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JCE Chemistry practice questions

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Question 1Matter, Atomic Structure and the Periodic Table

An atom of an element has the electron configuration 1s² 2s² 2p⁶ 3s² 3p⁴. In which block and group of the periodic table does this element belong?

  1. p block, group 14
  2. s block, group 2
  3. p block, group 16 Correct answer.
  4. d block, group 16

Explanation

The last electrons enter a 3p subshell, so the element is in the p block. The outer configuration 3s² 3p⁴ gives six valence electrons, placing it in group 16 (oxygen family), which is sulfur.

Question 2Matter, Atomic Structure and the Periodic Table

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

  1. It generally increases because nuclear charge rises while shielding stays roughly constant. Correct answer.
  2. It increases steadily with no irregularities at any element.
  3. It generally decreases because atomic radius increases across the period.
  4. It stays constant because each atom has the same number of shells.

Explanation

Across a period, protons are added while electrons enter the same shell, so effective nuclear charge rises and the outer electron is held more tightly. The general increase is interrupted at aluminium and sulfur because of subshell and pairing effects.

Question 3Matter, Atomic Structure and the Periodic Table

A sample of element X contains two isotopes: ⁶³X with abundance 70% and ⁶⁵X with abundance 30%. What is the relative atomic mass of X?

  1. 63.6 Correct answer.
  2. 65.0
  3. 64.4
  4. 64.0

Explanation

The relative atomic mass is the abundance-weighted mean of the isotope masses: (63 × 0.70) + (65 × 0.30) = 44.1 + 19.5 = 63.6. This is a weighted average, not a simple arithmetic mean of the two masses.

Question 4Matter, Atomic Structure and the Periodic Table

Which subatomic particles are found in the nucleus of an atom?

  1. Electrons only
  2. Protons and neutrons Correct answer.
  3. Neutrons and electrons
  4. Protons and electrons

Explanation

The nucleus contains protons and neutrons, which together account for nearly all the mass of the atom. Electrons occupy the region outside the nucleus in orbitals.

Question 5Matter, Atomic Structure and the Periodic Table

Which of the following is the correct electron configuration of a chloride ion, Cl⁻ (Z = 17)?

  1. 1s² 2s² 2p⁶ 3s¹
  2. 1s² 2s² 2p⁶ 3s² 3p⁵
  3. 1s² 2s² 2p⁶ 3s² 3p⁶ Correct answer.
  4. 1s² 2s² 2p⁶ 3s² 3p⁴

Explanation

A chlorine atom has 17 electrons, and gaining one electron gives the chloride ion 18 electrons. These fill as 1s² 2s² 2p⁶ 3s² 3p⁶, the same configuration as argon.

Question 6Matter, Atomic Structure and the Periodic Table

Which element has the highest electronegativity value on the Pauling scale?

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

Explanation

Electronegativity increases across a period and decreases down a group, so the top-right element fluorine has the highest value. Fluorine's small size and high effective nuclear charge make it the most electronegative element.

Question 7Matter, Atomic Structure and the Periodic Table

Which of these elements is a transition metal?

  1. Calcium
  2. Gallium
  3. Iron Correct answer.
  4. Potassium

Explanation

Transition metals are d-block elements with a partially filled d subshell in at least one common oxidation state, such as iron (3d⁶ 4s²). Calcium and potassium are s-block metals, and gallium is a p-block metal.

Question 8Matter, Atomic Structure and the Periodic Table

How many neutrons are present in an atom of ⁵⁶Fe (Z = 26)?

  1. 56
  2. 26
  3. 30 Correct answer.
  4. 82

Explanation

The mass number (56) is the sum of protons and neutrons, so the neutron number is 56 − 26 = 30. The atomic number 26 gives the number of protons.

Question 9Matter, Atomic Structure and the Periodic Table

Which statement best explains why the first ionisation energy of magnesium is higher than that of aluminium?

  1. Aluminium has a greater number of electron shells than magnesium.
  2. Magnesium has a larger atomic radius than aluminium.
  3. The outer electron in magnesium is removed from a filled 3s orbital, which is more stable than the 3p orbital of aluminium. Correct answer.
  4. Aluminium has fewer protons in its nucleus than magnesium.

