Syllabus notes
Calculations in chemistry


1.1 Mole concept & Avogadro's constant

1.1.1 : 1 Mole of something is equivalent to 6.023 x 1023 (<-- Avogadro's constant) units of it...ie lots of atoms, molecules etc...the periodic table gives molar masses...ie the number of grams of a substance required for 1 mol of atoms. This can be extrapolated to molecules of known molecular formula.

1.1.2 : Number of mols = mass / mass per mol (Usually found on periodic table)...The coefficients in chemical equations give the molar ratios of reactants and products...ie 2A + 3B -> C. There is 2/3 as much A as B, and 3 times more B than C involved in the reaction...Assuming the reaction goes to completion, there must be 3/2times as much B as A for neither to remain...If this ratio is not followed, one will be a limiting reactant, and so the reaction will have some of the other reactant left over when it completes.

1.2 Formulae

1.2.1 : Atomic mass, Molecular Mass, Formula Mass : All the mass per mol of a particular type of species...atoms, molecules or formula units. These can be found for the periodic table, and will give the mass for 1 mol of the species (or rather the average accounting for different isotopes and their relative abundance). Mr is the ratio between the molar masses of two species. Ar is the ratio of the number of atoms between two species. These two ratios will be equal.

1.2.2 : Moles vs Mass...Moles is a number of something...every mol being 6.023x1023 individual elements. Mass is the property which results in 'weight' in the presence of Gravity. Given a molar mass, Mr a mass m and a number of mols N then N=m/Mr

1.2.3 : An 'Empirical formula' is the formula describing the different atoms present in a molecules, and their ratios, but not the actual number present. ie AxByZc could be an empirical formula if x, y, and z are in lowest common terms. The molar mass can then be used to calculate the actual numbers of each atom present per molecule. The empirical formula can be determined by percentage composition, or anything else which gives the ratios of atoms present.

1.2.4 : A Molecular Formula is similar to an empirical formula except that it includes the the number of atoms present in each molecule, rather than the ratio. It will be an integer multiple of the empirical formula ie KAxByZc and can be calculated from the empirical formula and the molar mass of the substance.

1.3 Chemical Equations

1.3.1 : The mole ratio of two species in a chemical equation is the ratio of their coefficients...ie aX + bY -> cZ : The ratio of X/Yis a/b, Y/Z= b/cetc...

1.3.2 : Balancing equations...change only the coefficients, not the subscripts to make sure all atoms, and charge is conserved (half equations can be balanced by addition of electrons to either side...2 half equations can be added by making the number of electrons equal in each, then vertically adding.)

1.3.3 : State symbols -- (s)-Solid , (l)-liquid, (g)-gas, (aq)-aqueous solution...ie something dissolved in water. Should be included in all chemical reactions (but won't be penalized).

1.4 Mass relationships in chemical reactions

1.4.1 : Mass is conserved throughout reactions. This fact allows masses to be calculated based on other masses in the reaction eg burning Mg in air to produce MgO and so to find the mass or Mg present in the original sample (ie purity)...can be extended to concentrations...ie titration.

1.4.2 : When a reaction contains several reactants, some may be in excess...is more is present that can be used in the reaction. The first reactant to run out is the limiting reagent (or reactant). Knowing the number of mols of the limiting reagent allows all other species to be calculated, and so the yield, and remaining quantities of other reactants.

1.5 Solutions

1.5.1 : Solvent - the medium that you're dissolving in...ie water. Solute - the compound that you're dissolving...eg. an ionic compound ...Solution - the two of them when mixed together intimately so that the solid phase becomes indistinguishable from the liquid phase.

Concentration - the amount of solute per amount of solvent...in mols per dm3 (litre) or grams per litre.

1.5.2 : Apply the equation concentration = number/volume...rather obvious from the units of concentration, but remember to convert everything into the same units.

1.5.3 : Use chemical equations to relate the amount of one species to the amounts of others.