Memory tricks for reaction types, redox & equilibrium
From the 5 reaction types to OIL RIG to Le Chatelier's principle — chemical reactions follow predictable patterns. These memory tricks lock in the rules so you can identify, balance, and predict any reaction your exam throws at you.
Every reaction in general chemistry fits one of five categories. SDSDC: Synthesis (A+B→AB) · Decomposition (AB→A+B) · Single Replacement (A+BC→AC+B) · Double Replacement (AB+CD→AD+CB) · Combustion (fuel+O₂→CO₂+H₂O). Identify the type first — then predict the products.
Synthesis
Two or more substances combine to form one product. 2H₂ + O₂ → 2H₂O. Always A + B → AB pattern.
Decomposition
One compound breaks into two or more simpler substances. 2H₂O → 2H₂ + O₂. Opposite of synthesis.
Single replacement
One element displaces another from a compound. Zn + CuSO₄ → ZnSO₄ + Cu. Use activity series to predict if it occurs.
Double replacement
Ions of two compounds exchange partners. Often produces a precipitate, gas, or water. AgNO₃ + NaCl → AgCl↓ + NaNO₃.
Combustion
Hydrocarbon + O₂ → CO₂ + H₂O (complete) or CO + H₂O (incomplete). Always exothermic — releases energy.
Redox
OIL RIG
Oxidation Is Loss · Reduction Is Gain (of electrons)
Oxidation & Reduction — Never Mix Them Up
OIL RIG is the most tested mnemonic in all of chemistry. Oxidation Is Loss of electrons — the species being oxidized loses electrons and its oxidation number increases. Reduction Is Gain of electrons — the species being reduced gains electrons and its oxidation number decreases. They always happen together.
Oxidizing agent
The species that causes oxidation by accepting electrons — it gets REDUCED itself. "OA gets reduced."
Reducing agent
The species that causes reduction by donating electrons — it gets OXIDIZED itself. "RA gets oxidized."
LEO says GER
Alternative: Lose Electrons = Oxidized · Gain Electrons = Reduced. Both mnemonics work — pick your favorite.
Oxidation numbers
Increase = oxidized. Decrease = reduced. Free elements = 0. O usually = -2. H usually = +1. Monatomic ion = its charge.
Balancing
COACH (C=Count atoms, O=Only change Coefficients, A=Atoms conserved, C=Check both sides, H=Hydrogen and oxygen last)
Count atoms · Only change Coefficients · Atoms conserved · Check both sides · Hydrogen and oxygen last
Balancing Chemical Equations — The Strategy
COACH gives you the systematic approach to balancing any equation. Count atoms on each side · Only change Coefficients (never subscripts) · Atoms are Conserved — matter can't be created or destroyed · Check both sides when done · Balance Hydrogen and Oxygen last (they appear in the most compounds).
Never change subscripts
Changing H₂O to H₃O changes the compound entirely. Only the big numbers (coefficients) in front of formulas can change.
Start with metals
Balance elements that appear in only one reactant and one product first. Leave H and O for last — they appear everywhere.
Use fractions if needed
It's fine to use ½ or ¾ as intermediate coefficients — multiply everything by the denominator at the end to clear fractions.
Polyatomic ions
If a polyatomic ion appears unchanged on both sides, balance it as a unit — don't split it into atoms.
Equilibrium
LE CHAT
Le Chatelier's principle — stress the system, it shifts to relieve the stress
Le Chatelier's Principle — Predicting Shifts
LE CHAT: if a system at equilibrium is stressed, it shifts to counteract the stress. Add reactant → shifts right. Remove product → shifts right. Increase pressure → shifts toward fewer moles of gas. Increase temperature → shifts in the endothermic direction. Used to maximize yield in industrial reactions (Haber process).
Concentration
Add reactant or remove product → shifts right (forward). Remove reactant or add product → shifts left (reverse).
Pressure
Increase pressure (decrease volume) → shifts toward side with fewer moles of gas. Equal moles of gas = no shift.
Temperature
Treat heat as a reactant (endothermic) or product (exothermic). Increase temp → shifts away from heat. Decrease → shifts toward heat.
Catalyst
Speeds up both forward and reverse reactions equally — reaches equilibrium faster but does NOT shift the equilibrium position.
