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Question 1Atom structure and periodic law

Which subatomic particle determines the identity of an element?

  1. The neutron
  2. The positron
  3. The electron
  4. The proton Correct answer.

Explanation

The atomic number, which equals the number of protons in the nucleus, fixes the identity of an element. Changing the proton count converts the atom into a different element, whereas changing electron or neutron numbers gives ions or isotopes of the same element.

Question 2Atom structure and periodic law

In the notation ⁵⁶Fe³⁺, how many electrons does the ion contain?

  1. 56
  2. 23 Correct answer.
  3. 26
  4. 29

Explanation

Iron has atomic number 26, so the neutral atom has 26 electrons. The 3+ charge means three electrons have been lost, leaving 26 − 3 = 23 electrons. The mass number 56 gives 30 neutrons but does not affect the electron count.

Question 3Atom structure and periodic law

Which pair of species are isotopes of one another?

  1. ¹⁴N and ¹⁴C
  2. ¹⁶O and ¹⁶O²⁻
  3. ²³Na⁺ and ²⁴Mg⁺
  4. ³⁵Cl⁻ and ³⁷Cl⁻ Correct answer.

Explanation

Isotopes are atoms of the same element with different mass numbers, so they share a proton count but differ in neutrons. ³⁵Cl and ³⁷Cl both have 17 protons and differ by two neutrons, so these ions are isotopes.

Question 4Atom structure and periodic law

Which statement about the lines in an atomic emission spectrum is correct?

  1. The lines arise from electrons spiralling continuously into the nucleus
  2. Each line corresponds to a different isotope present in the sample
  3. The number of lines equals the number of neutrons in the atom
  4. Each line corresponds to an electron transition between quantised energy levels Correct answer.

Explanation

Emission lines are produced when excited electrons fall from a higher quantised level to a lower one, releasing a photon of fixed energy. Because the level spacings are fixed for each element, the line wavelengths act as a fingerprint of that element.

Question 5Atom structure and periodic law

Which subshell is filled immediately after the 4s subshell in the building-up order of ground-state atoms?

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

Explanation

After 4s the next subshell in energy order is 3d, which is why the first transition series runs from scandium to zinc. The 4s orbital has lower energy than 3d for these atoms, so it fills first despite the higher principal quantum number.

Question 6Atom structure and periodic law

Which block of the periodic table contains an element whose atoms have a partially filled p subshell in the ground state?

  1. The f-block
  2. The p-block Correct answer.
  3. The s-block
  4. The d-block

Explanation

The p-block consists of elements in which the outermost electrons occupy p orbitals, so these subshells are progressively filled across groups 13 to 18. The s-block fills s orbitals, and the d- and f-blocks fill d and f orbitals respectively.

Question 7Atom structure and periodic law

Which species has the ground-state electron configuration 1s² 2s² 2p⁶ 3s² 3p⁶?

  1. Cl⁻
  2. Ar⁺
  3. K⁺ Correct answer.
  4. S²⁻

Explanation

The configuration shown contains 18 electrons, matching the argon core. Potassium has atomic number 19, so K⁺ has 19 − 1 = 18 electrons and this configuration.

Question 8Atom structure and periodic law

Which property increases across period 3 from sodium to chlorine?

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

Explanation

Across a period the nuclear charge rises while electrons enter the same shell, so the effective nuclear charge felt by outer electrons increases. More energy is therefore needed to remove an outer electron, so first ionisation energy rises overall.

Question 9Atom structure and periodic law

Why is the first ionisation energy of magnesium higher than that of aluminium?

  1. Aluminium has a larger nuclear charge than magnesium, which attracts the outer electron more strongly
  2. Magnesium has a smaller atomic radius, so its outer electron is removed from a 3p orbital
  3. Aluminium has one more electron shell than magnesium, so its outer electron is further from the nucleus
  4. The electron removed from aluminium is in a 3p orbital, which is higher in energy and better shielded than the 3s orbital Correct answer.

Explanation

Magnesium loses a 3s electron, while aluminium loses a 3p electron that lies higher in energy and is shielded by the filled 3s subshell. This outweighs the increase in nuclear charge, so aluminium's first ionisation energy is lower.

