I85PREP
⌘ K
GCE A/L Chemistry study materials
GCE A/L Chemistry study resources

GCE A/L Chemistry practice questions

One question at a time. Build your GCE A/L Chemistry understanding.

First 50 of 199 questions. Each one lists the choices, the correct answer, and an explanation. The checker at the top of this page saves progress in this browser when it is available.

Continue in the practice tool above to load more questions, 50 at a time.

Question 1Atomic structure and periodicity

Which subatomic particles are present in the nucleus of an atom of chlorine-37 (atomic number 17)?

  1. 20 protons and 17 neutrons
  2. 17 protons and 20 neutrons Correct answer.
  3. 17 protons and 17 neutrons
  4. 17 protons and 37 neutrons

Explanation

The atomic number (17) gives the number of protons, and the mass number (37) is the sum of protons and neutrons. Neutrons = 37 − 17 = 20. Electrons are outside the nucleus and are not counted.

Question 2Atomic structure and periodicity

Which of the following is the electron configuration of a chromium atom in its ground state?

  1. 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁴ 4s²
  2. 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶
  3. 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹ Correct answer.
  4. 1s² 2s² 2p⁶ 3s² 3p⁶ 3d³ 4s² 4p¹

Explanation

Chromium (Z = 24) shows an anomalously stable configuration with a half-filled 3d subshell and a singly occupied 4s orbital. This gives extra exchange stability compared with the predicted 3d⁴ 4s² arrangement.

Question 3Atomic structure and periodicity

Which species has the same electron configuration as a neon atom?

  1. Cl⁻
  2. O²⁻
  3. Na⁺
  4. Mg²⁺ Correct answer.

Explanation

Mg²⁺ has lost two electrons from magnesium (Z = 12), leaving 10 electrons, identical to neon. Isoelectronic species share the same number of electrons, not the same nuclear charge.

Question 4Atomic structure and periodicity

Which statement correctly describes the trend in first ionisation energy across Period 3 from sodium to argon?

  1. It generally increases but dips only at argon.
  2. It decreases because atomic radius increases across the period.
  3. It generally increases but dips at aluminium and at sulfur. Correct answer.
  4. It increases steadily with no exceptions because nuclear charge increases.

Explanation

First ionisation energy generally rises across Period 3 as effective nuclear charge increases, but dips occur at Al (removing a 3p electron, which is higher in energy than 3s) and at S (pairing in a 3p orbital increases electron–electron repulsion).

Question 5Atomic structure and periodicity

Which property increases across Period 3 from sodium to chlorine?

  1. Atomic radius
  2. First ionisation energy Correct answer.
  3. Metallic character
  4. Number of occupied electron shells

Explanation

Across Period 3, effective nuclear charge increases while the number of shells stays the same, so electrons are held more tightly and first ionisation energy rises. Atomic radius and metallic character both decrease.

Question 6Atomic structure and periodicity

Which ion has the largest ionic radius?

  1. Al³⁺
  2. F⁻ Correct answer.
  3. Mg²⁺
  4. Na⁺

Explanation

Na⁺, Mg²⁺ and Al³⁺ are isoelectronic with 10 electrons but have increasing nuclear charge, so they shrink in the order Na⁺ > Mg²⁺ > Al³⁺. F⁻ also has 10 electrons but the lowest nuclear charge of the group, so it is the largest.

Question 7Atomic structure and periodicity

An element has successive ionisation energies (kJ mol⁻¹) of 590, 1150, 4940, 6480 and 8130. To which group of the periodic table does it most likely belong?

  1. Group 1
  2. Group 3
  3. Group 2 Correct answer.
  4. Group 14

Explanation

The large jump between the second and third ionisation energies indicates that the third electron is removed from a new, inner shell. Losing two outer electrons before a big jump is characteristic of a Group 2 element.

Question 8Atomic structure and periodicity

Which statement about the emission spectrum of hydrogen is correct?

