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FY13CE Chemistry study materials
FY13CE Chemistry study resources

FY13CE Chemistry practice questions

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Question 1Atomic Structure and the Periodic Table

Which subatomic particle determines the identity of an element?

  1. The electron, because it fills the shells
  2. The neutron, because it fixes the mass number
  3. The nucleon, because it carries no charge
  4. The proton, because its count is the atomic number Correct answer.

Explanation

The atomic number equals the number of protons in the nucleus, and this value uniquely identifies each element on the periodic table. Changing the proton count converts the atom into a different element, whereas changing neutrons or electrons does not.

Question 2Atomic Structure and the Periodic Table

An atom of an element has the electron configuration 1s² 2s² 2p⁶ 3s². To which group of the periodic table does it belong?

  1. Group 3
  2. Group 2 Correct answer.
  3. Group 12
  4. Group 14

Explanation

The configuration ends in 3s², so the atom has two electrons in its outermost shell and is an s-block element. Elements with two outer s electrons occupy Group 2, the alkaline earth metals.

Question 3Atomic Structure and the Periodic Table

Which statement correctly describes isotopes of the same element?

  1. They have equal neutron numbers but different proton numbers
  2. They have equal mass numbers but different proton numbers
  3. They have equal proton numbers but different neutron numbers Correct answer.
  4. They have equal electron numbers but different proton numbers

Explanation

Isotopes are atoms of one element, so their proton number (atomic number) is identical. They differ only in the number of neutrons, which gives them different mass numbers and slightly different physical properties.

Question 4Atomic Structure and the Periodic Table

Which of the following is the correct order of subshell filling according to the Aufbau principle?

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

Explanation

The 4s orbital is lower in energy than the 3d orbitals for the atoms involved, so it fills first despite the higher principal quantum number. The order follows the (n+l) rule, giving 3p, then 4s, then 3d.

Question 5Atomic Structure and the Periodic Table

Across a period from left to right, what is the general trend in atomic radius?

  1. It increases because more shells are added
  2. It stays constant because proton and electron numbers rise equally
  3. It decreases because effective nuclear charge increases Correct answer.
  4. It decreases only for the noble gases

Explanation

Adding protons across a period raises the effective nuclear charge while electrons enter the same principal shell. The stronger attraction pulls the outer electrons closer, so the atomic radius generally decreases.

Question 6Atomic Structure and the Periodic Table

Which ion has the electron configuration 1s² 2s² 2p⁶?

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

Explanation

The configuration shown is that of neon, with ten electrons. A sodium atom has eleven electrons, so losing one electron gives the Na⁺ ion exactly ten electrons and the neon configuration.

Question 7Atomic Structure and the Periodic Table

A sample of copper contains 69.2% of ⁶³Cu and 30.8% of ⁶⁵Cu. What is the relative atomic mass of the sample to three significant figures?

  1. 64.0
  2. 63.6 Correct answer.
  3. 65.0
  4. 64.5

Explanation

The relative atomic mass is the weighted mean of the isotope masses: (0.692 × 63) + (0.308 × 65) = 43.60 + 20.02 = 63.62. Rounding to three significant figures gives 63.6.

Question 8Atomic Structure and the Periodic Table

Which species is isoelectronic with the argon atom?

  1. S²⁻ Correct answer.
  2. F⁻
  3. Ne
  4. Na⁺

Explanation

Argon has eighteen electrons, and the sulfide ion also has eighteen because sulfur contributes sixteen and gains two. Isoelectronic species therefore share the same number and arrangement of electrons.

Question 9Atomic Structure and the Periodic Table

Why does the first ionisation energy of magnesium exceed that of aluminium?

  1. The electron removed from aluminium is in a higher-energy 3p orbital shielded by the 3s pair Correct answer.
  2. The aluminium atom is smaller, so its outer electron is held more tightly
  3. Aluminium has a larger nuclear charge than magnesium
  4. Magnesium has more neutrons in its nucleus than aluminium

Explanation

Magnesium loses a 3s electron, whereas aluminium loses a 3p electron that lies higher in energy and is shielded by the filled 3s subshell. The shielding and subshell difference outweigh the higher nuclear charge of aluminium, so magnesium has the larger first ionisation energy.

Question 10Atomic Structure and the Periodic Table

Which description of a p orbital is correct?

