International Baccalaureate IB Chemistry

Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton acceptor.

Deduce the Brønsted–Lowry acid and base in a reaction.

A pair of species differing by a single proton is called a conjugate acid–base pair.

Deduce the formula of the conjugate acid or base of any Brønsted–Lowry base or acid.

Some species can act as both Brønsted–Lowry acids and bases.

Interpret and formulate equations to show acid–base reactions of these species.

The pH scale can be used to describe the $[H^+]$ of a solution: $pH = -\log_{10}[H^+]$; $[H^+] = 10^{-pH}$.

Perform calculations involving the logarithmic relationship between pH and $[H^+]$.

The ion product constant of water, $K_w$, shows an inverse relationship between $[\mathrm{H}^+]$ and $[\mathrm{OH}^-]$: $K_w = [\mathrm{H}^+][\mathrm{OH}^-]$.

Recognise solutions as acidic, neutral and basic from the relative values of $[\mathrm{H}^+]$ and $[\mathrm{OH}^-]$.

Strong and weak acids and bases differ in the extent of ionisation.

Recognise that acid–base equilibria lie in the direction of the weaker conjugate.

Acids react with bases in neutralisation reactions.

Formulate equations for the reactions between acids and metal oxides, metal hydroxides, hydrogen‑carbonate and carbonates. 

Identify the parent acid and base of different salts.

pH curves for neutralization reactions involving strong acids and bases have characteristic shapes and features. 

Sketch and interpret the general shape of the pH curve. 

Interpretation should include the intercept with the pH axis and equivalence point. 

Only monoprotic neutralization reactions will be assessed.

The $pOH$ scale describes the $[\mathrm{OH}^-]$ of a solution. $pOH = -\log_{10}[\mathrm{OH}^-]$; $[\mathrm{OH}^-] = 10^{-\text{pOH}}$.

Interconvert $[\mathrm{H}^+]$, $[\mathrm{OH}^-]$, $pH$ and $pOH$ values.

The equations for $pOH$ are given in the data booklet.

The strengths of weak acids and bases are described by their Ka, Kb, pKa or pKb values.

Interpret the relative strengths of acids and bases from these data.

For a conjugate acid–base pair, the relationship $K_a \times K_b = K_w$ can be derived from the expressions for $K_a$ and $K_b$.

Solve problems involving these values.

The pH of a salt solution depends on the relative strengths of the parent acid and base.

Construct equations for the hydrolysis of ions in a salt, and predict the effect of each ion on the pH of the salt solution. 

Examples should include the ammonium ion $ \mathrm{NH_4^+}$, the carboxylate ion $ \mathrm{RCOO^-}$, the carbonate ion $ \mathrm{CO_3^{2-}}$, and the hydrogencarbonate ion $ \mathrm{HCO_3^-}$. The acidity of hydrated transition element ions and (aq) is not required. 

$$\mathrm{NH_4^+ + H_2O \rightleftharpoons NH_3 + H_3O^+}$$  

$$\mathrm{RCOO^- + H_2O \rightleftharpoons RCOOH + OH^-}$$  

$$\mathrm{CO_3^{2-} + H_2O \rightleftharpoons HCO_3^- + OH^-}$$  

$$\mathrm{HCO_3^- + H_2O \rightleftharpoons H_2CO_3 + OH^-}$$  

$$\mathrm{HCO_3^- + H_2O \rightleftharpoons CO_3^{2-} + H_3O^+}$$

pH curves of different combinations of strong and weak monoprotic acids and bases have characteristic shapes and features.

Interpret the general shapes of pH curves for all four combinations of strong and weak acids and bases. 

Interpretation should include: intercept with the pH 

Tool 1—When collecting data to generate a pH axis, equivalence point, buffer region, points where curve, when should smaller volumes of titrant be added between each measurement? $pH = pKa$ or $pOH = pKb$.

Acid–base indicators are weak acids, where the components of the conjugate acid–base pair have different colours. The $p\mathrm{H}$ of the end point of an indicator, where it changes colour, approximately corresponds to its $pK_a$ value.

Construct equilibria expressions to show why the colour of an indicator changes with $p\mathrm{H}$. 

The generalized formula $\mathrm{HInd_{(aq)}}$ can be used to represent the undissociated form of an indicator. 

The equilibrium can be written as 
$$\mathrm{HInd_{(aq)} \rightleftharpoons H^+ + Ind^-}$$ 
and the acid dissociation constant as 
$$K_a = \frac{[H^+][\mathrm{Ind^-}]}{[\mathrm{HInd}]}.$$ 

Examples of indicators with their $p\mathrm{H}$ range are given in the data booklet. 

Include universal indicator as a mixture of many indicators with a wide $p\mathrm{H}$ range of colour change.