  1. It shows lines only in the visible region.
  2. It is continuous because the atom can emit all wavelengths.
  3. It is identical to the absorption spectrum at all wavelengths.
  4. It consists of discrete lines because electron transitions involve fixed energy gaps. Correct answer.

Explanation

Electrons occupy discrete energy levels, so transitions emit photons of specific energies given by ΔE = hν. This produces a line, not continuous, spectrum. The lines occur in the ultraviolet, visible and infrared regions.

Question 9Atomic structure and periodicity

Which subshell is filled immediately after the 4s subshell according to the Aufbau principle?

  1. 4p
  2. 4d
  3. 5s
  4. 3d Correct answer.

Explanation

After 4s, the next lowest energy subshell is 3d, because the 3d orbitals lie slightly higher in energy than 4s for the first-row transition elements. Filling 4p or 5s would skip the 3d level.

Question 10Atomic structure and periodicity

An ion X²⁺ has the electron configuration 1s² 2s² 2p⁶ 3s² 3p⁶. What is the atomic number of element X?

  1. 22
  2. 18
  3. 20 Correct answer.
  4. 16

Explanation

The ion has 18 electrons, and because it carries a 2+ charge the neutral atom has 20 electrons. The atomic number equals the number of protons, which is 20 (calcium).

Question 11Atomic structure and periodicity

Which factor best explains why the first ionisation energy of magnesium is higher than that of sodium?

  1. Magnesium has one more occupied shell than sodium.
  2. Magnesium has a greater nuclear charge with the outer electron in the same shell. Correct answer.
  3. Magnesium has a larger atomic radius than sodium.
  4. Magnesium has a paired electron in its outer orbital.

Explanation

Mg has one more proton than Na while the outer electron is removed from the same third shell, so the greater effective nuclear charge holds it more strongly. The larger radius and extra shell claims are factually incorrect for this pair.

Question 12Atomic structure and periodicity

Which of the following correctly describes the shape of a p orbital?

  1. Spherical
  2. Tetrahedral
  3. A single ring
  4. Dumbbell-shaped with two lobes Correct answer.

Explanation

Each p orbital has two lobes on opposite sides of the nucleus, giving a dumbbell shape. The s orbital is spherical, and tetrahedral describes molecular geometry, not an orbital shape.

Question 13Atomic structure and periodicity

Which statement about the shielding effect is correct?

  1. Outer electrons shield inner electrons from the nucleus.
  2. Shielding is greatest for electrons in the same shell as the outer electron.
  3. Shielding increases the effective nuclear charge on outer electrons.
  4. Inner-shell electrons reduce the effective nuclear charge felt by outer electrons. Correct answer.

Explanation

Inner-shell electrons lie between the nucleus and the outer electrons and repel them, reducing the effective nuclear charge. Electrons in the same shell shield each other only weakly.

Question 14Atomic structure and periodicity

Which element has the highest electronegativity?

  1. Fluorine Correct answer.
  2. Nitrogen
  3. Oxygen
  4. Chlorine

Explanation

Electronegativity increases across a period and decreases down a group. Fluorine, at the top right of the periodic table (excluding noble gases), has the highest electronegativity of all elements.

Question 15Atomic structure and periodicity

Which pair of elements would be expected to show the most similar chemical properties?

  1. Li and Na Correct answer.
  2. Ne and Ar
  3. K and Ca
  4. Na and Mg

Explanation

Elements in the same group have the same number of outer-shell electrons and therefore similar chemical properties. Li and Na are both Group 1 alkali metals. Na/Mg and K/Ca are in different groups.

Question 16Atomic structure and periodicity

Which statement explains why the second ionisation energy of an element is always greater than its first ionisation energy?

  1. The second electron is removed from a more stable, positively charged ion. Correct answer.
  2. The second electron is always removed from a paired orbital.
  3. The second electron is removed from a closer shell in every case.
  4. The nucleus loses a proton after the first ionisation.