Question 10Atom structure and periodic law

Which ion has the electron configuration [Ar] 3d⁵?

  1. Ni²⁺
  2. Mn²⁺ Correct answer.
  3. Fe³⁺
  4. Cr³⁺

Explanation

Manganese has atomic number 25, so Mn²⁺ has 23 electrons: [Ar] 3d⁵. This half-filled d subshell is relatively stable and helps explain the stability of the Mn²⁺ ion.

Question 11Atom structure and periodic law

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

  1. The ion has gained an electron, increasing the electron–electron repulsion
  2. The ion has lost its only 3s electron, so the remaining shells are pulled in by the unchanged nuclear charge Correct answer.
  3. The ion has the same electron configuration as the atom but a lower mass
  4. The ion has one fewer proton than the atom, so the nucleus attracts electrons less strongly

Explanation

Forming Na⁺ removes the single 3s electron, leaving the neon-like 1s² 2s² 2p⁶ core. With the same nuclear charge but one fewer shell and less electron–electron repulsion, the remaining electrons are drawn closer to the nucleus.

Question 12Atom structure and periodic law

Which element has the ground-state configuration [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p²?

  1. Sn
  2. Bi
  3. Pb Correct answer.
  4. Tl

Explanation

The configuration adds four electrons beyond the xenon core, placing the element in group 14, period 6. That position corresponds to lead, whose total electron count is 82.

Question 13Atom structure and periodic law

Which observation provides the strongest evidence that electrons occupy discrete energy levels in atoms?

  1. The scattering of alpha particles by a thin gold foil
  2. The deflection of cathode rays by a magnetic field
  3. The sharp line spectrum emitted by excited hydrogen atoms Correct answer.
  4. The continuous spectrum emitted by a hot tungsten filament

Explanation

A line spectrum shows that excited atoms emit only photons of certain fixed energies, which implies that electrons can occupy only particular energy levels. A continuous spectrum would be expected if electron energies could take any value.

Question 14Atom structure and periodic law

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

  1. The second electron is removed from a shell closer to the nucleus in every element
  2. The nucleus gains a proton each time an electron is removed
  3. The second electron is removed from a cation, so it experiences a greater effective nuclear attraction Correct answer.
  4. The second electron is removed from a paired orbital, so extra repulsion lowers the energy needed

Explanation

After the first electron is lost, the remaining species is positively charged, so each remaining electron feels a stronger net pull from the nucleus. More energy is therefore needed to remove the next electron, regardless of which subshell it occupies.

Question 15Atom structure and periodic law

An element has first four ionisation energies of 590, 1150, 4940 and 6480 kJ mol⁻¹. To which group does it most likely belong?

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

Explanation

The large jump occurs between the second and third ionisation energies, showing that the third electron comes from a new, inner shell. This means the element has two valence electrons, placing it in group 2.

Question 16Atom structure and periodic law

Which property decreases down group 2 from magnesium to barium?

  1. First ionisation energy Correct answer.
  2. Number of occupied electron shells
  3. Atomic radius
  4. Reactivity with water

Explanation

Down the group each element has an extra filled shell, so the outer electron is further from the nucleus and better shielded. The attraction holding that electron is weaker, so first ionisation energy falls from magnesium to barium.

Question 17Atom structure and periodic law

Which species is isoelectronic with the sulfide ion, S²⁻?

  1. Ne
  2. P³⁻
  3. Ca²⁺ Correct answer.
  4. Cl⁻

Explanation

Sulfide has 16 + 2 = 18 electrons. Calcium has atomic number 20, so Ca²⁺ has 20 − 2 = 18 electrons and the same argon-type configuration, making the two species isoelectronic.

Question 18Atom structure and periodic law

The first ionisation energies of the period 2 elements show a dip at boron relative to beryllium and a dip at oxygen relative to nitrogen. Which pair of reasons correctly accounts for these two dips?

