S1.4 Counting particles, mass. The mole
Practice exam-style IB Chemistry questions for Counting particles, mass. The mole, aligned with the syllabus and grouped by topic.
Question 1
The number of elementary entities in 0.250 mol of atoms is best represented by which expression?
A.
0.250 × 6.022 × 10⁻²³
B.
0.250 × N_A
C.
0.250 ÷ N_A
D.
N_A ÷ 0.250
Question 2
The relative formula mass, M_r, of Al₂(SO₄)₃ is calculated using A_r values Al = 26.98, S = 32.06 and O = 16.00. What is M_r for Al₂(SO₄)₃?
A.
684.28
B.
342.14
C.
150.02
D.
123.04
Question 3
A sample contains 9.00 g of water. What amount of H₂O molecules is present? Use M(H₂O) = 18.02 g mol⁻¹.
A.
2.00 mol
B.
9.00 mol
C.
162 mol
D.
0.499 mol
Question 4
The empirical formula of a compound with molecular formula C₄H₈O₂ is:
A.
C₂H₄O
B.
C₄H₈O₂
C.
CH₂O
D.
C₂H₈O
Question 5
State the value and unit of the Avogadro constant, N_A.
[1]
Calculate the number of molecules in 0.125 mol of CO₂.
[1]
Question 6
Calculate the relative formula mass, M_r, of Ca(OH)₂ using A_r values Ca = 40.08, O = 16.00 and H = 1.01. [2]
Question 7
A solution contains 0.0200 mol of NaCl in a final volume of 250 cm³. What is the molar concentration of NaCl?
A.
5.00 mol dm⁻³
B.
80.0 mol dm⁻³
C.
0.0800 mol dm⁻³
D.
0.00500 mol dm⁻³
Question 8
Nitrogen and hydrogen react according to N₂
A.
40 cm³
B.
120 cm³
C.
13 cm³
D.
80 cm³
Question 9
The amount of chloride ions in 0.300 mol of MgCl₂ formula units is:
A.
0.150 mol
B.
0.300 mol
C.
0.600 mol
D.
0.900 mol
Question 10
A solution has a mass concentration of 5.85 g dm⁻³ NaCl. Using M(NaCl) = 58.44 g mol⁻¹, the molar concentration is:
A.
10.0 mol dm⁻³
B.
0.0100 mol dm⁻³
C.
342 mol dm⁻³
D.
0.100 mol dm⁻³
Question 11
The statement that correctly distinguishes relative formula mass from molar mass is:
A.
Relative formula mass is measured in g mol⁻¹; molar mass has no unit.
B.
Both relative formula mass and molar mass are unitless ratios.
C.
Both relative formula mass and molar mass have units g mol⁻¹.
D.
Relative formula mass has no unit; molar mass usually has units g mol⁻¹.
Question 12
A student weighs 2.50 g of sodium carbonate, Na₂CO₃.
Calculate the molar mass of Na₂CO₃. Use A_r values Na = 22.99, C = 12.01, O = 16.00.
[1]
Calculate the amount, in mol, of Na₂CO₃ in the sample.
[1]
Question 13
A compound contains 40.0% carbon, 6.7% hydrogen and 53.3% oxygen by mass.
State the assumption made about sample mass when using percentage composition to determine an empirical formula.
[1]
Determine the empirical formula.
[1]
Question 14
A 0.150 mol dm⁻³ solution of KNO₃ is prepared in a 250.0 cm³ volumetric flask.
Calculate the amount of KNO₃ required.
[1]
Calculate the mass of KNO₃ required. Use M(KNO₃) = 101.11 g mol⁻¹.
[1]
Question 15
Methane burns according to CH₄
- 2O₂
[1]
→ CO₂
[1]
- 2H₂O(g).
[1]
State the volume ratio of CH₄ to O₂ for gases measured under the same conditions.
[1]
Explain why this ratio can be used for gas volumes.
[2]
Question 16
The graph shows the mass of several pure samples of copper plotted against the amount of copper atoms in each sample.
State the relationship shown by the graph.
[1]
Use the gradient of the graph to determine the molar mass of copper.
[1]
Determine the mass of copper containing 0.250 mol of copper atoms.
