OCR Gateway B Module C5: How Much?
Note: you must be able to recall all formulas on this page, most questions will not give them and they are not at the front of the question paper
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:
Lead nitrate (aq) + sodium iodide (aq) → lead iodide (g) + sodium nitrate (aq)
Pb(NO3)2 + 2NaI → PbI2 + 2NaNO3
Pb2+ + 2I- → PbI2
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