Unit 3 – Chemical Quantities
6.1
- qualitative and quantitative analysis
6.3
- mole concept – count by units of Avogadro’s number (6.022 x 10 23)
- based on the 12C isotope – equal # of atoms on the periodic table
- equation ->
# particles = n x N; particles are: atoms, molecules, formula units or ions
6.4 and 6.5
- molar mass – calculate off the periodic table
- equation -> m = n x MM
- 6.5: moles link the two equations in 6.3 and 6.4
6.6
- percentage composition = mass of element/mass of compound, formulas to %
- eg. 2.80 g of CO2 contains 2.80 g x (12.0g/44.0g) = 0.764 g of carbon
6.7
- empirical formula -> opposite of 6.6, going from % to formulas
- convert from % --> mass ---> mol ---> atoms
- must be whole numbers, may have to multiply by a number to get a whole number (tbl 3 p. 291)
6.9
- molecular formula: EF x a = MF, need to know the MF mass and scale up the mass of the
EF by a whole number (a).
Lab - hydrates, anhydrous salts and how to calculate the amount of water in the formula
7.1
- mol-mol ratios in chemical reactions
7.2 Stoichiometry
- mass – mol relationships in chemical reactions
7.3 and 7.4 – Limiting Reactants
- limiting reagents and excess reagents – the L.R. determines the quantity of the product
- to find the L.R. calculate the mole of each reactant and divide by the mole ratio
7.5 Error analysis
- percent error and percent yield equations