Difficulty: Intermediate | Prerequisites: Chapters 5-7 (electron configurations, valence electrons, periodic trends)
Source: Chem 1 Final Study Guide, General Chemistry (Purdue University, Chhabra)
Tags: ionic bonding, covalent bonding, metallic bonding, Lewis structures, VSEPR, molecular geometry, hybridisation, polarity, chemical reactions, balancing equations, synthesis, decomposition, single replacement, double replacement, combustion, mole, Avogadro's number, molar mass, percent composition, empirical formula, molecular formula, stoichiometry, limiting reactant, percent yield
These five chapters cover why and how atoms join together, the types of chemical reactions they undergo, and the quantitative maths behind those reactions. Chapter 8 handles ionic bonds, Chapter 9 covers covalent bonds and molecular shape, Chapter 10 introduces reaction types, Chapter 11 is all about the mole concept, and Chapter 12 brings it together with stoichiometry. This is the quantitative heart of general chemistry. If you can master mole conversions and stoichiometry here, every later calculation-heavy topic (solutions, gas laws, thermochemistry, equilibrium) will be far more manageable.
Atoms bond by transferring electrons (ionic) or sharing electrons (covalent) to achieve stable electron configurations. Molecular shapes follow from VSEPR theory. Chemical reactions fall into five main types. The mole connects the atomic scale to the lab scale, and stoichiometry uses balanced equations to predict how much product a reaction yields.
Chemical bond
The force that holds two atoms together. In simple terms, the "glue" between atoms.
Octet rule
Elements tend to react to achieve the stable electron configuration of a noble gas (8 valence electrons).
Cation
A positively charged ion, formed when an atom loses one or more electrons.
Anion
A negatively charged ion, formed when an atom gains one or more electrons. Named with the suffix "-ide."
Ionic bond
The electrostatic force of attraction between oppositely charged ions in an ionic compound.
Lattice energy
The energy required to separate one mole of ions in an ionic compound. The more negative the lattice energy, the stronger the ionic bond. Smaller ions and higher charges produce more negative lattice energies.
Formula unit
The simplest ratio of ions represented in an ionic compound. The overall charge of a formula unit is zero.
Oxidation number
The charge of a monatomic ion; indicates how many electrons an atom has gained or lost.
Electrolyte
An ionic compound whose aqueous solution conducts electric current, because the dissolved ions carry charge.
Metallic bond
The attraction between metallic cations and the surrounding "sea" of delocalised electrons.
Delocalized electrons
Valence electrons in a metallic solid that are free to move throughout the entire structure, not bound to any one atom.
Alloy
A mixture of elements with metallic properties. Its properties differ from those of the individual component metals.
Covalent bond
A bond formed by the sharing of valence electrons between two atoms, typically nonmetals.
Lewis structure
A diagram showing how valence electrons are arranged among atoms in a molecule, using dots for lone pairs and lines for bonds.
Sigma bond (σ)
A single covalent bond formed by the head-on overlap of atomic orbitals.
Pi bond (π)
A bond formed when parallel p orbitals overlap sideways. Found in double and triple bonds (a double bond = 1σ + 1π; a triple bond = 1σ + 2π).
Bond length
The distance between two bonded nuclei. Shorter bond length = stronger bond.
Bond dissociation energy
The energy required to break a specific covalent bond. Always a positive value.
Diatomic molecules
Seven elements that exist naturally as pairs of atoms: Br₂, I₂, N₂, Cl₂, H₂, O₂, F₂ (mnemonic: BrINClHOF).
Resonance structures
Two or more valid Lewis structures for the same molecule or ion. The actual structure is an average (hybrid) of all resonance forms.
VSEPR model
Valence Shell Electron Pair Repulsion model. Predicts molecular shape by arranging electron pairs as far apart as possible.
Hybridisation
The mixing of atomic orbitals to form new, identical hybrid orbitals. The number of bonding zones around the central atom determines the hybridisation: 2 = sp, 3 = sp², 4 = sp³, 5 = sp³d, 6 = sp³d².
Polar covalent bond
A covalent bond with unequal sharing of electrons due to a difference in electronegativity between the atoms.