Explanation

After the first electron is removed, the resulting cation holds its remaining electrons more tightly because the same nuclear charge acts on fewer electrons. Removing a further electron therefore requires more energy.

Question 17Atomic structure and periodicity

A hydrogen atom emits a photon when its electron falls from the n = 3 level to the n = 2 level. To which series does this transition belong?

  1. Paschen series
  2. Brackett series
  3. Balmer series Correct answer.
  4. Lyman series

Explanation

Transitions ending at n = 2 form the Balmer series, which lies mainly in the visible region. The Lyman series ends at n = 1, while the Paschen and Brackett series end at n = 3 and n = 4 respectively.

Question 18Atomic structure and periodicity

Which statement about atomic orbitals is correct?

  1. An orbital can hold a maximum of two electrons with opposite spins. Correct answer.
  2. An orbital is a fixed circular path followed by an electron.
  3. All orbitals in a subshell have different energies in a free atom.
  4. An orbital can hold up to three electrons.

Explanation

The Pauli exclusion principle allows at most two electrons per orbital, and they must have opposite spins. Orbitals describe probability distributions, not fixed paths, and degenerate orbitals in a subshell have equal energy.

Question 19Atomic structure and periodicity

Which isoelectronic series is arranged in order of increasing ionic radius?

  1. Na⁺ < Mg²⁺ < Al³⁺ < F⁻
  2. F⁻ < Na⁺ < Mg²⁺ < Al³⁺
  3. Al³⁺ < Mg²⁺ < Na⁺ < F⁻ Correct answer.
  4. Mg²⁺ < Al³⁺ < F⁻ < Na⁺

Explanation

All four ions have 10 electrons. As nuclear charge increases (F < Na < Mg < Al), the electrons are pulled in more strongly, so radius decreases. Increasing radius order is therefore Al³⁺ < Mg²⁺ < Na⁺ < F⁻.

Question 20Atomic structure and periodicity

The first ionisation energies of the elements with atomic numbers 11, 12, 13 and 14 are approximately 496, 738, 578 and 786 kJ mol⁻¹. Which factor explains the lower value for the element with atomic number 13?

  1. Its outer electron is in a 3p orbital, which is higher in energy than 3s. Correct answer.
  2. It has a full outer shell.
  3. It has a larger number of occupied shells than element 12.
  4. It has fewer protons than the element with atomic number 12.

Explanation

Element 13 (Al) loses its outermost electron from a 3p orbital, which lies at higher energy than the 3s orbital used by Mg (element 12). This higher-energy orbital reduces the ionisation energy below Mg's value despite the greater nuclear charge.

Question 21Chemical bonding and intermolecular forces

Which type of bonding is present in solid sodium chloride?

  1. Covalent bonding between shared electron pairs
  2. Hydrogen bonding between polar molecules
  3. Metallic bonding through a delocalised electron sea
  4. Ionic bonding between oppositely charged ions Correct answer.

Explanation

Sodium chloride is built from Na⁺ and Cl⁻ ions held together by strong electrostatic attraction in a giant ionic lattice. Electron transfer from sodium to chlorine produces these ions, so the bonding is ionic rather than covalent, metallic or hydrogen bonding.

Question 22Chemical bonding and intermolecular forces

Which statement correctly describes the shape of a methane molecule, CH₄?

  1. Trigonal planar with bond angles of 120°
  2. Bent with bond angles of about 104.5°
  3. Linear with bond angles of 180°
  4. Tetrahedral with bond angles of about 109.5° Correct answer.

Explanation

The carbon atom in methane has four bonding pairs and no lone pairs, so electron-pair repulsion arranges them as far apart as possible in a tetrahedron. The resulting H–C–H bond angle is approximately 109.5°.

Question 23Chemical bonding and intermolecular forces

Which intermolecular force is primarily responsible for the relatively high boiling point of water compared with hydrogen sulfide, H₂S?

