SCH4U · Exam review

Chemistry Exam Review

A practice exam with questions from all 5 units, then the key facts from every lesson, unit by unit. For a deeper look at one unit, open its test review.

Unit test reviews

Practice exam

35 questions, the same number from each unit. Each one shows the lesson it comes from.

Unit 1: Hydrocarbons: Naming and Isomers
1
What is the IUPAC name of $C H_{3} CH (C H_{3}) C H_{2} C H_{2} C H_{3}$ ?
Unit 1: Functional Groups and Their Properties
2
Which functional group is present in $C H_{3} C H_{2} CHO$ ?
Unit 1: Stereoisomers and Chirality
3
Which compound contains a chiral carbon?
Unit 1: Organic Reactions
4
Classify the reaction $C H_{4} + C l_{2} \to C H_{3} Cl + HCl$ (in UV light).
Unit 1: Organic Reactions
5
Methyl ethanoate, $C H_{3} COOC H_{3}$ , is hydrolyzed with an acid catalyst. What are the products?
Unit 1: Polymers and Organic Compounds in Society
6
After a lab, a student has 5 mL of leftover hexane. What should they do with it?
Unit 1: Organic Compounds in Everyday Life: Risks, Benefits and Choices
7
What property of DDT and PCBs, besides persistence, allows them to be stored in animal tissue and biomagnify? (two or three words)
Unit 2: From Rutherford to the Quantum Atom
8
Which gold foil observation led most directly to the idea of a small, dense, positive nucleus?
Unit 2: From Rutherford to the Quantum Atom
9
Which statement describes a limitation of the Bohr model?
Unit 2: Electron Configurations
10
Match each element to the block of the periodic table it belongs to.
Calcium (Ca)
Iron (Fe)
Bromine (Br)
Uranium (U)
Unit 2: Molecular Shape and Polarity
11
Match each molecule to its molecular shape.
$CH_{4}$
$PCl_{5}$
$SF_{4}$
$ClF_{3}$
$BrF_{5}$
Unit 2: Intermolecular Forces and Types of Solids
12
Arrange the halogens from lowest to highest boiling point.
1. 1 $Cl_{2}$
2. 2 $I_{2}$
3. 3 $Br_{2}$
4. 4 $F_{2}$
Unit 2: Seeing Structure: Spectroscopy and Instrumental Analysis
13
An IR spectrum shows a strong, broad absorption centred near 3350 cm⁻¹ and no peak near 1700 cm⁻¹. Which compound fits best?
Unit 2: Seeing Structure: Spectroscopy and Instrumental Analysis
14
What medical imaging technique applies NMR to hydrogen nuclei in the body? (abbreviation or full name)
Unit 3: Enthalpy Changes and Calorimetry
15
A solid dissolves in water and the temperature of the solution drops from $22.0^{\circ} C$ to $17.5^{\circ} C$ . Which statement is correct?
Unit 3: Enthalpy Changes and Calorimetry
16
A student's measured molar enthalpy of combustion of a fuel is $- 720 kJ/mol$ , but the accepted value is $- 1367 kJ/mol$ . Which is the most likely cause?
Unit 3: Hess's Law and Enthalpies of Formation
17
Which equation is the correct formation equation for carbon dioxide?
Unit 3: Physical, Chemical and Nuclear Energy: Fuels Compared
18
What is the role of heavy water in a CANDU reactor?
Unit 3: Rates of Reaction and Collision Theory
19
In $2 NO (g) + O_{2} (g) \to 2 NO_{2} (g)$ , oxygen is consumed at $0.012 mol/ (L \cdot s)$ . At what rate is $NO_{2}$ formed?