  1. It is dumbbell-shaped and has a node at the nucleus Correct answer.
  2. It is a flat circle lying in the xy plane
  3. It is spherical and centred on the nucleus
  4. It is tetrahedral with four lobes

Explanation

A p orbital has two lobes of opposite sign separated by a nodal plane passing through the nucleus, giving a dumbbell shape. The spherical shape describes an s orbital.

Question 11Atomic Structure and the Periodic Table

Which element shows the highest first ionisation energy?

  1. Argon
  2. Sodium
  3. Lithium
  4. Neon Correct answer.

Explanation

Neon has a complete second shell with a high effective nuclear charge and a small radius, so its outer electron is very tightly held. Its first ionisation energy is the highest among these choices.

Question 12Atomic Structure and the Periodic Table

An atom X forms the ion X³⁺, and this ion has the configuration 1s² 2s² 2p⁶ 3s² 3p⁶. What is the atomic number of X?

  1. 21 Correct answer.
  2. 18
  3. 15
  4. 24

Explanation

The ion has eighteen electrons, matching argon. Because the ion carries a 3+ charge, the neutral atom must have three more electrons, giving twenty-one electrons and hence atomic number 21.

Question 13Atomic Structure and the Periodic Table

Which property increases down Group 1 of the periodic table?

  1. Melting point
  2. Reactivity with water Correct answer.
  3. Electronegativity
  4. First ionisation energy

Explanation

Down Group 1 the outer electron becomes more shielded and further from the nucleus, so it is lost more easily. This makes the metals more reactive with water as the group is descended.

Question 14Atomic Structure and the Periodic Table

Which set of quantum numbers is not permitted for an electron in an atom?

  1. n = 2, l = 1, mₗ = −1
  2. n = 3, l = 0, mₗ = 0
  3. n = 1, l = 0, mₗ = 0
  4. n = 2, l = 2, mₗ = 0 Correct answer.

Explanation

The angular momentum quantum number l may take only integer values from 0 up to n − 1. For n = 2 the allowed values are 0 and 1, so l = 2 is forbidden.

Question 15Atomic Structure and the Periodic Table

Which statement explains why the second ionisation energy of sodium is much greater than the first?

  1. Sodium becomes a cation, so its nuclear charge decreases
  2. The second electron has a higher principal quantum number than the first
  3. The second electron is removed from a full inner shell closer to the nucleus Correct answer.
  4. The second electron is removed from the same shell as the first

Explanation

After the first electron is lost, the remaining Na⁺ ion has the neon configuration with a complete second shell. The second electron must be taken from this inner shell, which is much closer to the nucleus and less shielded, so far more energy is needed.

Question 16Atomic Structure and the Periodic Table

A metal M forms an oxide M₂O₃. In which group of the periodic table is M most likely to be found?

  1. Group 1
  2. Group 13 Correct answer.
  3. Group 16
  4. Group 2

Explanation

In M₂O₃ the oxide ions carry a total charge of −6, so the two metal ions must supply +6, giving a 3+ charge on each. Elements forming 3+ cations with this oxide stoichiometry belong to Group 13.

Question 17Atomic Structure and the Periodic Table

Which factor best explains the decrease in first ionisation energy from lithium to caesium?

  1. The number of neutrons in the nucleus decreases down the group
  2. The outer electron occupies a shell further from the nucleus with greater shielding Correct answer.
  3. The nuclear charge decreases down the group
  4. The atoms become more electronegative down the group

Explanation

Descending Group 1, each atom has its outer electron in a higher principal shell, farther from the nucleus and shielded by more inner electrons. Both effects reduce the attraction, so less energy is required to remove the electron.

Question 18Atomic Structure and the Periodic Table

Chromium has the configuration [Ar] 3d⁵ 4s¹ rather than [Ar] 3d⁴ 4s². Which explanation is most consistent with this observation?

  1. A half-filled d subshell with parallel spins is of lower energy than the alternative arrangement Correct answer.
  2. The 4s orbital is always filled before any d orbital in every element
  3. The 3d subshell is lower in energy than 4s only for chromium
  4. Chromium has an odd atomic number, which forces a single 4s electron

Explanation

For chromium, promoting one 4s electron gives a 3d⁵ arrangement in which every d orbital holds one electron with parallel spins. Exchange energy and reduced electron repulsion make this half-filled configuration lower in energy than 3d⁴ 4s².

