CHEMICAL CHANGES

Acids and alkalis

• For example, stomach acid is around pH 2, and bleach is around pH 12

• Indicators are chemical dyes that change to laine with the pH colour

  • Universal indicator is the most common, and changes to the same colours as above 🙂

  • A pH probe and meter give a number reading, and are more accurate and precise

• Acids are any substance that forms an aqueous solution with a pH less than 7

• A base is any substance with a pH greater than 7

  • Alkalis are any substances that dissolves in water to form a solution with a pH more than 7

• Neutralisation reactions are when you react an acid and a base to form a solution CLOSER to pH 7 (pure water)

  • They form a salt and water

  • HCl + NaOH \rightarrow NaCl + H₂O

• Acids I need to know

  • Hydrochloric acid is HCl

  • Sulfuric acid is H₂SO₄

  • Nitric acid is HNO₃

• Alkalis I need to know

  • Sodium hydroxide is NaOH

  • Calcium carbonate is CaCO₃

Titration practical

• A titration is an experimental technique used to find an unknown concentration of an acid or alkali

• Neutralising an alkali with acid

  • Use a pipette to add 25cm³ of alkali to a conical flask

  • Add a few drops of an indicator (can’t be universal indicator) and place on a white tile

  • Fill a burette with acid and note the starting volume

  • Slowly add acid from the burette to the conical flask, swirling slowly, so that it is evenly distributed

  • Stop when the end point is reached - the acid has neutralised the alkali and the indicator has changed colour

  • Note the final volume reading and calculate the total acid added

  • Repeat until you get concordant results - within 0.10 cm³ of each other

  • Calculate a mean

• Indicators that can be used:

  • Litmus is RED in acids, and BLUE in alkaline solutions

  • Phenolphthalein is COLOURLESS in acids, and PINK in alkaline solutions

  • Methyl orange is RED in acid, and YELLOW in alkaline solutions

  • Universal indicator can’t be used as it gradually changes, and you can’t clearly see where it has reached the end point

Strong and weak acids

• Acids are ionised in aqueous solutions to release H⁺ ions

  • For example, HCl exists as H⁺ and Cl^- ions

• Strong acids ionise completely - all particles will dissociate

  • Like HCl or H₂SO₄

• Concentration is the amount of acid per amount of volume

  • The amount of H ions per unit of volume

• Weak acids don’t fully ionise - their reactions are reversible and stop reacting when they reach an equilibrium (net reactions)

  • Like CH₃COOH - ethanoic acid

• pH is a measure of concentration of H⁺ ions in a solution

  • The higher the concentration, the lower the pH

• HCl is a strong acid and could easily be pH0 at most concentrations

  • If it was a weak acid, to be pH0, it would have to be very very concentrated

Neutralisation reactions

• Metal oxide + acid \rightarrow salt + water

  • Metal oxides end in O

• Metal hydroxides + acid \rightarrow salt + water

  • Metal hydroxides end in OH

• Metal carbonates + acid \rightarrow salt + water + carbon dioxide

• To make a soluble salt, we have to react an insoluble base with an acid

  • We gently heat an acid, and slowly and a base a bit at a time until it stopes dissolving

    • This shows the acid is neutralised and that the base is in excess

  • Then filter the excess base, and the product is dissolved soluble salt

  • To form crystals from this solution, we can gently heat it to evaporate any water (at less than 60 degrees)

    • When crystals begin to form, stop heating and leave to cool

    • This will cause more crystals to form, which can be filtered and dabbed dry

The reactivity series and displacement

• Elements below carbon can be reduced with carbon (in metal oxides)

• Reactivity is how easily something forms a positive ion

• Metals + acid \rightarrow Salt + hydrogen

  • Reaction gets less violent down the reactivity series

  • Causes bubbles, explosions and temperature changes

• Metals + water \rightarrow Metal hydroxide + hydrogen

  • Only in the most reactive metals

• Displacement example:   

  • Mg + FeSO₄ \rightarrow MgSO₄ + Fe

    • Magnesium displaces iron, as it is more reactive

Separating metals from metal oxides

• Oxidisation is the process of losing electrons (gaining oxygen)

  • 2Mg + O₂ \rightarrow 2MgO

• Reduction is the process of gaining electrons (losing oxygen0

  • 2MgO \rightarrow 2Mg + O₂

• OILRIG - Oxidisation is loss, reduction is gain

• Most metals are fairly reactive

  • They reactive with oxygen (from the air) to form metal oxides

  • If they are really unreactive, they don’t react, like gold

• To get pure metals, we want to reduce them to remove the oxygen

• This is done by reacting metal oxides with carbon (displacement)

  • Oxygen will react with carbon and leave behind the pure metal

  • This only works with metals less reactive than carbon

  • It is cheap and quick

  • For example, 2Fe₂O₃ + 3C \rightarrow 4Fe + 3CO₂, when mining for iron

• More reactive metals are extracted via electrolysis, which uses lots of energy and is expensive

The process of electrolysis

• Electrolysis involves using to break down electrolytes to form elements

• - ions are attracted to the anode (+ electrode)

  • They are then discharged, and become atoms

• + ions are attracted to the cathode (- electrode)

  • They are then discharged to become atoms, and produce either a solid or a gas

• At the cathode, Pb²⁺ + 2e- \rightarrow Pb, reduction which forms solid lead

• At the anode, 2Br^- \rightarrow Br₂ + 2e-, oxidisation which pairs up and forms gas (covalent bonds)

• Electrons at the anode that are loss are passed through the wire to the cathode, where they are given off to the positive ions to form atoms

Electrolysis of aqueous solutions

• Electrolysis is used to split compounds into their elements

• Soluble compounds can be dissolved in water to become aqueous, so charged particles are free to flow

