Chapter 1: Stoichiometric Relationships

SOLID: - Arrangement: Regular, fixed shape, particles are close together - Movement: Vibrate slowly LIQUID: - Arrangement: Random, takes shape of container, particles are close together but further apart than solids - Movement: Vibrate faster around each other GAS: - Arrangement: Random, fills container, particles far apart - Movement: Moves fastest in all directions

HEATING CURVE: - As time increases, the amount of kinetic energy possessed by the particles increase, causing the temperature to increase - However, changes of state occur at a constant temperature, as heat is being supplied to overcome the intermolecular forces of attraction between particles

ELEMENT: pure substance containing only 1 type of atom ATOM: smallest part of an element that can still be recognised as that element COMPOUND: 2 or more elements combined chemically - Chemical change that elements undergo to form a compound allow compounds to have different properties from their constituent elements - Chemical equations show how elements combine in fixed ratios to form compounds (must be balanced to show conservation of mass –> same no. of atoms on both sides of equation)

MIXTURE: 2 or more substances (either elements or compounds) physically mixed together - Since components of a mixture do not undergo chemical change to form mixtures, mixtures have similar properties to their constituent elements - And since the components are physically combined, they can also be separated using physical methods • HOMOGENOUS MIXTURE: uniform composition, consisting only of 1 phase • HETEROGENOUS MIXTURE: non-uniform composition and consists of more than 1 phase

RELATIVE ATOMIC MASS (Aᵣ): Average of the masses of the isotopes found in a naturally occurring sample of the element, relative to 1/12 of a carbon-12 atom - Eg. Aᵣ of C: 12.01 RELATIVE MOLECULAR MASS (Mᵣ): Mass of a molecule of an element/compound, relative to 1/12 of a carbon-12 atom - Eg. Mᵣ of CO₂: 44.01 MOLAR MASS: Mass of 1 mole of atoms in an element/compound - Eg. Molar mass of C: 12.01 g mol⁻¹

1 mole of a substance: 6.02 x 10²³ particles (Avogrado’s Constant) No. of particles/6.02 x 10²³ = No. of mols

Mass (g) / No. of mols = Molar mass (g mol⁻¹)

% composition = (Aᵣ x no. of atoms of an element / Total Mᵣ) x 100%

EMPIRICAL FORMULA: Simplest whole number ratio of the elements in a compound Eg. CₓHᵧ MOLECULAR FORMULA: Total number of atoms of each element present in a molecule of the compound, and is a multiple of the empirical formula Eg. n(CₓHᵧ) Find empirical and molecular formula: - Divide masses by Mᵣ to find no. of mols of each - Express as whole number ratio (empirical formula) - Calculate Mᵣ of compound using empirical formula - Divide actual Mᵣ by calculated Mᵣ to find n - Find molecular formula

CₓHᵧ + O₂ –> x CO₂ + y/2 H₂O • Use given masses of CO₂ and H₂O to find no. of mols of each (x = No. of mols of CO₂) (y = no. of mols of H₂O x 2) • Express as empirical formula

1) Conservation of mass Total mass of reactants = Total mass of products 2) Converting to number of moles - Use molar ratio to determine number of moles of required quantity - Convert back to mass using formula

% yield = Actual yield/Theoretical yield x 100%

LIMITING REAGENT: reagent used up before the others - Find no. of mols of each reactant - Divide by coefficient in chemical equation - Smallest number: LR

IDEAL GAS: concept used by scientists to model the behaviour of real gases (under normal conditions, real gases behave similarly to ideal gases) Assumptions: - Particles themselves have no volume/volume of particles are negligible compared to the volume of the container - Intermolecular forces of attraction between gaseous particles are negligible Ideal conditions: 1) High temperature - High temp: High K.E, can overcome IMFA between particles - Low temp: Particles move slowly, making IMFA significant, and deviate from ideal behaviour 2) Low pressure - Low pressure: Particles are further apart, making volume of particles negligible compared to volume of container - High pressure: Particles are close together, making volume of particles and IMFA significant

AVOGRADO’S LAW: Equal volumes of gases measured at the same temperature and pressure have equal numbers of particles (no. of mols) Hence, Volume of gases ∝ Number of moles of gas –> volumes can be used to directly compare quantities of gases (instead of converting to mole)

• Volume of 1 mol of an ideal gas: 22.7 dm³ mol⁻¹ Volume/22.7 = No. of mols

1) Boyle’s law: when temperature is constant, pressure and volume are inversely proportional (P₁V₁ = P₂V₂) P∝ 1/V (P as y-axis, V as x-axis: inversely proportional graph) P = k/V; P=k x 1/v (P as y-axis, 1/V as x-axis: directly proportional graph) PV = k (P as y-axis, P as x-axis: horizontal graph) 2) Charles’ law: When pressure is constant, volume is directly proportional to its temperature (V₁/T₁ = V₂/T₂) V∝T; V=kT (V as y-axis, T as x-axis: directly proportional graph) 3) When volume is constant, pressure and temperature are directly proportional (P₁/T₁ = P₂/T₂) P∝T P=kT Therefore: PV = nRT OR P₁V₁/T₁ = P₂V₂/T₂ P: pa V: m³ n: no. of mols R: 8.31 J K⁻¹ mol⁻¹ T: K

Concentration x Volume = No. of mols

g dm⁻³/mol dm⁻³ = g mol⁻¹

Concentration in ppm = Mass of solute x 10⁶ / Mass of solution

Titration: technique to find the volumes of 2 solutions that react exactly with each other to find an unknown concentration • Find no. of mols of 1 substance using given concentration and volume • Use molar ratio to find no. of mols of unknown substance • Use volume obtained from titration to find concentration of unknown substance

• Find no. of mols of HCl reacted • Write out balanced chemical equation of reaction between HCl and Na₂CO₃ and find no. of mols of Na₂CO₃ reacted with 25.00cm³ solution using molar ratio • Calculate no. of mols of Na₂CO₃ reacted with 250.00cm³ solution • Find mass of Na₂CO₃ • Subtract mass of Na₂CO₃ from mass of Na₂CO₃* xH₂O to find mass of H₂O • Find no. of moles of H₂O • Find ratio of Na₂CO₃ to H₂O to find x

• Find no. of mols of NaOH • Find no. of mols of HCl in 25.00cm³ solution (unreacted HCI) • Find no. of mols of HCl in 100.00cm³ solution • Find total no. of mols of HCl added • Subtract unreacted HCl from total HCl to find reacted HCl • Find no. of mols of CaCO₃ • Find mass of CaCO₃ • Find percentage purity