Structure 1. Models of the particulate nature of matter

1.1 Introduction to the particulate nature of matter

1.2 The nuclear atom

1.3 Electron configurations

1.4 The mole

1.5 Ideal gases

Structure 2. Models of bonding and structure

2.1 The ionic model

2.2 The covalent model

2.3 The metallic model

2.4 From models to materials

Structure 3. Classification of matter

3.1 The periodic table

3.2 Functional groups: Organic compounds

Reactivity 1. What drives chemical reactions?

1.1 Measuring enthalpy changes

1.2 Energy cycles

1.3 Energy from fuels

Reactivity 2. How much, how fast and how far?

2.1 Amount of chemical change

2.2 Rate of chemical change

2.3 Extent of chemical change

Reactivity 3. What are the mechanisms of chemical change?

3.1 Proton transfer reactions

3.2 Electron transfer reactions

3.3 Electron sharing reactions (Radicals)

3.4 Electron-pair sharing reactions

IB Chemistry SL 1.4 The mole Question Bank

[Maximum number: 5]

A class studied the equilibrium established when ethanoic acid and ethanol react together in the presence of a strong acid, using propanone as an inert solvent. The equation is given below.

\mathrm{CH}_{3} \mathrm{COOH}+\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH} \rightleftharpoons \mathrm{CH}_{3} \mathrm{COOC}_{2} \mathrm{H}_{5}+\mathrm{H}_{2} \mathrm{O}

One group made the following initial mixture:

(a)

The density of ethanoic acid is $1.05 \mathrm{~g} \mathrm{~cm}^{-3}$ . Determine the amount, in mol, of ethanoic acid present in the initial mixture.

[ 3 ]

(b)

After one week, a $5.00 \pm 0.05 \mathrm{~cm}^{3}$ sample of the final equilibrium mixture was pipetted out and titrated with $0.200 \mathrm{~mol} \mathrm{dm}^{-3}$ aqueous sodium hydroxide to determine the amount of ethanoic acid remaining. The following titration results were obtained:

[ 2 ]

(i)

$23.00 \mathrm{~cm}^{3}$ of this $0.200 \mathrm{~mol} \mathrm{dm}^{-3}$ aqueous sodium hydroxide reacted with the ethanoic acid in the $5.00 \mathrm{~cm}^{3}$ sample. Determine the amount, in mol, of ethanoic acid present in the $50.0 \mathrm{~cm}^{3}$ of final equilibrium mixture.

[ 2 ]

[Maximum number: 2]

Reaction kinetics can be investigated using the iodine clock reaction. The equations for two reactions that occur are given below.

\begin{array}{ll} \text { Reaction } \mathrm{A}: & \mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})+2 \mathrm{I}^{-}(\mathrm{aq})+2 \mathrm{H}^{+}(\mathrm{aq}) \rightarrow \mathrm{I}_{2}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \\ \text { Reaction } \mathrm{B}: & \mathrm{I}_{2}(\mathrm{aq})+2 \mathrm{~S}_{2} \mathrm{O}_{3}{ }^{2-}(\mathrm{aq}) \rightarrow 2 \mathrm{I}^{-}(\mathrm{aq})+\mathrm{S}_{4} \mathrm{O}_{6}{ }^{2-}(\mathrm{aq}) \end{array}

Reaction B is much faster than reaction A , so the iodine, $\mathrm{I}_{2}$ , formed in reaction A immediately reacts with thiosulfate ions, $\mathrm{S}_{2} \mathrm{O}_{3}{ }^{2-}$ , in reaction B , before it can react with starch to form the familiar blue-black, starch-iodine complex.