Reaction Rates
CATS (C=Concentration, A=Area/surface area, T=Temperature, S=Surface area catalyst/catalyst)
Concentration · Area (surface) · Temperature · Surface area catalyst
4 Factors That Speed Up Reactions
Reaction rate depends on collision frequency and energy. CATS: Concentration (more particles = more collisions) · Area/surface area (powders react faster than chunks) · Temperature (faster molecules = more energetic collisions) · catalyst (lowers activation energy). All four increase the rate of successful collisions.
Collision theory
Reactions occur when particles collide with sufficient energy (≥ activation energy) and proper orientation. Rate = collision frequency × fraction with enough energy.
Temperature effect
Rule of thumb: every 10°C rise roughly doubles reaction rate. Higher temperature → more molecules exceed activation energy threshold.
Catalyst mechanism
Provides an alternative reaction pathway with lower activation energy. Is regenerated — not consumed. Enzymes are biological catalysts.
Activation energy
Minimum energy needed for a reaction to occur. Shown on energy diagrams as the "hill" between reactants and products. Exothermic: products lower than reactants.
Activity Series
Please Stop Calling Me A Zebra — I Like Copper's Silver Nature
K · Na · Ca · Mg · Al · Zn · Fe · Ni · Sn · Pb · H · Cu · Ag · Au · Pt
Activity Series of Metals — Most to Least Reactive
The activity series ranks metals by reactivity. Metals higher on the list displace metals lower on it from ionic solutions. Please Stop Calling Me A Zebra — I Like Copper's Silver Nature gives you K→Na→Ca→Mg→Al→Zn→Fe→Ni→Sn→Pb→H→Cu→Ag→Au→Pt from most to least reactive.
Rule
Metal A displaces Metal B from solution only if A is above B on the activity series. Zn displaces Cu; Cu cannot displace Zn.
Hydrogen's position
Metals above H react with acids to produce H₂ gas. Metals below H (Cu, Ag, Au, Pt) do NOT react with dilute acids.
Most reactive
K, Na, Ca react violently with water producing H₂ gas and metal hydroxide. Never store in water or open air.
Least reactive
Au, Pt are noble metals — resistant to corrosion and most chemical reactions. That's why gold jewelry lasts forever.
Precipitation
SNAP (N=Nitrates always soluble, A=Alkali metals always soluble, P=Perchlorates always soluble — the "S" frames the mnemonic)
SNAP covers the always-soluble categories. Beyond SNAP: most chlorides are soluble EXCEPT AgCl, PbCl₂, Hg₂Cl₂ · most sulfates soluble EXCEPT BaSO₄, PbSO₄, CaSO₄ · most carbonates, phosphates, and hydroxides are INSOLUBLE except with alkali metals or NH₄⁺. Insoluble = precipitate (↓).
Always soluble
All nitrates (NO₃⁻) · All alkali metal salts (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺) · All ammonium salts (NH₄⁺) · All perchlorates.
Carbonates · Phosphates · Hydroxides · Sulfides — all insoluble UNLESS combined with alkali metals or NH₄⁺.
Net ionic equation
Remove spectator ions (those that don't participate). The net ionic equation shows only the ions that actually react to form the precipitate.
Exo vs Endo
EXO exits · ENDO enters
Exothermic = heat exits system (ΔH negative) · Endothermic = heat enters system (ΔH positive)
Exothermic vs Endothermic Reactions
EXO exits — heat leaves the system, surroundings get warm, ΔH is negative. ENDO enters — heat flows into the system from surroundings, surroundings get cold, ΔH is positive. Combustion is always exothermic. Photosynthesis is endothermic. Hand warmers = exo. Ice packs = endo.
ΔH negative
Exothermic — products have LESS energy than reactants. Energy released. Feels hot. Combustion, neutralization, respiration.
ΔH positive
Endothermic — products have MORE energy than reactants. Energy absorbed. Feels cold. Photosynthesis, melting ice, cooking.
Energy diagram
Exothermic: products lower than reactants on the diagram. Endothermic: products higher. Activation energy is the hill in both cases.
Bond energy connection
Breaking bonds requires energy (endothermic). Forming bonds releases energy (exothermic). ΔH = energy to break bonds − energy released forming bonds.
Equilibrium Constant
K = Products over Reactants (raised to their coefficients)
K expression: products in the numerator, reactants in the denominator, each raised to the power of its coefficient. K > 1 → products favored (equilibrium lies to the right). K < 1 → reactants favored (lies left). K = 1 → roughly equal. Pure solids and liquids are NOT included in K expressions.