  1. Boron: removal of a paired 2s electron; oxygen: 2p is higher in energy than 2s
  2. Boron: smaller nuclear charge than beryllium; oxygen: larger nuclear charge than nitrogen
  3. Boron: extra shielding by the 1s shell; oxygen: extra shielding by the 2s shell
  4. Boron: 2p is higher in energy than 2s; oxygen: removal of a paired 2p electron relieves electron–electron repulsion Correct answer.

Explanation

Boron's outermost electron occupies a 2p orbital, which is higher in energy and better shielded than the 2s electron removed from beryllium. At oxygen the fourth 2p electron must pair in an already occupied orbital, and the resulting repulsion makes it easier to remove than nitrogen's unpaired 2p electron.

Question 19Atom structure and periodic law

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

  1. The 3d subshell must be filled before the 4s subshell in all atoms
  2. A half-filled 3d subshell with five parallel spins has exchange stabilisation that outweighs the 4s–3d energy gap Correct answer.
  3. Chromium has an anomalously small nuclear charge, so it cannot retain two 4s electrons
  4. The 4s orbital is always higher in energy than 3d in every element of the first transition series

Explanation

The 3d and 4s orbitals are close in energy in the first transition series, so small stabilising effects can decide the configuration. A 3d⁵ arrangement gives five unpaired electrons with parallel spins and extra exchange stabilisation, which favours the observed 3d⁵ 4s¹ configuration.

Question 20Atom structure and periodic law

Across the first transition series the atomic radii decrease only slightly, whereas across period 3 the decrease is much larger. Which explanation is correct?

  1. In period 3 the added electrons enter inner shells, which increases the radius
  2. In the transition series the 4s electrons are removed before the 3d electrons, keeping the radius constant
  3. In the transition series the added d electrons shield the outer 4s electrons effectively, so the increasing nuclear charge is largely offset Correct answer.
  4. In the transition series the nuclear charge does not increase across the series

Explanation

As the 3d subshell fills, the extra d electrons are inner electrons and shield the outer 4s electrons fairly well. The growing nuclear charge is therefore largely cancelled by the increased shielding, so the radius changes only slightly, unlike the pronounced contraction across a period of the s- and p-blocks.

Question 21Chemical bonding and molecular structure

Which description of a covalent bond is correct?

  1. A delocalised sea of electrons surrounding positive ions
  2. A shared pair of electrons attracted to both nuclei Correct answer.
  3. An electrostatic attraction between oppositely charged ions
  4. A transfer of electrons from one atom to another

Explanation

A covalent bond forms when two atoms share one or more pairs of electrons, and each shared pair is attracted to both nuclei, holding the atoms together. Electron transfer produces ionic bonding, while a sea of delocalised electrons describes metallic bonding.

Question 22Chemical bonding and molecular structure

Which molecule contains a polar covalent bond?

  1. CH4
  2. HCl Correct answer.
  3. Cl2
  4. NaCl

Explanation

A polar covalent bond arises when two bonded atoms have different electronegativities, so the shared electrons are pulled toward the more electronegative atom. In HCl, chlorine is more electronegative than hydrogen, giving a polar bond with partial charges.

Question 23Chemical bonding and molecular structure

What is the shape of a methane (CH4) molecule according to VSEPR theory?

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

Explanation

Carbon in methane has four bonding pairs and no lone pairs, so the four electron domains repel each other equally and adopt a tetrahedral arrangement with bond angles of about 109.5 degrees.

Question 24Chemical bonding and molecular structure

Which type of bond is present between the atoms within a polyatomic ion such as sulfate, SO4^2-?

  1. Metallic bonds
  2. Covalent bonds Correct answer.
  3. Ionic bonds only
  4. Hydrogen bonds

Explanation

Within the sulfate ion, sulfur and oxygen atoms share electrons, so the atoms are held together by covalent bonds. The overall charge makes the ion interact ionically with counter-ions, but that is between ions, not within the ion.

Question 25Chemical bonding and molecular structure

Why does sodium chloride conduct electricity when molten but not when solid?

  1. Ions become free to move when the lattice is broken by melting Correct answer.
  2. Electrons are delocalised in the solid lattice
  3. The covalent bonds break on melting to release electrons
  4. Melting converts ions into neutral atoms

Explanation

In solid NaCl the ions are held in a fixed lattice and cannot move, so no current flows. Melting breaks down the lattice and allows the ions to move freely, which carries charge through the liquid.