Reactants
The starting substances in a chemical reaction (left side of the equation).
Products
The substances formed during a chemical reaction (right side of the equation).
Skeleton equation
An unbalanced chemical equation showing only the chemical formulas of reactants and products.
Synthesis reaction
Two or more substances combine to form a single product: A + B → AB.
Decomposition reaction
A single compound breaks down into two or more simpler substances: AB → A + B.
Combustion reaction
A substance reacts with oxygen, releasing energy, water, and CO₂ (or CO).
Single-replacement reaction
An element replaces another element in a compound: A + BC → AC + B. Driven by the activity series.
Double-replacement reaction
The ions of two compounds exchange partners: AB + CD → AD + CB. Produces a precipitate, gas, or water.
Precipitate
A solid that forms during a reaction in solution. Will not form if the compound contains a group 1A element, ammonium, or is a nitrate.
Mole
The SI unit for amount of substance. One mole = 6.022 x 10²³ particles (Avogadro's number).
Avogadro's number
6.022 x 10²³, the number of atoms in exactly 12 g of carbon-12.
Molar mass
The mass in grams of one mole of a substance, numerically equal to the atomic or formula mass in amu.
Percent composition
The percentage by mass of each element in a compound: (mass of element in 1 mol compound / molar mass of compound) x 100.
Empirical formula
The simplest whole-number ratio of elements in a compound.
Molecular formula
The actual number of atoms of each element in one molecule. Found by multiplying the empirical formula by (experimental molar mass / empirical formula mass).
Hydrate
A compound with a specific number of water molecules bound to its structure (e.g. CuSO₄·5H₂O).
Stoichiometry
The study of quantitative relationships between reactants and products in a balanced chemical equation.
Mole ratio
The ratio of moles of any two substances in a balanced equation, derived directly from the coefficients.
Limiting reactant
The reactant that is completely consumed first, determining the maximum amount of product formed.
Excess reactant
The reactant that remains after the limiting reactant is used up.
Theoretical yield
The maximum amount of product that can be formed from a given amount of reactant, calculated from stoichiometry.
Actual yield
The amount of product actually obtained in an experiment.
Percent yield
(Actual yield / theoretical yield) x 100.
Forming ions
Metals lose electrons to form cations (positive); nonmetals gain electrons to form anions (negative)
The goal is to achieve a noble gas electron configuration
Transition metals commonly form 2+ ions (losing 2 s electrons) but can also form 3+ or higher by losing d electrons
Zinc and cadmium are always 2+; silver is always 1+
Ionic bond formation
Ionic bonds form between metals and nonmetals
Electrons lost by the metal are gained by the nonmetal
Ionic compounds form crystal lattices; a 1:1 ratio gives a cubic crystal
Ionic compounds are electrolytes when dissolved in water
Lattice energy
More negative = stronger bond
Affected by ion size (smaller = stronger) and charge (higher = stronger)
Writing formulas
Cation first, anion second
Charges must balance to zero
For polyatomic ions, the charge applies to the entire group
Naming ionic compounds
Cation name + anion name (with "-ide" suffix for monatomic anions)
Transition metals: include Roman numeral for the charge, e.g. iron(III) chloride
Polyatomic ions: use their standard names (oxyanions: "-ate" = more oxygen, "-ite" = less)
Metallic bonding
Electron sea model: metal atoms contribute valence electrons to a shared "sea"
Metallic bonds = attraction between cations and delocalised electrons
Explains conductivity, malleability, and lustre
Covalent bonds
Formed by sharing electrons between nonmetals
BrINClHOF elements exist as diatomic molecules
Single bond = 1 shared pair (σ bond); double = 2 pairs (1σ + 1π); triple = 3 pairs (1σ + 2π)
Bond strength: triple > double > single; bond length: triple < double < single
Drawing Lewis structures
Identify the central atom (least electronegative, never hydrogen)
Count total valence electrons
Divide by 2 to get bonding pairs
Place one pair between central and each terminal atom
Distribute remaining pairs as lone pairs or multiple bonds
Ensure all atoms satisfy the octet rule (where possible)
Resonance
Some molecules have multiple valid Lewis structures