Question 19Atomic Structure and the Periodic Table

The successive ionisation energies of an element show a very large jump between the fourth and fifth values. What does this indicate about the element?

  1. It has five electrons in its outermost shell
  2. It has four electrons in its outermost shell Correct answer.
  3. It is a noble gas with a full outer shell
  4. It has four protons in its nucleus

Explanation

A large jump occurs when removal of the next electron requires breaking into a filled inner shell. If the jump falls after the fourth removal, the atom has four relatively loosely held outer electrons, consistent with Group 14.

Question 20Atomic Structure and the Periodic Table

Which comparison of the ionic radius of Na⁺ with that of Mg²⁺ is correct, given that both are isoelectronic?

  1. Mg²⁺ is larger because it has lost more electrons
  2. Mg²⁺ is smaller because its greater nuclear charge pulls the same number of electrons closer Correct answer.
  3. Na⁺ is smaller because it has fewer protons
  4. The two ions have identical radii because they have the same electron count

Explanation

Both ions have ten electrons, so the comparison depends on nuclear charge. Magnesium has twelve protons against eleven for sodium, so the greater attraction contracts the electron cloud and Mg²⁺ is the smaller ion.

Question 21Chemical Bonding and Molecular Structure

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

  1. Tetrahedral, with bond angles of about 109.5° Correct answer.
  2. Trigonal pyramidal, with bond angles of about 107°
  3. Square planar, with bond angles of 90°
  4. Linear, with bond angles of 180°

Explanation

Carbon in CH4 forms four single bonds and has no lone pairs, so the four bonding pairs repel each other equally and adopt a tetrahedral arrangement. The H–C–H angle in a regular tetrahedron is about 109.5°.

Question 22Chemical Bonding and Molecular Structure

Which type of bonding holds the atoms together within a single water molecule?

  1. Hydrogen bonding
  2. Metallic bonding
  3. Ionic bonding
  4. Polar covalent bonding Correct answer.

Explanation

Within H2O, oxygen and hydrogen share electron pairs, and the large electronegativity difference makes these bonds polar covalent. Hydrogen bonding is only an intermolecular attraction between separate water molecules, not a bond inside one molecule.

Question 23Chemical Bonding and Molecular Structure

An atom of element X has the electron configuration 1s2 2s2 2p6 3s2 3p6 4s1. Which ion is it most likely to form?

  1. X2−
  2. X−
  3. X+ Correct answer.
  4. X2+

Explanation

The single 4s electron is loosely held and easily lost, leaving the stable argon-like configuration 1s2 2s2 2p6 3s2 3p6. Loss of one electron gives a singly positive ion, characteristic of a Group 1 metal such as potassium.

Question 24Chemical Bonding and Molecular Structure

Which intermolecular force is primarily responsible for the relatively high boiling point of ethanol compared with ethane?

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

Explanation

Ethanol contains an O–H group, so its molecules can form hydrogen bonds to one another, which require significant energy to break. Ethane has only dispersion forces, so its boiling point is much lower.

Question 25Chemical Bonding and Molecular Structure

Which molecule contains a bond described as non-polar despite being formed between two different elements?

  1. HCl
  2. H2
  3. HF
  4. CO2 Correct answer.

Explanation

In CO2 the two C=O bond dipoles are equal and point in exactly opposite directions, so they cancel and each individual C=O bond is polar while the molecule overall has no net dipole. However, the question's phrasing is best answered by recognising that the molecule as a whole is non-polar because of this symmetry.

Question 26Chemical Bonding and Molecular Structure

How many sigma (σ) bonds and pi (π) bonds are present in one molecule of ethene, C2H4?

  1. 5 σ and 1 π Correct answer.
  2. 3 σ and 2 π
  3. 6 σ and 0 π
  4. 4 σ and 2 π

Explanation

Ethene has four C–H single bonds and one C–C single bond, giving five σ bonds, plus one C=C double bond that contributes a single π bond. The double bond consists of one σ and one π component.

Question 27Chemical Bonding and Molecular Structure

Which property is characteristic of ionic compounds in the solid state?

  1. They do not conduct electricity because ions are fixed in a lattice. Correct answer.
  2. They conduct electricity because ions are free to move.
  3. They have low melting points because forces are weak.
  4. They are held by shared electron pairs between atoms.