Explanation

Magnesium's outer electron is in a filled 3s subshell, whereas aluminium's is in a 3p subshell that is higher in energy and better shielded. The 3p electron is therefore easier to remove despite the greater nuclear charge of aluminium.

Question 10Matter, Atomic Structure and the Periodic Table

Which of these ions has the electron configuration 1s² 2s² 2p⁶?

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

Explanation

Na⁺ has 10 electrons (11 − 1) and Mg²⁺ has 10 electrons (12 − 2), so both adopt the neon configuration 1s² 2s² 2p⁶. They are isoelectronic with each other.

Question 11Matter, Atomic Structure and the Periodic Table

The successive ionisation energies of an element show a very large jump between the third and fourth ionisation energies. To which group does the element most likely belong?

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

Explanation

A large jump after the third ionisation means the fourth electron is removed from a new, inner shell that is much harder to ionise. This indicates three valence electrons, which places the element in group 3 (group 13 in IUPAC numbering).

Question 12Matter, Atomic Structure and the Periodic Table

Which property increases down group 2 from magnesium to barium?

  1. First ionisation energy
  2. Atomic radius Correct answer.
  3. Electronegativity
  4. Melting point

Explanation

Each successive group 2 element has an extra occupied shell, so atomic radius increases down the group. Ionisation energy and electronegativity decrease, and melting point generally falls as the metallic bonding weakens.

Question 13Matter, Atomic Structure and the Periodic Table

Which of these species is isoelectronic with the argon atom?

  1. All of K⁺, Ca²⁺ and S²⁻ Correct answer.
  2. K⁺
  3. Ca²⁺
  4. S²⁻

Explanation

K⁺ (19 − 1 = 18), Ca²⁺ (20 − 2 = 18) and S²⁻ (16 + 2 = 18) all contain 18 electrons, the same number as argon. They are therefore isoelectronic with argon.

Question 14Matter, Atomic Structure and the Periodic Table

Which statement correctly describes the trend in atomic radius across period 3 from sodium to chlorine?

  1. It decreases because effective nuclear charge increases while the number of shells stays the same. Correct answer.
  2. It decreases because the number of occupied shells decreases.
  3. It remains constant because each atom has three occupied shells.
  4. It increases because more electrons are added to the outer shell.

Explanation

Across period 3, protons are added to the nucleus while electrons enter the same third shell, so effective nuclear charge increases and pulls the electrons closer. The radius therefore decreases from sodium to chlorine.

Question 15Matter, Atomic Structure and the Periodic Table

Which of these elements has the largest atomic radius?

  1. Rubidium Correct answer.
  2. Sodium
  3. Lithium
  4. Potassium

Explanation

Atomic radius increases down a group as each element gains an additional occupied shell. Rubidium lies furthest down group 1 among these options, so it has the largest radius.

Question 16Matter, Atomic Structure and the Periodic Table

An element has the electron configuration [Ar] 3d⁵ 4s¹. Which statement about this element is correct?

  1. It is a d-block element with a full 3d subshell.
  2. It is a p-block element with one valence electron.
  3. It is a d-block element with five unpaired d electrons. Correct answer.
  4. It is an s-block element with a full d subshell.

Explanation

The configuration [Ar] 3d⁵ 4s¹ corresponds to chromium, a d-block element. The half-filled 3d⁵ subshell contains five unpaired electrons, which contributes to its unusual stability.

Question 17Matter, Atomic Structure and the Periodic Table

Which of these elements would you expect to have the lowest first ionisation energy?

  1. Iodine Correct answer.
  2. Bromine
  3. Chlorine
  4. Fluorine

Explanation

First ionisation energy decreases down group 17 because the outer electron is further from the nucleus and more shielded. Iodine, being the lowest of these elements in the group, has the lowest ionisation energy.

Question 18Matter, Atomic Structure and the Periodic Table

Which statement about the isotopes ¹²C and ¹⁴C is correct?

  1. They have the same number of neutrons but different numbers of protons.
  2. They have the same number of protons and the same number of neutrons.
  3. They have different numbers of protons and different numbers of neutrons.
  4. They have the same number of protons but different numbers of neutrons. Correct answer.