The true structure is a hybrid (average) of all resonance forms
Octet rule exceptions
Free radicals: odd-electron species (e.g. NO)
Boron compounds: B can form only 3 bonds (6 electrons around B)
Expanded octets: elements in period 3+ can use d orbitals (e.g. PCl₅ with 5 bonds)
VSEPR and molecular geometry
2 bonding zones, 0 lone pairs: linear (180°)
3 bonding zones, 0 lone pairs: trigonal planar (120°)
4 bonding zones, 0 lone pairs: tetrahedral (109.5°), sp³
3 bonding zones, 1 lone pair: trigonal pyramidal (107.3°), sp³
2 bonding zones, 2 lone pairs: bent (104.5°), sp³
5 bonding zones, 0 lone pairs: trigonal bipyramidal (90°/120°), sp³d
6 bonding zones, 0 lone pairs: octahedral (90°), sp³d²
Hybridisation
Determined by the number of bonding zones (electron groups) around the central atom
2 zones = sp, 3 = sp², 4 = sp³, 5 = sp³d, 6 = sp³d²
Polar vs nonpolar molecules
Polar covalent bonds arise from electronegativity differences
A molecule with polar bonds can still be nonpolar if the bond dipoles cancel symmetrically (e.g. CO₂, CH₄)
Purely nonpolar bonds exist only between identical atoms
Naming binary molecular compounds
First element: full name; second element: root + "-ide"
Use prefixes (mono-, di-, tri-, tetra-, penta-, hexa-) to indicate atom counts
Naming acids
Anion ends in "-ide": hydro____ic acid (e.g. HCl = hydrochloric acid)
Anion ends in "-ite": ____ous acid (e.g. HNO₂ = nitrous acid)
Anion ends in "-ate": ____ic acid (e.g. HNO₃ = nitric acid)
Balancing equations
Write the skeleton equation
Count atoms on each side
Adjust coefficients to equalise each element
Use lowest whole-number ratios
Verify the balanced equation
Five reaction types
Synthesis: A + B → AB
Decomposition: AB → A + B
Combustion: substance + O₂ → CO₂ + H₂O + energy
Single replacement: A + BC → AC + B (use the activity series)
Double replacement: AB + CD → AD + CB (produces precipitate, gas, or water)
Precipitate rules (solubility shortcuts)
No precipitate forms if the product contains a group 1A cation, ammonium, or nitrate
The mole concept
1 mole = 6.022 x 10²³ particles (atoms, molecules, formula units)
Molar mass: the mass in grams of one mole, equal to atomic/formula mass in amu
Conversions: mass ↔ moles (divide or multiply by molar mass), moles ↔ particles (multiply or divide by Avogadro's number)
Molar mass of compounds
Add the molar masses of each element, multiplied by the number of atoms of that element in the formula
Percent composition
% by mass = (mass of element in 1 mol compound / molar mass of compound) x 100
Empirical formula from percent composition
Assume 100 g sample (percentages become grams)
Convert each mass to moles
Divide each by the smallest mole value
Round to nearest whole numbers (or multiply to eliminate fractions)
Molecular formula
Molecular formula = empirical formula x (experimental molar mass / empirical formula mass)
Hydrates
Named with Greek prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, deca-
To find the formula: subtract the anhydrous mass from the hydrated mass to get the mass of water, convert both to moles, divide moles of water by moles of anhydrous compound
Core principle
Stoichiometry is based on the law of conservation of mass
Coefficients in a balanced equation give mole ratios
Stoichiometric conversions
Mole-to-mole: use the mole ratio directly
Mole-to-mass: find moles of the unknown using the ratio, then multiply by molar mass
Mass-to-mass: convert given mass to moles, use ratio, convert back to mass
Limiting reactants
Write the balanced equation
Convert each reactant mass to moles
Determine which reactant runs out first by comparing the actual mole ratio to the required ratio
The limiting reactant determines the maximum product
The excess reactant is what is left over
Percent yield
Percent yield = (actual yield / theoretical yield) x 100
Theoretical yield comes from stoichiometric calculations; actual yield comes from the lab
Percent composition: (mass of element in 1 mol / molar mass of compound) x 100
Molecular formula: empirical formula x (experimental mass / empirical mass)
Percent yield: (actual yield / theoretical yield) x 100
Mole conversions: mass ÷ molar mass = moles; moles x 6.022 x 10²³ = particles
Stoichiometry is how industrial chemists calculate how much raw material they need and how much product they can expect. Every pharmaceutical, fertiliser, and polymer is manufactured using these calculations.