Writing Keq
For aA + bB ⇌ cC + dD: Keq = [C]^c[D]^d / [A]^a[B]^b. Use molar concentrations for Kc or partial pressures for Kp.
Excluded species
Pure solids (s) and pure liquids (l) have constant concentrations — they're excluded from K expressions. Only aqueous (aq) and gas (g) species are included.
Reaction quotient Q
Q uses same expression as K but with non-equilibrium concentrations. Q < K → reaction proceeds right. Q > K → reaction proceeds left. Q = K → at equilibrium.
Kp vs Kc
Kp uses partial pressures; Kc uses molar concentrations. Related by: Kp = Kc(RT)^Δn where Δn = moles gas products − moles gas reactants.
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🎓 Common Exam Questions
Q: What does SDSDC stand for and give an example of each reaction type?
A: SDSDC = Synthesis, Decomposition, Single replacement, Double replacement, Combustion. Synthesis (S): A + B → AB. Example: 2Na + Cl2 → 2NaCl. Decomposition (D): AB → A + B. Example: 2H2O2 → 2H2O + O2. Single replacement (S): A + BC → AC + B. Example: Zn + CuSO4 → ZnSO4 + Cu (Zn is above Cu in activity series). Double replacement (D): AB + CD → AD + CB. Example: AgNO3 + NaCl → AgCl↓ + NaNO3. Combustion (C): fuel + O2 → CO2 + H2O. Example: CH4 + 2O2 → CO2 + 2H2O. Complete combustion always gives CO2 and H2O.
Q: Explain OIL RIG and how to assign oxidation numbers.
A: OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain of electrons. They always occur together — you cannot have one without the other. Rules for oxidation numbers: pure element = 0; monatomic ion = its charge; O is usually -2 (except in peroxides: -1); H is usually +1 (except metal hydrides: -1); sum of oxidation numbers = charge of species. Example: in H2SO4, H is +1, O is -2, so S must be +6. Identify redox: the atom that increases in oxidation number is oxidized (loses electrons) — its compound is the reducing agent. The atom that decreases is reduced — its compound is the oxidizing agent.
Q: What does COACH stand for and walk through balancing a redox equation?
A: COACH = Count atoms, Only change Coefficients (never subscripts), Atoms are conserved, Check both sides, Hydrogen and oxygen last. For simple equations: balance non-H, non-O atoms first, then H, then O, then check. For redox in acidic solution: separate into half-reactions, balance atoms (add H2O for O, H⁺ for H), balance charge (add electrons), multiply half-reactions so electrons cancel, add and simplify. Example: MnO4⁻ + Fe²⁺ in acid → Mn²⁺ + Fe³⁺. Balance each half-reaction separately, then combine.
Q: Explain Le Chatelier's principle (LE CHAT) with examples of each type of stress.
A: LE CHAT: if a system at equilibrium is stressed, it shifts to relieve that stress. Stresses and responses: Add reactant → shifts RIGHT (makes more product). Remove product → shifts RIGHT. Add product → shifts LEFT. Remove reactant → shifts LEFT. Increase pressure/decrease volume → shifts toward FEWER moles of gas. Decrease pressure → shifts toward MORE moles of gas. Increase temperature → shifts in ENDOTHERMIC direction. Decrease temperature → shifts in EXOTHERMIC direction. Adding a catalyst does NOT shift equilibrium — it speeds up reaching equilibrium but doesn't change K or the equilibrium position. Adding an inert gas at constant volume also has no effect.
Q: What does SNAP cover in solubility rules, and what are the other key rules?
A: SNAP covers the always-soluble categories: N=Nitrates (NO3⁻) always soluble, A=Alkali metal ions (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺) always soluble, P=Perchlorates (ClO4⁻) always soluble — the S is the mnemonic frame holding it together. Beyond SNAP: most Chlorides soluble EXCEPT AgCl, PbCl2, Hg2Cl2. Most Sulfates soluble EXCEPT BaSO4, PbSO4, CaSO4. Carbonates (CO3²⁻), Phosphates (PO4³⁻), Hydroxides (OH⁻) generally insoluble EXCEPT with alkali metals or NH4⁺. Use these rules to predict whether a precipitate forms in double replacement reactions — if the product is insoluble, a precipitate (↓) forms.