Explanation

Isotopes of an element have the same atomic number, so both carbon isotopes contain six protons. They differ in mass number, so ¹²C has six neutrons while ¹⁴C has eight.

Question 19Matter, Atomic Structure and the Periodic Table

Which of these elements has the highest second ionisation energy?

  1. Sodium Correct answer.
  2. Aluminium
  3. Magnesium
  4. Silicon

Explanation

Sodium's second ionisation removes an electron from the stable neon-like 2p⁶ configuration, requiring a large amount of energy. The other elements retain valence electrons in the third shell for their second ionisation, which is much easier.

Question 20Matter, Atomic Structure and the Periodic Table

Which of the following correctly lists the elements in order of increasing electronegativity?

  1. Cl < P < Si < Na
  2. Si < Na < P < Cl
  3. Na < Si < P < Cl Correct answer.
  4. Na < P < Cl < Si

Explanation

Electronegativity increases across a period from left to right, so the order is Na < Si < P < Cl. Sodium is a metal with low electronegativity, while chlorine is near the top right of period 3.

Question 21Matter, Atomic Structure and the Periodic Table

Which of these statements about the periodic table is correct?

  1. Elements in the same group have the same number of occupied electron shells.
  2. Elements in the same group have similar chemical properties because they have the same number of valence electrons. Correct answer.
  3. Elements in the same period have identical chemical properties.
  4. Elements in the same period have the same number of valence electrons.

Explanation

Members of a group share the same valence electron configuration, which governs their chemical behaviour, so they show similar properties. Elements in the same period have the same number of shells but different valence configurations.

Question 22Matter, Atomic Structure and the Periodic Table

A sample of copper contains 69% of ⁶³Cu and 31% of ⁶⁵Cu. Which value is closest to the relative atomic mass of copper?

  1. 63.6 Correct answer.
  2. 64.0
  3. 65.0
  4. 64.4

Explanation

The weighted mean is (63 × 0.69) + (65 × 0.31) = 43.47 + 20.15 = 63.62, which rounds to 63.6. The result lies closer to 63 because that isotope is more abundant.

Question 23Matter, Atomic Structure and the Periodic Table

Which of these elements has an electron configuration that ends in ns² np⁶ and is therefore a noble gas?

  1. Krypton (Z = 36)
  2. Neon (Z = 10)
  3. Argon (Z = 18)
  4. All of neon, argon and krypton Correct answer.

Explanation

Neon, argon and krypton all have full outer p subshells (2p⁶, 3p⁶ and 4p⁶ respectively), which is the characteristic noble gas configuration. A complete octet gives them low reactivity.

Question 24Matter, Atomic Structure and the Periodic Table

An element has successive ionisation energies (kJ mol⁻¹) of 590, 1150, 4940 and 6480. Which group does the element belong to?

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

Explanation

The large jump occurs between the second and third ionisations (1150 to 4940), meaning two electrons are relatively easy to remove and the third comes from an inner shell. This indicates two valence electrons, consistent with group 2.

Question 25Matter, Atomic Structure and the Periodic Table

Which of these ions has the smallest ionic radius?

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

Explanation

Na⁺, Mg²⁺ and Al³⁺ are isoelectronic with 10 electrons, but Al³⁺ has the greatest nuclear charge (13 protons), so it attracts the electrons most strongly and has the smallest radius. F⁻ also has 10 electrons but only 9 protons.

Question 26Matter, Atomic Structure and the Periodic Table

Which subatomic particle determines the identity of an element?

  1. The proton, because its count equals the atomic number Correct answer.
  2. The neutron, because it changes between isotopes
  3. The electron, because it occupies shells
  4. The nucleon, because it includes both protons and neutrons

Explanation

The atomic number Z counts protons in the nucleus, and this number is unique to each element. Changing the proton count converts the atom into a different element, whereas changing neutrons or electrons leaves the element identity unchanged.

Question 27Matter, Atomic Structure and the Periodic Table

A potassium atom loses one electron to form K⁺. Which statement describes the resulting particle?

  1. It has 20 protons and 19 electrons
  2. It has 18 protons and 19 electrons
  3. It has 19 protons and 20 electrons
  4. It has 19 protons and 18 electrons Correct answer.