Question 26Chemical bonding and molecular structure

Which statement about the strength of a hydrogen bond is correct?

  1. It is stronger than an ionic bond
  2. It is stronger than a typical intermolecular force but weaker than a covalent bond Correct answer.
  3. It is stronger than a typical covalent bond
  4. It is the weakest of all intermolecular forces

Explanation

A hydrogen bond is a strong intermolecular attraction, typically around 10 to 40 kJ per mole, which is much stronger than dispersion forces but far weaker than a covalent bond. It forms when hydrogen bonded to N, O or F is attracted to a lone pair on a nearby electronegative atom.

Question 27Chemical bonding and molecular structure

The bond angle in a water molecule is approximately 104.5 degrees rather than the tetrahedral 109.5 degrees. What is the best explanation?

  1. Two lone pairs on oxygen repel the bonding pairs more strongly than bonding pairs repel each other Correct answer.
  2. The molecule is linear because oxygen has no lone pairs
  3. Oxygen uses sp2 hybrid orbitals, giving a trigonal planar arrangement
  4. Hydrogen atoms repel each other more than lone pairs do

Explanation

Oxygen in water has two bonding pairs and two lone pairs, giving four electron domains and an approximately tetrahedral electron geometry. Lone pairs occupy more space than bonding pairs, so they compress the H-O-H angle below the ideal tetrahedral value.

Question 28Chemical bonding and molecular structure

Which molecule is overall nonpolar despite containing polar bonds?

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

Explanation

In CO2 the two C=O bond dipoles are equal and point in opposite directions along the linear axis, so they cancel and the molecule has no net dipole. A molecule with polar bonds can still be nonpolar if its geometry makes the bond dipoles cancel.

Question 29Chemical bonding and molecular structure

Which species has a bond order of 2.5 according to molecular orbital theory?

  1. O2
  2. N2
  3. O2^2-
  4. NO Correct answer.

Explanation

For NO, the molecular orbital configuration gives eight bonding electrons and three antibonding electrons, so the bond order is (8 - 3)/2 = 2.5. Bond order equals half the difference between bonding and antibonding electron counts.

Question 30Chemical bonding and molecular structure

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

  1. London dispersion forces
  2. Dipole-dipole forces only
  3. Hydrogen bonding Correct answer.
  4. Ion-dipole forces

Explanation

Ethanol contains an O-H group, so its molecules can form hydrogen bonds with one another, which requires significant energy to overcome on boiling. Ethane has only dispersion forces, so it boils at a much lower temperature despite a similar molar mass.

Question 31Chemical bonding and molecular structure

Which element pair is most likely to form an ionic compound?

  1. Potassium and fluorine Correct answer.
  2. Nitrogen and chlorine
  3. Sulfur and oxygen
  4. Carbon and oxygen

Explanation

Ionic bonding occurs when a metal with low ionisation energy transfers electrons to a highly electronegative nonmetal. Potassium is a reactive alkali metal and fluorine is the most electronegative element, so the large electronegativity difference favours electron transfer and ionic bonding.

Question 32Chemical bonding and molecular structure

What is the hybridisation of the central carbon atom in ethene (C2H4)?

  1. sp3
  2. sp3d
  3. sp
  4. sp2 Correct answer.

Explanation

Each carbon in ethene is bonded to three other atoms with no lone pairs, giving three electron domains and sp2 hybridisation. The remaining unhybridised p orbitals overlap side-on to form the pi bond of the C=C double bond.

Question 33Chemical bonding and molecular structure

Which statement best explains why the bond angle in NH3 (about 107 degrees) is smaller than in NH4+ (109.5 degrees)?

  1. NH4+ contains a double bond that widens the angle
  2. NH3 has a lone pair that repels bonding pairs more strongly than bonding pairs repel each other Correct answer.
  3. NH3 has fewer electrons so the bonds are shorter and wider
  4. Nitrogen in NH3 is sp2 hybridised

Explanation

Both species have four electron domains around nitrogen, but NH4+ has four identical bonding pairs while NH3 has one lone pair. Lone-pair repulsion is greater than bonding-pair repulsion, so the lone pair squeezes the N-H bonds closer together in ammonia.

