Edexcel Chemistry - Topic 5: Formulae, Equations and Amounts

• The mole is the amount of substance in grams that has the same number of particles as there are atoms in 12 grams of Carbon-12
• One mole of any substance contains exactly 6.02x10^23 atoms or molecules of that substance - This is Avogadros Constant
• The Relative Atomic Mass is the average mass of one atom compared to one twelth of the mass of one atom of Carbon-12
• The Molar Mass is the mass in grams of 1 mole of a substance and is given the unit of g mol-1
• The Molar Mass for a compound can be calculated by adding up the mass numbers of each element in a compound.
  e.g. CaCO3 = 40.1 + 12.0 + (3 x 16.0) = 100.1 g mol-1

The number of moles of a pure solid, liquid or gaseous substance = Mass / Molar Mass

The number of moles of a gas at room temperature (25C) and pressure (1atm) = Volume / 24
The volume of a gas is measured in dm3

The number of moles in a solution = Concentration x Volume
Concentration is measured in mol dm-3 or M

The number of moles of atoms = Number of atoms / Avogadros Number

Empirical Formulas

• An empirical formula is the simplest ratio of atoms of each element in a compound

Deriving an Empirical Formula

• Divide each mass/percentage mass by the atomic mass of the element
• For each of the answers from step 1 divide by the smallest one of those numbers
• Multiply up to give whole numbers

Molecular Formulas

• A molecular formula is the actual number of atoms of each element in the compound
• From the relative molecular mass (Mr) work out how many times the mass of the empirical formula fits into the Mr

• A solution is a mixture formed when a solute dissolves in a solvent
• Molar concentrations can be measured for solutions by: no. moles / volume of solution
• When soluble ionic solids dissolve in water they will dissociate into seperate ions leading to differences in concentrations of ions and the concentration of the solute
  e.g. 0.1 mol of NaCl(s) dissolved in water will form 0.1 mol of Na+ ions and 0.1 mol of Cl- ions
• Diluting a solution involves pipetting a known volume of a solution into a volumetric flask and making up the solution to the mark with distilled water and mixing the solution
• Diluting a solution increases the volume of water in the solution and therefore lowers the concentration of the solution. The number of moles of solute doesn't change
• Therefore: diluted concentration = original concentration x (Original volume/new volume)
• The density of a substance is the number of grams per unit volume of a substance.
  Density = mass / volume

• Weigh the sample bottle containing the required mass of solid on a 2dp balance
• Transfer to a beaker and reweigh the sample bottle
• Record the difference in mass
• Add 100cm3 of distilled water to the beaker. Use a glass rod to stir to help dissolve the solid
• If the substance is not dissolving in cold water, gently heat the beaker until all the solid has dissolved
• Pour the solution into a 250cm3 volumetric flask via a funnel
• Rinse the beaker and funnel with distilled water and add the washings from the beaker and glass rod to the volumetric flask
• Make up the solution to the mark with distilled water using a dropping pipette for the last few drops. Ensure that the bottom of the meniscus sits on the line on the neck of the flask
• Invert the flask several times to ensure a uniform solution is achieved

• Rinse equipment (burette with acid, pipette with alkali, conical flask with distilled water)
• Pipette 25 cm3 of alkali into conical flask
• Touch surface of alkali with pipette( to ensure correct amount is added) 
• Add acid solution from burette
• make sure the jet space in the burette is filled with acid 
• Add a few drops of indicator and refer to colour change at end point
  Phenolphthalein - pink (alkali) to colourless (acid):use if NaOH is used
  Methyl orange - yellow (alkali) to red (acid): use if HCl is used
• Use a white tile underneath the flask to help observe the colour change 
• Add acid to flask whilst swirling the mixture and add acid dropwise near the end point 
• Note burette reading before and after addition of acid 
• Repeat titrations until at least 2 concordant results are obtained- two readings within 0.1 of each other
• Take a mean titre using only the concordant titre results
• Safety: Acids and Alkalis are corrosive (at low concs acids are irritants) - Wear eye protection and gloves. If there are any spillages, wash immidiately
• A conical flask is used instead of a beaker because it reduces spillages and is easier to swirl

• An Ideal Gas is one that occupies negligible space, have no interactions with one another and consequently obey the gas laws exactly
• The Ideal Gas equation combines several other gas laws (Boyles/Charles) into one equation

PV = nRT

P = Pressure (Pa)

V = Volume (m3)

n = Number of Moles (mol)

R = Ideal Gas Constant (8.31 JK-1mol-1)

T = Temperature (K)

• To convert from Celsius to Kelvin, you add 273

• Set up a gas syringe connected to a delivery tube and bung that can fit in a conical flask and ensure the syringe is set to zero.
• Pipette a known volume and concentration of hydrochloric acid into the conical flask.
• Record the temperature and pressure of the room
• Add a known mass of calcium carbonate to the flask and quickly add the bung to the flask.
• Record the final reading on the syringe
• Repeat the experiment at least three times and take an average reading
• Calculate the moles of gas produced by rearranging the ideal gas equation
• Potential errors in using the gas syinge include:
  Gas escaping before the bung is inserted
  The syringe gets stuck
  Some gases such as CO2 are soluble in water so the true amount of gas isn't measured
• On a diagram don't leave gaps where gas could escape and draw measurements on the syringe to show that measurements can be made.

