moles

Define the term molecular formula.

The actual number of atoms of each element present in a molecule.

Define the term empirical formula.

The simplest whole number ratio of atoms of each element present in a compound.

Define the term molar mass.

The mass, in grams, per mole of a substance; units are g mol⁻¹.

Define the term mole.

The amount of substance containing as many particles as there are carbon atoms in exactly 12g of Carbon-12.

Define the term Avogadro constant.

The number of particles per mole of a substance; value is 6.02 × 10²³ mol⁻¹.

Define the term molar gas volume.

The volume occupied per mole of a gas; 24.0 dm³ mol⁻¹ at room temperature and pressure (RTP).

Define the term water of crystallisation.

Water present in a compound giving it a crystalline appearance.

Define the term anhydrous.

A compound with all water of crystallisation removed.

Define the term hydrated.

A compound containing water of crystallisation.

Define the term standard solution.

A solution of known concentration.

State the equation linking moles, mass and molar mass.

moles = mass / molar mass; n = m / M.

State the equation linking moles, concentration and volume in dm³.

moles = concentration × volume; n = c × V.

State the equation linking moles, concentration and volume in cm³.

moles = (concentration × volume) / 1000; n = (c × V) / 1000.

State the equation linking moles and gas volume at RTP.

moles = volume / 24.0; n = V / 24.0 (volume in dm³); moles = volume / 24000 (volume in cm³).

State the ideal gas equation.

pV = nRT; where p = pressure in Pa, V = volume in m³, n = moles, R = 8.314 J mol⁻¹ K⁻¹, T = temperature in Kelvin.

State the equation for number of particles.

Number of particles = moles × Avogadro constant; N = n × Nₐ.

State the equation for concentration in g dm⁻³.

Concentration (g dm⁻³) = concentration (mol dm⁻³) × molar mass.

State the equation for percentage uncertainty in a burette reading.

Percentage uncertainty = (uncertainty × 2 × 100) / titre volume.

State the equation for percentage yield.

Percentage yield = (actual yield / theoretical yield) × 100.

State the equation for atom economy.

Atom economy = (Mᵣ of desired product / sum of Mᵣ of all products) × 100.

State the equation for the dilution factor.

Dilution factor = concentration of stock / concentration of diluted solution; or dilution factor = final volume / initial volume.

State the equation for relative atomic mass from isotopic abundance.

Aᵣ = (isotopic mass × percentage abundance) / 100.

State the units for molar mass.

g mol⁻¹.

State the units for concentration.

mol dm⁻³ or g dm⁻³.

State the units for molar gas volume.

dm³ mol⁻¹.

State the units for the gas constant R.

J mol⁻¹ K⁻¹.

State the units for pressure in the ideal gas equation.

Pascals (Pa).

State the units for volume in the ideal gas equation.

Cubic metres (m³).

State the units for temperature in the ideal gas equation.

Kelvin (K).

State the standard conditions for gas volume measurements.

Room temperature and pressure (RTP): 20°C (293 K) and 101 kPa.

State the value of the molar gas volume at RTP.

24.0 dm³ mol⁻¹ (or 24,000 cm³ mol⁻¹).

State the value of the Avogadro constant.

6.02 × 10²³ mol⁻¹.

State the value of the gas constant R.

8.314 J mol⁻¹ K⁻¹.

State how to convert cm³ to m³.

Divide by 1,000,000 (×10⁻⁶).

State how to convert dm³ to m³.

Divide by 1000 (×10⁻³).

State how to convert °C to K.

Add 273.

State how to convert kPa to Pa.

Multiply by 1000 (×10³).

State how to convert kg to g.

Multiply by 1000.

State how to convert tonnes to g.

Multiply by 10⁶.

State how to convert mg to g.

Divide by 1000.

Explain the four-step method for reacting mass calculations.

Step 1: identify the known substance and the unknown substance; Step 2: calculate moles of the known substance using mass divided by molar mass; Step 3: use the molar ratio from the balanced equation to find moles of the unknown; Step 4: convert moles of the unknown to the required quantity such as mass, volume or concentration.

Explain the four-step method for titration calculations.

Step 1: identify the known solution and the unknown solution; Step 2: calculate moles of the known using concentration multiplied by volume divided by 1000; Step 3: use the molar ratio from the balanced equation to find moles of the unknown; Step 4: convert moles of the unknown to concentration using moles multiplied by 1000 divided by volume.