mol/(L·s)
Unit 3: Potential Energy Diagrams and Mechanisms
20
A reaction has $E_{a} (forward) = 65 kJ$ and $Δ H = + 25 kJ$ . What is $E_{a} (reverse)$ in kJ?
kJ
Unit 3: Potential Energy Diagrams and Mechanisms
21
Put the steps for analyzing a proposed reaction mechanism in order.
1. 1Label intermediates, catalysts and the slow step
2. 2Check that the result is the overall balanced equation
3. 3Cancel species that appear on both sides
4. 4Write out each elementary step
5. 5Add all the steps together
Unit 4: Dynamic Equilibrium and the Equilibrium Constant
22
Which is the correct equilibrium expression for $CaCO_{3} (s) ⇌ CaO (s) + CO_{2} (g)$ ?
Unit 4: Le Chatelier's Principle and the Reaction Quotient
23
For $2 SO_{2} (g) + O_{2} (g) ⇌ 2 SO_{3} (g),$ $K = 40.0$ . A mixture has $[SO_{3}] = 0.500$ , $[SO_{2}] = 0.100$ and $[O_{2}] = 0.200 mol/L$ . What happens?
Unit 4: Acid-Base Equilibria
24
In $HCO_{3}^{-} + OH^{-} ⇌ CO_{3}^{2 -} + H_{2} O,$ which species is the conjugate base of $HCO_{3}^{-}$ ?
Unit 4: Solubility Equilibria, Buffers and Titration
25
A slightly soluble salt AB has a molar solubility of $2.0 \times 1 0^{- 5} mol/L$ . Calculate its Ksp (2 significant figures, for example 3.4e-10).
Unit 4: Solubility Equilibria, Buffers and Titration
26
Match each term to its description.
Buffer
Common ion effect
Equivalence point
Endpoint
Unit 4: Equilibrium at Work: Industry, Body and Environment
27
Which ion's concentration falls in seawater as it absorbs CO₂, making it harder for corals and shellfish to build calcium carbonate? (name or formula)
Unit 4: Measuring K in the Lab
28
What is the name of the law $A = ε b c$ ?
Unit 5: Oxidation Numbers and Redox Reactions
29
What is the oxidation number of chlorine in the chlorate ion, $ClO_{3}^{-}$ ? Include the sign.
Unit 5: Oxidation Numbers and Redox Reactions
30
In the balanced equation $MnO_{4}^{-} + 5 Fe^{2 +} + 8 H^{+} \to Mn^{2 +} + 5 Fe^{3 +} + 4 H_{2} O,$ how many electrons are transferred in total?
Unit 5: Balancing Redox with Half-Reactions
31
A student balances a reaction in acidic solution and gets $\dots + 6 H^{+} \to \dots$ . To convert it to basic solution, what should they do next?
Unit 5: Galvanic Cells and Cell Potential
32
Using $E^{\circ} (Ag^{+} / Ag) = + 0.80 V$ and $E^{\circ} (Mg^{2 +} / Mg) = - 2.37 V$ , calculate $E_{cell}^{\circ}$ for the spontaneous magnesium-silver cell. Give your answer to 2 decimal places.
V
Unit 5: Galvanic Cells and Cell Potential
33
At which electrode of a galvanic cell does reduction occur? (one word)
Unit 5: Electrolytic Cells, Corrosion and Batteries
34
Copper is refined using a current of 1.50 A for 1.00 h. What mass of copper, in grams, is deposited at the cathode? Use $F = 96485 C/mol$ and $M (Cu) = 63.55 g/mol$ . Give 3 significant figures.
g
Unit 5: Electrolytic Cells, Corrosion and Batteries
35
Match each technology to the electrochemical idea behind it.
Magnesium rod in a water heater tank
Recharging a lithium-ion battery
Hydrogen fuel cell
Refining copper
Painting a steel bridge