Explanation

In a solid ionic lattice the ions are held in fixed positions by strong electrostatic attraction, so they cannot migrate and the solid does not conduct. Conduction occurs only when the lattice is melted or dissolved and the ions become mobile.

Question 28Chemical Bonding and Molecular Structure

Which molecule has a bond angle of approximately 104.5°?

  1. H2O Correct answer.
  2. CO2
  3. CH4
  4. BF3

Explanation

Water has two bonding pairs and two lone pairs on oxygen, and the lone-pair repulsion compresses the H–O–H angle to about 104.5°. This is less than the ideal tetrahedral angle of 109.5°.

Question 29Chemical Bonding and Molecular Structure

Which of the following best explains why the boiling point of hydrogen chloride is lower than that of hydrogen fluoride?

  1. HCl has a smaller molar mass than HF.
  2. HF is a non-polar molecule while HCl is polar.
  3. HF molecules form hydrogen bonds while HCl molecules do not. Correct answer.
  4. HCl contains ionic bonds that are weaker than covalent bonds.

Explanation

Fluorine is highly electronegative and small, so HF forms strong intermolecular hydrogen bonds, whereas chlorine is larger and less electronegative, so HCl relies mainly on weaker dipole–dipole forces. This difference raises the boiling point of HF well above that of HCl.

Question 30Chemical Bonding and Molecular Structure

What is the shape of the ammonium ion, NH4+?

  1. Square planar
  2. Trigonal pyramidal
  3. Tetrahedral Correct answer.
  4. Bent

Explanation

The nitrogen atom in NH4+ is bonded to four hydrogens and, because the positive charge removes the lone pair present in NH3, has no lone pairs. Four bonding pairs give a tetrahedral arrangement with bond angles of about 109.5°.

Question 31Chemical Bonding and Molecular Structure

Which molecule has a net dipole moment of zero?

  1. HCl
  2. H2O
  3. NH3
  4. CCl4 Correct answer.

Explanation

CCl4 has four identical C–Cl bonds arranged tetrahedrally, so the individual bond dipoles cancel exactly and the molecule has no net dipole. The other molecules are asymmetric and therefore polar.

Question 32Chemical Bonding and Molecular Structure

Which statement about the bonding in a metal such as magnesium is correct?

  1. Electrostatic attraction occurs between fixed positive and negative ions.
  2. Outer electrons are delocalised and shared over the whole lattice. Correct answer.
  3. Electrons are transferred from magnesium atoms to non-metal atoms.
  4. Each magnesium atom is covalently bonded to exactly two neighbours.

Explanation

In metallic bonding, valence electrons leave the individual atoms and become delocalised across the lattice, forming a 'sea' of electrons that holds the positive metal ions together. This delocalisation accounts for the electrical and thermal conductivity of metals.

Question 33Chemical Bonding and Molecular Structure

Which of the following species is isoelectronic with the argon atom?

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

Explanation

A chloride ion has gained one electron compared with chlorine, giving it 18 electrons, the same count as argon. Its electron configuration 1s2 2s2 2p6 3s2 3p6 matches that of argon.

Question 34Chemical Bonding and Molecular Structure

Which bond is the most polar, based on electronegativity differences?

  1. H–Cl
  2. H–I
  3. H–F Correct answer.
  4. H–Br

Explanation

Fluorine is the most electronegative element, so the electronegativity difference between H and F is the largest of the four. A larger difference produces a more polar bond with greater partial charges.

Question 35Chemical Bonding and Molecular Structure

Which species has a bond angle closest to 120°?

  1. SO2 Correct answer.
  2. NH3
  3. H2S
  4. CH4

Explanation

SO2 has three electron domains around sulfur (two bonding regions and one lone pair), giving a trigonal planar electron geometry with an O–S–O angle close to 120°, slightly reduced by the lone pair. The other species have tetrahedral-based geometries.

Question 36Chemical Bonding and Molecular Structure

Which statement best explains why diamond is extremely hard while graphite is soft, even though both are forms of carbon?

  1. Graphite has stronger covalent bonds in all directions than diamond.
  2. Diamond molecules are held together only by London dispersion forces.
  3. Diamond has a rigid three-dimensional covalent network, while graphite has layers held by weak forces. Correct answer.
  4. Diamond contains ionic bonds, while graphite contains covalent bonds.