In one experiment the reaction mixture contained:
$5.0 \pm 0.1 \mathrm{~cm}^{3}$ of $2.00 \mathrm{~mol} \mathrm{dm}^{-3}$ hydrogen peroxide $\left(\mathrm{H}_{2} \mathrm{O}_{2}\right)$ $5.0 \pm 0.1 \mathrm{~cm}^{3}$ of 1 % aqueous starch
$20.0 \pm 0.1 \mathrm{~cm}^{3}$ of $1.00 \mathrm{~mol} \mathrm{dm}^{-3}$ sulfuric acid $\left(\mathrm{H}_{2} \mathrm{SO}_{4}\right)$ $20.0 \pm 0.1 \mathrm{~cm}^{3}$ of $0.0100 \mathrm{~mol} \mathrm{dm}{ }^{-3}$ sodium thiosulfate $\left(\mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3}\right)$ $50.0 \pm 0.1 \mathrm{~cm}^{3}$ of water with $0.0200 \pm 0.0001 \mathrm{~g}$ of potassium iodide (KI) dissolved in it.
After 45 seconds this mixture suddenly changed from colourless to blue-black.

(a)

Calculate the amount, in mol , of KI in the reaction mixture.

[ 1 ]

(b)

Calculate the amount, in mol , of $\mathrm{H}_{2} \mathrm{O}_{2}$ in the reaction mixture.

[ 1 ]

[Maximum number: 5]

A student decided to determine the molecular mass of a solid monoprotic acid, HA, by titrating a solution of a known mass of the acid.

The following recordings were made.

(a)

This known mass of acid, HA, was then dissolved in distilled water to form a $100.0 \mathrm{~cm}^{3}$ solution in a volumetric flask. A $25.0 \mathrm{~cm}^{3}$ sample of this solution reacted with $12.1 \mathrm{~cm}^{3}$ of a $0.100 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{NaOH}$ solution. Calculate the molar mass of the acid.

[ 3 ]

(b)

The percentage composition of HA is 70.56 % carbon, 23.50 % oxygen and 5.94 % hydrogen. Determine its empirical formula.

[ 2 ]

[Maximum number: 5]

Iron tablets are often prescribed to patients. The iron in the tablets is commonly present as iron(II) sulfate, $\mathrm{FeSO}_{4}$ .
Two students carried out an experiment to determine the percentage by mass of iron in a brand of tablets marketed in Cyprus.
Experimental Procedure:
- The students took five iron tablets and found that the total mass was 1.65 g .
- The five tablets were ground and dissolved in $100 \mathrm{~cm}^{3}$ dilute sulfuric acid, $\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq})$ . The solution and washings were transferred to a $250 \mathrm{~cm}^{3}$ volumetric flask and made up to the mark with deionized (distilled) water.
- $25.0 \mathrm{~cm}^{3}$ of this $\mathrm{Fe}^{2+}(\mathrm{aq})$ solution was transferred using a pipette into a conical flask. Some dilute sulfuric acid was added.
- A titration was then carried out using a $5.00 \times 10^{-3} \mathrm{moldm}^{-3}$ standard solution of potassium permanganate, $\mathrm{KMnO}_{4}(\mathrm{aq})$ . The end-point of the titration was indicated by a slight pink colour.
The following results were recorded.

(a)

(i)

Determine the amount, in mol , of $\mathrm{MnO}_{4}^{-}$ (aq), used in each accurate titre.

[ 2 ]

(ii)

Calculate the amount, in mol, of $\mathrm{Fe}^{2+}(\mathrm{aq})$ ions in $250 \mathrm{~cm}^{3}$ of the solution.

[ 1 ]

(iii)

Determine the total mass of iron, in g , in the $250 \mathrm{~cm}^{3}$ solution.

[ 1 ]

(iv)

Determine the percentage by mass of iron in the tablets.

[ 1 ]

[Maximum number: 6]

Two groups of students (Group A and Group B) carried out a project* on the chemistry of some group 7 elements (the halogens) and their compounds.

(a)

In the first part of the project, the two groups had a sample of iodine monochloride (a corrosive brown liquid) prepared for them by their teacher using the following reaction.

\mathrm{I}_{2}(\mathrm{~s})+\mathrm{Cl}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{ICl}(\mathrm{l})

The following data were recorded.

[ 3 ]

(i)

The iodine used in the reaction was in excess. Determine the theoretical yield, in g , of ICl(l).