  1. Hydrogen bonding Correct answer.
  2. London dispersion forces
  3. Dipole–dipole attractions
  4. Ionic attractions

Explanation

Oxygen is highly electronegative and small, so water molecules form strong hydrogen bonds through O–H···O interactions. Sulfur is larger and less electronegative, so H₂S has only weaker dipole–dipole and dispersion forces, giving it a much lower boiling point.

Question 24Chemical bonding and intermolecular forces

Which molecule contains a polar covalent bond but has no net dipole moment?

  1. CO₂ Correct answer.
  2. H₂O
  3. NH₃
  4. HCl

Explanation

Carbon dioxide has two polar C=O bonds, but the molecule is linear, so the two bond dipoles are equal and opposite and cancel exactly. The other molecules are bent or linear with unequal dipoles, giving a permanent net dipole moment.

Question 25Chemical bonding and intermolecular forces

Which property is characteristic of a substance with a giant covalent (macromolecular) structure such as diamond?

  1. It conducts electricity when molten
  2. It is soluble in water
  3. It has a very high melting point Correct answer.
  4. It has a low melting point

Explanation

In a giant covalent structure every atom is joined to its neighbours by strong covalent bonds extending throughout the lattice. Breaking the structure requires breaking many strong bonds, so the melting point is extremely high.

Question 26Chemical bonding and intermolecular forces

Which statement about metallic bonding is correct?

  1. Positive ions are attracted to a sea of delocalised electrons Correct answer.
  2. Molecules are held together by hydrogen bonds
  3. Atoms share pairs of electrons in fixed directions
  4. Electrons are transferred from metal atoms to non-metal atoms

Explanation

Metallic bonding is modelled as a lattice of positive metal ions held together by electrostatic attraction to delocalised valence electrons. This delocalisation explains properties such as electrical conductivity and malleability.

Question 27Chemical bonding and intermolecular forces

Which pair of species is isoelectronic with the neon atom?

  1. Na⁺ and F⁻ Correct answer.
  2. Mg²⁺ and K⁺
  3. Na⁺ and Cl⁻
  4. O²⁻ and Ne²⁺

Explanation

Na⁺ has lost one electron from sodium (10 electrons) and F⁻ has gained one electron (10 electrons), so both have the neon configuration. Isoelectronic species must contain the same total number of electrons.

Question 28Chemical bonding and intermolecular forces

Which molecule has a trigonal pyramidal shape?

  1. CH₄
  2. CO₂
  3. BF₃
  4. NH₃ Correct answer.

Explanation

Nitrogen in ammonia has three bonding pairs and one lone pair, giving four electron domains arranged tetrahedrally. The lone pair compresses the H–N–H angles, producing a trigonal pyramidal molecular shape.

Question 29Chemical bonding and intermolecular forces

Which statement about the bonding in a molecule of ethene, C₂H₄, is correct?

  1. The two carbon atoms are joined by one σ bond and one π bond Correct answer.
  2. The two carbon atoms are joined by two σ bonds
  3. The two carbon atoms are joined by a single σ bond only
  4. The two carbon atoms are joined by three π bonds

Explanation

The carbon–carbon double bond consists of one σ bond formed by end-on overlap of hybrid orbitals and one π bond formed by sideways overlap of unhybridised p orbitals. The π bond is weaker and is the one broken in addition reactions.

Question 30Chemical bonding and intermolecular forces

Which substance has the weakest forces between its particles?

  1. Magnesium oxide
  2. Iodine Correct answer.
  3. Sodium chloride
  4. Diamond

Explanation

Solid iodine consists of I₂ molecules held together only by London dispersion forces, which are weak. Diamond has a giant covalent network, and sodium chloride and magnesium oxide have strong ionic lattices, all giving much stronger interparticle forces.

Question 31Chemical bonding and intermolecular forces

Which property of water is best explained by hydrogen bonding between water molecules?