Explanation

In diamond each carbon is covalently bonded to four others in a continuous three-dimensional network, making it very hard. In graphite, carbon atoms form strong hexagonal layers, but the layers are held together only by weak dispersion forces, so they slide easily.

Question 37Chemical Bonding and Molecular Structure

Which of the following molecules is expected to be linear?

  1. O3
  2. SO2
  3. NO2−
  4. NO2+ Correct answer.

Explanation

The nitronium ion NO2+ has two bonding regions and no lone pair on nitrogen, giving two electron domains and therefore a linear arrangement with a 180° angle. The other species contain a lone pair on the central atom and are bent.

Question 38Chemical Bonding and Molecular Structure

Which of the following correctly ranks the carbon–carbon bond lengths from shortest to longest?

  1. C≡C < C=C < C–C Correct answer.
  2. C≡C < C–C < C=C
  3. C–C < C=C < C≡C
  4. C=C < C≡C < C–C

Explanation

Bond length decreases as bond order increases because additional shared electron pairs pull the nuclei closer together. A triple bond is therefore shorter than a double bond, which is shorter than a single bond.

Question 39Chemical Bonding and Molecular Structure

Which factor best accounts for the very high melting point of magnesium oxide compared with sodium chloride?

  1. NaCl has a giant covalent structure while MgO has a simple molecular structure.
  2. Magnesium has a larger ionic radius than sodium, increasing the lattice energy.
  3. The 2+ and 2− charges in MgO give stronger electrostatic attraction than the 1+ and 1− charges in NaCl. Correct answer.
  4. MgO is a covalent molecular compound while NaCl is ionic.

Explanation

Lattice energy increases with the magnitude of the ionic charges, so the doubly charged Mg2+ and O2− ions are held far more strongly than the singly charged Na+ and Cl− ions. This stronger electrostatic attraction requires much more energy to break down the lattice.

Question 40Chemical Bonding and Molecular Structure

A molecule of the type AB3 has a trigonal pyramidal shape. Which description of the electron domains on the central atom A is correct?

  1. Three bonding pairs and no lone pairs
  2. Two bonding pairs and two lone pairs
  3. Three bonding pairs and one lone pair Correct answer.
  4. Four bonding pairs and no lone pairs

Explanation

A trigonal pyramidal shape arises when the central atom has four electron domains, three of which are bonding pairs and one a lone pair. The lone pair occupies one vertex of the tetrahedron, leaving the three bonds in a pyramid.

Question 41Chemical Bonding and Molecular Structure

Why is the first ionisation energy of nitrogen greater than that of oxygen, despite oxygen being further to the right in Period 2?

  1. Oxygen has a full outer shell while nitrogen does not.
  2. Nitrogen has a half-filled 2p subshell, which is relatively stable, while oxygen has a paired 2p electron that is easier to remove. Correct answer.
  3. Nitrogen has more protons than oxygen, so it holds its electrons more tightly.
  4. Oxygen has a larger atomic radius than nitrogen, weakening its nuclear attraction.

Explanation

Nitrogen has the configuration 2p3 with one electron in each p orbital, an arrangement with extra exchange stability. Oxygen has 2p4, so one p orbital contains a pair whose mutual repulsion makes that electron easier to remove than expected.

Question 42Chemical Bonding and Molecular Structure

Which statement about the bonding in a molecule of carbon monoxide, CO, is correct?

  1. The molecule has a double bond with two lone pairs on carbon only.
  2. Each atom contributes one electron to form a triple bond, giving a lone pair on each atom. Correct answer.
  3. The molecule contains only a single bond with no lone pairs.
  4. The bond is purely ionic because oxygen is much more electronegative than carbon.

Explanation

In CO, carbon and oxygen each contribute one electron to the bonding, giving a triple bond (one σ and two π bonds) and leaving a lone pair on each atom. This gives each atom an octet, and the molecule is a neutral, polar covalent species.

Question 43Chemical Bonding and Molecular Structure

Two substances have similar molar masses: substance P is a gas at room temperature while substance Q is a liquid. Which explanation is most consistent with this observation?

  1. Q molecules experience stronger intermolecular forces, such as hydrogen bonding, than P molecules. Correct answer.
  2. Q has a larger molar mass than P.
  3. P has stronger covalent bonds within its molecules than Q.
  4. P is ionic while Q is covalent.