Key facts by unit

Unit 1: Organic Chemistry

Hydrocarbons: Naming and Isomers

Alkanes $C_{n} H_{2n + 2}$ (saturated); alkenes $C_{n} H_{2n}$ ; alkynes $C_{n} H_{2n - 2}$ ; cycloalkanes $C_{n} H_{2n}$ .
Name = branch locants and names (alphabetical) + parent prefix + ending, e.g. 3-ethyl-2,2-dimethylpentane.
The parent chain is the longest continuous chain, and it must contain any double or triple bond, which gets the lowest number.
Structural isomers share a molecular formula but differ in connections; more branching means weaker London forces and a lower boiling point.
Cis-trans isomers need restricted rotation (C=C or a ring) and two different groups on each double-bond carbon.
Benzene, $C_{6} H_{6}$ , has delocalized ring electrons, so it is more stable and less reactive than an alkene.

Functional Groups and Their Properties

Endings: alcohol -ol, aldehyde -al, ketone -one, carboxylic acid -oic acid, ester alkyl -oate, amine -amine, amide -amide; ethers use alkoxy-.
Aldehydes, carboxylic acids and amides always have the functional-group carbon as carbon 1, so they need no locant for it.
Hydrogen bonding needs H bonded directly to O or N; ethers, aldehydes, ketones and esters cannot hydrogen bond with each other.
At similar molar mass, boiling point rises: alkane < ether < ester, aldehyde, ketone < alcohol < carboxylic acid < amide.
Solubility in water falls as the nonpolar carbon chain gets longer.

Stereoisomers and Chirality

Isomers: structural (chain, position, functional group) vs stereoisomers (cis-trans and enantiomers).
A chiral carbon is bonded to four different groups; a molecule with one chiral carbon has a non-superimposable mirror image (an enantiomer).
Enantiomers have identical physical and chemical properties except for rotating plane-polarized light in opposite directions and interacting differently with chiral molecules.
A racemic mixture is 50:50 and shows no optical rotation. Specific rotation: $[α] = \frac{α}{l \times c}$ .
Receptors and enzymes are chiral, so enantiomers can smell different (spearmint vs caraway carvone) or act differently as drugs (thalidomide, ibuprofen).

Organic Reactions

Addition: atoms add across C=C or C≡C to give one saturated product; with HX or H₂O, H adds to the carbon that already has more H (Markovnikov's rule).
Substitution: one atom or group replaces another, e.g. $C H_{4} + C l_{2} \to C H_{3} Cl + HCl$ in UV light; benzene substitutes rather than adds.
Elimination: removing $H_{2} O$ (from an alcohol) or HX (from a haloalkane) from neighbouring carbons forms a C=C.
Oxidation: primary alcohol → aldehyde → carboxylic acid; secondary alcohol → ketone; tertiary alcohol → no reaction.
Carboxylic acid + alcohol ⇌ ester + water (acid catalyst, heat); hydrolysis reverses it.
Complete combustion of a hydrocarbon gives $C O_{2}$ and $H_{2} O$ ; too little oxygen also gives CO and soot.

Polymers and Organic Compounds in Society

Addition polymers form from monomers with C=C; nothing is released, so the polymer has the monomer's empirical formula.
Condensation polymers form from monomers with two functional groups; each link releases a small molecule, usually $H_{2} O$ .
Polyester: diacid + diol (ester links). Polyamide: diacid + diamine (amide links). Proteins are natural polyamides of amino acids joined by peptide bonds.
Number of repeating units: $n = \frac{M ( polymer )}{M ( repeating unit )} .$
A chain of $n$ amino acids has $n - 1$ peptide bonds, so $n - 1$ water molecules are released.
WHMIS 2015: read the label and 16-section SDS; handle volatile solvents in a fume hood away from ignition sources, and dispose of them as hazardous waste.

Organic Compounds in Everyday Life: Risks, Benefits and Choices

Aspirin is made by esterification: salicylic acid + ethanoic anhydride → acetylsalicylic acid + ethanoic acid. Moisture hydrolyzes it back.
Hydrogenation is an addition reaction: $H_{2}$ adds across C=C. Partial hydrogenation creates trans fats, which raise LDL and lower HDL.
Small esters give fruity flavours (ethyl butanoate: pineapple; 3-methylbutyl ethanoate: banana).
Like dissolves like: mostly nonpolar vitamins (A, D, E, K) are fat-soluble and stored; vitamins with many –OH groups (C, B) are water-soluble and excreted.
Persistent, nonpolar, chlorinated compounds (DDT, PCBs) dissolve in fat and biomagnify; polar compounds move with water.
Soap = fatty acid salt (forms scum with $C a^{2 +}$ and $M g^{2 +}$ ); detergents use sulfonate heads and work in hard water.