• The uncertainty of a reading is +/- 0.5 of the smallest scale reading
• % uncertainty is calculated by: uncertainty/Measurement made on apparatus x 100
• To calculate the maximum % uncertainty, the sum of all the apparatus % uncertainties are added together
• The uncertainty for a burette is double because two measurements are made on it (an initial and a final reading)
• Uncertainties can be reduced by:
  Using a more accurate balance or larger mass - when weighing a solid
  Weighing the sample before and after addition and calculating a difference to ensure a more accurate measurement of mass being added
• The percentage difference can be calculated to determine how accurate results are
  (Experimental Value - Actual Value)/Actual Value x 100

% Yield = Actual Yield / Theoretical Yield x 100

• Percentage yield can be lowered by:
  Incomplete reactions
  Side reactions
  Losses on transfers of substances
  Losses during purification stage

% Atom Economy = Mass of Useful Products / Mass of all Reactants x 100

• Sustainable chemistry requires chemists to design processes with high atom economies that minimise production of waste products
• Reactions with only one product are most desired as they have atom economies of 100% e.g. Hydrogenation of ethene to form ethane
• If there are waste products, the economics of the process can be improved by selling the bi-product for other uses

• All Salts containing Group 1 elements and Ammonium are soluble
• All Salts containing Nitrates are soluble
• Most Salts containing halides (Cl-, Br- or I-) Except those containing Ag+, Pb2+ or (Hg2)2+
• Most Salts containing sulphate ions are soluble Except Ca, Ba, Pb, Ag and Sr
• Most Salts containing hydroxides are slightly soluble Except Transition Metals and Aluminium
• Most Salts containing carbonates are insoluble Except for Group 1 Metals and Ammonium
• Most Salts containing phosphates are insoluble Except for Group 1 Metals and Ammonium
• Most Salts containing sulpides are insoluble Except for Group 1 Metals, Ammonium and Ca, Sr and Ba

ACID + BASE --> SALT + WATER

• HCl + NaOH --> NaCl + H2O

ACID + METAL --> SALT + HYDROGEN

• 2HCl + Mg --> MgCl2 + H2
• Observations: effevescence and squeaky pop test

ACID + CARBONATE --> SALT + WATER + CARBON DIOXIDE

• H2SO4 + CaCO3 --> CaSO4 + H2O + CO2
• Observations: effervescence and limewater turns cloudy

• The most common strong acids are Hydrochloric, Sulphuric and Nitric
• The most common bases are Metal Oxides, Metal Hydroxides and Ammonia

• More reactive metals will displace less reactive from their compounds
  Mg + CuSO4 --> Cu + MgSO4
  Mg + Cu2+ --> Cu + Mg2+
• A halogen that is a strong oxidising agent will displace a halogen that has a lower oxidising power from one of its compounds
  Cl2(aq) + 2Br-(aq) --> 2Cl-(aq) + Br2(aq)
• A precipitate is a solid insoluble salt
  Pb(NO3)2(aq) + 2NaCl(aq) --> PbCl2(s) + 2NaNO3 (aq)
• An ionic equation can ahow the ions that are reacting and leave out spectator ions. Spectator ions are ions that neither change state nor change oxidation number
  Pb2+(aq) + 2Cl-(aq) --> PbCl2(s)
  H+(aq) + OH-(aq) --> H2O(l)

Ionic Substances

• Positive Ions come first
• If there are multiple charges, use roman numerals to specify e.g. Iron(II)
• If the negative ion is monatomic, add the suffix -ide to the nonmetal e.g. Chloride
• If the negative ion is polyatomic and contains oxygen, the suffix -ate is used e.g. Sulphate
  CuSO4 = Copper(II) Sulphate

Covalent Substances

• If the substance contains multiple atoms of one element, the prefixes mono-, di-, tri- and the suffix -ide are used e.g. Carbon Dioxide
• If the substance is an acid (starts with a H) and contains a halogen, then nomenclature follows Hydro___ic Acid e.g. Hydrochloric Acid
• If the substance is an acid and contains a polyatomic ion, the nomenclature follows
  ____ate = _____ic Acid e.g. Sulphate = Sulphuric Acid
• Exceptions to the nomenclature are H2O (water) and NH3 (Ammonia)

• A Hazard is a substance or procedure that has the potential to do harm
• Risk is the probability or chance that harm will result from the use of a hazardous substance or procedure
• Hazardous substances in low concentrations or amounts will not pose the same risks as the pure substance
• Irritants - Cause irritation of skin/eyes - dilute acids and alkalis - wear goggles and gloves
• Corrosive - Burns through skin or surfaces - stronger acids and alkalis - wear goggles and gloves
• Flammable - Easily catches fire - Organic substances - keep away from naked flames
• Toxic - a substance that is poisonous - Some Inorganic substances e.g. lead or mercury - wash hands after use
• Oxidising - Provides oxygen to make other substances burn more fiercely - Potassium Manganate(VII) - Keep away from flammable or easily oxidised substances