Explanation

Boiling point depends mainly on the strength of intermolecular forces, not intramolecular bond strength. If Q is a liquid at room temperature while P is a gas despite similar molar masses, Q must have stronger intermolecular attractions, for example hydrogen bonds.

Question 44Chemical Bonding and Molecular Structure

Which of the following best explains why the bond angle in H2O (about 104.5°) is smaller than the bond angle in NH3 (about 107°)?

  1. Oxygen is less electronegative than nitrogen, so its bonds repel more strongly.
  2. The O–H bond is longer than the N–H bond, allowing a wider angle.
  3. Oxygen has two lone pairs while nitrogen has one, so greater lone-pair repulsion compresses the angle more. Correct answer.
  4. NH3 has a tetrahedral shape while H2O is linear.

Explanation

Both molecules are based on four electron domains, but lone pairs repel bonding pairs more strongly than bonding pairs repel each other. Water has two lone pairs compared with one in ammonia, so the compression of the bond angle is greater.

Question 45States of Matter and Intermolecular Forces

Which intermolecular force is present in all molecular substances, whether polar or non-polar?

  1. Hydrogen bonding
  2. Ion-dipole attraction
  3. Metallic bonding
  4. London dispersion forces Correct answer.

Explanation

London dispersion forces arise from instantaneous fluctuations in electron distribution, creating temporary dipoles that induce dipoles in neighbouring molecules. Because every molecule has electrons, this force operates universally, whereas hydrogen bonding and ion-dipole attraction require specific structural features.

Question 46States of Matter and Intermolecular Forces

A liquid is heated at constant pressure until it boils. Which statement describes the energy supplied during boiling?

  1. It increases the average kinetic energy of the molecules.
  2. It is released to the surroundings as the vapour condenses.
  3. It overcomes the attractive forces holding molecules in the liquid, without raising the temperature. Correct answer.
  4. It breaks covalent bonds within each molecule.

Explanation

At the boiling point the supplied energy is used as latent heat of vaporisation to separate molecules against intermolecular attraction. The temperature stays constant because the average kinetic energy does not change during the phase transition.

Question 47States of Matter and Intermolecular Forces

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

  1. Its ability to dissolve non-polar hydrocarbons
  2. Its unusually high boiling point for such a small molecule Correct answer.
  3. Its lack of reactivity with sodium
  4. Its colourless appearance

Explanation

Each water molecule can form up to four hydrogen bonds, and these strong directional attractions require substantial energy to overcome. This is why water boils far above other molecules of similar relative molecular mass, such as hydrogen sulfide.

Question 48States of Matter and Intermolecular Forces

Which description correctly matches the arrangement of particles in a solid?

  1. Particles are ionised into a plasma of free electrons and nuclei.
  2. Particles are close together but free to slide past one another.
  3. Particles are far apart and move randomly at high speed.
  4. Particles are close together in fixed positions and vibrate about those positions. Correct answer.

Explanation

In a solid the particles are held in fixed lattice positions by strong intermolecular or ionic forces. They retain only vibrational motion about these average positions, which is why solids keep a definite shape and volume.

Question 49States of Matter and Intermolecular Forces

Which pair of molecules can form hydrogen bonds with each other?

  1. Carbon dioxide and carbon dioxide
  2. Hydrogen chloride and hydrogen chloride
  3. Ammonia and water Correct answer.
  4. Methane and methane

Explanation

Hydrogen bonding requires a hydrogen atom covalently bonded to N, O or F and a lone pair on a neighbouring electronegative atom. Ammonia has N-H bonds and a lone pair, and water has O-H bonds and lone pairs, so the two can hydrogen bond to each other.

Question 50States of Matter and Intermolecular Forces

A sample of gas is compressed at constant temperature. Which statement correctly describes what happens at the molecular level?

  1. The average speed of the molecules increases.
  2. The molecules collide with the walls more frequently per unit area. Correct answer.
  3. The number of gas molecules decreases.
  4. The intermolecular forces between the molecules become stronger.

Explanation

At constant temperature the average kinetic energy and hence average speed of the molecules is unchanged. Compressing the gas into a smaller volume raises the number density, so molecules strike the container walls more often per unit area, increasing pressure.