An appropriate indicator for a titration has an end point range that coincides with the $pH$ at the equivalence point.

Identify an appropriate indicator for a titration from the identity of the salt and the $pH$ range of the indicator. 

Distinguish between the terms “end point” and “equivalence point”.

A buffer solution is one that resists change in pH on the addition of small amounts of acid or alkali.

Describe the composition of acidic and basic buffers and explain their actions. 

Why must buffer solutions be composed of weak acid or base conjugate systems, not of strong acids or bases?

The $pH$ of a buffer solution depends on both: the $pKa$ or $pKb$ of its acid or base; the ratio of the concentration of acid or base to the concentration of the conjugate base or acid.

Solve problems involving the composition and $pH$ of a buffer solution, using the equilibrium constant. Include explanation of the effect of dilution of a buffer.

Oxidation and reduction can be described in terms of electron transfer, change in oxidation state, oxygen gain/loss or hydrogen loss/gain.

Deduce oxidation states of an atom in a compound or an ion. 

Identify the oxidized and reduced species and the oxidizing and reducing agents in a chemical reaction.

Half-equations separate the processes of oxidation and reduction, showing the loss or gain of electrons.

Deduce redox half-equations and equations in acidic or neutral solutions. 

Tool 1, Inquiry 2—Why are some redox titrations described as “self-indicating”.

The relative ease of oxidation and reduction of an element in a group can be predicted from its position in the periodic table.

The reactions between metals and aqueous metal ions demonstrate the relative ease of oxidation of different metals. 

Predict the relative ease of oxidation of metals. Predict the relative ease of reduction of halogens. Interpret data regarding metal and metal ion reactions.

Acids react with reactive metals to release hydrogen.

Deduce equations for reactions of reactive metals with dilute and .

Oxidation occurs at the anode and reduction occurs at the cathode in electrochemical cells.

Identify electrodes as anode and cathode, and identify their signs/polarities in voltaic cells and electrolytic cells, based on the type of reaction occurring at the electrode.

A primary (voltaic) cell is an electrochemical cell that converts energy from spontaneous redox reactions to electrical energy.

Explain the direction of electron flow from anode to cathode in the external circuit, and ion movement across the salt bridge.

Secondary (rechargeable) cells involve redox reactions that can be reversed using electrical energy.

Deduce the reactions of the charging process from given electrode reactions for discharge, and vice versa.

An electrolytic cell is an electrochemical cell that converts electrical energy to chemical energy by bringing about non-spontaneous reactions.

Explain how current is conducted in an electrolytic cell. Deduce the products of the electrolysis of a molten salt.

Functional groups in organic compounds may undergo oxidation.

Deduce equations to show changes in the functional groups during oxidation of primary and secondary alcohols, including the two-step reaction in the oxidation of primary alcohols.

$$\text{Primary alcohol:}\quad RCH_2OH \xrightarrow{\text{[O]}} RCHO \xrightarrow{\text{[O]}} RCOOH$$

$$\text{Secondary alcohol:}\quad R_2CHOH \xrightarrow{\text{[O]}} R_2C=O$$

Functional groups in organic compounds may undergo reduction.

Deduce equations to show reduction of carboxylic acids to primary alcohols via the aldehyde, and reduction of ketones to secondary alcohols. 

$$\ce{R-COOH + 2\,\text{[H]} -> R-CHO + H2O}$$
$$\ce{R-CHO + \text{[H]} -> R-CH2OH}$$
Overall: $$\ce{R-COOH + 4\,\text{[H]} -> R-CH2OH + H2O}$$

$$\ce{R2C=O + \text{[H]} -> R2CHOH}$$

Include the role of hydride ions in the reduction reaction. 

The hydride ion $\ce{H-}$, delivered by reagents such as $\ce{LiAlH4}$ or $\ce{NaBH4}$, acts as the nucleophilic reducing agent that transfers a hydride to the carbonyl carbon in each step.

Reduction of unsaturated compounds by the addition of hydrogen lowers the degree of unsaturation.

Deduce the products of the reactions of hydrogen with alkenes and alkynes.

The hydrogen half-cell $ \mathrm{H}^+ (aq) + e^- \rightleftharpoons \frac{1}{2}\,\mathrm{H}_2 (g) $ is assigned a standard electrode potential of zero by convention. It is used in the measurement of standard electrode potential, $E^\circ$.

Interpret standard electrode potential data in terms of ease of oxidation/reduction.

Standard cell potential, $E^\circ_{\text{cell}}$, can be calculated from standard electrode potentials. $E^\circ_{\text{cell}}$ has a positive value for a spontaneous reaction.

Predict whether a reaction is spontaneous in the forward or reverse direction from $E^\circ$ data.

