another, and the opposite charges keep them stuck together.
Covalent bond
A bond formed by the sharing of electron pairs between two nonmetal atoms. In a nonpolar covalent bond, electrons are shared equally; in a polar covalent bond, they are shared unequally.
Metallic bond
A bond in which metal atoms share a "sea" of delocalised electrons. This accounts for electrical conductivity and malleability in metals.
Lewis structure (Lewis dot structure)
A diagram showing the bonding between atoms in a molecule and the lone pairs of electrons that may exist. Each line represents a shared pair; dots represent lone pairs.
Octet rule
Most main-group atoms tend to form bonds until they are surrounded by eight valence electrons (or two for hydrogen). There are exceptions, notably for elements in period 3 and beyond that can expand their octet.
VSEPR theory (Valence Shell Electron Pair Repulsion)
A model used to predict the geometry of a molecule based on the number of electron groups (bonding and lone pairs) around the central atom. Electron groups arrange themselves to be as far apart as possible.
Electron geometry
The arrangement of all electron groups (bonding pairs and lone pairs) around a central atom.
Molecular geometry
The arrangement of only the atoms (not lone pairs) in a molecule. Molecular geometry and electron geometry are the same when there are no lone pairs on the central atom.
Lattice energy
The energy released when gaseous ions come together to form one mole of an ionic solid. It increases with higher ion charges and smaller ion sizes.
In simple terms, lattice energy measures how strongly the ions in a crystal are held together.
Molar mass
The mass of one mole of a substance, in g/mol. For a compound, it is the sum of the atomic masses of every atom in the formula.
Law of conservation of mass
In a chemical reaction, matter is neither created nor destroyed. The total mass of reactants equals the total mass of products, and a balanced equation reflects this.
Ionic bonds form between metals and nonmetals. Electrons transfer completely. Example: NaCl.
Covalent bonds form between nonmetals. Electrons are shared. Example: H₂O, CO₂.
Metallic bonds form between metal atoms. Delocalised electrons are shared across a lattice. Example: Cu, Fe.
The complete transfer of electrons from a metal to a nonmetal defines an ionic bond.
Count the total valence electrons. For polyatomic ions, add electrons for negative charges and subtract for positive charges.
Place the least electronegative atom in the centre (usually not hydrogen or fluorine).
Draw single bonds from the central atom to each surrounding atom.
Distribute remaining electrons as lone pairs, starting with the outer atoms, to satisfy octets.
If the central atom lacks an octet, convert lone pairs on surrounding atoms into double or triple bonds.
Valence electrons: N has 5, each O has 6, plus 1 for the negative charge. Total = 5 + 3(6) + 1 = 24 electrons.
N is the central atom. Draw three N-O single bonds (uses 6 electrons, leaving 18).
Distribute remaining 18 electrons as lone pairs on the three oxygens (6 each gives each O an octet, using all 18).
Check N: it only has 6 electrons (3 single bonds). Convert one lone pair from one O into a double bond. Now N has 8 electrons.
The structure has one N=O double bond and two N-O single bonds, with lone pairs on all three oxygens. Resonance structures exist where the double bond can be on any of the three oxygens.
Molecular geometry: trigonal planar (3 bonding groups, 0 lone pairs on N). The octet rule is satisfied for all atoms.
For a central atom with no lone pairs:
2 electron groups: linear (180°). Example: CO₂.
3 electron groups: trigonal planar (120°). Example: BF₃.
4 electron groups: tetrahedral (109.5°). Example: CH₄, CBr₄.
When lone pairs are present on the central atom, the electron geometry stays the same but the molecular geometry changes:
4 electron groups, 1 lone pair: trigonal pyramidal. Example: NH₃.
4 electron groups, 2 lone pairs: bent. Example: H₂O, H₂S.
2 electron groups, 0 lone pairs: linear. Example: CO₂, CS₂.
CBr₄: 4 bonding pairs, 0 lone pairs on C. Electron geometry = tetrahedral. Molecular geometry = tetrahedral. These match because there are no lone pairs.
CS₂: C is the central atom with 2 double bonds to S and no lone pairs. Electron geometry = linear. Molecular geometry = linear.
H₂S: S has 2 bonding pairs and 2 lone pairs. Electron geometry = tetrahedral. Molecular geometry = bent.
CO₂: 2 double bonds, no lone pairs on C. Electron geometry = linear. Molecular geometry = linear.
To count the total number of ions in one formula unit of an ionic compound, count every distinct ion produced:
NaOH: Na⁺ + OH⁻ = 2 ions per formula unit.
MgCl₂: Mg²⁺ + 2 Cl⁻ = 3 ions per formula unit.
Al₂O₃: 2 Al³⁺ + 3 O²⁻ = 5 ions per formula unit.
Mg₃(PO₄)₂: 3 Mg²⁺ + 2 PO₄³⁻ = 5 ions per formula unit.
MgCl₂ contains exactly three ions per formula unit.
Binary covalent compounds (two nonmetals): use Greek prefixes for both elements, and the second element gets an -ide ending.
N₂O₅: dinitrogen pentoxide
NCl₃: nitrogen trichloride
Common prefixes: mono- (1, usually dropped for the first element), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6).