Explain the four-step method for gas volume calculations.

Step 1: identify the known substance; Step 2: calculate moles of the known; Step 3: use molar ratio to find moles of the unknown gas; Step 4: convert moles to volume using V = n × 24.0 at RTP or V = nRT over p for non-standard conditions.

Explain the method for calculating empirical formula from mass data.

Divide each mass by the relative atomic mass to find moles; divide each mole value by the smallest mole value; convert to whole numbers by multiplying if necessary.

Explain the method for calculating empirical formula from percentage composition.

Treat percentages as masses in 100g; divide each mass by relative atomic mass to find moles; divide by the smallest mole value; convert to whole numbers by multiplying if necessary.

Explain the method for calculating molecular formula from empirical formula.

Calculate the empirical formula mass; divide the given molecular mass by the empirical mass; multiply the empirical formula by this factor.

Explain the method for calculating the value of x in a hydrated salt.

Find moles of anhydrous salt and moles of water lost by heating; divide moles of water by moles of anhydrous salt; the result is the value of x.

Explain how a standard solution is prepared from a solid.

Calculate mass required; weigh accurately by difference; dissolve in a beaker with distilled water; transfer to volumetric flask; rinse beaker and add washings; make up to the mark; stopper and invert to mix.

Explain how a standard solution is prepared by dilution.

Calculate volume of stock solution required; pipette this volume into a volumetric flask; add distilled water to the mark; stopper and invert to mix.

Explain the key steps in an acid-base titration.

Pipette a known volume of solution into a conical flask; add indicator; fill burette with other solution; record initial burette reading; add solution dropwise until indicator changes colour; record final burette reading; calculate titre; repeat for concordant results.

Explain the key equipment needed for a titration.

Burette, pipette, conical flask, white tile, stand and clamp, indicator, volumetric flask.

Explain the key equipment needed to prepare a standard solution.

Weighing balance, beaker, glass rod, funnel, volumetric flask, dropping pipette, weighing boat.

Explain what is meant by concordant titres.

Titres that agree within 0.10 cm³ of each other.

Explain how to calculate the mean titre from titration results.

Select only concordant titres; calculate the average of these values; do not include the trial titre.

Explain the purpose of a trial titration.

To get an approximate endpoint; to determine the rough volume needed; saves time in subsequent titrations.

Explain the purpose of a white tile in a titration.

To provide a white background; makes the colour change of the indicator easier to see; helps identify the endpoint accurately.

Explain the purpose of a pipette in a titration.

To deliver a very accurate and reproducible volume of solution.

Explain the purpose of a burette in a titration.

To allow precise addition of solution; readings can be taken to nearest 0.05 cm³.

Explain why a volumetric flask is used to prepare standard solutions.

It is calibrated to contain a very accurate volume; the solution can be made up precisely to the graduation mark.

Explain the effect of not rinsing the beaker when preparing a standard solution.

The solution would be less concentrated because some solute remains in the beaker.

Explain the effect of filling above the graduation mark when preparing a standard solution.

The solution would be less concentrated because too much water has been added and the volume is too large.

Explain the effect of filling below the graduation mark when preparing a standard solution.

The solution would be more concentrated because not enough water has been added and the volume is too small.

Explain the effect of washing out the pipette with water before use.

The solution would be more dilute because residual water dilutes the sample and there are fewer moles of solute in the pipette volume.

Explain why burette readings have twice the uncertainty of a single reading.

Two readings are taken: initial and final; the titre is calculated from both; the uncertainty doubles.

Explain how to calculate percentage error in a mass measurement.

Percentage error equals maximum error multiplied by number of readings multiplied by 100 divided by the measured value.

Explain the effect of using a more precise balance on percentage error.

Percentage error decreases because the maximum error is smaller.

Explain the effect of using a larger sample mass on percentage error.

Percentage error decreases because the measured mass is larger.

Explain how to ensure all water of crystallisation is removed when heating a hydrated salt.

Heat until constant mass is achieved; repeated heating and weighing until the mass no longer changes.

Explain why the salt must be heated to constant mass.

To ensure all water of crystallisation has been removed; no further mass loss occurs.

Explain why the crucible must be cooled before reweighing.

Hot objects are less massive due to convection currents; weighing a hot crucible gives an inaccurate lower mass.

