C5 OCR Gateway

• Created by: allie_99
• Created on: 30-03-15 14:55

The Mole

• one mole of atoms or molecules of any substance will have a mass in grams equal to the relative formula mass for that substance (e.g. Carbon Ar = 12. so 1mole of carbon is 12g)
• the relative atomic mass of an element is the average mass of an atom of the element compared to the mass of 1/12 of an atom of carbon-12.

finding moles in given mass Mr - relative formula mass, Ar - atomuc mass:

• Number of moles = mass in g/Mr
• write down balanced symbol equation
• write down mass under formula (x if unknown)
• write down Mr of substances under that
• fill in to formula above and work out
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Reacting masses

You can use Moles to calculate masses in reactions:

• write down balanced equation
• calculate number of moles (moles=mass/mr)
• look at the ratio, these will be the bigger number infront in the symbol equation
• calculate the mass of unknown
• e.g. calculate mass of aluminium oxide formed when 135g of aluminium is burned in the air.
• 4Al + 3O2 - 2Al2O3
• 135/27 = 5
• 4Al:2Al2O3 - so half the number of moles produced. so 5 moles of Al will react to get 2.5 moles of Al2O3
• 2.5 x 102 = 255g

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Empirical Formula

• Empirical formula - simplest whole number ratio of each type of atom in a compound
• For example, the empirical formula of ethane is CH3
• The empirical formula and molecular formula can be the same for some compounds. For example, they are the same for carbon dioxide CO2 and methane CH4 because the numbers in their molecular formulae are already in their simplest whole number ratios.

Method:

• write each element's symbol
• write each mass in g(or%)
• write the Ar
• find the number of moles (mass/Mr)
• divide by the smallest number (to get ratio)
• write down the formula e.g.. 1:1 = MgO 2:1 = Mg2O
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Concentration

• the more solute you dissolve in a given volume, the more crowded the solute molecules are and the more concentrated the solution.
• measured in mol/dm3 or g/dm3(1 litre=1000cm3 = 1dm3)
• concentration = number of moles ÷ volume
• use the volume and concentration of one reactant to calculate the moles
• use the chemical equation to find the moles of the other reactant
• calculate the volume or concentration as required in the question
• converting: conc (g/dm3) = conc(mol/dm3) x Mr
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Percentage Composition

you can work out how much of one particular element in a substance

Percentage of element in substance = (mass of element / mass of whole compound) x 100

• e.g. work out the percentage by mass of nitrogen in ammonium nitrate (NH4NO3)
• calculate the Mr of compound using Ar for the elements
• multiply the Ar of the element in question by the number of its atoms in the compound
• work out the percentage using numbers these two answers
• Mr of NH4NO3 =80
• mass of N = 28 (appears twice in formula NH4NO3)
• percentage = 28/80 x 100 = 35%

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Dilution and GDA

scientists need to dilute accurately so that they can carry out tests on them.

• if too low or high, machines may give an innacurate reading or may be unable to detect
• dilute by adding water then mixing. this can be calculated by:
• volume of water to add = (starting concentration / target concentration) x starting volume

GDA (guideline daily amounts)

• manufacturers often include GDA tables on their food lables to show amount of energy, sugar, fat, sturated fat and salt provided
• amount listed may not always be amount you eat as
• amounts are given per 100g and you may eat more/less
• may add other things (e.g. milk to cereals)
• can estimate how much salt by mass of sodium
• find ratio of NaCl Mr to Na Ar
• multiply by amount of Na
• sodium present may not all come from NaCl, there may be other compounds (preservative sodium nitrate) so may be overestimate
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Titrations

method used to find the concentration of an acid or alkali using a neutralisation reaction

• Use the pipette and pipette filler to add 25 cm3 of alkali to a clean conical flask.
• Add a few drops of indicator and put the conical flask on a white tile (so you can see the colour of the indicator more easily).
• Fill the burette with acid and note the starting volume.
• Slowly add the acid from the burette to the alkali in the conical flask, swirling to mix.
• Stop adding the acid when the end-point is reached (the appropriate colour change in the indicator happens).
• Note the final volume reading.
• Repeat steps 1 to 5 until you get consistent readings (because):
• increases level of accuracy of the titre and to irradicate anomalous results so the mean titre can be calculated
• The difference between the reading at the start and the final reading gives the volume of acid (or alkali) added. This volume is called the titre.