\footnotetext{
* Adapted from J Derek Woollins, (2009), Inorganic Experiments and Open University, (2008), Exploring the Molecular World.
}

[ 3 ]

(b)

The students reacted ICl(l) with CsBr(s) to form a yellow solid, $\mathrm{CsICl}_{2}(\mathrm{~s})$ , as one of the products. $\mathrm{CsICl}_{2}(\mathrm{~s})$ has been found to produce very pure CsCl(s) which is used in cancer treatment.

To confirm the composition of the yellow solid, Group A determined the amount of iodine in 0.2015 g of $\mathrm{CsICl}_{2}(\mathrm{~s})$ by titrating it with $0.0500 \mathrm{moldm}^{-3} \mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3}(\mathrm{aq})$ . The following data were recorded for the titration.

[ 3 ]

(i)

Calculate the percentage of iodine by mass in $\mathrm{CsICl}_{2}(\mathrm{~s})$ , correct to three significant figures.

[ 1 ]

(ii)

Determine the amount, in mol, of $0.0500 \mathrm{moldm}^{-3} \mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3}(\mathrm{aq})$ added in the titration.

[ 1 ]

(iii)

Calculate the mass of iodine, in g , present in the sample of $\mathrm{CsICl}_{2}(\mathrm{~s})$ .

[ 1 ]

[Maximum number: 2]

A student carried out an experiment to determine the concentration of a hydrochloric acid solution and the enthalpy change of the reaction between aqueous sodium hydroxide and this acid by thermometric titration.

She added $5.0 \mathrm{~cm}^{3}$ portions of hydrochloric acid to $25.0 \mathrm{~cm}^{3}$ of $1.00 \mathrm{moldm}^{-3}$ sodium hydroxide solution in a glass beaker until the total volume of acid added was $50.0 \mathrm{~cm}^{3}$ , measuring the temperature of the mixture each time. Her results are plotted in the graph below.

The initial temperature of both solutions was the same.

(a)

(i)

Determine the concentration of the hydrochloric acid, including units.

[ 2 ]

[Maximum number: 1]

Phosphine (IUPAC name phosphane) is a hydride of phosphorus, with the formula $\mathrm{PH}_{3}$ .

(a)

2.478 g of white phosphorus was used to make phosphine according to the equation:

\mathrm{P}_{4}(\mathrm{~s})+3 \mathrm{OH}^{-}(\mathrm{aq})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{PH}_{3}(\mathrm{~g})+3 \mathrm{H}_{2} \mathrm{PO}_{2}^{-}(\mathrm{aq})

[ 1 ]

(i)

Calculate the amount, in mol, of white phosphorus used.

[ 1 ]

[Maximum number: 1]

0.040 mol of $\left(\mathrm{NH}_{4}\right)_{2} \mathrm{Ni}\left(\mathrm{SO}_{4}\right)_{2} \cdot 6 \mathrm{H}_{2} \mathrm{O}$ is dissolved in water to give $200 \mathrm{~cm}^{3}$ of aqueous solution. What is the concentration, in $\mathrm{mol} \mathrm{dm}^{-3}$ , of ammonium ions?

0.00040

0.0080

0.20

0.40

(a)

Magnesium sulfate can exist in either the hydrated form or in the anhydrous form. Two students wished to determine the enthalpy of hydration of anhydrous magnesium sulfate. They measured the initial and the highest temperature reached when anhydrous magnesium sulfate, $\mathrm{MgSO}_{4}(\mathrm{~s})$ , was dissolved in water. They presented their results in the following table.

[ 1 ]

(i)

Calculate the amount, in mol, of anhydrous magnesium sulfate.

[ 1 ]

[Maximum number: 2]

This question is about a mug made of a lead alloy.

The rate of lead dissolving in common beverages with various pH values was analysed.

Lead dissolving in beverages at various times and temperatures

(a)

(i)

A mean daily lead intake of greater than $5.0 \times 10^{-6} \mathrm{~g}$ per kg of body weight results in increased lead levels in the body.

Calculate the volume, in $\mathrm{dm}^{3}$ , of tap water from experiment 8 which would exceed this daily lead intake for an 80.0 kg man.

[ 2 ]

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