[1]
Question 17
A class heated different masses of magnesium in air. The graph shows final mass of magnesium oxide against initial mass of magnesium.
Identify one anomalous point on the graph.
[1]
Use the best-fit line to determine the mass of oxygen combined with a given mass of magnesium shown on the graph.
[1]
Use the graph to determine the empirical formula of magnesium oxide.
[1]
Question 18
The diagram shows gas volumes measured at the same temperature and pressure for the reaction between carbon monoxide and oxygen:
2CO
| Mixture | CO initial / cm³ | O₂ initial / cm³ |
|---|---|---|
| A | 40 | 20 |
| B | 30 | 25 |
| C | 60 | 20 |
- O₂
[1]
→ 2CO₂
[1]
State the volume ratio CO:O₂:CO₂.
[1]
Use the data to determine the limiting gas in one mixture.
[1]
Calculate the volume of CO₂ formed in that mixture.
[2]
Question 19
A 20.0 cm³ sample of 0.500 mol dm⁻³ CuSO₄(aq) is diluted to 100.0 cm³. What is the concentration after dilution?
A.
2.50 mol dm⁻³
B.
0.500 mol dm⁻³
C.
0.0200 mol dm⁻³
D.
0.100 mol dm⁻³
Question 20
A hydrocarbon contains 85.6% carbon by mass. The empirical formula is best represented by:
A.
C₂H
B.
CH₄
C.
CH₂
D.
C₃H₈
Question 21
The number of oxygen atoms in 0.100 mol of Al₂O₃ formula units is:
A.
0.200N_A
B.
0.100N_A
C.
0.500N_A
D.
0.300N_A
Question 22
A compound has empirical formula NO₂ and molar mass 92.02 g mol⁻¹. Its molecular formula is:
A.
NO₂
B.
N₂O₄
C.
N₂O₂
D.
NO₄
Question 23
C₂H₂
A.
30 cm³
B.
75 cm³
C.
60 cm³
D.
15 cm³
Question 24
In determining the empirical formula of a metal oxide by heating the metal in air, the calculated oxygen content is too low. The most likely cause is:
A.
The crucible was heated to constant mass.
B.
The balance had more decimal places than required.
C.
Some metal did not react before the final weighing.
D.
The oxygen mass was found from the final mass minus the initial metal mass.
Question 25
A 10.0 cm³ aliquot of 0.200 mol dm⁻³ solution is diluted to 250.0 cm³, then 25.0 cm³ of this diluted solution is taken. The amount of solute in the 25.0 cm³ portion is:
A.
2.00 × 10⁻³ mol
B.
5.00 × 10⁻³ mol
C.
2.00 × 10⁻⁴ mol
D.
8.00 × 10⁻⁴ mol
Question 26
A drink contains glucose, C₆H₁₂O₆, at a mass concentration of 18.0 g dm⁻³.
Calculate the molar mass of glucose.
[1]
Calculate the molar concentration of glucose.
[1]
State the meaning of [glucose].
[1]
Question 27
A compound has empirical formula CH₂O and molar mass 180.18 g mol⁻¹.
Calculate the empirical formula mass.
[1]
Determine the molecular formula.
[1]
Question 28
A sample contains 4.50 × 10²² sulfate ions, SO₄²⁻.
Calculate the amount, in mol, of sulfate ions.
[1]
Calculate the amount, in mol, of oxygen atoms in the sample.
[1]
Calculate the number of oxygen atoms in the sample.
[1]
Question 29
On complete combustion, 0.420 g of a compound containing only carbon, hydrogen and oxygen produces 0.616 g CO₂ and 0.252 g H₂O.
Determine the amount of carbon atoms in the original sample.
[1]
Determine the amount of hydrogen atoms in the original sample.
[1]
Determine the empirical formula of the compound.
[1]
Question 30
A standard solution is prepared by dissolving a solid and making the solution up to the mark in a volumetric flask.
State why the final volume, rather than the volume of water first added, is used in n = CV.
[1]
Explain why the beaker and funnel are rinsed into the volumetric flask.
[1]
Question 31
Propane burns according to C₃H₈
- 5O₂
[1]
→ 3CO₂
[1]
- 4H₂O(g). All gas volumes are measured at the same temperature and pressure.
[1]
Calculate the volume of O₂ needed for 25.0 cm³ of C₃H₈.