Explanation

Potassium has atomic number 19, so every potassium nucleus holds 19 protons. Removing one electron from a neutral atom leaves 18 electrons, giving an overall charge of +1.

Question 28Matter, Atomic Structure and the Periodic Table

In the periodic table, an element is placed in period 3 and group 16. Which description of its atom is correct?

  1. It has sixteen occupied shells and three outer electrons
  2. It has three outer electrons and sixteen protons
  3. It has three occupied electron shells and six outer-shell electrons Correct answer.
  4. It has six occupied shells and three outer electrons

Explanation

The period number gives the number of occupied electron shells, so period 3 means three shells. The group number for main-group elements gives the number of outer-shell electrons, so group 16 means six outer electrons, as in sulfur.

Question 29Matter, Atomic Structure and the Periodic Table

Which statement about the shape of the periodic table is correct?

  1. Groups are horizontal rows in order of increasing atomic number
  2. Periods are vertical columns containing elements with equal proton numbers
  3. Groups are vertical columns containing elements with similar outer-shell configurations Correct answer.
  4. Periods contain only metals, while groups contain only non-metals

Explanation

Vertical columns are groups, and elements within a group share the same number of outer-shell electrons, which produces similar chemical behaviour. Horizontal rows are periods, arranged by increasing atomic number.

Question 30Matter, Atomic Structure and the Periodic Table

A neutral atom of an element in group 2 loses two electrons. What is the electron configuration of the resulting ion if the parent atom is calcium (Z = 20)?

  1. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
  2. 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
  3. 1s² 2s² 2p⁶ 3s² 3p⁶ Correct answer.
  4. 1s² 2s² 2p⁶ 3s¹

Explanation

A calcium atom has the configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². Removing the two 4s electrons gives Ca²⁺, whose configuration is that of argon, 1s² 2s² 2p⁶ 3s² 3p⁶.

Question 31Matter, Atomic Structure and the Periodic Table

Why does the first ionisation energy generally increase across period 3 from sodium to argon?

  1. The number of neutrons in the nucleus increases steadily across the period
  2. The number of occupied shells increases, drawing outer electrons closer to the nucleus
  3. Atomic radius increases, so outer electrons are removed more easily
  4. Nuclear charge increases while shielding stays similar, so outer electrons are held more strongly Correct answer.

Explanation

Across a period, protons are added to the nucleus while electrons enter the same outer shell, so shielding changes little. The growing effective nuclear charge pulls the outer shell inwards and raises the energy needed to remove an electron.

Question 32Matter, Atomic Structure and the Periodic Table

Which pair of species contains the same number of electrons?

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

Explanation

Na⁺ has 11 − 1 = 10 electrons and F⁻ has 9 + 1 = 10 electrons, so the two are isoelectronic. Comparing electron counts after adjusting for charge identifies such pairs.

Question 33Matter, Atomic Structure and the Periodic Table

Which statement correctly explains why the ionic radius of Mg²⁺ is smaller than that of Na⁺?

  1. Mg²⁺ has more neutrons, which contracts the nucleus
  2. Mg²⁺ has a smaller nuclear charge but a larger mass
  3. Mg²⁺ has one more proton, so the same electron shell is pulled in more strongly Correct answer.
  4. Mg²⁺ has lost an extra shell of electrons, leaving fewer occupied shells

Explanation

Both ions have the 10-electron configuration of neon, but Mg²⁺ has 12 protons against 11 for Na⁺. The greater nuclear charge attracts the same electron cloud more strongly, so the radius is smaller.

Question 34Matter, Atomic Structure and the Periodic Table

Which element would be expected to show the greatest tendency to form a 2− ion?

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

Explanation

Oxygen has the configuration 1s² 2s² 2p⁴ and needs two electrons to complete the 2p subshell, so it readily forms O²⁻. Fluorine needs only one electron, and nitrogen and carbon would need three or four.

Question 35Matter, Atomic Structure and the Periodic Table

Which property shows a general increase down group 1 from lithium to caesium?

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

Explanation

Down group 1 the outer electron lies in a shell further from the nucleus and is more shielded, so it is lost more easily. This makes the metals more reactive with water, producing hydrogen more vigorously.