Explain why the lid should be left off during heating.

Allows water vapour to escape; prevents water condensing back into the crucible.

Explain the effect of incomplete heating of a hydrated salt.

Some water of crystallisation remains; calculated mass of water lost is too small; the value of x calculated is too small.

Explain the effect of the salt decomposing during heating.

Additional mass loss occurs such as CO₂ or SO₂; calculated mass of water lost is too large; the value of x calculated is too large.

Explain the effect of the sample spitting from the crucible during heating.

Some solid is lost; calculated mass of water lost is too large; the value of x calculated is too large.

Explain the effect of soot deposits from the Bunsen burner.

Soot adds mass to the crucible; calculated mass of water lost is too small; the value of x calculated is too small.

Explain the assumptions of the ideal gas model.

Random motion of particles; elastic collisions between particles; negligible size of particles; no intermolecular forces between particles.

Explain how real gases deviate from ideal gas behaviour.

At high pressure or low temperature molecules are closer together; intermolecular forces become significant; molecular volume is no longer negligible.

State the conditions under which ideal gas behaviour is approached.

Low pressure; high temperature.

Explain the relationship between volume and moles of gas at constant temperature and pressure.

Volume is directly proportional to moles; equal volumes of gases contain equal numbers of moles; this is Avogadro's Law.

Explain the term limiting reagent.

The reactant that is completely used up in a reaction; it determines the maximum amount of product formed.

Explain how to identify the limiting reagent in a reaction.

Calculate moles of each reactant; compare mole ratio from balanced equation; the reactant with fewer moles than required is the limiting reagent.

Explain why the actual yield is usually less than the theoretical yield.

Reaction may not go to completion; product may be lost during purification; side reactions may occur.

Explain why a high percentage yield is beneficial.

Efficient conversion of reactants to products; less waste; cost-effective in industry.

Explain why a high atom economy is beneficial.

Less waste is produced; more efficient use of resources; reduces environmental impact.

Explain why addition reactions have 100% atom economy.

Only one product is formed; all atoms from reactants are incorporated into the product.

Explain why substitution reactions often have lower atom economy.

A by-product is formed; some atoms are wasted in the by-product.

Suggest how atom economy can be improved.

Find a use for the by-product; design reactions with fewer by-products; use addition reactions instead of substitution.

Explain the importance of considering both percentage yield and atom economy.

Percentage yield shows efficiency of conversion; atom economy shows efficiency of atom usage; both are important for sustainability.

Explain the term stoichiometry.

The molar ratio between reactants and products as shown by the coefficients in a balanced equation.

Explain why a balanced equation is needed for reacting mass calculations.

The coefficients give the molar ratio; this allows moles of one substance to be converted to moles of another.

Explain why the empirical formula is used instead of molecular formula for ionic compounds.

Ionic compounds do not exist as molecules; the empirical formula shows the simplest ratio of ions.

Explain the difference between empirical and molecular formula for covalent compounds.

Covalent compounds exist as molecules; they can have different empirical and molecular formulae; the molecular formula is a multiple of the empirical formula.

Explain why the Avogadro constant is so large.

Atoms and molecules are extremely small; one mole contains 6.02 × 10²³ particles.

Explain what is meant by one mole of a substance.

The amount of substance containing 6.02 × 10²³ particles; for an element, one mole is the relative atomic mass in grams; for a compound, one mole is the relative molecular mass in grams.

Explain what is meant by one mole of gas at RTP.

The amount of gas that occupies 24.0 dm³; contains 6.02 × 10²³ molecules.

Explain why the molar mass of an element equals its relative atomic mass in grams.

Relative atomic mass is measured compared to 1/12 of Carbon-12; one mole of atoms has this mass in grams.

Explain why the molar mass of a compound equals its relative molecular mass in grams.

Relative molecular mass is the sum of relative atomic mass values; one mole of molecules has this mass in grams.

Explain the use of mass by difference in weighing.

Weigh container plus solid; transfer the solid; reweigh the container; the difference gives accurate mass of solid used; avoids inaccuracies from solid sticking to the container.

Explain why solutions should be inverted after making up to volume.

To ensure the solution is homogeneous; concentration is uniform throughout the volumetric flask.

Explain why distilled water should be used when preparing solutions.

Tap water contains impurities such as dissolved ions that could react with the solute; impurities affect the concentration.