Chemical changes C4

Reactivity Series definition

• A list of metals arranged in order of their reactivity based on how easily they lose electrons from their outer shell to form positive ions. Metals form positive ions most easily= most reactive

Reactivity Series order (Zebra acronym)

• Potassium, Sodium, Lithium (group 1), Calcium, Magnesium (group 2), [Carbon], Zinc, Iron, [Hydrogen], Copper (transition metals)

Metal reaction with Water (or acid) to see how fast/violent reactions are to figure out reactivity series list

Only most reactive metals can react with water to produce metal hydroxide and hydrogen. (Metal + Water → Metal hydroxide + Hydrogen).

Potassium, Sodium, Lithium, and Calcium react vigorously, Magnesium very slightly, Zinc/ iron/ copper do not react with water.

Metal reaction with dilute Acid (or water) to see how fast/violent reactions are to figure out reactivity series list

• Metals above hydrogen react with acids to produce a salt and hydrogen gas (Metal + Acid → Salt + Hydrogen).

Reactions get gradually less violent down reactivity series: Potassium reacts the most explosively with lots of hydrogen, Magnesium lots of bubbles as metal disappears, Zinc/ iron fewer bubbles. Fizzing will stop when reaction has completed.

• Metals below hydrogen do not react with dilute acids.

Other than watching the reactions to see how violent, measure temperature change of the reactions. Most reactive metals= most heat. Fair test: metal same mass/surface area and acid same type/concentration.

Displacement reaction definition

• A reaction where a more reactive metal displaces a less reactive metal from an aqueous solution of its salt e.g Calcium + Iron sulphate→ Calcium sulphate + Iron (as Calcium more reactive than Iron).

Observations: solution fading in colour, new solid forming and temperature rise (displacement reactions are exothermic, the greater the difference in reactivity the larger temperature rise)

Oxidation and Reduction (acronym OILRIG)

• Oxidation is the gain of oxygen and loss of electrons e.g metal + oxygen→ metal oxides

• Reduction is the loss of oxygen and gain of electrons

Unreactive/ reactive metal extraction

• Unreactive metals like gold and platinum are found as pure metals and do not need chemical extraction, as they do not react with oxygen to form metal oxides.

More reactive metals that react with oxygen to form metal oxides and need chemical extractions (with carbon or electrolysis) to get the pure metal back through reduction.

Example: Magnesium + Oxygen → Magnesium oxide. This burns with white flame to form a white powder.

Metal extraction methods

• 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 through a displacement reaction (Metal Oxide + Carbon → Metal + Carbon Dioxide).

For metals more reactive than carbon to extract them from their oxides, use process called electrolysis (requiring lots of energy and expensive)

Redox reaction definition

• A reaction where both oxidation and reduction take place at the same time, as if one substance loses electrons another gains electrons.

Ionic equations example

Only show the particles that take part in the reaction and change in someway.

Example for displacement reaction of Calcium + Iron sulphate→ Calcium sulphate + Iron. The sulphate ions don’t change and are called spectator ions so are removed: Ca + Fe²⁺ + SO₄²⁻ → Ca²⁺ + SO₄²⁻ + Fe

Ca + Fe²⁺ → Ca²⁺ + Fe

Half equations and example

• Show the individual gain and loss of electrons for each of the elements involved.

Example for displacement reaction of Calcium + Iron sulphate→ Calcium sulphate + Iron. Copy the ionic equation and split them into seperate ones for each element: Ca + Fe²⁺ → Ca²⁺ + Fe

Oxidation: Ca → Ca²⁺ then balance charges Ca → Ca²⁺ + 2e⁻

Reduction: Fe²⁺ → Fe then balance charges Fe²⁺ + 2e⁻ → Fe

(Make sure the positive ion and the negative electrons are written on the same side of the arrow so they cancel out to make 0 for metals)

Acids

• Acid is a substance with a pH less than 7 and release hydrogen ions (H⁺) when they dissolve in aqueous solutions (water).

Bases and alkalis

• Base is a substance with a pH greater than 7 (can be insoluble or soluble- alkali). Alkalis (subgroup of bases that are soluble) release hydroxide ions (OH⁻) when they dissolve in aqueous solutions.

pH Scale values

• How acidic or alkaline a solution is from a scale of 0 to 14. pH 0 to 6 is acidic, pH 7 is completely neutral, and pH 8 to 14 is alkaline.

Ways to measure pH

Indictors- chemical dyes that change colour depending on pH e.g Universal indicator

pH probes- dipping probe into a solution the connected pH meter measures the pH. Benefit compared to indicator is more precise and reliable as remove human judgement guessing shades/ colours

Neutralisation reaction

• The chemical reaction between an acid and a base (alkali) Acid + Base→ Salt + Water e.g Hydrochloric acid + Sodium hydroxide→ Sodium chloride + Water

• Ionic equation: H⁺(aq) + OH⁻(aq) → H₂O(l).

Most common acids: Hydrochloric acid (HCl), Sulphuric acid (H₂SO₄), Nitric acid (HNO₃)

Most common bases: Sodium hydroxide (NaOH), Calcium carbonate (CaCO₃)

Neutralisation reaction: with Acid and Metal oxide (base)

• Acid + Metal Oxide → Salt + Water e.g Hydrochloric acid + Sodium oxide→ Sodium chloride + Water (hydrogen and sodium ions combine and hydrogen and oxygen combine. 2HCl + Na₂O → 2NaCl + H₂O

Neutralisation reaction: with Acid and Metal hydroxide (base)

• Acid + Metal Hydroxide → Salt + Water e.g Sulphuric acid + Potassium hydroxide→ Potassium sulphate + Water. H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O

Neutralisation reaction: with Acid and Metal carbonate (base)

• Acid + Metal Carbonate → Salt + Water + Carbon Dioxide e.g Nitric acid + Calcium carbonate→ Calcium nitrate + Water + Carbon dioxide. 2HNO₃ + CaCO₃ → Ca(NO₃)₂ + H₂O + CO₂

Fizzing is observed due to the release of carbon dioxide gas.

Strong Acid definition

• An acid that releases hydrogen ions (H+) by completely ionising (dissociates) in an aqueous solution (e.g. hydrochloric, sulphuric, and nitric acids).

Weak Acid definition

• An acid that releases hydrogen ions (H+) by only partially ionises (dissociates) in an aqueous solution e.g. ethanoic, citric, and carbonic acids (due to its a reversible reaction so more undissociated acid at equilibrium).

Strong vs Weak Acid pH link

• Concentration is how much acid with hydrogen ions there are in a certain volume.

pH depends on the concentration of hydrogen ions in a solution:

At the same concentration, a strong acid will have a lower pH than a weak acid because it has a higher concentration of hydrogen ions (H⁺) that ionise (dissociate) fully.

However, highly concentrated weak acid can have a lower pH than a dilute (less concentrated) strong acid because it releases more hydrogen ions into the volume overall.

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 as it become a stronger acid e.g pH 5 to 3 the hydrogen ion concentration increases by x100 (as 10×10)

Required Practical: Making a Soluble Salt (by a neutralisation reaction copper oxide- insoluble base) Step 1

• Measure a set volume of dilute hydrochloric acid into a beaker and heat it gently over a Bunsen burner.

• Add the insoluble copper oxide in excess, stirring constantly until no more dissolves (neutralised the acid)

Use an insoluble salt instead of soluble as it does not dissolve in water and unreacted solid filtered out easily when we add excess to make sure the acid is completely used up.

Required Practical: Making a Soluble Salt Step 2

• Filter the mixture using a funnel and filter paper to physically remove the excess unreacted copper oxide.

• Pour the remaining pure filtrate solution of copper chloride 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 until it becomes saturated with crystals forming.

• Leave the solution to cool slowly through crystallisation so large salt crystals grow, then filter again 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 (e.g with one known measured volume of an alkali to find out the other unknown volume of an acid for neutralisation).

Then use volume and concentration already known from the alkali and the volume of the acid from the titration, to work out the unknown concentration of the acid.

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 methyl orange 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 (neutralisation).

Required Practical: Titration method Step 3 (Triple Only)

• Record the final volume on the burette and calculate the volume of acid added (titre by the final volume- initial volume)

• Repeat the experiment multiple times until concordant results (readings within 0.10 cm³ of each other) to calculate a mean.

Concordant matters: First one is a rough titration not added into the mean, reduce random errors and anomalies with a mean closer to true value.

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 rather than a gradual one.

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 through electrolysis e.g Aluminium Oxide

• Aluminium oxide is solid and its ions are fixed. A ore bauxite is purified into aluminium oxide, then mixed with cryolite to lower the melting point (also reduce cost by saving energy) and melted into molten aluminium oxide so the ions are free to move and act as an electrolyte.

Aluminium extraction: Carbon electrodes

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

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 reaction

• 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 reaction

• 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 and hydrogen ions are attracted to the negative cathode, since copper is less reactive than hydrogen ions it gains 2 electrons through reduction and form pure copper that is discharged, since it’s a solid it accumulates around cathode.

Aqueous Copper Sulfate: Anode reaction & pH change

• Sulphate and hydroxide ions are attracted to the positive anode, since sulphate is not a halide the hydroxide ions lose an electron through oxidation to form oxygen gas and extra water (4OH⁻ → 2H₂O + O₂ + 4e⁻). Since the electrolyte loses OH⁻ ions and leaves H⁺ ions, the solution becomes more acidic, remaining as sulphuric acid.

Aqueous Sodium Chloride: Cathode reaction

• Sodium and hydrogen ions are attracted to the negative cathode, since hydrogen is less reactive than sodium ions, it gains an electron through reduction, that is discharged to from hydrogen gas that floats off.

Aqueous Sodium Chloride: Anode reaction & pH change

• Chloride ions ans hydroxide ions are attracted to the positive anode, since chloride is a halide it loses an electron through oxidation, that is discharged to from chlorine gas. Since the electrolyte loses H⁺ ions and leaves OH⁻ ions, the solution becomes more alkaline, remaining as sodium hydroxide.

Tips for electrolysis

Use acronyms PANIC (Positive anode and Negative cathode)

Non-metals (start with n) so start as negative ions, metals are positive ions until both are neutral

For half equations the left side should include the gain of electrons, whilst right side should include the loss of electrons.

The total charge on left should equal total charge on right

To help with balancing equations halogens always come in pairs so write them as e.g Br₂

Check on periodic table for the charges or transition metals exam will state with Roman numerals e.g Iron (II) means Fe²⁺. Always write the amount of electrons lost/ gain in written questions