[1]
Calculate the volume of CO₂ produced.
[1]
Explain why the same volume method must not be used to calculate a volume of liquid water.
[2]
Question 32
The table shows absorbance data for standard solutions of nickel(II) sulfate and one diluted unknown.
| Solution | c(NiSO₄) / mol dm⁻³ | Absorbance |
|---|---|---|
| Standard 1 | 0.0000 | 0.000 |
| Standard 2 | 0.0010 | 0.082 |
| Standard 3 | 0.0020 | 0.160 |
| Standard 4 | 0.0030 | 0.241 |
| Standard 5 | 0.0040 | 0.319 |
| Standard 6 | 0.0050 | 0.401 |
| Diluted unknown | — | 0.280 |
Describe the relationship between absorbance and concentration for the standards.
[1]
Use the calibration data to determine the concentration of the diluted unknown.
[1]
The original unknown was diluted by transferring 10.0 cm³ to a 100.0 cm³ volumetric flask. Determine the concentration of the original unknown.
[1]
Question 33
A student prepared sodium chloride solutions using different glassware. The table compares the intended concentration with concentrations determined by an independent method.
| Glassware used | Intended concentration / mol dm⁻³ | Determined concentration / mol dm⁻³ |
|---|---|---|
| Volumetric flask | 0.1000 | 0.0996 |
| Measuring cylinder | 0.1000 | 0.1028 |
| Beaker | 0.1000 | 0.0935 |
Identify which preparation gives a concentration closest to the intended value.
[1]
Calculate the percentage error for one preparation.
[1]
Evaluate which glassware is most suitable for preparing a standard solution.
[1]
Question 34
A stock solution has concentration 0.800 mol dm⁻³. A 5.00 cm³ portion is diluted to 100.0 cm³. Then 10.0 cm³ of the diluted solution is further diluted to 250.0 cm³.
Calculate the concentration after the first dilution.
[1]
Calculate the concentration after the second dilution.
[1]
Question 35
In an experiment to determine the empirical formula of magnesium oxide, a student heats magnesium in a crucible.
State why the crucible is heated, cooled and weighed repeatedly until constant mass is obtained.
[1]
Suggest two experimental errors that could make the calculated Mg:O mole ratio too high.
[1]
Question 36
A compound contains 52.14% carbon, 13.13% hydrogen and 34.73% oxygen by mass.
Determine its empirical formula.
[1]
The molar mass is 92.14 g mol⁻¹. Determine the molecular formula.
[1]
Question 37
A student prepares a calibration curve using five standard solutions of a coloured ion. The unknown solution gives an absorbance higher than all standards.
State one problem with using this calibration curve directly.
[1]
Suggest a suitable procedure to obtain a reliable concentration for the unknown.
[1]
State one glassware choice that improves reliability when preparing the standards.
[1]
Question 38
Combustion analysis data for several oxygen-containing organic compounds are shown. Each compound contains only C, H and O.
| Compound | Sample mass / g | CO₂ mass / g | H₂O mass / g |
|---|---|---|---|
| A | 0.600 | 0.880 | 0.360 |
| B | 0.460 | 0.880 | 0.540 |
| C | 0.580 | 1.320 | 0.540 |
| D | 0.880 | 1.760 | 0.720 |
For compound A, determine the amount of carbon from the mass of CO₂ produced.
[1]
Determine the amount of hydrogen from the mass of H₂O produced.
[1]
Use the data to determine the empirical formula of compound A.
[1]
Question 39
The graph shows a calibration curve for a coloured complex ion. An unknown was measured before and after dilution.
Determine why the undiluted unknown measurement is unsuitable.
[1]
Use the diluted unknown absorbance to determine its concentration.
[1]
Determine the concentration of the original unknown and evaluate one limitation of the result.
[1]
Question 40
The table shows reacting gas volumes for the complete combustion of a gaseous hydrocarbon, C_xH_y. Water is gaseous under the conditions used.
| Gas | Volume / cm³ |
|---|---|
| Hydrocarbon, CxHy | 20.0 |
| O₂ consumed | 100.0 |
| CO₂ produced | 60.0 |
| H₂O(g) produced | 80.0 |
Use the volume of CO₂ produced to determine x for the hydrocarbon formula.
