Chemical Analysis

• Usual definition of purity = when a substance has had nothing added to it (it's in its natural state)
• Chemical definition of purity = how close a substance is to being a pure substance
• Pure substance = a substance that only contains one compound or chemical throughout - it is not mixed with anything else

• Testing for purity:
• a chemically pure substance will boil or melt at a specific temperature
• you can test the purity of a sample by measuring its boiling or melting point and comparing it with the boiling or melting point of a pure substance (can be found in a data book)
• the closer the measured value is to the true value, the purer the sample
• Impurities may lower the melting point and increase the melting range
• Impurities may increase the boiling point and increase the boiling range

• Formulations = useful mixtures with a precise purpose made by following a formula ('recipe')
• every component in a formulation is in a carefully measured quantity as part of a fixed ratio and contributes to the properties of the formulation so it can perform its intended function
• e.g. fuels, cleaning agents, paints, medicines, alloys, fertilisers and food

Separation method used to separate the substances (typically dyes) in a mixture - the substances can the be identified

All chromatography has 2 phases:

• Mobile phase - the solvent - this moves up the paper
• Stationary phase - the paper - this does not move

• During chromatography, the susbtances in the mixture move continuously between the mobile and stationary phases - an equilibrium is formed between them
• The mobile phase (the solvent) moves through the stationary phase (the paper), carrying anything dissolved in it through the paper as well
• How quickly a chemical moves through the stationary phase depends on how it's 'distributed' between the two phases
• Substances that spend more time in the mobile phase will move further throught the paper
• A mixture will normally separate out into its components on the stationary phase
• A pure substance will only ever form one dot - it contains no other substances
• The amount of time the molecules of the substances spend in each phase depends on their solubility in the solvent and their attraction to the paper
• There is more on chromatography in the practicals booklet
  

Chlorine - put damp litmus paper into the gas - if it is bleached white, the gas is chlorine

Oxygen - put a glowing splint into a test tube of the gas - if the splint relights, the gas is oxygen

Carbon Dioxide - the gas should be bubbled through or shaken with aqueous calcium hydroxide solution (limewater) - if the limewater turns milky/cloudy, the gas is carbon dioxide

Hydrogen - hold a lit splint at the open end of a test tube of the gas - if the gas burns quickly with a 'squeaky pop', the gas is hydrogen

Carbonate ions = CO­₃^2-

When a metal carbonate reacts with an acid, it produces a salt, water and carbon dioxide

The carbon dioxide produced turns limewater (calcium hydroxide) cloudy

Example Equations:

• calcium carbonate + hydrochloric acid à calcium chloride + water + carbon dioxide
• CaCO₃ + 2HCl à  CaCl₂ + H₂O + CO₂
• metal carbonate + acid à salt +water + carbon dioxide

Equation in limewater:

• calcium hydroxide + carbon dioxide à calcium carbonate + water
• calcium hydroxide + carbon dioxide à calcium carbonate + water
• Calcium carbonate is always formed, and it is insoluble making the solution cloudy

Sulphate ions = SO₄^2-

HCl must first be added to the sulphate solution to remove any unwanted ions (like carbonate ions)

The reaction between a sulphate solution and barium chloride solution forms barium sulphate - which can be identified as a white precipitate (has to be in the presence of HCl)

Example Equations:

• sodium sulphate (aq) + barium chloride (aq) à sodium chloride (aq) + barium sulphate (s)
• Na₂SO₄ + BaCl₂ à 2NaCl + BaSO₄

Ionic Equation:

• SO₄^2- + Ba²⁺ à BaSO₄

Halide ions = Cl^-, Br^-, I^-

Nitric acid must be added to the solutions first, to remove any unwanted ions (like carbonates)

When a halide ion solution reacts with silver nitrate solution, a precipitate is formed - its colour depends on the halide anion in it (chlorine is white, bromide is cream, iodide is yellow)

Ionic Equations:

• Chloride: Cl^-(aq) + Ag⁺ (aq) à AgCl (s) // silver chloride is a white precipitate
• Bromide: Br^- (aq) + Ag⁺ (aq) à AgBr (s) // silver bromide is a cream precipitate
• Iodide: I^- (aq) + Ag⁺ (aq) à AgI (s) // silver iodide is a yellow precipitate

Some positive metal ions burn with a characteristic colour - so it is possible to identify various metal ions by heating them and seeing what colour flame they burn with:

• Lithium ions, Li⁺, burn with a crimson/red flame
• Sodium ions, Na⁺, burn with a yellow flame
• Potassium ions, K⁺, burn with a lilac flame
• Calcium ions, Ca²⁺, burn with an orange-red flame
• Copper ions, Cu²⁺, burn with a green flame

In a sample containing a mixture of metal ions, some of the flame colours may be masked - this can make it less useful than flame emission spectroscopy which can be used to test for multiple ions in mixtures

• Many metal hydroxides are insoluble and precipitate out of solutions when formed - some of these hydroxides have identifiable characteristic colours
• If you add sodium hydroxide solution to a metal ion solution, a coloured precipitate may be formed which will allow you to identify the metal ion in the solution:

• Calcium, Ca²⁺, white precipitate, Ca²⁺ (aq) + 2OH^- (aq) à  Ca(OH)₂ (s)
• Copper (II), Cu²⁺, blue precipitate, Cu²⁺ (aq) + 2OH^- (aq) à Cu(OH)₂ (s)
• Iron (II), Fe²⁺, green precipitate, Fe²⁺ (aq) + 2OH^- (aq) à Fe(OH)₂ (s)
• Iron (III), Fe³⁺, brown precipitate, Fe³⁺ (aq) + 3OH^- (aq) à Fe(OH)₃ (s)
• Magnesium, Mg²⁺, white precipitate, Mg²⁺ (aq) + 2OH^- (aq) à  Mg(OH)₂ (s)
• Aluminium, Al³⁺, white precipitate (redissolves), Al³⁺ (aq) + 3OH^- (aq) à Al(OH)₃ (s)

• Calcium and magnesium both form white precipitates, but can be distinguished using a flame test, where calcium will burn with an orange-red flame and magnesium with a white flame
• The aluminium precipitate begins white, but redissolves in excess NaOH to form a colourless solution

• When an ion is placed in a flame, the electrons heat up and gain energy.
• The ion transfers energy as light as its electrons drop back to their original energy levels.
• The light given out is usually not visible to the human eye, but it can be detected by a spectroscope.
• A prism is used to split the light into a pattern/combination of wavelengths, forming a spectrum, and each ion has its own distinctive line spectrum - as the combination of wavelengths depends on the the charge and the electron arrangement of the ion, and no two ions share both of these characteristics
• The intensity of the spectrum indicates the concentration of that ion in the solution it is in
• So line spectrums can be used to identify ions and calculate their concentrations

• Pros:
• very fast (tests can be automated)
• very accurate
• very sensitive (can work with very small samples)

• Cons:
• expensive to carry out (cost of machinery)
• requires trained operators
• can only identify ions when compared to known data