Unit 2: Structure and Properties of Matter

From Rutherford to the Quantum Atom

Gold foil: most alpha particles pass through (atom is mostly empty space); a few deflect sharply (tiny, dense, positive nucleus).
Photon energy: $E = h f = \frac{h c}{λ}$ , with $h = 6.63 \times 1 0^{- 34} J \cdot s$ and $c = 3.00 \times 1 0^{8} m/s$ .
Bohr hydrogen levels: $E_{n} = - \frac{2.18 \times 1 0 ^{- 18} J}{n ^{2}}$ ; an emitted photon carries $∣ E_{final} - E_{initial} ∣$ .
Line spectra prove energy is quantized: drops to $n = 1$ give ultraviolet (Lyman), to $n = 2$ give visible (Balmer), to $n = 3$ give infrared (Paschen).
An orbital is a region of high probability of finding an electron, holding at most 2 electrons; s orbitals are spherical, p orbitals are dumbbell-shaped.

Electron Configurations

Filling order: $1 s 2 s 2 p 3 s 3 p 4 s 3 d 4 p 5 s 4 d 5 p 6 s 4 f 5 d 6 p .$
Capacities: s holds 2, p holds 6, d holds 10, f holds 14; shell $n$ holds at most $2 n^{2}$ electrons.
Pauli: at most 2 electrons per orbital, with opposite spins. Hund: fill equal-energy orbitals singly with parallel spins before pairing.
Exceptions: $Cr = [Ar] 4 s^{1} 3 d^{5}$ and $Cu = [Ar] 4 s^{1} 3 d^{10}$ .
Transition metals lose their 4s electrons before their 3d electrons when forming cations, e.g. $Fe^{3 +} = [Ar] 3 d^{5}$ .

Molecular Shape and Polarity

Electron groups (lone pairs plus bonded atoms, with multiple bonds counting once) spread out as far apart as possible.
2 groups linear (180°); 3 trigonal planar (120°); 4 tetrahedral (109.5°); 5 trigonal bipyramidal (90°, 120°); 6 octahedral (90°).
Lone pairs repel more than bonding pairs and shrink bond angles: $CH_{4}$ 109.5°, $NH_{3}$ about 107°, $H_{2} O$ about 104.5°.
Expanded octets: 5 groups with 1, 2, 3 lone pairs give seesaw, T-shaped, linear; 6 groups with 1, 2 lone pairs give square pyramidal, square planar.
A molecule is polar only if it has polar bonds and its shape does not let the bond dipoles cancel.

Intermolecular Forces and Types of Solids

Strength of intermolecular forces for molecules of similar size: London dispersion < dipole-dipole < hydrogen bonding, all far weaker than chemical bonds.
Dispersion forces act between all molecules and increase with the number of electrons (molar mass) and with surface contact.
Hydrogen bonding needs H bonded directly to N, O or F, attracted to a lone pair on N, O or F of another molecule.
Ionic: high melting point, brittle, conducts only molten or dissolved. Molecular: low melting point, non-conducting. Covalent network: very high melting point, very hard. Metallic: conducts as a solid, malleable and ductile.
Metal + nonmetal gives an ionic solid; nonmetals alone give molecular solids, except the network solids C (diamond, graphite), Si, $SiO_{2}$ and SiC.

Seeing Structure: Spectroscopy and Instrumental Analysis

Atoms absorb or emit only photons that match their energy-level gaps; AAS measures one element by shining that element's own lines through atomized sample.
Absorbance $A = lo g_{10} (I_{0} / I)$ is proportional to concentration; a calibration line from standards converts absorbance to concentration.
IR makes bonds vibrate: O–H 3200–3550 (broad), C–H 2850–3000, C=O 1680–1750 (strong, sharp); wavenumber $\tilde{ν} = 1/ λ$ .
Mass spectrometry: the molecular ion peak gives molar mass; one Cl gives M : M+2 = 3 : 1; one Br gives 1 : 1.
NMR: number of signals = number of H environments; signal area ∝ number of H atoms. MRI images hydrogen nuclei in the body.
X-ray crystallography maps electron density in crystals: Hodgkin (penicillin, B12, insulin), Franklin (DNA, Photo 51).