The same method works for adding an alkali to an acid - just swap around the liquids that go into the conical flask and burette.

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Titration Indicators

• If universal indicator is used, the colour changes gradually through a range of colours.
• single indicator like litmus or phenolphthalein gives a sharp end-point where the colour changes suddenly so more appropriate to use
• litmus - red in acidic solutions and blue in alkaline solutions
• phenolphthalein - pink in alkaline solutions and colourless in acidic solutions
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pH curves

neutralisation - acid = alkali = salt +water

Acid to Alkali - slowly at first as acid is added to the alkali; rapidly at the end-point (the point where the alkali is completely neutralised); slowly again once excess acid is being added

Alkali to acid - slowly at first as alkali is added to the acid; rapidly at the end-point (the point where the acid is completely neutralised; slowly again once excess alkali is being added

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Concentration of Titrations

e.g. titration carried out with 0.1 mol/dm3 HCl and 25cm3 of an unknown concentration of NaOH. the mean titre of acid is 24cm3. what is the concentration of NaOH?

• convert all volumes to dm3 (i.e. divide by 1000)
• 24cm3 = 0.024 dm3
• 25cm3 = 0.025dm3
• amount of HCl = concentration x volume
• 0.1 x 0.024 = 0.0024 mol
• Looking at balanced equation, 1 mol of HCl reacts with 1 mol of NaOH. so 0.0024 mol of HCl will react with 0.0024 mol of NaOH
• HCl + NaOH - NaCl + H2O
• use formula (concentration = mol / volume) to work out concentration
• 0.0024/0.025 = 0.096 mol/dm3
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Measuring Gas

• Gas Syringe
• Syringe connected to the reaction container by tubing. As the reaction takes place, gas fills the syringe and pushes plunger out.
• Accurate – nearest cm 3
• If reaction is too vigorous the plunger could blow out the end of syringe
• Upturned Measuring Cylinder/ burette
• Measuring cylinder upside-down filled with water. Delivery tube is led from reaction container into cylinder. As reaction occurs, gas fills cylinder (upward displacement is using a burette instead)
• Burette – more accurate (nearest 0.1 cm3)
• Not good for collecting substances that dissolve in water (ammonia)
• Mass of gas produced (scales)
• Reaction container placed on the balance. As the gas is released the mass will decrease
• Really accurate (0.01 g)
• Released into the room
• If only small change, may not be picked up on school scales
• one mole of gas occupies a volume of 24 dm3 at rtp (room temperature and pressure) -  volume = moles x 24
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Following Reactions

• reactions stop when one reactant is used up (limiting reactant)
• any other reactats are in excess
• amount of product formed is directly proportional to amount of limiting reactant
• double the limiting reactant means double the volume of gas

Faster Rates of reaction are shown by steeper curves

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Equilibrium

• A reversible reaction can go in both directions - it can go forwards and backwards.
• at equilibrium the rate of the forward reaction is equal to the backwards reaction.
• only reached if it takes place in a "closed" system so none of the products or reactants can escape (e.g. stoppered flask or beaker of liquid where reactants stay in solution)
• Position of equlibrium can be on the right or the left
• left = lots of reactants and not a lot of products
• right = lots of products and not a lot of reactant
• three things can change position of equilibrium
• temperature - if exothermic reaction - moves to left. if endothermic - moves to right
• concentration - increase moves to opposite side. decrease moves to decreasing side
• pressure - increase moves to side with less moles
• adding a catalyst speed up forward and backward reaction by same amount, so reaching equlibrium faste but you end up with same amount of product as you would without catalyst
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the contact process

