كيف يحسب الرقم الهيدروجيني pH للمحاليل؟ How to calculate pH and pOH?

يحسب الرقم الهيدروجيني بالعلاقة: pH = -log[H⁺]، وللمحاليل المائية عند 25°C: pH + pOH = 14. إذا كان pH 7 فالمحلول قاعدي. [الصيغة الرمزية الموزونة: pH = -log[H⁺] | pOH = -log[OH⁻] | pH + pOH = 14 | [H⁺][OH⁻] = 10⁻¹⁴]

R3.2Reactivity 3 — What are the mechanisms of chemical change?

Electron transfer reactions — IB Chemistry SL Revision Notes & Topical Questions

30 min read 46 parts Lesson 2 of 12 1 quiz

📚 IB Chemistry SL Hub / Topical Revision

Curriculum Equation Key قوانين الرقم الهيدروجيني والاتزان الأيوني - pH & Acid-Base Equilibrium FormulasFast Exam Solution

pH = -log[H⁺]  |  pOH = -log[OH⁻]  |  pH + pOH = 14  |  [H⁺][OH⁻] = 10⁻¹⁴

Word Equation & Conditions: الرقم الهيدروجيني = -لوغاريتم تركيز أيون الهيدروجين | pH + pOH = 14

Location & Specifications: الاتزان الأيوني، كيمياء 3 ثانوي، التحصيلي، و AP Chemistry.

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Lesson summary

R3.2 Electron transfer reactions — IB CHEMISTRY SL notes: Identify oxidising and reducing agents, write redox equations, compare cells and electrolysis, and describe organic oxidation and reduction.Oxidising agent = a species that accepts electrons, is itself reduced and has a de…

Redox reaction
Dual behaviour
Zn + Cu2+ → Zn2+ + Cu
Manganate(VII) with iron(II)
Self-indicating
Iodine-thiosulfate

Part 1 of 10Oxidising and reducing agents

Oxidising agent

Oxidising agent a species that accepts electrons, is itself reduced and has a decreasing oxidation number.

Reducing agent

Reducing agent a species that donates electrons, is itself oxidised and has an increasing oxidation number.

Redox reaction

Redox reaction requires both an oxidising agent and a reducing agent.

Dual behaviour

Dual behaviour some species act as either, depending on the partner and conditions, e.g. H₂O₂ oxidises Fe²⁺ but reduces Fe³⁺ (O from −1 to 0).
- (O from −1 to −2)

Zn + Cu²⁺ → Zn²⁺ + Cu

Zn is oxidised and Cu²⁺ is reduced (the oxidising agent).
- (the reducing agent)

Naming

Roman numerals show the oxidation state, as in iron(II) sulfate and iron(III) sulfate.

Part 2 of 10Balancing half-equations and redox titrations

Balance atoms other than O and H.

Balance O with H₂O.

Balance H with H⁺ (acidic solution).

Balance charge with electrons.

Multiply so the electrons cancel, add and simplify.

Balancing half-equations and redox titrations

Manganate(VII) with iron(II)

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O and Fe²⁺ → Fe³⁺ + e⁻ give MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O.

Self-indicating

Self-indicating acidified MnO₄⁻ is purple until the Fe²⁺ is used up,
- then the first permanent pale pink shows the end point; it is used to analyse iron tablets.

Iodine-thiosulfate

I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻, with starch added near the end point (pale yellow) and the blue-black colour vanishing at the end;
- it is used to analyse chlorine in bleach.

Part 3 of 10Relative ease of oxidation and reduction

Relative ease of oxidation and reduction

Groups 1 and 2

Groups 1 and 2 metals are oxidised more easily down the group,
- so they become stronger reducing agents: K reacts more violently than Na,
- and Ca much faster than Mg.

Group 17

Group 17 oxidising power decreases down the group, F₂ > Cl₂ > Br₂ > I₂.

Halogen displacement

Halogen displacement a more reactive halogen displaces a less reactive halide, e.g. Cl₂ + 2KBr → 2KCl + Br₂ (orange) and Cl₂ + 2I⁻ → 2Cl⁻ + I₂ (brown).

No reaction

No reaction iodine cannot oxidise Br⁻ or Cl⁻.

Aqueous colours

Aqueous colours chlorine very pale green, bromine orange or yellow, iodine brown.
- (often colourless)

Reactivity series

Mg, Al, Zn, Fe, Pb, H, Cu, Ag from the strongest to the weakest reducing agent,
- and a more reactive metal displaces a less reactive metal ion, e.g. Mg + CuSO₄ → MgSO₄ + Cu.

Part 4 of 10Acids with reactive metals

General reaction

General reaction acid + metal → salt + hydrogen, e.g. 2HCl + Zn → ZnCl₂ + H₂.

Ionic equation

2H⁺(aq) + Zn(s) → Zn²⁺(aq) + H₂(g), with Cl⁻ as a spectator ion.

Redox roles

Zn is oxidised (0 to +2) and acts as the reducing agent;
- H⁺ is reduced (+1 to 0) and acts as the oxidising agent.

Very reactive metals

Very reactive metals react dangerously and are not usually used.

Unreactive metals

Unreactive metals such as copper, silver and gold do not react with dilute non-oxidising acids.

Strong versus weak acid

Strong versus weak acid a strong acid reacts faster because [H⁺] is higher, with more effervescence, though both give the same final volume of H₂.

Part 5 of 10Voltaic cells

Voltaic (galvanic) cell

Voltaic (galvanic) cell generates a potential difference (EMF) from a spontaneous redox reaction using two half-cells joined by a wire and a salt bridge.