Unit 3: Energy Changes and Rates of Reaction

Enthalpy Changes and Calorimetry

$Δ H = H_{products} - H_{reactants} :$ negative means exothermic (surroundings warm), positive means endothermic (surroundings cool).
$q = m c Δ T$ , with $c = 4.18\text{ J/{(g}\cdot{}^\circ\text{C)}}$ for water and dilute aqueous solutions.
In a calorimeter, $q_{reaction} = - q_{surroundings} .$
Molar enthalpy: $Δ H_{molar} = \frac{q _{reaction}}{n}$ , and $Δ H = n Δ H_{molar}$ .
Reversing a thermochemical equation flips the sign of $Δ H$ ; multiplying the coefficients multiplies $Δ H$ by the same factor.

Hess's Law and Enthalpies of Formation

Hess's law: $Δ H$ depends only on the initial and final states, so $Δ H_{target} = \sum Δ H_{steps}$ .
Reverse an equation: change the sign of $Δ H$ . Multiply an equation by a factor: multiply $Δ H$ by the same factor.
$Δ H_{r x n}^{\circ} = \sum n Δ H_{f}^{\circ} (products) - \sum n Δ H_{f}^{\circ} (reactants)$
$Δ H_{f}^{\circ} = 0$ for any element in its standard state, e.g. $O_{2} (g)$ or $C (s, graphite)$ .
Bond energies: $Δ H \approx \sum (broken) - \sum (formed)$ ; breaking bonds absorbs energy, forming bonds releases it.

Physical, Chemical and Nuclear Energy: Fuels Compared

Energy scale: phase changes (tens of kJ/mol) < combustion (hundreds to thousands of kJ/mol) < fission and fusion (about $1 0^{9}$ to $1 0^{10}$ kJ/mol).
Nuclear energy comes from the mass defect: $E = m c^{2}$ ; 1 MeV $= 1.602 \times 1 0^{- 13}$ J; fission of uranium-235 releases about 200 MeV per nucleus.
CANDU reactors use natural uranium with a heavy-water ( $D_{2} O$ ) moderator and coolant and can be refuelled while running (Pickering, Bruce, Darlington).
Energy per gram $= \frac{Δ H _{c} ( kJ/mol )}{M ( g/mol )}$ ; CO₂ per kJ $= \frac{mass of CO _{2} per mol of fuel}{Δ H _{c}} .$
Efficiency $= \frac{useful energy out}{energy in} \times 100%;$ thermal plants reach about 30–40%, combined-cycle gas plants about 55–60%.
Fuels with more H per C give more energy per gram and less CO₂ per kJ: hydrogen > methane > octane > coal.

Rates of Reaction and Collision Theory

Average rate of consumption $= - \frac{Δ [ reactant ]}{Δ t}$ ; average rate of formation $= \frac{Δ [ product ]}{Δ t}$ , in mol/(L·s).
For $a A + b B \to c C + d D$ , the rate of each species divided by its coefficient gives the same rate of reaction.
An effective collision needs the correct orientation and at least the activation energy $E_{a}$ .
Rate increases with concentration (or gas pressure), temperature, surface area, and a catalyst; the nature of the reactants sets the activation energy.
Higher temperature increases the fraction of particles above $E_{a}$ ; a catalyst lowers $E_{a}$ without changing the energy distribution.

Potential Energy Diagrams and Mechanisms

$E_{a} (forward) = E_{activated complex} - E_{reactants}$ and $E_{a} (reverse) = E_{activated complex} - E_{products} .$
$Δ H = E_{products} - E_{reactants},$ so $E_{a} (reverse) = E_{a} (forward) - Δ H .$
A catalyst lowers $E_{a}$ for both the forward and reverse reactions by the same amount and does not change $Δ H$ .
Elementary steps must add up to the overall equation; intermediates are made then used up, catalysts are used up then regenerated.
The rate-determining step is the slowest step (highest activation energy) and limits the overall rate.