  1. Its lack of colour
  2. Its low molar mass
  3. Its unusually high surface tension Correct answer.
  4. Its ability to dissolve non-polar substances

Explanation

Hydrogen bonds between water molecules are strong and extensive, so significant energy is needed to disrupt them at the surface. This produces the unusually high surface tension of water compared with other liquids of similar molar mass.

Question 32Chemical bonding and intermolecular forces

Which electron-pair repulsion order is used in the VSEPR model?

  1. Lone pair–bond pair > lone pair–lone pair > bond pair–bond pair
  2. All electron-pair repulsions are exactly equal in strength
  3. Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair Correct answer.
  4. Bond pair–bond pair > lone pair–bond pair > lone pair–lone pair

Explanation

A lone pair is held by only one nucleus and spreads out more than a bonding pair, so lone pair–lone pair repulsions are strongest and bond pair–bond pair repulsions weakest. This ordering explains why lone pairs compress bond angles below ideal values.

Question 33Chemical bonding and intermolecular forces

Which of the following shows the correct trend in the strength of London dispersion forces?

  1. CH₄ > C₂H₆ > C₃H₈
  2. C₃H₈ > C₂H₆ > CH₄ Correct answer.
  3. CH₄ > C₃H₈ > C₂H₆
  4. C₂H₆ > C₃H₈ > CH₄

Explanation

Dispersion forces increase with the number of electrons and the surface area available for contact. Propane has the most electrons and largest surface, so its dispersion forces are strongest, followed by ethane and then methane.

Question 34Chemical bonding and intermolecular forces

Which statement explains why magnesium oxide has a much higher melting point than sodium chloride?

  1. Magnesium oxide has weaker intermolecular forces
  2. Magnesium oxide has a covalent network structure
  3. Magnesium oxide contains ions with higher charges, giving stronger electrostatic attraction Correct answer.
  4. Magnesium oxide contains smaller molecules

Explanation

MgO is an ionic lattice of Mg²⁺ and O²⁻ ions, whereas NaCl contains singly charged Na⁺ and Cl⁻ ions. The higher charges in MgO produce much stronger electrostatic attraction, so more energy is needed to melt the lattice.

Question 35Chemical bonding and intermolecular forces

Which molecule would be expected to have the largest bond angle at the central atom?

  1. H₂O
  2. CO₂ Correct answer.
  3. CH₄
  4. NH₃

Explanation

Carbon dioxide has only two electron domains around the carbon, so the bonding regions lie opposite each other and the O–C–O angle is 180°. The other molecules have three or four domains, giving angles of about 104.5°, 107° and 109.5° respectively.

Question 36Chemical bonding and intermolecular forces

Which explanation correctly describes why graphite conducts electricity while diamond does not?

  1. Graphite has weaker covalent bonds than diamond
  2. Graphite has free ions between its layers, while diamond has no ions
  3. Graphite contains delocalised electrons between its layers, while diamond has all electrons localised in σ bonds Correct answer.
  4. Graphite contains small molecules that can move

Explanation

Each carbon atom in graphite forms three σ bonds, leaving one p electron delocalised in a π system within each layer. These mobile electrons carry current, whereas every valence electron in diamond is held in a localised covalent bond.

Question 37Chemical bonding and intermolecular forces

Which species has a bond angle of approximately 104.5°?

  1. H₃O⁺
  2. SO₂
  3. H₂O Correct answer.
  4. BF₄⁻

Explanation

Water has two bonding pairs and two lone pairs on oxygen, giving a bent shape with an H–O–H angle of about 104.5°. The two lone pairs repel more strongly than bonding pairs, compressing the angle below the tetrahedral value.

Question 38Chemical bonding and intermolecular forces

Which statement correctly compares the melting points of the period 3 chlorides NaCl and SiCl₄?