Half-cell

Half-cell a metal in a solution of its own ions with an equilibrium such as Zn ⇌ Zn²⁺ + 2e⁻.

Anode

Anode oxidation, negative; cathode — reduction, positive.

Electron flow

Electron flow from the anode to the cathode through the external circuit only.

Salt bridge

Salt bridge filter paper soaked in an inert electrolyte such as KNO₃ lets anions drift to the anode and cations to the cathode, keeping each half-cell neutral.

Voltmeter

Voltmeter a high-resistance voltmeter draws almost no current, so the reading is close to the maximum EMF.

Cell diagram

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s), with | for a phase boundary, || for the salt bridge and the anode on the left.

Zn-Cu cell

E_cell = +1.10 V, and a larger EMF results when the two metals are further apart in the reactivity series.

Part 6 of 10Primary, secondary and fuel cells

Primary cell

Primary cell single use, e.g. zinc-carbon and alkaline cells, cheap and light with long shelf life but low current.

Secondary cell

Secondary cell recharged by applying a voltage greater than the cell voltage, which reverses the reaction.

Lead-acid

Pb negative and PbO₂ positive in H₂SO₄, Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O, about 2 V per cell and 12 V from six,
giving high current for starting engines but heavy and containing toxic lead and corrosive acid.

NiCd

NiCd about 1.2 V, with toxic cadmium and a memory effect.

Lithium-ion

Li⁺ moves through a polymer electrolyte between a carbon negative electrode and lithium cobalt oxide positive electrode, about 3.5–4.0 V, light with no memory effect,
but capacity fades, lithium is limited and damaged cells can catch fire.

Fuel cell

Fuel cell an alkaline hydrogen cell has anode H₂ + 2OH⁻ → 2H₂O + 2e⁻, cathode O₂ + 2H₂O + 4e⁻ → 4OH⁻ and overall 2H₂ + O₂ → 2H₂O (E = +1.23 V), running continuously without storing energy.

Part 7 of 10Electrolytic cells

Electrolysis

Electrolysis electrical energy drives a non-spontaneous reaction, splitting an ionic compound that is molten or in solution.

Electrodes

Electrodes oxidation at the anode (positive) and reduction at the cathode (negative), the opposite polarity to a voltaic cell.

Memory aid

RED CATREDuction at the CAThode.

Ion movement

Ion movement cations move to the cathode and anions to the anode, while electrons move only in the external circuit.

Molten binary compounds

Molten binary compounds the metal forms at the cathode and the non-metal at the anode, e.g. molten PbBr₂ gives Pb at the cathode (Pb²⁺ + 2e⁻ → Pb) and Br₂ at the anode (2Br⁻ → Br₂ + 2e⁻).

Aqueous solutions

Aqueous solutions water may be reduced to H₂ or oxidised to O₂ instead, depending on E^⦵ and concentration, e.g. concentrated NaCl gives Cl₂.

Uses

Uses purifying copper, electroplating, extracting aluminium, and making chlorine, hydrogen and sodium hydroxide.

Part 8 of 10Oxidation of alcohols

1Oxidising agent

Oxidising agent acidified potassium dichromate(VI), where orange Cr₂O₇²⁻ is reduced to green Cr³⁺.

2Primary alcohol → aldehyde

Primary alcohol → aldehyde distil the aldehyde off as it forms, because it has a lower boiling point than the alcohol.

3Aldehyde → carboxylic acid

Aldehyde → carboxylic acid heat the primary alcohol with excess oxidant under reflux.

4Secondary alcohol

Secondary alcohol oxidised to a ketone, which resists further oxidation.

5Tertiary alcohol

Tertiary alcohol not oxidised because the carbon bearing OH has no hydrogen to lose.

6Equations

Equations use [O], e.g. CH₃CH₂OH + [O] → CH₃CHO + H₂O.

Part 9 of 10Reduction of carbonyls, acids and alkenes

LiAlH₄

LiAlH₄ in dry ether, then dilute acid, reduces carboxylic acids, aldehydes and ketones.

NaBH₄

NaBH₄ in aqueous or alcoholic solution, reduces only aldehydes and ketones.

Products

Products aldehydes and carboxylic acids give primary alcohols, and ketones give secondary alcohols.

Equation

CH₃COOH + 4[H] → CH₃CH₂OH + H₂O.

No aldehyde from an acid

LiAlH₄ goes all the way to the alcohol, so the aldehyde cannot be isolated.

Hydrogenation

Hydrogenation alkenes add H₂ over a nickel catalyst at about 150–200 °C to give alkanes, and alkynes need two H₂ per triple bond.

Margarine

Margarine partial hydrogenation of vegetable oils gives straighter chains with stronger London forces and a higher melting point.

Part 10 of 10More from the specification

Key terms

Term	Meaning
Oxidation	Loss of electrons or an increase in oxidation state.
Reduction	Gain of electrons or a decrease in oxidation state.
Oxidising agent	A species that accepts electrons and is itself reduced.
Reducing agent	A species that donates electrons and is itself oxidised.
Voltaic cell	A cell that generates electricity from a spontaneous redox reaction.
Electrolysis	The use of electricity to drive a non-spontaneous reaction in a molten or dissolved ionic compound.
Salt bridge	A strip of inert electrolyte that completes the circuit between half-cells.

Avoid these mistakes

Reversing oxidising and reducing agent — the oxidising agent is reduced and the reducing agent is oxidised.

Labelling the cathode of a voltaic cell as negative — it is positive, whereas in electrolysis it is negative.

Using NaBH₄ to reduce a carboxylic acid — LiAlH₄ is needed.

Saying electrons flow through the salt bridge — ions carry the charge there, electrons flow in the wire.

Last updated: 18 October 2026