Unit 4: Chemical Systems and Equilibrium

Dynamic Equilibrium and the Equilibrium Constant

At dynamic equilibrium, forward rate = reverse rate, so concentrations stay constant while both reactions continue.
For $a A + b B ⇌ c C + d D$ : $K_{e q} = \frac{[ C ] ^{c} [ D ] ^{d}}{[ A ] ^{a} [ B ] ^{b}} .$ Pure solids and pure liquids are left out.
$K ≫ 1$ means mostly products at equilibrium; $K ≪ 1$ means mostly reactants. Only temperature changes K.
Reversing a reaction inverts K ( $1/ K$ ); multiplying coefficients by $n$ gives $K^{n}$ .
ICE table: Initial + Change = Equilibrium, with the Change row in the same ratio as the coefficients.

Le Chatelier's Principle and the Reaction Quotient

Reaction quotient Q has the same form as K but uses current concentrations.
$Q < K$ : shifts right (toward products). $Q > K$ : shifts left (toward reactants). $Q = K$ : at equilibrium.
Le Chatelier's principle: a system at equilibrium shifts to partly counteract any change imposed on it.
Decreasing volume shifts a gas equilibrium toward fewer moles of gas; a catalyst or an inert gas at constant volume causes no shift.
Only temperature changes K: raising temperature lowers K for an exothermic reaction and raises K for an endothermic one.

Acid-Base Equilibria

Brønsted-Lowry: an acid donates $H^{+}$ , a base accepts $H^{+}$ ; conjugate pairs differ by one $H^{+}$ .
$K_{w} = [H_{3} O^{+}] [OH^{-}] = 1.0 \times 1 0^{- 14}$ and $pH + pOH = 14.00$ at 25 °C.
$pH = - lo g [H_{3} O^{+}]$ , $[H_{3} O^{+}] = 1 0^{- pH}$ .
Weak acid: $K_{a} = \frac{x ^{2}}{c - x} \approx \frac{x ^{2}}{c}$ if percent ionization $= \frac{x}{c} \times 100% < 5%$ .
For a conjugate pair: $K_{a} K_{b} = K_{w}$ .

Solubility Equilibria, Buffers and Titration

AB salt: $K_{s p} = s^{2}$ . AB2 salt: $K_{s p} = 4 s^{3}$ .
$Q > K_{s p}$ : a precipitate forms. $Q < K_{s p}$ : no precipitate. Use concentrations after mixing.
Common ion effect: an ion already in solution lowers the solubility of a salt containing that ion.
Buffer: weak acid + conjugate base; $pH = p K_{a} + lo g \frac{[ A ^{-} ]}{[ HA ]}$ , and pH = pKa at half-equivalence.
Equivalence pH: strong-strong = 7, weak acid-strong base > 7, strong acid-weak base < 7.

Equilibrium at Work: Industry, Body and Environment

Contact process: $2 SO_{2} + O_{2} ⇌ 2 SO_{3}$ , $Δ H = - 198$ kJ; about 400–450 °C, $V_{2} O_{5}$ catalyst, excess air, 1–2 atm; a rate-versus-yield compromise.
Kidney stones form when $Q > K_{s p}$ for calcium oxalate in urine; dilution, less oxalate and citrate lower Q.
CO binds hemoglobin about 200 times more strongly than $O_{2}$ ; high $[O_{2}]$ shifts $HbO_{2} + CO ⇌ HbCO + O_{2}$ to the left.
Heavy metals are removed by precipitating hydroxides or (more completely) sulfides with very small Ksp values.
Heating hard water drives off $CO_{2}$ and precipitates $CaCO_{3}$ scale.
Ocean acidification: extra $H^{+}$ converts $CO_{3}^{2 -}$ to $HCO_{3}^{-}$ , lowering Q for $CaCO_{3}$ and making shells harder to build.

