Amount of substance

Relative atomic mass (Ar)

Average mass of an atom of an element

1/12 mass of one atom of C12

Contains 1 mole of atoms

Formula for moles

Moles = mass / mr

Formula for concentration

Conc= mass/volume Or conc= moles/volume

Mass= g

Volume= dm³

Empirical formula

Simplest whole number ratio of atoms of each element in a compound

Molecular formula

Actual number of atoms of each element in a compound

A hydrocarbon contains 1.2g of carbon and 0.4g of hydrogen. What is the empirical formula?

Carbon Hydrogen

Mass 1.2 0.4

Ar 12.0 1.0

Moles 1.2/12=0.1 0.4/1.0=0.4

Mole ratio 1 : 4

=CH4

Water of crystallisation

Water found in a crystalline framework of an ionic compound

CuSO4 . XH2O

Hydrated copper sulfate mass= 6.864g

Anhydrous copper sulfate mass= 4.389g

Find X

Mass of water loss= 6.864-4.389=2.475

Moles of CuSO4:

Moles=mass/mr

4.389/159.6=0.0275

Moles of water:

Moles=mass/mr

2.475/18.0=0.138

0.1375/0.0275=5

1:5

So CuSO4 . 5H2O

Law of conservation of mass

Atoms cannot be created or destroyed, only reorganised

Charges on NH4 , OH , NO3 , SO4, CO3 , PO4

Ammonium- NH4 1+

Hydroxide- OH-

Nitrate- NO3 1-

Sulfate SO4 2-

Carbonate- CO3 2-

Phosphate- PO4 3-

Ideal gas equation

pV= n R T

Pressure is Pa (pascals)

Volume is m³

R (universal gas constant) 8.31 J K-1 Mol-1

How many kPa is 1 atmosphere (atm)

101 kPa = 1 atmosphere (atm)

Convert Celsius to kelvin

Celsius + 273 = K

Volume conversion from cm³ to m³

Cm³ → m³ Divide by 1,000,000 or 10^6

Limiting factor

Chemical that gets used up

Excess

More than enough of this chemical, will be some left over

Theoretical yield (max yield)

Largest amount of product that could be made (assuming no problems)

Actual yield

Amount of product you make in an experiment (usually less than theoretical yield, CANNOT be bigger)

Percentage yield

Actual yield/ theoretical yield x 100

Mass produced/ max mass possible x 100

Moles produced/ theoretical moles x 100