Master chemical reactions and stoichiometry — the mathematical heart of chemistry — with this comprehensive flashcard deck. Stoichiometry allows chemists to predict exactly how much of each substance is needed or produced in a reaction, making it essential for...
Master chemical reactions and stoichiometry — the mathematical heart of chemistry — with this comprehensive flashcard deck. Stoichiometry allows chemists to predict exactly how much of each substance is needed or produced in a reaction, making it essential for every branch of chemistry, from laboratory work to industrial manufacturing.
This deck covers the five main types of chemical reactions, how to balance chemical equations, mole calculations, molar mass, limiting and excess reagents, theoretical and percent yield, concentration calculations, and gas stoichiometry. Essential for AP Chemistry, GCSE, A-Level, and university general chemistry.
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1. Synthesis: A + B → AB. 2. Decomposition: AB → A + B. 3. Single displacement: A + BC → AC + B. 4. Double displacement: AB + CD → AD + CB. 5. Combustion: Fuel + O₂ → CO₂ + H₂O.
Two or more substances combine to form a single product. Example: 2H₂ + O₂ → 2H₂O. 2Na + Cl₂ → 2NaCl.
A single compound breaks down into two or more simpler substances. Example: 2H₂O₂ → 2H₂O + O₂. Often requires heat, light, or electricity.
A more reactive element displaces a less reactive element from a compound. Example: Zn + CuSO₄ → ZnSO₄ + Cu. Determined by the activity series.
Ions of two compounds exchange partners. Example: AgNO₃ + NaCl → AgCl↓ + NaNO₃. Often forms a precipitate, gas, or water.
A substance (usually hydrocarbon) reacts rapidly with oxygen, releasing heat and light. Complete combustion: Fuel + O₂ → CO₂ + H₂O. Incomplete combustion produces CO and soot (C).
Matter cannot be created or destroyed in a chemical reaction. The total mass of reactants = total mass of products. This is why we balance equations.
Adjust coefficients (numbers in front of formulas) so the number of each type of atom is equal on both sides. NEVER change subscripts in formulas. Example: H₂ + O₂ → H₂O becomes 2H₂ + O₂ → 2H₂O.
The mass in grams of one mole of a substance. Equal to the sum of atomic masses of all atoms in the formula. Example: H₂O = 2(1) + 16 = 18 g/mol.
Moles = Mass (g) ÷ Molar Mass (g/mol). Rearranged: Mass = Moles × Molar Mass. Molar Mass = Mass ÷ Moles.
The calculation of quantities (moles, mass, volume) of reactants and products in a chemical reaction, based on the balanced equation's molar ratios.
The coefficients in a balanced equation give the mole ratio. Example: N₂ + 3H₂ → 2NH₃. If 1 mol N₂ reacts, it uses 3 mol H₂ and produces 2 mol NH₃.
The reactant that is completely consumed first and determines the maximum amount of product that can be formed. The other reactant is in excess.
1. Convert each reactant to moles. 2. Divide moles by each reactant's coefficient. 3. The smallest result is the limiting reagent. Then use limiting reagent moles to calculate product.
The maximum amount of product that could be produced if the reaction goes to completion with no losses. Calculated from stoichiometry using the limiting reagent.
The amount of product actually obtained from an experiment. Always ≤ theoretical yield due to incomplete reactions, losses during collection, and side reactions.
% Yield = (Actual Yield ÷ Theoretical Yield) × 100%. A 100% yield is ideal but rarely achieved. A yield > 70% is generally considered good in organic chemistry.
Molarity = concentration of a solution in moles per litre (mol/L). Formula: M = moles of solute ÷ volume of solution (L). Example: 2 mol NaCl in 0.5 L = 4 M NaCl.
Adding solvent to decrease a solution's concentration. Formula: M₁V₁ = M₂V₂ (moles before = moles after dilution). Example: Diluting 1 L of 6M HCl to 3 L gives 2 M HCl.
PV = nRT. P = pressure (atm), V = volume (L), n = moles, R = 0.0821 L·atm/mol·K, T = temperature (Kelvin). Describes behavior of an ideal gas.
Temperature = 273 K (0°C) and Pressure = 1 atm. At STP, one mole of any ideal gas occupies 22.4 L (molar volume).
A reaction that releases energy (heat) to the surroundings. ΔH is negative. The products have lower energy than the reactants. Example: combustion, neutralization reactions.
A reaction that absorbs energy (heat) from the surroundings. ΔH is positive. Products have higher energy than reactants. Example: photosynthesis, dissolving ammonium nitrate in water.
A double displacement reaction that forms an insoluble solid (precipitate) when two aqueous solutions are mixed. Example: Pb(NO₃)₂ + 2KI → PbI₂↓ (yellow precipitate) + 2KNO₃.
A reaction involving the transfer of electrons. Oxidation = loss of electrons (OIL). Reduction = gain of electrons (RIG). OIL RIG — the substance oxidized is the reducing agent; reduced = oxidizing agent.