[1]
Use the volume of H₂O produced to determine y.
[1]
Write a balanced equation for the combustion reaction.
[1]
Question 41
The graph compares the number of specified entities with amount of substance for atoms, molecules and formula units.
State what the gradient of each line represents.
[1]
Explain why the lines have the same gradient even though different substances are shown.
[1]
For the MgCl₂ line, determine the amount of chloride ions present when the graph shows a stated amount of MgCl₂ formula units.
[1]
Question 42
A student must prepare 250.0 cm³ of 0.100 mol dm⁻³ sodium carbonate solution, Na₂CO₃(aq), from solid Na₂CO₃.
Calculate the mass of Na₂CO₃ required. Use M(Na₂CO₃) = 105.99 g mol⁻¹.
[1]
Explain the procedure and glassware choices needed to prepare the solution accurately.
[1]
Question 43
A compound contains 24.3% carbon, 4.1% hydrogen and 71.6% chlorine by mass. Its molar mass is 98.96 g mol⁻¹.
Determine the empirical formula of the compound.
[1]
Determine the molecular formula and explain the difference between empirical and molecular formulas.
[1]
Question 44
A 1.20 g sample of hydrated copper(II) sulfate, CuSO₄·xH₂O, is heated until all water is removed. The mass of anhydrous CuSO₄ remaining is 0.768 g.
Calculate the amount of anhydrous CuSO₄ remaining. Use M(CuSO₄) = 159.61 g mol⁻¹.
[1]
Determine the value of x in CuSO₄·xH₂O.
[1]
Question 45
Two groups determined the empirical formula of copper oxide by heating copper in air. Their mass measurements are shown.
| Measurement | Group A mass / g | Group B mass / g |
|---|---|---|
| Empty crucible | 25.67 | 27.12 |
| Crucible + copper before heating | 26.95 | 28.40 |
| After heat-cool-weigh cycle 1 | 27.20 | 28.62 |
| After heat-cool-weigh cycle 2 | 27.26 | 28.67 |
| After heat-cool-weigh cycle 3 | 27.27 | 28.72 |
| After heat-cool-weigh cycle 4 | 27.27 | 28.77 |
Calculate the mass of oxygen gained for one group.
[1]
Determine the empirical formula obtained by that group.
[1]
Discuss which group’s data are more reliable.
[1]
Question 46
Ammonia is produced by the reaction N₂
- 3H₂
[1]
→ 2NH₃(g). A mixture initially contains 50.0 cm³ N₂ and 120.0 cm³ H₂ at the same temperature and pressure.
[1]
Determine the limiting reactant and the volume of excess gas remaining.
[1]
Explain, using Avogadro’s law, how the volume of NH₃ produced can be predicted and calculate its value.
[1]
Question 47
A student determines the empirical formula of tin oxide by heating tin powder in a crucible. The initial mass of tin is 1.19 g and the final mass of oxide is 1.51 g.
Determine the empirical formula from these data. Use A_r(Sn) = 118.71 and A_r
[1]
= 16.00.
[1]
Evaluate the experimental method and suggest improvements to increase the reliability of the empirical formula.
[1]
Question 48
A volatile liquid contains only carbon, hydrogen and oxygen. A 0.735 g sample is completely combusted, producing 1.47 g CO₂ and 0.901 g H₂O. The molar mass of the liquid is 88.12 g mol⁻¹.
Determine the empirical formula of the liquid.
[1]
Determine the molecular formula and justify any approximation used in the calculation.
[1]
Question 49
A laboratory technician prepares a series of standard solutions for a calibration curve from a 0.500 mol dm⁻³ stock solution of a coloured salt.
Calculate the volume of stock solution required to prepare 100.0 cm³ of a 0.0800 mol dm⁻³ standard.
[1]
Discuss how the standards and calibration curve should be prepared and used to determine the concentration of an unknown solution accurately.
[1]
Question 50
A student compares three samples: 0.200 mol Ar atoms, 0.200 mol O₂ molecules and 0.200 mol CaCl₂ formula units.
Calculate the total number of specified entities in one of the samples.
[1]
Compare and contrast the numbers of atoms and ions present in the three samples, and explain why specifying the elementary entity is essential when using the mole.
[1]
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