• Because the compounds are dissolved in water, both the ions of the compound, and H and HO ions will be present

• For example:

• The cathode attracts positive ions - Na⁺ and H⁺

  • Only the less reactive will be discharged at the cathode

  • As hydrogen is less reactive than chlorine, hydrogen atoms will be formed, and given off as hydrogen gas

• The anode attracts negative ions - OH^- and Cl^-

  • If a halide is present, it will be discharged. If not, OH^- will be discharged

  • Halides are any of group 7 - F, Cl, Br, I and At

  • Chloride is a halide, so chloride is discharged, and forms chlorine gas at the anode

• This leaves NaOH in the beaker as a solution

Oxidation and reduction in terms of electrons

• Oxidation is loss, reduction is gain

• Reduction and oxidisation reactions both take at the same time - redox reactions

  • This is because the electrons that are loss also need to be gained

• Ionic equations are used to show displacement

  • They only show the particles that take part in the reaction and change

  • Ca + Fe²⁺SO₄^2- \rightarrow Ca²⁺SO₄^2- + Fe

    • Ca + Fe²⁺ \rightarrow Ca²⁺ + Fe

    • SO₄^2- is left out as it doesn’t change

• Half equations show the gain and loss of electrons for each element involved

  • Ca \rightarrow Ca²⁺ + 2e-

  • Fe²⁺ + 2e- \rightarrow Fe

  • The charged particles are always placed on the same side

Reactions of acids with metals

• The products of reactions of acids with metals are always a salt and hydrogen

• Acid + metal \rightarrow salt + hydrogen

• They are redox reactions

  • 2H⁺ + Mg \rightarrow Mg²⁺ + H²

  • Magnesium loses electrons - is oxidised

  • Hydrogen gains electrons - is reduced

Soluble salts

• Soluble salts can be prepared by reacting an acid with a suitable insoluble reactant

  • The reactant can be a metal, metal oxide, carbonate or a metal hydroxide

  • The reactant depends on the salt required

    • For example, copper doesn’t react with dilute acids, so can’t be used, and sodium is too reactive to use

• The reaction between an acid and a metal produces hydrogen

  • Hydrogen is flammable, so we usually use metal oxides or carbonates and an acid instead

• Making a salt

  • Add the powdered insoluble reactant to an acid in a beaker, 1 spatula at a time

    • The acid can be gently heated by a Bunsen burner to increase the reaction speed (particles have more energy)

  • Stir between each addition and continue until the powder is in excess and no longer reacts

  • Filter the mixture in the beaker with filter paper and a funnel to remove excess solid - this means the filtrate now only contains the salt and water

  • Heat the solution in a evaporating dish over a water bath

  • Stop when small crystals begin to appear - the solution is now saturated and most of the water has evaporated

  • Leave the solution for 1-2 days at room temperature to allow the rest of the water to evaporate and leave large crystals

  • Dry by gently dabbing with filter paper

Electrolysis of molten ionic compounds

• This is the process of separating elements in insoluble ionic compounds (splitting with electricity)

• In electrolysis, the electrolyte has to be a liquid or aqueous solution that contains an ionic compound - so the ions can move freely

  • CuSO₄ is soluble, and can be dissolve in water to create this

  • However, lead bromide is insoluble and has to be melted and made molten to allow ions to move freely

• PbBr₂ is heated to become molten

• It splits into it’s two ions - Br^- and Pb²⁺

• Br^- is attracted to the positive anode and is discharged as bromine gas

  • Br^- \rightarrow Br + 2e-

  • Two bromines react covalently to form 2Br as a gas

• Pb²⁺ is attracted to the negative anode and is discharged as solid lead

  • Pb²⁺ + 2e- \rightarrow Pb

  • This creates a layer of molten lead in the beaker

• Ions are being oxidised and reduced at the electrodes

  • Cathode = reduction

  • Anode = oxidisation

• Electrons that are lost when Br^- become Br are transported via the wire with current to the anode, where Pb²⁺ gains two electrons

Extracting metals with electrolysis

• Reactive metals are extracted from metal oxides by melting and making them molten compounds

• The cheapest way to reduced a metal from an oxide is with carbon, but it only works with metals that are less reactive than carbon

  • So we use electrolysis instead, which is expensive and needs lots of energy

• Electrolysis only works in ions can freely move through the solution

• For example, Al₂O₃ \rightarrow Al + O₂

  • Aluminium is more reactive than carbon, so electrolysis is used

  • Aluminium oxide is found as a solid, and found mixed with bauxite when mined

    • This needs to become molten 🔥

• First, we need to purify Al₂O₃ from bauxite

• Then, we mix aluminium oxide with cryolite as it lowers the melting point

  • Aluminium oxide has a very high melting point (it was ~ 2000 degrees C), but is lowered by cryolite

  • We then melt this mixture to become molten - lots of energy is required as the melting point is still high

• The electrodes are made of graphite, and the cathode is found around the outside of the steel case

• O^2- is attracted to the anode in the centre

  • At the anode, it is discharge and becomes oxygen gas (makes a pair and covalently bonds once atoms0

  • 2 electrons are lost per oxygen, and they are transferred through the wire to the cathode for reduction

  • 2O^2- \rightarrow O₂ + 4e-

• Al³⁺ is attracted to the cathode around the edge

  • It is discharged and becomes Al

    • 3 electrons are gained from the electrons lost at the anode

    • The aluminium formed pools at the bottom, and leaves via a channel at the bottom to be collected

    • Al³⁺ + 3e- \rightarrow Al

• 2Al₂O₃ (l) \rightarrow 4Al (l) + 3O₂ (g)

DONE!!!