Measuring K in the Lab

Beer's law: $A = ε b c$ ; a calibration curve of A against c for standards has slope $ε b$ .
FeSCN²⁺ standards use a large excess of $Fe^{3 +}$ so nearly all $SCN^{-}$ is converted; test mixtures plus an ICE table give K.
Ksp of $Ca (OH)_{2}$ : titrate filtered saturated solution with HCl; $[Ca^{2 +}] = \frac{1}{2} [OH^{-}]$ and $K_{s p} = [Ca^{2 +}] [OH^{-}]^{2}$ .
Ka from pH: $[H_{3} O^{+}] = 1 0^{- pH} = [A^{-}]$ , $[HA] = c_{0} - [H_{3} O^{+}]$ ; at half-equivalence $pH = p K_{a}$ .
Cobalt chloride: pink $[\text{Co{(H}_2\text{O)}}_6]^{2+}$ ⇌ blue $[CoCl_{4}]^{2 -}$ ; added $Cl^{-}$ or heat turns it blue (forward reaction endothermic).
Percent error $= \frac{∣ exp - acc ∣}{acc} \times 100%$ ; look for the direction of systematic errors.

Unit 5: Electrochemistry

Oxidation Numbers and Redox Reactions

Oxidation is loss of electrons (oxidation number goes up); reduction is gain of electrons (oxidation number goes down).
The oxidizing agent is reduced; the reducing agent is oxidized.
Sum of oxidation numbers = 0 for a neutral compound, or = the charge for a polyatomic ion.
Usual values: O is −2 (−1 in peroxides), H is +1 (−1 in metal hydrides), F is −1, uncombined elements are 0.
In a balanced redox equation, total electrons lost = total electrons gained.

Balancing Redox with Half-Reactions

Order in acidic solution: other atoms, then O with $H_{2} O$ , then H with $H^{+}$ , then charge with $e^{-}$ .
Oxidation half-reaction: electrons are products. Reduction half-reaction: electrons are reactants.
Multiply the half-reactions so the electrons lost equal the electrons gained, then add and cancel.
In basic solution, add $OH^{-}$ to both sides to match each $H^{+}$ , form water, and cancel.
A correctly balanced redox equation balances in both atoms and total charge.

Galvanic Cells and Cell Potential

Anode: oxidation (negative electrode in a galvanic cell). Cathode: reduction (positive electrode).
Electrons flow through the wire from anode to cathode; in the salt bridge, anions move toward the anode and cations toward the cathode.
$E_{cell}^{\circ} = E_{cathode}^{\circ} - E_{anode}^{\circ},$ using reduction potentials for both.
$E_{cell}^{\circ} > 0$ : spontaneous as written. $E_{cell}^{\circ} < 0$ : not spontaneous (the reverse is).
Standard hydrogen electrode: $2 H^{+} + 2 e^{-} \to H_{2}$ , $E^{\circ} = 0.00 V$ by definition.
Cell notation: anode $∣$ anode solution $∥$ cathode solution $∣$ cathode.

Electrolytic Cells, Corrosion and Batteries

Electrolytic cell: electrical energy drives a non-spontaneous reaction ( $E_{cell}^{\circ} < 0$ ). Anode is positive, cathode is negative, but oxidation is still at the anode.
Minimum voltage needed for electrolysis is at least the size of the negative $E_{cell}^{\circ}$ (e.g. about 4.07 V for molten NaCl).
In electroplating, the object to be plated is the cathode and the plating metal is the anode.
Corrosion of iron: $Fe \to Fe^{2 +} + 2 e^{-}$ and $O_{2} + 2 H_{2} O + 4 e^{-} \to 4 OH^{-}$ ; needs both water and oxygen.
Galvanizing and sacrificial anodes protect iron by giving it a more active metal to be oxidized instead.
$Q = I t$ and $n (e^{-}) = \frac{I t}{F}$ , where $F \approx 96485 C/mol$ .