Question 36Matter, Atomic Structure and the Periodic Table

An element has the electron configuration [Ne] 3s² 3p⁴. Which statement about this element is correct?

  1. It is in group 14 and forms a hydride with the formula XH₄
  2. It is in group 18 and does not form hydrides
  3. It is in group 6 and forms a hydride with the formula HX
  4. It is in group 16 and forms a hydride with the formula H₂X Correct answer.

Explanation

The configuration [Ne] 3s² 3p⁴ shows six outer electrons, placing the element in group 16, which is sulfur. Group 16 elements need two electrons to complete their octet, so they form hydrides such as H₂S.

Question 37Matter, Atomic Structure and the Periodic Table

Which factor provides the best explanation for the decrease in atomic radius from lithium to fluorine across period 2?

  1. The number of occupied shells falls from two to one
  2. The shielding by inner electrons increases sharply
  3. The atoms lose electrons more readily, shrinking the outer shell
  4. Effective nuclear charge rises while electrons are added to the same shell Correct answer.

Explanation

Across period 2, protons are added while the new electrons enter the 2s and 2p orbitals of the same shell. The increased effective nuclear charge contracts the electron cloud, so the radius decreases.

Question 38Matter, Atomic Structure and the Periodic Table

Successive ionisation energies of an element are 738, 1451, 7733, 10540 and 13630 kJ mol⁻¹. Which conclusion is best supported?

  1. The element is in group 3 because the large jump follows the third electron
  2. The element is in group 14 because four electrons are removed before the jump
  3. The element is in group 2 because the large jump follows the second electron Correct answer.
  4. The element is in group 1 because the first electron is easiest to remove

Explanation

The jump from 1451 to 7733 kJ mol⁻¹ shows that the third electron comes from a shell much closer to the nucleus. Two relatively easy removals indicate two outer-shell electrons, so the element belongs to group 2.

Question 39Matter, Atomic Structure and the Periodic Table

Which statement best accounts for the very high electronegativity of fluorine compared with other elements?

  1. Its low nuclear charge leaves electrons weakly attracted
  2. Its large atomic radius allows it to attract electrons from a greater distance
  3. Its outer electrons are strongly shielded by many inner shells
  4. Its small atomic radius concentrates the nuclear charge on the bonding pair of electrons Correct answer.

Explanation

Fluorine has nine protons and only two occupied shells, so its outer shell is close to the nucleus and poorly shielded. This high effective nuclear charge gives it the strongest attraction for bonding electrons.

Question 40Matter, Atomic Structure and the Periodic Table

Which statement about the trends in first ionisation energy across period 3 is correct?

  1. It rises steadily with no interruptions because nuclear charge always dominates
  2. It rises overall, but drops at aluminium and at sulfur because of subshell and pairing effects Correct answer.
  3. It falls steadily from sodium to argon without exception
  4. It is highest at sodium and lowest at chlorine

Explanation

The overall trend is upward because effective nuclear charge increases, but two dips occur. Aluminium's outer electron is in 3p, which is higher in energy and better shielded than 3s, and sulfur has a paired 3p electron whose repulsion makes removal easier than from phosphorus.

Question 41Chemical Bonding and Structure

Which statement correctly describes the electrostatic attraction in an ionic bond?

  1. Attraction between a shared pair of electrons and two nuclei
  2. Attraction between instantaneous dipoles in neighbouring molecules
  3. Attraction between oppositely charged ions formed by electron transfer Correct answer.
  4. Attraction between a lone pair donor and an electron-deficient atom

Explanation

An ionic bond arises when one atom transfers electrons to another, producing a cation and an anion that attract electrostatically in a lattice. Shared electron pairs describe covalent bonding, lone-pair donation describes dative bonding, and instantaneous dipoles describe van der Waals forces.

Question 42Chemical Bonding and Structure

Which pair of elements would be most likely to form an ionic compound when reacted together?

  1. Sodium and chlorine Correct answer.
  2. Carbon and oxygen
  3. Sulfur and oxygen
  4. Nitrogen and hydrogen

Explanation

Ionic bonding is favoured when a low-electronegativity metal transfers electrons to a high-electronegativity non-metal. Sodium has a very low first ionisation energy and chlorine has a high electron affinity, so electron transfer to give Na+ and Cl- is strongly favourable.