  1. SiCl₄ has a higher melting point because it has a giant covalent structure
  2. SiCl₄ has a higher melting point because silicon is more electronegative than sodium
  3. NaCl has a higher melting point because it is a giant ionic lattice with strong electrostatic forces Correct answer.
  4. NaCl has a higher melting point because it contains hydrogen bonds

Explanation

NaCl forms a giant ionic lattice held by strong electrostatic attraction between ions, so it melts at a high temperature. SiCl₄ is a simple molecular liquid with only weak London dispersion forces between molecules, so it melts at a much lower temperature.

Question 39Chemical bonding and intermolecular forces

In a molecule of ammonia, the H–N–H bond angle is about 107°, whereas in the ammonium ion, NH₄⁺, it is about 109.5°. Which statement best explains this difference?

  1. The ammonium ion has a lone pair that increases repulsion
  2. The nitrogen atom in the ammonium ion has a different number of electron shells
  3. The ammonium ion has fewer electron domains than ammonia
  4. The lone pair in ammonia repels bonding pairs more strongly, compressing the angle Correct answer.

Explanation

Ammonia has one lone pair and three bonding pairs, and the lone pair repels the bonding pairs more strongly, compressing the angle below the tetrahedral value. In NH₄⁺ the lone pair is used to bond the fourth hydrogen, so four equal bonding pairs give the ideal tetrahedral angle of 109.5°.

Question 40Gases, kinetic theory and the mole concept

A sealed rigid container holds 0.50 mol of an ideal gas at 300 K. The gas is heated to 600 K. What happens to the pressure inside the container?

  1. It quadruples, because pressure depends on the square of the temperature
  2. It doubles, because pressure is directly proportional to absolute temperature at constant volume Correct answer.
  3. It halves, because the molecules move faster and collide less often with the walls
  4. It stays the same, because the number of gas molecules has not changed

Explanation

For a fixed amount of gas in a fixed volume, the ideal gas equation pV = nRT reduces to p proportional to T. Doubling the absolute temperature from 300 K to 600 K therefore doubles the pressure. The average kinetic energy of the molecules rises with T, so they strike the walls harder and more frequently.

Question 41Gases, kinetic theory and the mole concept

Which statement about the molar volume of an ideal gas is correct under standard conditions?

  1. It is larger for denser gases at the same temperature
  2. It is the same for all gases at the same temperature and pressure Correct answer.
  3. It depends on the molar mass of the gas
  4. It is only defined for gases that behave perfectly at all pressures

Explanation

Avogadro's law states that equal volumes of gases at the same temperature and pressure contain equal numbers of moles. The molar volume therefore depends only on T and p, not on the identity of the gas. At standard temperature and pressure it is approximately 22.4 dm3 per mole.

Question 42Gases, kinetic theory and the mole concept

A sample of gas contains 0.30 mol of nitrogen and 0.20 mol of oxygen in a total volume of 5.0 dm3 at constant temperature. What is the mole fraction of nitrogen?

  1. 0.60 Correct answer.
  2. 0.50
  3. 0.40
  4. 0.30

Explanation

Mole fraction is the moles of the component divided by the total moles. The total is 0.30 + 0.20 = 0.50 mol, so the mole fraction of nitrogen is 0.30/0.50 = 0.60. The volume is not needed for this calculation.

Question 43Gases, kinetic theory and the mole concept

Which quantity is directly proportional to the absolute temperature of an ideal gas?

  1. The average kinetic energy of its molecules Correct answer.
  2. The average momentum of its molecules
  3. The average distance between its molecules
  4. The average speed of its molecules

Explanation

Kinetic theory gives the average translational kinetic energy of a molecule as (3/2)kT, which is directly proportional to absolute temperature. Average speed depends on the square root of T, so it is not directly proportional.

Question 44Gases, kinetic theory and the mole concept

Which of the following is a postulate of the kinetic theory of gases as applied to an ideal gas?

