C5 - How much?

• 1 mole = 6.023 X 10²³

• One mole of any substance will have a mass in grams equal to the relative formula mass for that substance
• Molar mass = mass of one mole  measured in grams per mole

Finding the number of moles with a given mass:

Number of moles = 

.

.

.

Relative Atomic Mass of an element is the average mass of an atom of the element compared to 1/12th of an atom of carbon-12

Using moles to calculate masses:

• write out balanced equation
• calculate number of moles of element with given mass
• look at ratio of moles in equation
• use mass = moles x rfm of compound

Percentage composition by mass of compounds:

...........

Finding empirical formulas:

• list all elements in compound and write experimental masses
• divide each mass by rfm of particular element
• multiply and/or divide by well chosen number to find a simple ratio
• get ratio in its simplest form

• measured in moles per dm3
• 1 litre = 1000cm3 = 1dm3

Concentration = 

.......

........

Converting mol/dm3 into g/dm3:

• work out relative formula mass for solute
• convert concentration in moles into concentration in grams using mass=no. of moles x rfm

Diluting solutions:

• Work out ratio of two concentrations by dividing the smaller concentration by the bigger one to get anumber less than one
• multiply ratio by volume of solution you want to end up with to find out the volume to dilute
• work out the volume of water you need by subtracting the volume to dilute from the total volume

Using sodium content to estimate mass of salt:

• find ratio of sodium chloride's relative formula mass to sodium's relative atomic mass by dividing sodium's ram with sodium chloride's rfm
• multiply this ratio by amount of sodium
• this is probably an overestimate as there are other sodium compounds

Titrations: allow you to find out exactly how much acid is needed to neutralise an alkali

• using pipette and pipette filler, add some alkali (25cm3) to a conical flask and some drops of indicator solution
• fill a burette with acid and do this below eye level
• using the buretter, add acid to alkali small quantities at a time giving the conical flask a regular swirl, do this slowly when you think you're about to reach the end point
• record volume of acid used to neutralise alkali and repeat many times until you get the same result 
• the first titration is a rough titration

Indicators:

• universal indicator estimates the pH of different values
• a single indicator shows sudden colour change

When interpreting graphs:

• there's a sudden change in the pH as the alkali is added which shows the exact amount of acid needed to neutralise the alkali 

Calculating concentration of acid (or alkali) when given the volume of acid and volume and concentration of alkali:

• work out how many moles of known substance - no. of moles = concentration x volume
• write balanced equation of reaction and work out how many moles of th eunknown substance you have
• work out concentration of the unknown substance using formula

Collection method:

• gas syringe - gives volumes of gases accurate to nearest 1cm3
• upturned measuring cylinder - use delivery tube to bubble gas into upside down measuring cylinder - not good for gases that dissolve in water

Measure mass of gas = most accurate method:

• carry out experiment on mass balance
• as gas released, disappearing mass = mass of gas
• has a disdvantage of releasing gas straight into room

One mole of gas:

• one mole of any gas always occupies 24dm3 / 24000cm3 at room temperature and pressure (RTP = 25 degrees and 1 atmosphere)
• number of moles (dm3) = volume of gas / 24

Reactions stop when one reactant is used up:

• reactant that's used up in a reaction that causes the reaction to stop
• amount of product formed is directly proportional to amount of limiting reactant
• more reactant = more reactant particles to take part in reaction = more product particles

Faster rate of reaction:

• rate of reaction increases if you have:
  • more acid
  • more of limiting reactant

A reversible reaction is one where the products of the reaction can themselves react to produce the original reactant

Reversible reactions will reach equilibrium:

• equilibrium = the forward and backward reactions are going at exactly the same rate 
  • both reactions are still hppening but there is no overall effect (dynamic equilibrium) so the concetrations of reactants have reached a balance and won't change
• equilibrium is only reached if the reversible reaction takes place in a closed system where none of the reactants or products can escape
• position of equilibrium can change:
  • to the right - concentration of product is greater than the concentration of reactant
  • to the left - concentration of reactant is greater than the concentration of product
• adding a catalyst does not change the position of equilibrium because it only speeds up the reaction

The position of equilibrium changes with:

• Temperature:
  • increase in temperature = equilibrium tries to decrease temperature and moves in endothermic direction
  • decrease in temperature = equilibrium tries to increase temperature and moves in exothermic direction
• Pressure: - only affects equilibrium involving gases
  • increase pressure = equilibrium tries to reduce it and moves in direction where there are fewer moles of gas
  • decrease pressure = equilibrium tries to increase it and moves in direction where there are more moles of gas
• Concentration:
  • increase concentration of reactants = equilibrium tries to decrease it and shifts to right, making more products
  • increase concentration of products = equilibrium tries to increase it and shifts to left, making more reactants

Used to make sulfuric acid:

• raw materials needed = air, water and sulfur
• sulfur is burned in air to make sulfur dioxide: sulfur + oxygen → sulfur dioxide
  • S(l) + O₂(g) → SO₂(g)
• sulfur dioxide then oxidised to make sulfur trioxide: sulfur dioxide + oxygen ⇌ sulfur trioxide
  • 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)  (reversible)
  • conditions needed are atmospheric pressure, 450 degree temperature and vanadium oxide catalyst
• sulfur trioxide reacted with water to make sulfuric acid: sulfur trioxide + water → sulfuric acid
  • H₂O(l) + SO₃(g) → H₂SO₄(aq)

Conditons of the contact process are chosen carefully:

• reaction: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)    (forward reaction = exothermic)
• temperature:
  • reducing the temperature slows reaction down so compromise temperature of 450 degrees is used to get high yield quite quickly
• pressure:
  • increasing pressure expensive and is not necessary as equilibrium is already on right side
  • atmospheric pressure (1 atmosphere) is used
• catalyst:
  • vanadium pentoxide increases rate of reaction but DOESN'T change the position of equilibrium

With farily high temperature, a low pressure and vanadium pentoxide catalyst, the reaction goes fairly quickly and you get a good yield of aout 99%.

• Strong acids (e.g. hydrchloric acid):
  • ionise completely in water - every hydrogen atom releases a hydrogen ion so there are lots of H+ ions
• Weak acids (e.g. ethanoic acid):
  • do not fully ionise - some hydrogen atoms in compound release hydrogen ions so small number of H+ ions are formed
  • ionisation of weak acid is reversible - equilibrium lies to the left
• pH of an acid or alkali = a measure of the concentration of H+ ions in the solution

Strong acids are better electrical conductors:

• strong acids have a higher concentration of hydrogen ions and so more charge is carried therefore they are better electrical conductors

Strong and weak acids react with reactive metals and with carbonates in the same way:

• both hydrochloric acid and ethanoic acid will react with magnesium to give hydrogen
• both hydrochloric and ethanoic acid with calcium carbonate to give carbon dioxide
• difference = rate of reaction:
  • ethanoic acid will react more slowly than hydrochloric acid of same concentration
  • a weak acid releases few H+ ions but concentration of H+ ions is low compared to a strong acid so collision frequency between reactants is low
  • in a weak acid, all of acid molecules are ionised so the collision frequency is really high

Volume of gas produced depends on amount of acid:

• strong acids react faster than weak acids but amount of product formed will be the same
• the end volume of gas produced is the same however the reaction with weak acids is slower

Precipitation reactions:

• two solutions react together making an insoluble substance
• reactions are extremely quick because there is a high collision frequency
• in a precipitation reaction, the reactants are both solutions but you end up with a solid precipitate which turns the water cloudy

Testing for sulfate ions:

• add dilute hydrochloric acid and barium chloride
• forms a white precipitate of barium sulfate which shw that original compound was a sulfate

Testing for halides (Cl-, Br- I- ions):

• add dilute nitric acid and lead nitrate:
  • chloride gives white precipitate of lead chloride
  • bromide gives cream precipitate of lead bromide
  • iodide gives yellow precipitate of lead iodide

Insoluble salt reactions:

• need ions in solutions so they can move about

Preparing the salt:

• add 1 spatula of lead nitrate to a test tube and fill with distilled water and shake thoroughly and then do same with potassium iodide
• tip two solutions in small beaker and stir to form a salt precipitate
• put folded piece of filter paper into filter funnel and put in conical flask
• pour contents of beaker into middle of filter paper 
• rinse contents of filter paper with distilled water and scrape lead iodide on to some fresh filter paper and leave to dry