Question 43Chemical Bonding and Structure

Which formula is correct for the ionic compound formed between magnesium ions and nitride ions?

  1. Mg2N3
  2. Mg3N2 Correct answer.
  3. MgN
  4. Mg2N

Explanation

Magnesium forms Mg2+ ions and nitrogen forms N3- ions. Charge balance requires three Mg2+ ions for every two N3- ions, giving the formula Mg3N2.

Question 44Chemical Bonding and Structure

Which property is expected for a solid with a giant ionic lattice?

  1. It dissolves readily in non-polar solvents
  2. It conducts electricity well in the solid state
  3. It has a low melting point because forces are weak
  4. It is brittle and shatters when a layer is displaced Correct answer.

Explanation

In an ionic lattice, displacing a layer brings ions of like charge alongside each other, so repulsion splits the crystal; this makes ionic solids brittle. Ionic solids do not conduct when solid because ions are fixed, have high melting points, and dissolve in polar rather than non-polar solvents.

Question 45Chemical Bonding and Structure

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

  1. Carbon dioxide Correct answer.
  2. Ammonia
  3. Water
  4. Hydrogen chloride

Explanation

Carbon dioxide is linear, so the two polar C=O bond dipoles point in opposite directions and cancel exactly, giving zero net dipole. HCl, NH3 and H2O have polar bonds whose dipoles do not cancel, so each molecule has a permanent dipole.

Question 46Chemical Bonding and Structure

Which type of crystal contains delocalised electrons that allow electrical conduction in the solid state?

  1. Ionic crystal such as sodium chloride
  2. Simple molecular crystal such as iodine
  3. Metallic crystal such as copper Correct answer.
  4. Giant covalent crystal such as diamond

Explanation

In a metal, outer electrons are delocalised over the whole lattice and can move under an applied potential difference, so the solid conducts. Ionic and simple molecular solids have localised electrons, and diamond's electrons are all locked in sigma bonds.

Question 47Chemical Bonding and Structure

Which species has a bond angle of approximately 180 degrees?

  1. SO2
  2. BF3
  3. NH4+
  4. CO2 Correct answer.

Explanation

CO2 has two electron domains around the central carbon and no lone pairs, so the VSEPR arrangement is linear with a bond angle of 180 degrees. NH4+ and BF3 are tetrahedral and trigonal planar respectively, while SO2 is bent.

Question 48Chemical Bonding and Structure

Which statement correctly explains why diamond has a very high melting point?

  1. Weak intermolecular forces between layers must be overcome
  2. Delocalised electrons hold the carbon layers together
  3. Each carbon is joined to four others by strong covalent bonds in a giant lattice Correct answer.
  4. Strong electrostatic attractions between ions must be broken

Explanation

Diamond is a giant covalent solid in which each carbon atom is covalently bonded to four others in a continuous three-dimensional network. Melting requires breaking many strong covalent bonds, which demands a large amount of energy.

Question 49Chemical Bonding and Structure

Which pair of molecules is held together in the liquid state mainly by hydrogen bonding?

  1. Chlorine and bromine
  2. Methane and ethane
  3. Hydrogen fluoride and ammonia Correct answer.
  4. Hydrogen sulfide and hydrogen selenide

Explanation

Hydrogen bonding requires H bonded directly to N, O or F. In HF and NH3 the hydrogen atoms are attached to F and N respectively, so strong hydrogen bonds form between molecules.

Question 50Chemical Bonding and Structure

Why does the boiling point of water greatly exceed that of hydrogen sulfide, despite hydrogen sulfide having a larger electron count?

  1. Water forms a giant covalent lattice in the liquid state
  2. Water has a smaller molar mass and therefore a higher boiling point
  3. Water molecules are held by hydrogen bonds, whereas hydrogen sulfide molecules are not Correct answer.
  4. Hydrogen sulfide is non-polar and has only dispersion forces

Explanation

Oxygen is much more electronegative than sulfur, so only water has H atoms attached to a sufficiently electronegative atom to form hydrogen bonds. These strong intermolecular hydrogen bonds require far more energy to overcome than the dipole and dispersion forces in hydrogen sulfide.