MOLES AND EQUATIONS
I) COUNTING CHEMICAL SUBSTANCES IN BULK
Chemical formulae show how many atoms of each element are in one molecule, or one formula unit, of a compound.
Adding the relative atomic masses of elements and taking into account their quantities in a compound, allows Mr to be calculated.
The ratio of elements in terms of atomic mass in compounds tells us how much element makes up a certain mass of a compound.
For example, water has a ratio of hydrogen : oxygen of 2 : 16; therefore in 18 g of water, 2 g will be hydrogen and 16 g will be oxygen.
Any unit can be used as long as the ratios are kept the same.
The actual number of atoms would be a very large number, so chemists use a quantity called the mole.
A mole of a substance is its mass in g which is equal to its relative atomic, relative molecular or relative formula mass. Therefore a mole has the same number of particles as 12 g of carbon-12.
The number of atoms, molecules or ions in one mole is a constant called Avogadro’s number.
NA = 6.02 x 1023 mol-1
One mole of a substance is its molar mass, M.
The relative atomic mass’ unit could be defined as g mol-1.
To find the amount of moles in given mass of a substance we use the following formula:
n = m / Mr
Where n is the number of moles, m is the mass of the substance and Mr is the molecular/formula/molar mass of the substance.
If we are given a mass of one of the reactants used in a reaction, and we know the ratios of reactants and products, we can predict the exact amount products that will be produced.
For example:
2H2 + O2 à 2H2O
It can be seen there is a ratio of 2 mol of hydrogen to 2 mol of water. Therefore we can calculate the amount of water produced if we react 5g of hydrogen in excess oxygen. Ar of hydrogen is 1
2.5 mol of hydrogen 2.5 mol of water 2.5 x (16 + 1 + 1) = 45 g
II) CALCULATION OF EMPIRICAL AND MOLECULAR FORMULAE
The empirical formula of a compound shows the simplest whole-number ratio of the elements present.
For many simple compounds it is identical to the molecular formula.
The molecular formula is most useful, as it enables balanced chemical equations to be written.
The empirical formula can be found from data which gives a percentage composition of a compound.
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Step 1: Find the composition in grams of an element in a compound.
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Step 2: Find the ratio of the number of moles for each element by dividing the masses by the element’s relative atomic mass.
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Step 3: Divide by the smallest amount to give whole number ratios.
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Step 4: Write the ratios as an empirical formula.
Instead of masses in grams, percentage compositions can be used.
Chemical formulae may be deduced from the charge on ions in ionic compounds, so that the sum of the ionic charges is equal to zero.
In molecules, chemical formulae may be deduced from the numbers of electrons required to complete the outer electron shell.
Metals in compounds rarely change their name when in a compound, however non-metals change their names by become –ides. When oxygen is bonded to a non-metal in an ion, it becomes an –ate, for example sulphate.
III) BALANCING CHEMICAL EQUATIONS
When writing chemical equations, it is important that there is the same number of atoms of each element on both sides of the equation.
This is done by writing the formulae of all the reactants on one side and all the products on the other side of the arrow, with their state symbols as a subscript.
Then the total number of each element should be counted up in the reactants and compared to the products. Numbers can be added in front of the compounds so that the amount of each element balances on each side of the equation.
IV) BALANCING IONIC EQUATIONS
In some situations, ionic equations are useful as they show the changes to ions in a reaction. They can show how an ion becomes an element by gaining or losing electrons off or to another ion or element.
As with normal equations, the overall charge on the ions must balance on each side. Also use of state symbols is required.
They are useful in redox and precipitation reactions, as they show the transfer of electrons and the formation of solid compounds from two ions in solution.
V) CALCULATIONS INVOLVING CONCENTRATIONS AND GAS VOLUMES
Similar to the calculations of mass, the number of moles can be used in calculations involving solutions and gases.
Calculations involving solutions
When using moles, it important that volumes are measured in dm3, since concentrations are measured in mol dm-3 or M (molar). Therefore the equation for concentration is defined as:
c = n / v
A solution with a comparatively high concentration of solute is called a concentrated solution. A solution with a low concentration of solute is a dilute solution.
Titrations are an experimental technique used to accurately calculate the concentration of a solution of unknown concentration, using a known concentration (known as a standard solution).
To use titrations, there are four of five things that need to be known:
· The balanced reaction equation
· The volume of the solution of known concentration (first reagent)
· The concentration of the solution of first reagent
· The volume of the solution of the unknown concentration (second reagent)
· The concentration of the solution of the second reagent.
Using four of these quantities, we can calculate the fifth quantity, which is usually the concentration of the second reagent. Using the above formula and the balanced equation, we can calculate the number of moles in a particular volume.
This number of moles will be in a ratio with the solution of unknown concentration. Therefore if we know the number of moles of one solution, we can calculate the concentration of the other solution if we know its volume.
The process of calculating the ratios of reacting moles is called stoichiometry.
To convert between g dm-3 and mol dm-3 divide by the Mr of the substance.
Calculations involving gas volumes
Avogadro discovered that 1 mol of any gas at room temperature and pressure occupies 24 dm3. Therefore from the number of moles of a gas n, its volume V can be calculated using:
V = 24 x n
GLOSSARY
Empirical formula: the simplest whole-number ratio of the elements present in a compound.
Mole: the unit of amount of substance (abbreviation: mol). One mole of a substance is the mass that has the same number of particles (atoms, molecules, ions or electrons) as there are atoms in exactly 12 g of carbon-12.
Molecular formula: shows the total number of atoms present in a molecule of the compound.
Precipitate: an insoluble solid formed when two solutions react, e.g. white silver chloride formed when silver nitrate is added to sodium chloride solutions.
Redox: reactions which involve reduction and oxidation processes.
Titration: measurement of the exact amount of one solution needed to react with a fixed amount of another solution.
Stoichiometry: the stoichiometry (or stoichiometric ratio) for a reaction shows the mole ratio of reactants and products in the balanced equation for the reaction.
