Pros and Cons of Batch and Continuous Processes

INTERMEDIATE 2 CHEMISTRY SUMMARY NOTES Unit Three: Acids, Bases & Metals

1. Acids & Alkalis pH Scale pH is measured using Universal indicator, pH paper or a pH meter. The scale can go from -1 to 15 but most substances lie between 1 and 14. Acids have a pH < 7 Alkalis have a pH > 7 Neutral substances have a pH = 7.

The pH scale is dependent on the concentration of H+(aq) ions and is based on the fact that water is partially ionised. + - H2O(l) H (aq) + OH (aq) Acids and alkalis affect the pH when added to water, so change the H+ ion concentration. In water: Concentration of H+ = Concentration of OH- In acids: Concentration of H+ > Concentration of OH- In alkalis: Concentration of H+ < Concentration of OH-

Soluble metal oxide form alkalis. Soluble non-metal oxides form acids.

All fizzy drinks are acidic because of the carbon dioxide gas dissolved in them. Most cleaning compounds tend to be alkaline. Strong acids are those which are completely dissociated into ions (ie. Every molecule is broken into ions). Weak acids are those which are only partially dissociated into ions. All alkanoic acids (eg. Ethanoic) are weak acids. Hydrochloric, sulphuric and nitric acids are all strong.

Strong alkalis are those which are completely dissociated into ions (ie. Every molecule is broken into ions). Weak alkalis are those which are only partially dissociated into ions. Ammonia solution (ammonium hydroxide) is a weak alkali. Ammonia is also very water-soluble and has a pungent smell (wet nappies). Sodium, potassium and calcium hydroxides are all strong. Strong vs Weak Strong acids have a higher ion concentration than a weak acid of the same concentration. This means they:  React faster  Have a higher electrical conductivity  Have a lower pH However, they are neutralised by the same amount of a base.

Effect of Dilution Diluting an acid increases its pH towards 7. Diluting an alkali decreases its pH towards 7. Diluting both reduces their electrical conductivity.

Concentration and The Mole For solutions: N = cV where N = number of moles (mol) c = concentration in moles per litre (mol l-1) V = volume in litres (l) 2. Salt Preparation

Reactions of Acids Neutralisation reaction: moves the pH of an acid to 7, producing water. Base: a substance which neutralises an acid. Alkali: a soluble base. Salt: ionic compound formed from the negative ion of the acid and the positive ion from the neutraliser (usually metal). Acid Type of salt Hydrochloric … chloride Nitric … nitrate Sulphuric … sulphate Ethanoic … ethanoate

1. acid + alkali (metal hydroxide) → salt + water 2. acid + metal oxide → salt + water 3. acid + metal carbonate → salt + water + carbon dioxide 4. acid + metal (MAZINTL) → salt + hydrogen 1 → 3 are neutralisations, while 4 is a displacement.

Spectator ions: Ions which do not change during the reaction so don’t actually take part in the reaction.

Acid rain: caused by sulphur dioxide, nitrogen dioxide and carbon dioxide. Affect buildings, structures and wildlife. Farmers use lime to neutralise the effects on soils and lakes.

Fertilisers: water-soluble compounds containing one of the 3 essential elements needed for healthy plants – nitrogen, phosphorus and potassium (NPK). Can be made by neutralisation of acids by ammonia. Precipitation: formation of an insoluble salt from 2 solutions. How you know a precipitation reaction will take place:

1. Use the solubility data on page 5 of the databook. 2. Take the names of the 2 soluble chemicals you start with: For example: lead nitrate + potassium iodide 3. Swap the 2 positive ions to make 2 new substances: That is: lead iodide + potassium nitrate 4. If one of these compounds is insoluble, then a precipitation reaction will occur! How to spot a precipitation reaction equation: 1. Check the state symbols. 2. The 2 reactants must both be (aq) and one product must be (s).

Example: Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq)

Titration Calculations At the end point (neutralisation) of a titration: Number of H+ ions = Number of OH- ions So the calculation is: pcV (acid) = pcV (alkali) where p(acid) = number of H ions in formula p(alkali) = number of OH ions in formula c = concentration and V = volume

3. Metals

A cell is the simplest form of battery and consists of 2 different metals joined in a circuit, separated by an electrolyte.

An electrolyte is used to complete the circuit and is usually a solution of an ionic compound. This allows charge to flow. The voltage produced in a cell varies as follows: 1. The further apart the 2 metals are in the Electrochemical Series, the higher the voltage. 2. The higher the concentration of the electrolyte, the higher the voltage.

Electrons flow through the wires of the circuit from the metal higher up the Electrochemical Series to the one lower down, eg. Zn → Cu.

Displacement Reactions These occur if the solid metal is higher up the Electrochemical Series than the metal in the salt solution. The metals will then swap places. The metal higher up would rather be in the ionic form so gives its electrons to the metal in solution.

Oxidation and Reduction OILRIG Oxidation Is Loss (of electrons) Reduction Is Gain (of electrons)

A Redox reaction is one where reduction and oxidation occur at the same time, eg. Displacements, electrolysis and making electricity in a cell. Redox Equations The number of electrons in the 2 ion-electron equations must be the same, so that they cancel each other out.

1 Zn(s) → Zn2+(aq) + 2e- 2 Ag+(aq) + e- → Ag(s)

Equation 2 has only one electron moving so we must multiply this by 2 before we can add the 2 equations together to get the redox equation.

1 Zn(s) → Zn2+(aq) + 2e- 2x2 2Ag + (aq) + 2e - → 2Ag(s) add Zn(s) + 2Ag+(aq) → Zn2+(aq) + 2Ag(s)

Reactions of Metals With Oxygen With Water With Acid Magnesium Potassium Magnesium Zinc Sodium Aluminium Iron Calcium Zinc Copper Magnesium (when hot) Iron Nickel Tin Lead Extraction of Metals Metals are extracted from their ores. This is a reduction process. Metal Method of Extraction potassium Electrolysis aluminium zinc Heating with carbon or carbon monoxide copper silver Heat alone

Extraction of iron is done in a blast furnace. As well as the iron ore, carbon is added (turned into carbon monoxide gas for more efficient extraction) and limestone removes impurities.

Corrosion This is a reaction where the surface of a metal is turned into a compound. This is an oxidation reaction of the metal.

Rusting is the corrosion of iron. This process can be investigated using ferroxyl indicator. This turns blue in the presence of Fe2+ ions (the first stage of rusting)

2 things are needed for rusting to occur: 1. Oxygen (from air) 2. Water Rusting is speeded up by: 1. Salt 2. Acid 3. Attaching a metal lower in the Electrochemical Series

Preventing Corrosion: 1. Physical protection (paint, plastic, metal stops air and water getting to the metal) 2. Sacrificial protection (a metal higher in the Electrochemical Series is attached) 3. Direct Electrical Protection (attached to a negative terminal supplying electrons)

In sacrificial protection, the metal higher up gives its electrons up to the iron. It is corroded to save the iron.

Iron can be covered by another metal by electroplating. It is attached to the negative terminal and dipped into a solution containing ions of the other metal.

Galvanising is the covering of iron by zinc (can be done by electroplating or just dipping in liquid zinc). This offers BOTH physical and sacrificial protection as zinc is higher in the Electrochemical Series.

On the other hand, tin-plating offers only physical protection and if cracked would then speed up rusting as tin is below iron.