Chemical changes C4

Reactivity Series definition

• A list of metals arranged in order of their reactivity based on how easily they lose electrons to form positive ions.

Reactivity Series order

• Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, [Carbon], Zinc, Iron, [Hydrogen], Copper.

Metal reaction with Water

• Potassium, sodium, lithium, and calcium react vigorously with cold water to produce a metal hydroxide solution and hydrogen gas.

Metal reaction with dilute Acid

• Metals above hydrogen react with acids to produce a salt and hydrogen gas (Metal + Acid → Salt + Hydrogen). • Metals below hydrogen do not react with dilute acids.

Displacement reaction definition

• A reaction where a more reactive metal physically takes the place of a less reactive metal from an aqueous solution of its salt.

Oxidation and Reduction in terms of Oxygen

• Oxidation is the gain of oxygen by a chemical substance. • Reduction is the loss of oxygen from a chemical substance.

Unreactive metal extraction

• Unreactive metals like gold and platinum are found in the Earth's crust as pure native elements and do not need chemical extraction.

Metal extraction using Carbon

• Metals less reactive than carbon (zinc, iron, copper) are extracted from their oxides by heating them with carbon. • The carbon reduces the metal oxide by removing its oxygen (Metal Oxide + Carbon → Metal + Carbon Dioxide).

Oxidation and Reduction in terms of Electrons (OIL RIG)

• Oxidation Is Loss of electrons. • Reduction Is Gain of electrons.

Redox reaction definition

• A reaction where both oxidation (loss of electrons) and reduction (gain of electrons) occur simultaneously at the exact same time.

Half equation definition

• An equation that shows what happens to electrons for a single species during an oxidation or reduction process.

Acids ionisation in water

• Acids produce hydrogen ions (H⁺) when they dissolve in aqueous solutions.

Alkalis ionisation in water

• Alkalis produce hydroxide ions (OH⁻) when they dissolve in aqueous solutions.

pH Scale values

• A scale from 0 to 14 measuring acidity or alkalinity. • pH 0 to 6 is acidic, pH 7 is completely neutral, and pH 8 to 14 is alkaline.

Neutralisation ionic equation

• The chemical reaction between an acid and an alkali. • Ionic equation: H⁺(aq) + OH⁻(aq) → H₂O(l).

Acid + Metal Oxide reaction

• Acid + Metal Oxide → Salt + Water.

Acid + Metal Hydroxide reaction

• Acid + Metal Hydroxide → Salt + Water.

Acid + Metal Carbonate reaction

• Acid + Metal Carbonate → Salt + Water + Carbon Dioxide. • Effervescence (fizzing) is observed due to the release of carbon dioxide gas.

Strong Acid definition

• An acid that completely ionises (splits into ions) in an aqueous solution (e.g. hydrochloric, sulfuric, and nitric acids).

Weak Acid definition

• An acid that only partially ionises in an aqueous solution (e.g. ethanoic, citric, and carbonic acids).

Strong vs Weak Acid pH link

• For a given concentration, a strong acid will have a lower pH than a weak acid because it has a higher concentration of hydrogen ions (H⁺).

pH and H+ concentration scale factor

• As the pH decrease by 1 unit, the hydrogen ion concentration of the solution increases by a factor of 10.

Required Practical: Making a Soluble Salt Step 1

• Measure a set volume of dilute acid into a beaker and heat it gently over a Bunsen burner. • Add the insoluble metal oxide or carbonate in excess, stirring constantly until no more dissolves.

Required Practical: Making a Soluble Salt Step 2

• Filter the mixture using a funnel and filter paper to physically remove the excess unreacted solid residue. • Pour the remaining pure filtrate solution into an evaporating dish.

Required Practical: Making a Soluble Salt Step 3

• Heat the solution gently using a water bath to evaporate some of the water solvent until it becomes saturated. • Leave the solution to cool slowly so large salt crystals grow, then filter and pat dry.

Titration definition & purpose (Triple Only)

• A highly accurate laboratory technique used to measure the exact volumes of acid and alkali that react together to achieve complete neutralisation.

Required Practical: Titration method Step 1 (Triple Only)

• Use a volumetric pipette and pipette filler to safely measure 25 cm³ of alkali into a conical flask. • Add a few drops of a single indicator like phenolphthalein to the flask and place it on a white tile.

Required Practical: Titration method Step 2 (Triple Only)

• Fill a burette with acid, record the initial volume, and slowly add the acid to the conical flask while swirling constantly. • As the end-point approaches, add the acid drop by drop until the indicator undergoes a sharp, permanent colour change.

Required Practical: Titration method Step 3 (Triple Only)

• Record the final volume on the burette and calculate the volume of acid added. • Repeat the experiment multiple times until you achieve concordant results (readings within 0.10 cm³ of each other) to calculate a mean.

Titration Indicator rules (Triple Only)

• You must use a single indicator (like phenolphthalein or methyl orange) rather than Universal Indicator because they show a sharp, sudden colour change at the exact point of neutralisation.

Electrolysis definition

• Splitting apart ionic substances into simpler substances using electricity.

Soluble ionic compounds treatment

• Dissolve it in water to turn into an aqueous solution then add into trough as the water breaks ions apart so they are free to move.

Insoluble ionic compounds treatment

• Melt it to turn into a molten liquid then add into trough as ions are free to move. An insoluble substance is too strong to be broken apart by water so extra heat is needed.

Electrolyte definition

• A liquid or solution that contains an ionic compound where ions are free to move and conduct electricity between two electrodes.

Electrodes definition

• Solid electrical conductors submerged into the electrolyte, made of metal or carbon.

Molten Lead Bromide: Cathode reaction

• Positive lead ions are attracted to the negative cathode. They are discharged, going from a charged ion (Pb²⁺) to a metal atom (Pb) by gaining two electrons, falling to the bottom as a layer of pure molten lead. They are reduced. Half equation: Pb²⁺ + 2e⁻ → Pb.

Molten Lead Bromide: Anode reaction

• Negative bromide ions are attracted to the positive anode. They are discharged from a charged ion (Br⁻) to a neutral atom (Br), by each bromide ion losing one electron, then pair up to form bromine gas and float off. They are oxidised. Half equation: 2Br⁻ → Br₂ + 2e⁻.

Extracting metals more reactive than carbon

• Extracting metal to form pure metals by melting them into molten ionic compounds for electrolysis, used for the metals that we have more reactive than carbon.

Extracting metals less reactive than carbon

• If less reactive than carbon, you can make metals with carbon via reduction with carbon (e.g. 2CuO + C → 2Cu + CO₂).

Aluminium Oxide state issue

• Aluminium oxide is solid and its ions are fixed. It must be melted as bauxite and turned into molten aluminium oxide so the ions are free to move and act as an electrolyte.

Aluminium extraction: Use of Cryolite

• Purified aluminium oxide (Al₂O₃) is mixed with cryolite, which lowers the melting point and reduces cost by saving energy.

Aluminium extraction: Carbon electrodes

• Carbon is used because it can conduct electricity at a cheaper cost and has a high melting point.

Aluminium extraction: Anodes wearing away

• The carbon anodes must be constantly replaced because the oxygen gas can react with the carbon to form CO₂ gas, so it wears away.

Aluminium extraction: Cathode half equation

• Positive aluminium ions are attracted to the negative cathode where they gain three electrons to form pure aluminium atoms and are reduced. Half equation: Al³⁺ + 3e⁻ → Al.

Aluminium extraction: Anode half equation

• Negative oxygen ions are attracted to the positive anode. Each oxide ion loses two electrons to give oxygen atoms in the molecule and are oxidised. Half equation: 2O²⁻ → O₂ + 4e⁻.

Aluminium extraction: Overall equation

• The overall balanced chemical equation is: 2Al₂O₃(l) → 4Al(l) + 3O₂(g).

Aqueous Solution components

• A solution dissolved in water has hydroxide (OH⁻) and hydrogen (H⁺) ions present as water splits them apart (H₂O → H⁺ + OH⁻), along with the metal and non-metal ions.

Aqueous Electrolysis: Cathode Rules

The negative cathode attracts positive metal and hydrogen ions.

• Hydrogen ions are discharged if the metal is more reactive than hydrogen (e.g. sodium). The metal ions are discharged if the metal is less reactive than hydrogen (e.g. copper).

Aqueous Electrolysis: Anode Rules

• The positive anode attracts negative non-metal and hydroxide ions.

If a halide is present, the halide gets discharged.

If a halide isn't present, it is always the hydroxide that gets discharged to form oxygen and water.

Aqueous Copper Sulfate: Cathode reaction

• Copper ion is less reactive than hydrogen ions, so copper ions gain electrons from the cathode via reduction to form pure copper, which is discharged and accumulates around the cathode as a solid copper layer.

Aqueous Copper Sulfate: Anode reaction & pH change

• Sulfate is not a halide, so hydroxide ions lose electrons to the anode via oxidation to form oxygen gas and extra water (4OH⁻ → 2H₂O + O₂ + 4e⁻). Since it loses OH⁻ ions and leaves H⁺ ions, the solution becomes more acidic, remaining as sulfuric acid.

Aqueous Sodium Chloride: Cathode reaction

• Sodium ions are more reactive than hydrogen ions, so hydrogen ions gain electrons from the cathode via reduction to form hydrogen gas, which is discharged and floats off.

Aqueous Sodium Chloride: Anode reaction & pH change

• Chloride is a halide, so chloride ions lose electrons to the anode via oxidation to form chlorine gas, which is discharged and floats off. Since it loses H⁺ ions and leaves OH⁻ ions, the solution becomes more alkaline, remaining as sodium hydroxide.