  1. Gas molecules move with the same speed at a given temperature
  2. Gas molecules attract each other strongly at ordinary pressures
  3. Collisions between gas molecules are perfectly elastic Correct answer.
  4. The volume occupied by the molecules themselves is a large fraction of the container volume

Explanation

The kinetic theory assumes that intermolecular collisions and collisions with the walls are perfectly elastic, so no kinetic energy is lost overall. It also assumes negligible intermolecular forces and negligible molecular volume compared with the container.

Question 45Gases, kinetic theory and the mole concept

2.8 g of a gas occupies 2.24 dm3 at standard temperature and pressure, where the molar volume is 22.4 dm3 mol-1. What is the relative molecular mass of the gas?

  1. 2.8
  2. 56
  3. 14
  4. 28 Correct answer.

Explanation

The number of moles is 2.24/22.4 = 0.100 mol. The molar mass is mass divided by moles, giving 2.8/0.100 = 28 g mol-1, so the relative molecular mass is 28. This is consistent with a gas such as nitrogen or carbon monoxide.

Question 46Gases, kinetic theory and the mole concept

A fixed mass of ideal gas is compressed at constant temperature to one quarter of its original volume. What happens to the pressure?

  1. It becomes one quarter of the original pressure
  2. It is unchanged because the temperature is constant
  3. It becomes twice the original pressure
  4. It becomes four times the original pressure Correct answer.

Explanation

Boyle's law states that for a fixed mass of gas at constant temperature, pressure is inversely proportional to volume. Reducing the volume to one quarter therefore raises the pressure to four times its original value, since the molecules strike the walls more frequently.

Question 47Gases, kinetic theory and the mole concept

0.020 mol of a gas is collected over water at a temperature where the saturated vapour pressure of water is 2.0 kPa. The total pressure is 102 kPa. What is the pressure of the dry gas?

  1. 102 kPa
  2. 100 kPa Correct answer.
  3. 104 kPa
  4. 51 kPa

Explanation

Dalton's law of partial pressures gives the dry gas pressure as the total pressure minus the water vapour pressure. Hence 102 - 2.0 = 100 kPa. The collected gas is saturated with water vapour, which contributes its own partial pressure.

Question 48Gases, kinetic theory and the mole concept

A mixture of gases contains 0.40 mol of hydrogen and 0.10 mol of helium at a total pressure of 200 kPa. What is the partial pressure of hydrogen?

  1. 40 kPa
  2. 200 kPa
  3. 160 kPa Correct answer.
  4. 100 kPa

Explanation

The mole fraction of hydrogen is 0.40/(0.40 + 0.10) = 0.80. The partial pressure is the mole fraction multiplied by the total pressure, giving 0.80 x 200 = 160 kPa.

Question 49Gases, kinetic theory and the mole concept

Which change increases the rate of diffusion of a gas through a small hole into a vacuum?

  1. Replacing the gas with one of lower molar mass at the same temperature Correct answer.
  2. Replacing the gas with one of higher molar mass at the same temperature
  3. Reducing the pressure difference across the hole to zero
  4. Lowering the temperature while keeping the same gas

Explanation

Graham's law states that the rate of diffusion is inversely proportional to the square root of molar mass at a given temperature. A lighter gas therefore diffuses faster because its molecules have a higher root-mean-square speed. A larger pressure difference also increases the rate, but the option reduces it to zero.

Question 50Gases, kinetic theory and the mole concept

Which observation provides the strongest evidence that gas molecules are in continuous random motion?

  1. The increase in pressure when a gas is heated at constant volume
  2. The uniform colour of a gas in a sealed flask
  3. The liquefaction of a gas when it is cooled and compressed
  4. The random zig-zag paths of smoke particles seen under a microscope Correct answer.

Explanation

Brownian motion of smoke particles arises because they are struck unevenly by rapidly moving gas molecules, revealing the random motion of the molecules themselves. Uniform colour, pressure changes on heating and liquefaction are consistent with kinetic theory but do not directly show random molecular motion.