• used to make sulfuric acid in 3 steps:
• sulfur + oxygen → sulfur dioxide S(l) + O2(g) → SO2¬(g)
• sulfur dioxide + oxygen ⇌ sulfur trioxide 2SO2(g) + O2(g) ⇌ 2SO3(g)
• sulfur trioxide + water → sulfuric acid H2O(l) + SO3(g) → H2SO4(aq)

Conditions are carefully chosen to get a higher yield (more product) for equilibrium reaction (2SO2(g) + O2(g) ⇌ 2SO3(g)

• Temperature
• Forward reaction is exothermic, so reducing temperature increases yield.
• Reaction will be slow as of low temperature, so 450­­oC is used (high enough to give reasonable rate without decreasing rate too much)
• Pressure
• High pressure would increase yield of sulphur trioxide
• High cost to high pressures so not worth is as only small increase so Atmospheric pressure (1 atmosphere) is used as its already to the right
• Catalyst
• Increases rate of reaction by using vanadium pentoxide (V2O5)

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Strong and Weak Acids

• Acids ionise in water to produce hydrogen ions, H+
• strong acids ionise completely in water
• HCl → H+ + Cl hydrochloric acid
• weak acids do not fully ionise in water. this is a reversible reaction. since only a few Has released, the equilibrium lies on the left
• CH3COOH ⇌ CH3COO + H+ ethanoic acid
• pH is a measure of concentration of H+ ions in solution (measured with universal indicator or by how fast it reacts)
• acid strength tells you what proportion of acid molecules ionise in water whereas concentration is how many moles of acid there are in a given volume
• Strong acids are better electrical conductors
• more H+ ions to carry the charge through solution.
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Reaction of Acids

strong acids react faster than weak acids

Magnesium reacts with acids to produce a magnesium salt and hydrogen:

• magnesium + hydrochloric acid → magnesium chloride + hydrogen
• magnesium + ethanoic acid → magnesium ethanoate + hydrogen

reaction with calcium carbonate:

• calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide
• calcium carbonate + ethanoic acid → calcium ethanoate + water + carbon dioxide
• weak acids react slower then strong acids as is releases fewer H+ ions so there is a lower collision frequency betwee the reactants.
• strong acids have all the H+ ions so there is a higher collision frequency
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Precipitation Reactions

• two solutions reacting together to make an insoluble substance (the precipitate)
• involves ions reacting so has to be in solution or molten so they can move to collide
• quick as high collision frequency between the ions
• spectator ions are in the solution and theyre still dissolved afterward but they dont change during the reaction
• barium chloride + sodium sulfate→barium sulfate + sodium chloride
• A white precipitate of barium sulfate forms.
• BaCl2(aq)  + Na2SO4(aq) → BaSO4(s) + 2NaCl(aq) - symbol equation
• Ba2+(aq) + 2SO42-(aq) → BaSO4(s) - ionic equation
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Test for Sulfates and Halides

Test for sulfate ions

• add dilute hydrochloric acid and barium chloride to sulfate ion. white precipitate means original compound was a sulfate
• Ba2+(aq)SO42–(aq) →BaSO4 (s)
• adding hydrochloric acid and barium chloride to potassium sulfate or magnesium sulfate produces a white precipitate.

Test for Halide (chloride, bromide, iodide) ions

• chloride gives white precipitate of lead chloride Pb2+(aq) + 2Cl(aq) → PbCl2(s)
• bromide gives cream precipitate of lead chloride Pb2+(aq) + 2Br(aq) → PbBr2(s)
• iodide gives yellow precipitate of lead chloride Pb2+(aq) + 2I(aq) → PbI2(s)
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Preparing Insoluble Salts

Three main stages are involved:

1. Mixing the required reactant solutions.

2. Filtration to remove soluble impurities.

3. Washing and drying the residue (the insoluble compound that remains in the filter paper).

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