Quantitative Chemistry Summary

Atoms are extremely small, so scientists use relative atomic mass instead of actual masses.
Originally, masses were compared to hydrogen, but now carbon-12 is used for more accuracy.
Relative atomic mass is defined as: \text{Relative atomic mass} = \frac{\text{Mass of one atom of an element}}{\frac{1}{12} \text{Mass of atom of carbon-12}}
Carbon-12 has a relative atomic mass of 12.000000.
Relative atomic mass is also known as the mass number of the element.

The relative molecular mass of a compound is calculated by adding the relative atomic masses of all the atoms in a molecule.
Example: For CO2, the relative molecular mass is calculated as (1 carbon atom × 12) + (2 oxygen atoms × 16) = 44.

Percentage composition by mass of each element in a compound can be found using the chemical formula.
The general equation is: \% \text{ of Element E} = \frac{\text{Relative atomic mass of element E} \times \text{No. of atoms E in chemical formula}}{\text{Relative molecular mass of the compound}} \times 100
Example: For MgO, % of Mg = \frac{24 \times 1}{40} \times 100 = 60\%, % of O = \frac{16 \times 1}{40} \times 100 = 40\%.

Chemical equations provide information about chemical reactions, including the amounts of substances reacting.
Avogadro's concept: If the mass in grams is equal to the relative atomic mass, there will be 6.02 × 10^{23} atoms (Avogadro’s Number), defined as one mole.
Example: 24 g of Magnesium, 16 g of Oxygen, 12 g of Carbon, and 14 g of Nitrogen each contain 6.02 × 10^{23} atoms.

To find the number of atoms in a given mass of an element use the following:
Determine the number of moles: x = \frac{\text{given mass} \times 1}{\text{relative atomic mass}}
Multiply the number of moles by Avogadro's Number (6.02 × 10^{23}) to find the number of atoms.

The mole concept can be used to determine the amounts of reactants needed in a chemical reaction.
56 g of iron and 32 g of sulfur contain the same number of atoms (1 mole).

Molar mass is the mass of one mole of a compound, expressed in grams per mole (g/mol).
To find the molar mass, calculate the relative molecular mass of the compound.
Example: The molar mass of sodium hydroxide (NaOH) is 40 g/mol.

Chemical equations indicate the amounts of each substance involved in a reaction.
Stoichiometry refers to the relationship between the amounts of reactants and products.
Example: 2KHCO3(s) → K2CO3(s) + CO2(g) + H2O(g) indicates that 2 moles of KHCO3 decompose into 1 mole of K2CO3, 1 mole of CO2, and 1 mole of H2O.

The empirical formula shows the simplest whole-number ratio of atoms in a compound.

The molecular formula is a multiple of the empirical formula.
To find the molecular formula: Divide the molar mass of the compound by the relative molecular mass of the empirical formula, then multiply the subscripts in the empirical formula by this factor.
Example: If a compound has an empirical formula of CH2O and a molecular mass of 360 g/mol, the molecular formula is C12H24O12.

To determine the formula of a hydrated salt, find the ratio of moles of the anhydrous salt to moles of water.
Example: For MgSO4 • xH2O, heat the compound to find the mass of water lost, then calculate the moles of MgSO4 and H2O. The ratio of these moles gives the value of x.

Percentage yield compares the actual mass of product obtained to the calculated theoretical mass.
\text{Percentage Yield (%) = (Actual Yield / Theoretical Yield)} \times 100

Gases are measured by volume at standard temperature and pressure (s.t.p: 0 °C and 1 atmosphere).
According to Avogadro, 1 mole of any gas at s.t.p occupies 22.4 dm3.