Percent yield tells chemical engineers how efficient a process is. A low yield means wasted reagents and higher costs, which is why catalysts and optimised conditions matter commercially.
Ionic bonding explains why table salt dissolves in water and conducts electricity: the ions separate and move freely in solution, carrying charge.
Students sometimes think that a molecule with polar bonds must be a polar molecule. That is only true if the bond dipoles do not cancel. CO₂ has polar C=O bonds, but the linear shape makes them cancel, so the molecule is nonpolar.
A frequent error in stoichiometry: using the molar mass of one substance as the mole ratio. The mole ratio comes from the balanced equation coefficients, not from molar masses.
Students often confuse empirical and molecular formulas. CH₂O (empirical) and C₆H₁₂O₆ (molecular, glucose) have the same elemental ratio but very different actual compositions.
In limiting reactant problems, students sometimes assume the substance present in the smallest mass is limiting. Mass does not determine the limiting reactant; the mole ratio does.
⚠️ Lewis structures and VSEPR geometry questions are on virtually every exam. Practise drawing structures and predicting shapes until it is fast.
⚠️ Balancing equations is assumed skill from this point forward. Every stoichiometry problem, every equilibrium problem, every electrochemistry problem starts with a balanced equation.
⚠️ Limiting reactant and percent yield problems are staples of the final exam. Work through several multi-step examples.
⚠️ Know the five reaction types and be able to predict products, especially for double-replacement (precipitate or no precipitate).
⚠️ Acid naming rules (-ide → hydro___ic acid, -ite → ___ous acid, -ate → ___ic acid) are a common quick test question.
True or False: A double bond consists of one sigma bond and one pi bond.
Fill in the blank: The mnemonic for diatomic elements is __________.
True or False: In a balanced equation, the coefficients represent mass ratios.
Fill in the blank: The empirical formula of a compound with 40% carbon, 6.7% hydrogen, and 53.3% oxygen is __________.
True or False: Percent yield can exceed 100% in a properly conducted experiment.
Q: What is the molar mass of Ca(OH)₂?
A: Ca = 40.08, O = 16.00 x 2, H = 1.008 x 2. Total = 40.08 + 32.00 + 2.016 = 74.10 g/mol.
Q: If 10.0 g of hydrogen reacts with 80.0 g of oxygen to form water, which is the limiting reactant?
A: 2H₂ + O₂ → 2H₂O. Moles H₂ = 10.0 / 2.016 = 4.96 mol. Moles O₂ = 80.0 / 32.00 = 2.50 mol. The ratio requires 2 mol H₂ per 1 mol O₂, so 4.96 mol H₂ needs 2.48 mol O₂. We have 2.50 mol O₂, which is just enough. H₂ is the limiting reactant (it would need slightly more O₂ if the ratio were exact, but it runs out first at scale).
Q: A student obtains 15.0 g of product in a reaction whose theoretical yield is 20.0 g. What is the percent yield?
A: (15.0 / 20.0) x 100 = 75.0%.
Q: What is the molecular geometry of water (H₂O)?
A: Bent. Oxygen has 4 electron groups (2 bonding pairs, 2 lone pairs), giving a tetrahedral electron geometry but a bent molecular shape with a bond angle of about 104.5°.
Q: Name the compound N₂O₅.
A: Dinitrogen pentoxide.
Ionic and covalent bonding directly determine intermolecular forces (Ch. 13), which control physical properties like boiling point and solubility.
Stoichiometry is the computational backbone of gas law calculations (Ch. 14), solution concentration (Ch. 15), and thermochemistry (Ch. 16).
Reaction types reappear in acids and bases (Ch. 19, neutralisation is a double-replacement reaction) and redox chemistry (Ch. 20-21).
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