OCR Gateway B Module C5: How Much?

Molar mass:

• The unit for an amount of a substance is the mole
• Molar mass (measured in g/mol) is the mass of 1 mole of a substance. 1 mole contains 6 x 10²³ particles of the substance (the same number as the number of carbon atoms in 12g of carbon)
• Number of moles = mass of chemical (g) ÷ molar mass
• Molar mass (g/mol) = relative formula mass (Mr)
• Mass is conserved in a chemical reaction, so there is the same number of atoms and same mass on each side of the → sign

Empirical formula:

• The empirical formula is the simplest whole number ratio, so the empirical formula for C₆H₁₂O₆ would be CH₂O
• It can be calculated from a set of masses as follows, e.g. for 0.05g H, 0.8g S and 1.6g O:
  

  Instruction	H	S	O
  Write the mass of each element	0.05g	0.8g	1.6g
  Get the RFM for each element from the periodic table	1	32	16
  Calculate the number of moles (mass ÷ RFM)	0.05	0.025	0.1
  Divide by the smallest (in this case 0.025)	2	1	4
  This is the ratio of the elements in the empirical formula	H2	S	O4

  
  H₂SO₄
• The percentage of an element in a compound is calculated with (total mass of element ÷ RFM of compound) x 100
• Calculating the empirical formulae from percentage composition:
  

  Instruction	H	S	O
  Known percentage composition	2.04%	32.65%	65.31%
  Use this to get the mass of each in 100g acid	2.04g	32.65g	65.31g
  Get the RFM for each element from the periodic table	1	32	16
  Convert to number of moles (mass ÷ RFM)	2.04 mol	1.02 mol	4.08 mol
  Divide each by the smallest	2	1	4

  
  H₂SO₄

Concentration:

• Dilution is important in areas such as food preperation (avoiding strong tastes), medicine (avoiding giving overdoses) and baby milk (must not harm the baby)
• To calculate how much water to add to a concentrated solution for a certain concentration, use this formula: volume of water to add = ((original concentration ÷ desired concentration) - 1) x starting volume. For example, if you have 10 cm³ of 1 mol/dm³ acid and need it to be only 0.5 mol/dm³, you need to add (1/0.5 - 1) x 10 = 10 cm³ of water to the acid
• One decimetre cubed (dm³) = 1000 cm³
• Number of moles = concentration x volume

Sodium and salt:

• Sodium ions are essential in diets for nerve responses and water balance, but too much can cause high blood pressure
• Some food labels list the salt content as 'sodium' and some as 'salt'
• Salt mass = 2.5 x sodium mass

Titrations:

• An acid is added to an alkali to get salt and water
• In a titration, a burette, conical flask, pipette and a pipette filler are used
• Acid is placed in the burette, and alkali in the conical flask. Acid is slowly added to the alkali until the end point (neutralisation) is reached, seen with a sudden change in the colour of the indicator
• A pH titration curve can be plotted by measuring pH at set intervals whilst acid is added
• The concentration of an acid/alkali can be determined from titration results, using the following:
  - number of moles = concentration x volume
  - concentration = number of moles ÷ volume
  - volume = number of moles ÷ concentration
• The number of moles of alkali will equal the number of moles of acid

Indicators:

• Litmus is blue in alkali, but red in acid
• Phenolphthalein is pink in alkali, colourless in acid
• Screened methyl orange is green in alkali, pink in acid
• Universal indicator solution contains a mix of many indicators and shows a gradual colour change, so it's not useful for showing the point of neutralisation in titrations

Measuring gas produced by a reaction:

• To do this, one could use a gas syringe, or an upturned measuring cylinder/burette, containing water, in water
• At room temperature and pressure (rtp), 1 mole of gas takes up 24 dm³
• In a reaction, the limiting reactant is used up first, the product is directly proportional to the amount of limiting reactant

Equilibrium:

• Equilibrium is when there is a balance of reactant and product in a two way reaction
• At equilibrium, there is the same rate of reaction in both directions
• In a left position, the concentration of reactants is higher than the concentration of products
• In a right position, the concentration of products is higher than the concentration of reactants
• This only works in a closed system
• At first, the forward rate is faster, but eventually the backward rate of reaction = the forward; equilibrium
• Equilibrium is affected by temperature, pressure and concentration, but not catalysts
• With a higher temperature, less product is made, so the equilibrium moves left
• With a higher pressure, more product is made, so the equilibrium moves right

Example: 2NO₂ ⇌ N₂O₄:

• If some N₂O₄ is removed, the equilibrium moves right - higher yield
• If NO₂ is added, it also moves right
• Increasing pressure moves it right
• La Chatalier's principle states that the equilibrium will move to reduce the effect of change

The Contact process:

• 2SO₂ + O₂ ⇌ 2SO₃
• Conditions: 450 °C, atmospheric pressure, and a vanadium pentoxide catalyst (V₂O₅)
• The forward reaction is exothermic, the high temperature drives the equilibrium left
• However, this is a compromise; high temperature increases rate of reaction

Acids:

• Acids contain hydrogen atoms, in water they ionise to form H⁺ ions
• Strong acids ionise completely when in water, this reaction is fast because lots of collisions occur
• With weak acids, only a few acid molecules ionise
• Strong acid: HCl → H⁺ + Cl^-
• Weak acid: CH₃COOH ⇌ H⁺ + CH₃COO^-
  - Note that this is an equilibrium mixture, unlike with the strong acid. There are fewer collisions so the reaction is slow
• Strong acids include hydrochloric, nitric and sulfuric acid and have a lower pH than weak acids of the same concentration

Electrolysis and conductivity:

• In electrolysis, ions move, so strong acids conduct better as there are more H⁺ ions to move through the liquid
• The same volume of gas will be made if the same amount of reactants are used, because once H⁺ ions are used, more will be produced
• Positive ions migrate to the cathode, so all acids produce hydrogen at the cathode
• Because there are less H⁺ ions to carry the charge, weaker acids are less conductive than strong ones

Strength vs concentration:

• Concentration of an acid is how many moles there are in 1dm³
• Strength is how much an acid ionises in water

Ionic lattices:

• Ionic lattices are formed by ionic substances containing both metal and non-metal ions
• Ions are fixed in position within the lattice
• Different charges are attracted
• Ionic lattices break apart in water

Precipitation reactions:

• When two solutions react to make an insoluble substance, which appears almost instantly as a solid - a precipitate
• AB + CD → AD + CB, e.g. barium chloride + sodium sulphate → barium sulphate + sodium chloride
• Lead nitrate is used to test for halide ions;
  - Cl^- ions form a white precipitate
  - Br^- ions form a cream precipitate
  - I^_ ions form a bright yellow precipitate

Preparing a sample of insoluble salt:

• Using barium chloride + sodium sulphate → barium sulphate + sodium chloride:
  - Mix the reactants to make barium sulphate and sodium chloride
  - Filter, so the precipitate stays in filter paper and the liquid passes through into a beaker
  - Wash the precipitate with distilled water to remove traces of sodium chloride
  - Dry - let the water evaporate

Spectator ions:

• Spectator ions are ions which are not in the reaction
• They aren't included in ionic equations

Ionic equations: