2 Inorganic chemistry

Exam analysis

Most tested topics

(a) Group 1 (alkali metals) – lithium, sodium and potassium

Lithium, sodium and potassium form one family because each reacts with water in the same chemical pattern: the products are hydrogen gas and a solution of the metal hydroxide.

\ce{2M(s) + 2H2O(l) -> 2MOH(aq) + H2(g)}

Shared evidence	What it shows
bubbles or fizzing	hydrogen gas is produced
the metal moves and gets smaller	the metal is reacting and being used up
the final colourless solution is alkaline	a soluble metal hydroxide has formed

The shared products and reaction pattern identify the family; the speed and intensity do not have to be identical. Melting into a ball or producing a flame is not a reliable observation for all three metals.

Reactivity increases down Group 1: lithium is the least reactive of these three metals, sodium is more reactive, and potassium is the most reactive.

Metal	Reaction with water	Evidence from air
lithium	fizzes and moves; usually keeps its solid shape	tarnishes, but least rapidly of the three
sodium	reacts faster and usually melts into a moving ball	tarnishes faster than lithium
potassium	reacts very vigorously and may ignite with a lilac flame	reacts most readily, so it must be kept away from air and moisture

The comparison is valid because the metals undergo similar types of reaction but at different rates or intensities. Faster fizzing, quicker disappearance, easier melting or ignition and more rapid reaction with air are evidence of greater reactivity.

A flame colour alone identifies a particular metal; it does not define reactivity. Compare how readily and vigorously the metals react under similar conditions. These reactive metals are stored under oil to prevent contact with oxygen and water vapour.

Use the established Group 1 pattern to predict an unfamiliar alkali metal: keep the family's shared chemical behaviour, then extend the trend in the correct direction.

Position of unfamiliar metal	Prediction
below potassium, such as rubidium or caesium	more reactive than potassium; reacts with water even faster and more violently
anywhere in Group 1	forms a $+1$ ion and reacts with water to form hydrogen and a metal hydroxide

For caesium, predict the same broad observations as potassium—fizzing, movement and an alkaline solution—but a still faster and more violent reaction. The general equation remains $\ce{2M + 2H2O -> 2MOH + H2}$ .

A trend supports a comparative prediction, not an invented exact value. State 'more reactive than potassium' or 'reacts more vigorously', unless numerical data are supplied; do not guess a precise rate, temperature or flame colour.

Every Group 1 atom has one electron in its outer shell. It reacts by losing that electron to form a $+1$ ion, so the ease of electron loss controls its reactivity.

Atom	Electronic configuration	Occupied shells	Relative ease of losing the outer electron
lithium	2,1	2	hardest of these three
sodium	2,8,1	3	easier
potassium	2,8,8,1	4	easiest

Down the group, atoms have more occupied electron shells. The outer electron is farther from the nucleus and is more shielded by inner electrons, so its attraction to the nucleus is weaker. It is therefore lost more easily, making reactions more rapid and vigorous down the group.

Although nuclear charge also increases down the group, increased distance and shielding outweigh it for the outer electron. The explanation must end with easier electron loss; simply saying that potassium has more electrons does not explain greater reactivity.

(b) Group 7 (halogens) – chlorine, bromine and iodine

At room temperature, chlorine, bromine and iodine show a clear physical trend down Group 7: their colours become darker and their melting and boiling points increase.

Element	Colour	Physical state at room temperature
chlorine, $\ce{Cl2}$	pale green	gas
bromine, $\ce{Br2}$	red-brown	liquid
iodine, $\ce{I2}$	dark grey	solid

The rising melting and boiling points explain the state sequence gas → liquid → solid at the same room temperature. Each element consists of diatomic molecules, so the elemental formula is $\ce{X2}$ , not a single atom or a halide ion.

State the colour of the element in the named state. Solid iodine is dark grey; purple describes iodine vapour, and an aqueous bromine solution may appear orange-brown rather than matching the pure liquid exactly.

Predict an unfamiliar halogen by extending the Group 7 trends in the correct direction and keeping the shared diatomic formula $\ce{X2}$ .

Position	Physical prediction	Reactivity prediction
above chlorine, such as fluorine	paler, lower melting and boiling points; gas at room temperature	more reactive
below iodine, such as astatine	darker, higher melting and boiling points; solid at room temperature	less reactive

The trends support fluorine as a pale yellow gas and astatine as a dark grey or black solid. Their molecular formulae are $\ce{F2}$ and $\ce{At2}$ . Displacement evidence establishes the reactivity direction chlorine > bromine > iodine, which can then be extended upward or downward.

A trend supports a range or comparison, not an invented exact value. Predict 'higher boiling point than iodine' or 'less reactive than iodine' unless numerical data are supplied; do not guess a precise boiling point or shade.

A more reactive halogen displaces a less reactive halogen from a solution containing its halide ions. A less reactive halogen cannot displace a more reactive one.

Halogen added	Chloride ions	Bromide ions	Iodide ions
chlorine	no reaction (same element)	reaction	reaction
bromine	no reaction	no reaction (same element)	reaction
iodine	no reaction	no reaction	no reaction (same element)

\ce{Cl2 + 2Br^- -> 2Cl^- + Br2}\qquad\ce{Br2 + 2I^- -> 2Br^- + I2}

Chlorine displaces bromide and iodide, while bromine displaces only iodide. Iodine displaces neither chloride nor bromide. This evidence gives the order chlorine > bromine > iodine, so reactivity decreases down Group 7. For example, chlorine added to colourless potassium bromide solution forms orange bromine.

Keep the species distinct: chlorine is the halogen $\ce{Cl2}$ , whereas chloride is the ion $\ce{Cl^-}$ . A halogen cannot displace itself, and a colour change is evidence only after the product halogen has been identified.

Every Group 7 atom has seven outer-shell electrons and reacts by gaining one electron to form a $-1$ halide ion. Reactivity therefore depends on how strongly the atom attracts an incoming electron.

Atom	Electronic configuration	Occupied shells	Relative reactivity
fluorine	2,7	2	highest
chlorine	2,8,7	3	lower
bromine	4 occupied shells, 7 outer electrons	4	lower again
iodine	5 occupied shells, 7 outer electrons	5	lowest of these four

Down the group, atoms have more occupied shells. The outer shell is farther from the nucleus and inner electrons provide more shielding, so the nucleus attracts an incoming electron less strongly. Gaining that electron becomes harder, so reactivity decreases down Group 7.

Do not apply the Group 1 explanation here: halogens gain an electron rather than lose one. Nuclear charge increases down the group, but the increased distance and shielding outweigh it for the incoming electron.

(c) Gases in the atmosphere

Dry air is a mixture dominated by nitrogen and oxygen. Its four most abundant gases have approximate percentages by volume that add to about 100%.

Gas	Approximate percentage by volume in dry air
nitrogen, $\ce{N2}$	78%
oxygen, $\ce{O2}$	21%
argon, $\ce{Ar}$	0.9% (about 1%)
carbon dioxide, $\ce{CO2}$	0.04%

In a $100\text{ cm}^3$ sample of dry air, this corresponds approximately to $78\text{ cm}^3$ nitrogen and $21\text{ cm}^3$ oxygen, with less than $1\text{ cm}^3$ argon and a much smaller carbon dioxide volume.

'Dry air' excludes variable water vapour. The percentages are approximate: 78% nitrogen and 21% oxygen are not 79% nitrogen and 20% oxygen unless the question explicitly permits that rounding.

To determine the oxygen percentage, trap a known volume of air and react its oxygen completely with a substance while the other main gases remain unreacted. The decrease in gas volume represents the oxygen removed.

Reactant	How oxygen is removed	Evidence reaction is complete
wet iron wool or filings	iron rusts and binds oxygen in a solid oxide	gas volume stops decreasing
heated copper	air is passed repeatedly over hot copper, forming black copper(II) oxide	repeated volume readings become constant
burning phosphorus	phosphorus consumes oxygen and forms a solid oxide	burning stops and the apparatus cools before the final reading

\text{oxygen percentage}=\frac{\text{initial gas volume}-\text{final gas volume}}{\text{initial gas volume}}\times100

If the trapped air starts at $100\text{ cm}^3$ and ends at $79\text{ cm}^3$ , the decrease is $21\text{ cm}^3$ , so oxygen is $21/100\times100=21\%$ by volume.

Use the total initial air volume, including connecting tubes, when required. Compare readings at the same temperature and pressure after cooling; leaks, incomplete reaction or a hot final gas volume make the result unreliable.

Combustion in oxygen forms an oxide. Magnesium, hydrogen and sulfur give different observations and products, so each reaction must be identified precisely.

Element	Observation in oxygen	Product
magnesium	intense bright white flame; white solid forms	magnesium oxide, $\ce{MgO}$
hydrogen	pale blue flame; water vapour forms and may condense	water, $\ce{H2O}$
sulfur	blue flame; colourless choking gas forms	sulfur dioxide, $\ce{SO2}$

\ce{2Mg + O2 -> 2MgO}\qquad\ce{2H2 + O2 -> 2H2O}\qquad\ce{S + O2 -> SO2}

A white magnesium product is a solid or ash, not a precipitate. Hydrogen combustion forms water, not hydrogen peroxide; sulfur normally forms sulfur dioxide under these conditions.

Thermal decomposition breaks one compound into simpler substances by heating it. Many metal carbonates form a metal oxide and carbon dioxide.

\ce{metal carbonate ->[heat] metal oxide + carbon dioxide}

When green copper(II) carbonate is heated, it forms black copper(II) oxide and carbon dioxide: $\ce{CuCO3(s) ->[heat] CuO(s) + CO2(g)}$ .

Evidence	Conclusion
green solid becomes black	copper(II) oxide has formed
gas turns limewater milky/cloudy	carbon dioxide has formed

The gas is released because the carbonate decomposes; it is not combustion, because oxygen is not a reactant in the equation. 'Thermal' means heat is supplied, not that the reaction necessarily gives heat out.

Carbon dioxide is a greenhouse gas. It absorbs some infrared radiation emitted by Earth's surface and re-emits it, reducing the rate at which energy escapes to space.

If the amount of atmospheric carbon dioxide increases, more outgoing infrared radiation can be absorbed. This strengthens the greenhouse effect and may raise average global temperatures, contributing to climate change.

Climate change can alter rainfall patterns and increase the likelihood of some extreme conditions; warming also contributes to melting land ice and sea-level rise. These are consequences of a changing climate, not the definition of a greenhouse gas.

The natural greenhouse effect keeps Earth warm enough for life. The concern is an enhanced effect caused by increasing greenhouse-gas concentrations. Carbon dioxide may contribute to climate change; it is not correct to describe it as the sole influence on climate.

A practical way to estimate oxygen in air is to trap a measured column of air with wet iron wool. Water is required for rusting, and oxygen is removed from the trapped gas as solid rust forms.

Stage	Action
1	Push wet iron wool into one end of a graduated tube and record the initial trapped-air length or volume.
2	Invert the open end in water so the air remains sealed; leave the apparatus until the reading becomes constant.
3	Record the final air length or volume at the same temperature and pressure.
4	Subtract final from initial; divide by the initial value and multiply by 100.

\text{oxygen percentage}=\frac{84-69}{84}\times100=17.9%\approx18%

Use excess finely divided iron to provide enough surface area, make airtight connections and wait for a constant reading. A value below 21% may mean not all oxygen reacted; leaks, temperature or pressure changes, or inaccurate meniscus readings can shift the result.

Wear eye protection and handle rusty or sharp iron wool with tools. Keep the apparatus stable and do not force a sealed plunger; the volume change is caused by oxygen removal, not by nitrogen being produced.

(d) Reactivity series

A more reactive metal reacts faster and more vigorously under the same conditions. Compare like with like: use equal-sized clean metal samples, the same volume and concentration of water or dilute acid, and the same temperature.

Test	Evidence of greater reactivity	Typical product pattern
cold water	faster bubbles, faster disappearance and a larger temperature rise	metal hydroxide + hydrogen
steam	faster hydrogen production and oxide formation	metal oxide + hydrogen
dilute hydrochloric or sulfuric acid	faster hydrogen bubbles and faster metal loss	salt + hydrogen

\ce{Mg + 2HCl -> MgCl2 + H2}\qquad\ce{Zn + H2SO4 -> ZnSO4 + H2}

If magnesium gives hydrogen faster than zinc in the same dilute acid, magnesium is more reactive. A metal below hydrogen, such as copper, does not normally release hydrogen from dilute hydrochloric or sulfuric acid.

Do not compare results from different acid concentrations, temperatures or surface areas. 'More bubbles in total' is not by itself a fair rate comparison, and no visible reaction with cold water does not prove that a metal cannot react with steam.

A metal displaces another metal from its compound only if the added metal is more reactive. A reaction therefore places the added metal above the displaced metal; no reaction places it below, provided the test conditions are suitable.

System	Example	Reactivity conclusion
metal + metal oxide	$\ce{2Al + Fe2O3 -> Al2O3 + 2Fe}$	aluminium is more reactive than iron
metal + aqueous metal salt	$\ce{Zn + CuSO4 -> ZnSO4 + Cu}$	zinc is more reactive than copper

For zinc in copper(II) sulfate, a brown copper coating forms and the blue solution becomes paler as copper(II) ions are removed. Observations identify that a reaction occurred; the equation identifies which metal displaced which.

A metal cannot displace itself. In a salt solution, compare the metals—not the whole salts or ions as if they were metals. Some metal-oxide reactions need heating to overcome activation energy; heating does not reverse the reactivity rule.

The reactivity series runs from metals that lose electrons most readily at the top to the least reactive metals at the bottom.

Position	Metal	Symbol
1	potassium	K
2	sodium	Na
3	lithium	Li
4	calcium	Ca
5	magnesium	Mg
6	aluminium	Al
7	zinc	Zn
8	iron	Fe
9	copper	Cu
10	silver	Ag
11	gold	Au

Use the order to predict displacement: zinc displaces copper from copper(II) sulfate because zinc is above copper; silver does not displace copper because silver is below it. The same order explains why gold is often found uncombined.

Hydrogen and carbon are often inserted as useful reference points, but they are not metals in the specified eleven-metal list. Preserve the exact order of potassium, sodium and lithium at the top, and copper, silver and gold at the bottom.

Iron rusts only when both oxygen and water are present. Rust is hydrated iron(III) oxide; it is a corrosion product rather than a simple layer of pure iron oxide.

Tube	Available conditions	Result	What it shows
iron + water + air	water and oxygen	rust forms	both together permit rusting
iron + boiled water under oil	water, but oxygen excluded	no rust	oxygen is required
iron + dry air with drying agent	oxygen, but water removed	no rust	water is required

Boiling removes dissolved air from water and the oil layer prevents oxygen re-entering. A drying agent removes water vapour from air. Each control changes one required condition while keeping iron present.

Air alone is not a complete explanation: dry air contains oxygen but lacks water. Water alone is also insufficient when dissolved oxygen has been removed and kept out.

Rust prevention either keeps oxygen or water away from iron, or makes a more reactive metal oxidise instead of the iron.

Method	How it protects	What happens if scratched
paint, oil, grease or plastic	barrier blocks oxygen and water	exposed iron can rust
galvanising with zinc	zinc coating is a barrier and zinc is more reactive than iron	zinc still oxidises in preference to iron
sacrificial protection	attached magnesium or zinc loses electrons instead of iron	protection continues while sacrificial metal remains connected

Choose barriers for surfaces that can remain continuously coated. Use galvanising or sacrificial protection where damage is possible, because the more reactive metal continues to protect exposed iron.

Zinc does not protect iron because it is unreactive; it protects because it is more reactive and is oxidised first. Rusting applies specifically to iron and steel, although other metals can corrode.

Oxidation and reduction occur together in a redox reaction. Oxidation is gain of oxygen or loss of electrons; reduction is loss of oxygen or gain of electrons.

Term	Oxygen description	Electron description
oxidation	gains oxygen	loses electrons
reduction	loses oxygen	gains electrons
oxidising agent	supplies oxygen to, or accepts electrons from, another substance	is itself reduced
reducing agent	removes oxygen from, or donates electrons to, another substance	is itself oxidised

\ce{Zn -> Zn^{2+} + 2e^-}\qquad\ce{Cu^{2+} + 2e^- -> Cu}

In $\ce{Zn + Cu^{2+} -> Zn^{2+} + Cu}$ , zinc loses electrons and is oxidised, so zinc is the reducing agent. Copper(II) ions gain electrons and are reduced, so $\ce{Cu^{2+}}$ is the oxidising agent.

The agent causes the other species to change and undergoes the opposite change itself. Do not call an electron donor an oxidising agent, and do not describe reduction only as 'removing a substance'—state oxygen loss or electron gain.

Investigate magnesium, zinc and iron with dilute hydrochloric acid or dilute sulfuric acid by comparing hydrogen-production rates under controlled conditions.

Stage	Action
1	Place equal volumes of the same dilute acid at the same concentration and temperature in labelled vessels.
2	Clean and add equal moles, masses or surface areas of magnesium, zinc and iron; keep the chosen measure consistent.
3	Start timing immediately and measure gas volume at regular intervals with a gas syringe, or time collection of a fixed hydrogen volume.
4	Repeat and compare initial gradients or the time to the fixed volume: faster hydrogen production means greater reactivity.

\ce{Mg + 2HCl -> MgCl2 + H2}\qquad\ce{Fe + H2SO4 -> FeSO4 + H2}

Expected order is magnesium faster than zinc, and zinc faster than iron. Confirm hydrogen only on a small collected sample with a lighted splint: a squeaky pop is positive.

Wear eye protection, use dilute acids and small metal samples, keep flames away while hydrogen is being produced, and point apparatus away from people. Do not stopper a vessel unless gas can escape into a correctly fitted syringe.

(e) Extraction and uses of metals

Most metals in the Earth's crust occur chemically combined with other elements in ores. Extraction separates the useful metal from its compounds and the rest of the ore.

Metal behaviour	How it is commonly found	Why
reactive metal	as a compound in an ore	it has readily reacted with substances in the environment
unreactive metal	sometimes as the uncombined, native element	it is less likely to have formed a compound

Gold is very unreactive, so it may be found native. A metal being present in an ore does not mean the ore is made only of that metal; it contains metal compound and other material.

Unreactive metals are often—not always—found uncombined, and most metals still require extraction from ores. 'Native' means the element is uncombined, not that it is pure enough to use without separation.

The reactivity series determines whether carbon can remove oxygen from a metal oxide. Metals below carbon can be extracted by reduction with carbon or carbon monoxide; metals above carbon require electrolysis of a molten compound.

Example	Position relative to carbon	Suitable method	Reason
iron	below carbon	carbon extraction	carbon removes oxygen from iron oxide
aluminium	above carbon	electrolysis	carbon cannot displace aluminium from aluminium oxide

\ce{Fe2O3 + 3CO -> 2Fe + 3CO2}\qquad\ce{Al^{3+} + 3e^- -> Al}

Electrolysis must use a molten aluminium compound so its ions can move; an aqueous solution would not produce aluminium metal. This objective uses iron and aluminium to illustrate the rule, not to require detailed knowledge of a blast furnace or industrial cell.

To comment on an unfamiliar extraction process, convert the supplied information into linked advantages, disadvantages and a qualified judgment. Detailed recall of a named industrial process is not required.

Supplied information	Evidence-based comment
energy or temperature	lower energy demand may reduce operating cost and fuel use
yield or metal recovered	a higher yield gives more useful metal from the same input
purity	higher purity may suit demanding uses but may need extra processing
raw-material or equipment cost	abundant inputs or simpler equipment may make the process cheaper
emissions, waste, land or water effects	less pollution or waste reduces environmental impact
recycling data	recycling may use less ore and energy, but collection and separation still have impacts

A strong conclusion states the deciding criterion and uses comparative data when supplied: for example, one method may use less energy but recover less metal, so the preferred method depends on whether cost, yield or environmental impact has priority.

Do not invent missing temperatures, costs or emissions. A process is not automatically 'best' because it wins on one measure; acknowledge any material trade-off shown by the information.

A use is explained only when the relevant property is linked to what the object must do. Different steels have different carbon or alloy content, so they should not be treated as one material.

Material	Relevant properties	Property-linked uses
aluminium	low density, corrosion resistant, malleable, conducts heat and electricity	aircraft need low mass; cans and foil need shaping; pans and cables need conduction
copper	excellent electrical and thermal conductor, ductile, corrosion resistant	wires need conduction and drawing; water pipes need corrosion resistance; cookware needs heat transfer
iron	relatively soft and malleable but rusts unless protected	gates, railings and shaped ironwork use its formability
low-carbon (mild) steel	strong, tough and relatively malleable	car bodies, nails and structural shapes need strength with shaping
high-carbon steel	harder, stronger and less malleable	cutting tools, blades and drill bits need a hard edge
stainless steel	hard and corrosion resistant	cutlery, sinks and chemical equipment must resist wear and corrosion

Do not give a property without explaining its relevance: 'aluminium is used for aircraft because it is low density' is linked; 'because it is a metal' is not. Hardness, strength, toughness and malleability are different properties.

An alloy is a mixture of a metal with one or more other elements. The added elements are usually other metals or carbon.

Alloy	Main metal	Other element or elements
steel	iron	carbon
stainless steel	iron	chromium and often nickel, plus carbon
magnalium	aluminium	magnesium

Because an alloy is a mixture, its elements are not combined in one fixed chemical formula. Changing their proportions can change the alloy's properties.

An alloy is not necessarily a mixture of metals only: steel contains the non-metal carbon. It is not a compound, and 'two or more elements' is incomplete unless at least one is a metal.

In a pure metal, equal-sized atoms or positive ions form regular layers that can slide over one another when a force is applied. This makes many pure metals relatively soft and malleable.

Structure	Arrangement under force	Result
pure metal	regular layers of similarly sized particles slide more easily	softer and easier to shape
alloy	differently sized particles distort the regular layers and obstruct sliding	harder to deform

Different-sized particles → disrupted regular arrangement → layers cannot slide as easily → a larger force is needed to change shape. Metallic bonding remains; hardness does not arise because the alloy becomes an ionic or molecular substance.

Hardness means resistance to scratching or permanent shape change. It is not identical to strength or toughness, and the explanation must mention disrupted layers and reduced sliding rather than simply saying that atoms are 'packed tighter'.

(f) Acids, alkalis and titrations

An acid–base indicator changes colour according to whether a solution is acidic or alkaline. Add only a few drops to a small sample and identify the solution from the resulting colour.

Indicator	In an acidic solution	In an alkaline solution
litmus solution	red	blue
phenolphthalein	colourless	pink
methyl orange	red	yellow

When acid is added to alkaline phenolphthalein until the acid is in excess, the colour changes from pink to colourless. The direction matters because the starting solution determines the first colour.

These indicators distinguish acidic from alkaline solutions but do not give a numerical pH. Use the named indicator's own colour pair—methyl orange is not blue in alkali, and phenolphthalein is not pink in acid.

The pH scale from 0 to 14 classifies aqueous solutions. Values below 7 are acidic, 7 is neutral and values above 7 are alkaline.

pH range	Classification
0–3	strongly acidic
4–6	weakly acidic
7	neutral
8–10	weakly alkaline
11–14	strongly alkaline

As pH decreases below 7, a solution is more acidic; as pH increases above 7, it is more alkaline. Thus pH 2 is strongly acidic, pH 5 weakly acidic, pH 9 weakly alkaline and pH 13 strongly alkaline.

Do not reverse the scale: a high pH is alkaline, not strongly acidic. The classification describes the solution at that pH; it does not by itself identify which acid or alkali is present.

Universal indicator gives an approximate pH because it produces a sequence of colours across the pH scale rather than one colour change.

Stage	Action
1	Place a small sample of the aqueous solution in a clean container, or use universal-indicator paper.
2	Add a few drops of universal indicator, or touch the paper with the solution.
3	Compare the final colour immediately with the supplied colour chart.
4	Report the closest pH value or range as an approximation.

Typical progression is red/orange for acidic values, green near neutral and blue/purple for alkaline values. The exact pH must be read from the chart supplied with that indicator.

Universal indicator estimates pH; it does not provide the precision of a calibrated pH meter. It is also unsuitable for an accurate titration end-point because its gradual sequence of colours does not give one sharp change.

In aqueous solution, acids are sources of hydrogen ions, $\ce{H+}$ , while alkalis are sources of hydroxide ions, $\ce{OH-}$ .

Solution	Ions formed in water	Ion responsible for behaviour
hydrochloric acid	$\ce{H+}$ and $\ce{Cl-}$	$\ce{H+}$ makes it acidic
sodium hydroxide solution	$\ce{Na+}$ and $\ce{OH-}$	$\ce{OH-}$ makes it alkaline

\ce{HCl(aq) -> H+(aq) + Cl-(aq)}\qquad\ce{NaOH(aq) -> Na+(aq) + OH-(aq)}

The aqueous condition matters. Dry hydrogen chloride consists of covalent molecules and does not show acidic behaviour without water; hydrogen chloride dissolved in water supplies mobile ions. Write hydroxide as $\ce{OH-}$ , including its charge.

Neutralisation occurs when an acid reacts with an alkali. Hydrogen ions from the acid combine with hydroxide ions from the alkali to form water.

\ce{H+(aq) + OH-(aq) -> H2O(l)}

The complete reaction forms a salt and water. For example, $\ce{HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)}$ : sodium and chloride ions remain in solution as the salt while $\ce{H+}$ and $\ce{OH-}$ form water.

Mixing an acid and alkali does not guarantee a neutral final solution: if either reactant is in excess, the mixture remains acidic or alkaline. Neutralisation describes the reaction, while exact neutral conditions require suitable reacting amounts.

A titration measures the volume of one solution needed to react exactly with a fixed volume of another. A pipette delivers the fixed volume; a burette measures the variable volume accurately.

Stage	Accurate action
1	Rinse the pipette with the solution it will deliver, then transfer $25.0\text{ cm}^3$ to a conical flask. Rinse the flask only with distilled water.
2	Add a few drops of a suitable indicator such as methyl orange or phenolphthalein; place the flask on a white tile.
3	Rinse and fill the burette with the other solution, fill the jet, remove the funnel and record the initial reading at eye level from the bottom of the meniscus.
4	Add from the burette while swirling. Near the end-point, add dropwise until one permanent indicator colour change occurs; record the final reading.
5	Calculate titre = final − initial reading. Repeat with fresh portions until concordant titres within $0.20\text{ cm}^3$ are obtained, then average only concordant accurate titres.

Use a rough titration to locate the end-point, then approach it slowly in accurate repeats. A few drops of indicator do not materially change the volume; excess indicator or universal indicator makes the end-point less reliable.

Wear eye protection, clamp the burette vertically and fill it below eye level. Remove air bubbles from the jet and never read the scale from above or below, because parallax changes the measured titre.

(g) Acids, bases and salt preparations

An ionic compound is soluble if enough of it dissolves in water to form an aqueous solution. Apply the cation rule first when it gives an unconditional result, then check the anion and its exceptions.

Compound group	General rule	Exceptions or limits
sodium, potassium and ammonium compounds	soluble	none in the required rules
nitrates	soluble	none in the required rules
chlorides	soluble	silver chloride and lead(II) chloride are insoluble
sulfates	soluble	barium, calcium and lead(II) sulfates are insoluble
carbonates	insoluble	sodium, potassium and ammonium carbonates are soluble
hydroxides	insoluble	sodium and potassium hydroxides are soluble; calcium hydroxide is slightly soluble

Potassium carbonate is soluble because all common potassium compounds are soluble. Barium sulfate is insoluble because barium is a sulfate exception, so mixing solutions that supply $\ce{Ba^{2+}}$ and $\ce{SO4^{2-}}$ forms a precipitate.

Do not apply only the broad anion rule and ignore an exception. 'Slightly soluble' is the specified classification for calcium hydroxide; it is not grouped with fully soluble sodium and potassium hydroxides.

An acid–base reaction involves transfer of a proton, $\ce{H+}$ , from one particle to another. The proton is not released and left unaccounted for: one species loses it as another gains it.

\ce{H+(aq) + OH-(aq) -> H2O(l)}

In neutralisation, the acid supplies $\ce{H+}$ and the hydroxide ion gains that proton to become $\ce{H2O}$ . Tracking the same hydrogen and its positive charge shows exactly what has transferred.

Proton transfer is not electron transfer. A proton is $\ce{H+}$ ; oxidation and reduction instead track electrons, so an acid–base reaction is not automatically redox.

An acid is a proton donor and a base is a proton acceptor. Identify each role by comparing a species before and after the reaction.

\ce{HCl + NH3 -> NH4+ + Cl-}

Species	Change	Role
$\ce{HCl}$	loses $\ce{H+}$ and becomes $\ce{Cl-}$	acid: proton donor
$\ce{NH3}$	gains $\ce{H+}$ and becomes $\ce{NH4+}$	base: proton acceptor

Do not decide from a formula alone when a reaction is supplied: follow the proton. The acid loses $\ce{H+}$ ; the base does not donate it and does not need to contain $\ce{OH-}$ .

Hydrochloric, sulfuric and nitric acids form different salt families: chlorides, sulfates and nitrates. The other products depend on whether the acid reacts with a metal, a base or a metal carbonate.

Reactants	Products	Example
acid + metal	salt + hydrogen	$\ce{2HCl + Mg -> MgCl2 + H2}$
acid + base	salt + water	$\ce{H2SO4 + CuO -> CuSO4 + H2O}$
acid + metal carbonate	salt + water + carbon dioxide	$\ce{2HNO3 + CaCO3 -> Ca(NO3)2 + H2O + CO2}$

Metal reactions usually show hydrogen bubbles as the metal disappears. Carbonates effervesce because carbon dioxide escapes. With a solid base such as copper(II) oxide, the solid disappears as a salt solution forms when the base is not in excess.

Reactions between nitric acid and metals are explicitly excluded here, so do not apply the salt-plus-hydrogen pattern to them. Preserve formulas, coefficients and the acid-derived salt name.

A base neutralises an acid by accepting protons. Metal oxides, metal hydroxides and ammonia can act as bases; an alkali is specifically a base that is soluble in water.

Substance	Acts as a base?	Alkali?	Reason
copper(II) oxide	yes	no	neutralises acid but is insoluble in water
sodium hydroxide	yes	yes	dissolves in water and supplies $\ce{OH-}$
calcium hydroxide	yes	slightly soluble alkali	its limited solubility still gives an alkaline solution
ammonia	yes	aqueous ammonia is alkaline	accepts $\ce{H+}$ and produces an alkaline solution in water

All alkalis are bases because they neutralise acids, but not all bases are alkalis because many bases are insoluble. Solubility—not whether the name contains 'hydroxide'—controls the alkali label.

A metal oxide can be a base without containing $\ce{OH-}$ in its formula, and ammonia can be a base without being a metal hydroxide. Do not use 'base' and 'alkali' as exact synonyms.

Use an excess insoluble oxide, hydroxide or carbonate—or a suitable metal with hydrochloric or sulfuric acid—to consume all the acid, then remove the unused solid before crystallising the soluble salt.

Stage	Purpose
1. Warm dilute acid gently.	increases reaction rate without boiling away acid
2. Add the insoluble solid in small portions while stirring until some remains.	excess solid shows all acid has reacted
3. Filter.	removes the unreacted excess solid; the salt is in the filtrate
4. Heat the filtrate until near saturation.	evaporates some water without drying the salt completely
5. Cool to crystallise, filter the crystals and dry them.	obtains a pure, dry sample

Choose reactants that form the required soluble salt. For magnesium nitrate, warm nitric acid with excess magnesium oxide, then filter and crystallise the magnesium nitrate solution.

Do not evaporate the solution to dryness when hydrated crystals are required. Filtering before crystallisation is essential: otherwise excess reactant contaminates the crystals.

An acid and an alkali are both soluble, so excess reactant cannot be removed by filtration. First use a titration to find the exact reacting volumes, then remake the salt solution without indicator.

Stage	Action
1	Pipette a fixed alkali volume into a conical flask, add a suitable indicator and titrate with acid to the end-point.
2	Repeat to obtain a reliable reacting volume.
3	Mix the same measured acid and alkali volumes again, but add no indicator.
4	Heat the pure salt solution to near saturation, then leave it to cool and crystallise.
5	Filter the crystals, rinse with a little cold distilled water and dry them.

Sulfuric acid and sodium hydroxide form sodium sulfate solution: $\ce{H2SO4 + 2NaOH -> Na2SO4 + 2H2O}$ . The indicator-free repeat prevents coloured indicator contaminating the crystals.

Do not add excess acid or alkali: both remain dissolved and cannot be filtered off. The first titration finds the proportions; it is not normally the solution crystallised because it contains indicator.

Prepare an insoluble salt by mixing two soluble salt solutions whose ions combine to form the required precipitate.

Stage	Action and reason
1	Select two soluble compounds that supply the required cation and anion; dissolve them separately if starting from solids.
2	Mix and stir the solutions so the insoluble salt precipitates.
3	Filter to collect the precipitate as the residue.
4	Wash the residue with distilled water to remove soluble impurities.
5	Dry between filter papers, in a warm oven or in a desiccator.

\ce{BaCl2(aq) + Na2SO4(aq) -> BaSO4(s) + 2NaCl(aq)}

The desired insoluble salt is the filter residue, not the filtrate. Do not use evaporation or crystallisation: those methods recover a dissolved soluble salt, whereas the precipitate has already formed as a solid.

Prepare hydrated copper(II) sulfate by reacting warm dilute sulfuric acid with excess insoluble copper(II) oxide, then crystallising the blue solution.

\ce{CuO(s) + H2SO4(aq) -> CuSO4(aq) + H2O(l)}

Stage	Observation or purpose
1	Warm dilute sulfuric acid in a beaker; do not boil.
2	Add black copper(II) oxide a little at a time with stirring until black solid remains.
3	Filter off excess $\ce{CuO}$ ; collect the blue copper(II) sulfate filtrate.
4	Heat gently until a cooled drop forms crystals, then leave the solution to cool.
5	Filter the blue crystals, rinse with a little cold distilled water and dry between filter papers.

Excess copper(II) oxide removes all acid and is later filtered off. Do not strongly heat the final crystals or evaporate to dryness, because the target is hydrated copper(II) sulfate.

Lead(II) sulfate is insoluble, so prepare it as a precipitate by mixing two soluble solutions that supply $\ce{Pb^{2+}}$ and $\ce{SO4^{2-}}$ ions.

\ce{Pb(NO3)2(aq) + Na2SO4(aq) -> PbSO4(s) + 2NaNO3(aq)}

Stage	Action
1	Mix measured lead(II) nitrate and sodium sulfate solutions and stir to complete precipitation.
2	Filter the mixture; retain the solid lead(II) sulfate residue.
3	Wash the residue with distilled water to remove soluble sodium nitrate and excess ions.
4	Dry the solid between filter papers or in a warm oven.

Lead compounds are toxic: wear eye protection and gloves, avoid skin contact and dust, use small quantities, wash hands and place all lead-containing residues and liquids in the designated hazardous-waste container.

(h) Chemical tests

A gas test is reliable only when both the test procedure and its distinctive positive result are stated. Use a fresh sample and expose it to the reagent or splint as described; do not identify a gas from colour or smell.

Gas	Test	Positive result
hydrogen	place a lighted splint at the mouth of the container	burns with a squeaky pop
oxygen	insert a glowing splint	the glowing splint relights
carbon dioxide	bubble the gas through limewater	limewater turns milky or cloudy
ammonia	hold damp red litmus paper in the gas	the paper turns blue
chlorine	hold damp blue litmus paper in the gas	it turns red, then is bleached white

Test only a small quantity. Chlorine and ammonia are harmful to inhale, so keep the gas contained and use suitable ventilation; never use smell as the test.

A lighted splint tests hydrogen, whereas a glowing splint tests oxygen. Litmus must be damp for ammonia and chlorine because the gases must dissolve before affecting the indicator.

A flame test identifies certain metal ions from the colour they produce in a non-luminous Bunsen flame. A clean wire and uncontaminated sample are essential because traces of another ion can mask the colour.

Stage	Action and purpose
1	Dip a platinum or nichrome wire loop in hydrochloric acid, then heat it until no flame colour is seen; this removes contamination.
2	Dip the clean loop into the sample, using a little hydrochloric acid to help a solid sample adhere if needed.
3	Place the loop at the edge of a non-luminous blue Bunsen flame.
4	Observe and record the flame colour, then clean the loop before testing another sample.

Wear eye protection, keep the acid away from skin and point the wire away from people. Use a blue flame: a yellow safety flame would hide the test colour.

Cleaning is complete only when heating the loop produces no colour. Reusing an unclean loop can give a false result, especially because sodium contamination produces an intense yellow flame.

After carrying out a clean flame test, compare the observed colour with the required reference colours. The colour identifies the metal ion, not its accompanying negative ion.

Metal ion	Flame colour
lithium, $\ce{Li+}$	red
sodium, $\ce{Na+}$	yellow
potassium, $\ce{K+}$	lilac
calcium, $\ce{Ca^{2+}}$	orange-red
copper(II), $\ce{Cu^{2+}}$	blue-green

For example, an orange-red flame supports the presence of calcium ions, while a lilac flame supports potassium ions. Record the specified colour precisely rather than only calling it red or blue.

Do not confuse lithium red with calcium orange-red, or potassium lilac with copper(II) blue-green. A yellow result may be caused by sodium contamination, so repeat the test with a freshly cleaned loop if the result is unexpected.

Aqueous sodium hydroxide can identify ammonium ions by releasing ammonia and can identify three metal ions by the colour of the hydroxide precipitate formed. Use a fresh sample for each test.

Ion	Procedure	Positive result
ammonium, $\ce{NH4+}$	add aqueous sodium hydroxide and warm gently; test any gas with damp red litmus	ammonia is released and turns the paper blue
copper(II), $\ce{Cu^{2+}}$	add aqueous sodium hydroxide	blue precipitate
iron(II), $\ce{Fe^{2+}}$	add aqueous sodium hydroxide	green precipitate
iron(III), $\ce{Fe^{3+}}$	add aqueous sodium hydroxide	brown or red-brown precipitate

For ammonium ions, warming drives the reaction that releases ammonia; the gas test completes the identification. For the metal ions, the named precipitate colour is the identifying observation.

Do not report only that a precipitate forms: copper(II), iron(II) and iron(III) are distinguished by blue, green and brown respectively. Do not try to identify ammonia by smell.

An anion test requires the correct reagents in the correct order and the expected observation. Acidification removes interfering ions, but the acid must not introduce the ion being tested.

Ion	Procedure	Positive result
chloride, $\ce{Cl-}$	acidify with dilute nitric acid, then add aqueous silver nitrate	white precipitate
bromide, $\ce{Br-}$	acidify with dilute nitric acid, then add aqueous silver nitrate	cream precipitate
iodide, $\ce{I-}$	acidify with dilute nitric acid, then add aqueous silver nitrate	pale-yellow precipitate
sulfate, $\ce{SO4^{2-}}$	acidify with dilute hydrochloric acid, then add aqueous barium chloride	white precipitate
carbonate, $\ce{CO3^{2-}}$	add dilute hydrochloric acid and pass the gas through limewater	effervescence; the carbon dioxide turns limewater milky

For halides, nitric acid is used before silver nitrate because hydrochloric acid would add chloride ions and could create a false white precipitate. A carbonate result is completed by identifying the evolved gas as carbon dioxide.

A white precipitate alone is not enough to distinguish chloride from sulfate: the reagent sequence identifies which test was performed. Preserve the cream result for bromide and pale-yellow result for iodide.

Anhydrous copper(II) sulfate tests whether water is present. The anhydrous solid is white and becomes blue when it is hydrated by water.

Stage	Action or observation
1	Place a small amount of white anhydrous copper(II) sulfate on a dry surface.
2	Add the liquid being tested.
3	A change from white to blue is a positive result for water.

Keep the reagent dry before use because moisture from the air can turn it blue and spoil the test. A known dry sample can be used to confirm the starting colour.

This test shows that water is present; it does not show that the liquid is pure water. An aqueous solution containing dissolved impurities would also turn anhydrous copper(II) sulfate blue.

A pure substance changes state at a sharp, fixed temperature. At standard atmospheric pressure, pure water boils at $100\,^{\circ}\mathrm{C}$ and freezes at $0\,^{\circ}\mathrm{C}$ .

Stage	Action and interpretation
1	Measure the atmospheric conditions and choose either boiling point or freezing point.
2	Heat or cool the sample while measuring temperature with a suitable thermometer.
3	A sharp boiling point of $100\,^{\circ}\mathrm{C}$ or freezing point of $0\,^{\circ}\mathrm{C}$ at standard pressure supports that the sample is pure water.
4	A shifted temperature or a change over a range indicates dissolved impurities.

Boiling point depends on atmospheric pressure, so compare the measurement with the expected value for the conditions. Take repeated readings and avoid reading the thermometer while its bulb touches the container.

The anhydrous copper(II) sulfate test cannot establish purity because it responds to water in both pure water and solutions. Purity requires a physical-property measurement with a sharp value.