The equation $ \Delta G^\circ = -nF E^\circ_{\text{cell}} $ shows the relationship between standard Gibbs energy and standard cell potential for a reaction.

Determine the value for $ \Delta G^\circ $ from $ E^\circ $ data. 

The equation and the value of $ F $ in C mol$^{-1}$ are given in the data booklet.

During electrolysis of aqueous solutions, competing reactions can occur at the anode and cathode, including the oxidation and reduction of water.

Deduce from standard electrode potentials the products of the electrolysis of aqueous solutions.

Electroplating involves the electrolytic coating of an object with a metallic thin layer.

Deduce equations for the electrode reactions during electroplating.

Radicals are produced by homolytic fission, e.g. of halogens, in the presence of ultraviolet (UV) light or heat.

Explain, including with equations, the homolytic fission of halogens, known as the initiation step in a chain reaction.

Radicals take part in substitution reactions with alkanes, producing a mixture of products.

Explain, using equations, the propagation and termination steps in the reactions between alkanes and halogens.

A nucleophile is a reactant that forms a bond to its reaction partner (the electrophile) by donating both bonding electrons.

Recognize nucleophiles in chemical reactions. 

Both neutral and negatively charged species should be included.

In a nucleophilic substitution reaction, a nucleophile donates an electron pair to form a new bond, as another bond breaks producing a leaving group.

Deduce equations with descriptions and explanations of the movement of electron pairs in nucleophilic substitution reactions. 

$$\text{SN2: } \mathrm{Nu^{-}} + \mathrm{R\!-\!X} \rightarrow \mathrm{R\!-\!Nu} + \mathrm{X^{-}} \quad \text{(backside attack; nucleophile donates a pair, C–X bond electrons go to }X^{-}\text{)}$$

$$\text{SN1 (ionization): } \mathrm{R\!-\!X} \rightarrow \mathrm{R^{+}} + \mathrm{X^{-}} \quad \text{(heterolytic cleavage forming a carbocation)}$$

$$\text{SN1 (nucleophilic attack): } \mathrm{R^{+}} + \mathrm{Nu^{-}} \rightarrow \mathrm{R\!-\!Nu} \quad \text{(nucleophile donates a pair to the planar carbocation)}$$

Heterolytic fission is the breakage of a covalent bond when both bonding electrons remain with one of the two fragments formed.

Explain, with equations, the formation of ions by heterolytic fission.

An electrophile is a reactant that forms a bond to its reaction partner (the nucleophile) by accepting both bonding electrons from that reaction partner.

Recognize electrophiles in chemical reactions. 

Both neutral and positively‑charged species should be included.

Alkenes are susceptible to electrophilic attack because of the high electron density of the carbon–carbon double bond. These reactions lead to electrophilic addition.

Deduce equations for the reactions of alkenes with water, halogens, and hydrogen halides. 

$$\mathrm{R\!-\!CH=CH_2 + H_2O \xrightarrow{H^+} R\!-\!CH(OH)\!-\!CH_3}$$

$$\mathrm{R\!-\!CH=CH_2 + X_2 \rightarrow R\!-\!CHX\!-\!CHX\!-\!R'}$$

$$\mathrm{R\!-\!CH=CH_2 + HX \rightarrow R\!-\!CHX\!-\!CH_3}$$

A Lewis acid is an electron‑pair acceptor and a Lewis base is an electron‑pair donor.

Apply Lewis acid–base theory to inorganic and organic chemistry to identify the role of the reacting species.

When a Lewis base reacts with a Lewis acid, a coordination bond is formed. Nucleophiles are Lewis bases and electrophiles are Lewis acids.

Draw and interpret Lewis formulas of reactants and products to show coordination bond formation in Lewis acid–base reactions.

Coordination bonds are formed when ligands donate an electron pair to transition element cations, forming complex ions.

Deduce the charge on a complex ion, given the formula of the ion and ligands present.

Nucleophilic substitution reactions include the reactions between halogenoalkanes and nucleophiles.

Describe and explain the mechanisms of the reactions of primary and tertiary halogenoalkanes with nucleophiles.

The rate of the substitution reactions is influenced by the identity of the leaving group.

Predict and explain the relative rates of the substitution reactions for different halogenoalkanes.

Alkenes readily undergo electrophilic addition reactions.

Describe and explain the mechanisms of the reactions between symmetrical alkenes and halogens, water and hydrogen halides.

The relative stability of carbocations in the addition reactions between hydrogen halides and unsymmetrical alkenes can be used to explain the reaction mechanism.

Predict and explain the major product of a reaction between an unsymmetrical alkene and a hydrogen halide or water.

Electrophilic substitution reactions include the reactions of benzene with electrophiles.

Describe and explain the mechanism of the reaction between benzene and a charged electrophile, $E^+$.