Ionic compounds: name the cation first, then the anion with an -ide ending (for monatomic anions) or the polyatomic ion name.
Add the atomic masses of every atom in the formula.
Worked example: Aluminum carbonate, Al₂(CO₃)₃
Al: 2 x 26.98 = 53.96
C: 3 x 12.01 = 36.03
O: 9 x 16.00 = 144.00
Total = 53.96 + 36.03 + 144.00 = 233.99 g/mol
In any chemical reaction, the total mass of the reactants equals the total mass of the products. This is why chemical equations must be balanced: the number of atoms of each element must be the same on both sides of the arrow.
Example: 2H₂ + O₂ → 2H₂O. Four hydrogen atoms and two oxygen atoms appear on each side.
Lattice energy is governed by Coulomb's law: it is directly proportional to the product of the ion charges and inversely proportional to the sum of the ionic radii.
NaCl: charges of +1 and -1, with Na⁺ (95 pm) and Cl⁻ (181 pm). Lattice energy is approximately 787 kJ/mol.
MgO: charges of +2 and -2, with Mg²⁺ (65 pm) and O²⁻ (140 pm). Lattice energy is approximately 3850 kJ/mol.
MgO has a much larger lattice energy because both the charges are doubled (the charge product goes from 1 to 4) and the ions are smaller, so the distance between their centres is shorter. Higher lattice energy means a higher melting point and a harder, more stable crystal.
Molar mass = Σ (number of atoms of each element x atomic mass of that element)
Lattice energy ∝ (q₊ x q₋) / (r₊ + r₋) (Coulomb's law relationship)
Lewis structures and molecular geometry determine whether a molecule is polar, which in turn determines its solubility, boiling point, and biological activity. Drug molecules must have the right shape to fit into receptor proteins, which is why VSEPR geometry matters far beyond the exam. Lattice energy explains why ceramics (ionic compounds with high lattice energies like MgO and Al₂O₃) are used in heat-resistant applications, from furnace linings to engine components.
Students often confuse electron geometry with molecular geometry. They are only the same when there are no lone pairs on the central atom. H₂S has tetrahedral electron geometry but bent molecular geometry.
Students sometimes place an atom other than the least electronegative in the centre of a Lewis structure. Hydrogen is never a central atom.
When counting ions per formula unit, students sometimes count each atom individually rather than each ion. In Mg₃(PO₄)₂, PO₄³⁻ is one polyatomic ion, not five separate ions.
Students forget to add an electron when drawing the Lewis structure of an anion. For NO₃⁻, you must include 24 valence electrons, not 23.
⚠️ Identifying the correct molecular geometry of CO₂ (linear) and H₂S (bent) are standard exam questions. Know the common shapes by heart.
⚠️ Lewis structure questions for polyatomic ions (especially NO₃⁻, CO₃²⁻, SO₄²⁻) are commonly tested. Remember to account for the charge when counting valence electrons.
⚠️ Nomenclature questions require you to distinguish between ionic naming (no prefixes) and covalent naming (Greek prefixes). N₂O₅ is "dinitrogen pentoxide," not "nitrogen oxide."
⚠️ Molar mass calculations require careful attention to subscripts, especially in compounds with polyatomic ions inside parentheses like Al₂(CO₃)₃.
⚠️ Lattice energy comparisons are tested conceptually: higher charges and smaller ions mean greater lattice energy.
True or false: CO₂ has a bent molecular geometry. False (it is linear).
Fill in the blank: The total number of valence electrons in NO₃⁻ is ___. 24.
True or false: MgCl₂ produces 3 ions per formula unit when dissolved. True (Mg²⁺ + 2 Cl⁻).
Fill in the blank: The name of NCl₃ is ___. Nitrogen trichloride.
True or false: MgO has a higher lattice energy than NaCl. True.
Q: What type of bond involves the complete transfer of electrons from a metal to a nonmetal?
A: An ionic bond.
Q: What is the molecular geometry of CO₂?
A: Linear.
Q: Calculate the molar mass of Al₂(CO₃)₃, given Al = 26.98, C = 12.01, O = 16.00.
A: 2(26.98) + 3(12.01) + 9(16.00) = 53.96 + 36.03 + 144.00 = 233.99 g/mol.
Q: Name the compound N₂O₅.
A: Dinitrogen pentoxide.
Q: In the Lewis structure of NO₃⁻, what is the molecular geometry and are all octets satisfied?
A: Trigonal planar geometry. All octets are satisfied (one N=O double bond and two N-O single bonds in one resonance structure, with resonance delocalising the double bond across all three positions).
Q: Explain why MgO has a much higher lattice energy than NaCl.
A: MgO has ions with charges of +2 and -2 (vs +1 and -1 in NaCl), and the ions are smaller. Both factors increase lattice energy according to Coulomb's law.
The type of bond (ionic vs covalent) determines physical properties like melting point, conductivity, and solubility, which connect to the solutions and gas laws material. Molecular geometry connects to polarity, which determines intermolecular forces (not covered in depth in this exam but crucial for CHM 11600). Molar mass is the bridge between grams and moles, which underpins every stoichiometry and solution concentration problem in the course. Lattice energy connects back to periodic trends (ion size) and forward to thermochemistry (energy changes in forming ionic compounds).
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