Question 34Chemical bonding and molecular structure

Which property of a metal is best explained by the delocalised electron model of metallic bonding?

  1. Its low melting point in all cases
  2. Its brittleness when struck
  3. Its inability to conduct heat
  4. Its high electrical conductivity Correct answer.

Explanation

In the metallic model, outer electrons are delocalised over a lattice of positive ions. When a potential difference is applied, these mobile electrons drift through the structure, giving metals their high electrical conductivity.

Question 35Chemical bonding and molecular structure

Rank the following bonds in order of increasing polarity: H-H, H-Cl, H-F, H-Br. Which ordering is correct?

  1. H-H < H-Br < H-Cl < H-F Correct answer.
  2. H-Br < H-H < H-F < H-Cl
  3. H-H < H-Cl < H-Br < H-F
  4. H-F < H-Cl < H-Br < H-H

Explanation

Bond polarity increases with the electronegativity difference between the bonded atoms. Hydrogen and hydrogen are identical, so H-H is nonpolar; among the halogens electronegativity decreases down the group from F to Cl to Br, so H-F is the most polar.

Question 36Chemical bonding and molecular structure

Which molecule has a trigonal pyramidal shape?

  1. BF3
  2. SiCl4
  3. PCl3 Correct answer.
  4. CO2

Explanation

Phosphorus in PCl3 has three bonding pairs and one lone pair, giving four electron domains arranged tetrahedrally. The lone pair occupies one vertex, so the atoms form a trigonal pyramid with bond angles slightly below 109.5 degrees.

Question 37Chemical bonding and molecular structure

Which statement about sigma and pi bonds is correct?

  1. A pi bond results from sideways overlap of p orbitals and is generally weaker than a sigma bond Correct answer.
  2. A double bond consists of two identical sigma bonds
  3. A sigma bond results from sideways overlap of p orbitals and is weaker than a pi bond
  4. Pi bonds allow free rotation about the internuclear axis

Explanation

Sigma bonds form by head-on overlap along the internuclear axis and concentrate electron density between the nuclei, making them stronger. Pi bonds form by sideways overlap of parallel p orbitals, with electron density above and below the axis, and they are weaker and restrict rotation.

Question 38Chemical bonding and molecular structure

Which species has the same shape as the carbonate ion, CO3^2-?

  1. H2O
  2. SO2
  3. SO3 Correct answer.
  4. NH4+

Explanation

Carbonate has three bonding regions and no lone pairs on carbon, giving a trigonal planar shape with 120-degree angles. Sulfur trioxide likewise has three bonding regions and no lone pair on sulfur, so it is also trigonal planar.

Question 39Chemical bonding and molecular structure

Using molecular orbital theory, why is O2 paramagnetic while N2 is diamagnetic?

  1. O2 has two unpaired electrons in its degenerate pi* antibonding orbitals, whereas N2 has all electrons paired Correct answer.
  2. Both molecules contain unpaired electrons, but N2's are paired by hydrogen bonding
  3. O2 is paramagnetic because of unpaired electrons in its sigma bonding orbital
  4. N2 contains unpaired electrons in its sigma* orbital that are too weakly coupled to show magnetism

Explanation

Filling the molecular orbitals of O2 places the last two electrons singly into the two degenerate pi* antibonding orbitals, giving two unpaired electrons and paramagnetism. In N2 the highest occupied orbitals are fully paired, so the molecule is diamagnetic.

Question 40Chemical bonding and molecular structure

Which factor best explains why the ionic radius of Mg2+ is smaller than that of Na+?

  1. Na+ forms stronger covalent bonds, which expand its radius
  2. Mg2+ has a larger number of electrons, increasing screening
  3. Mg2+ has a greater nuclear charge with the same electron configuration, pulling electrons inward more strongly Correct answer.
  4. Mg2+ has an extra electron shell compared with Na+

Explanation

Both Mg2+ and Na+ are isoelectronic with neon, having the same ten electrons in the same shells. Magnesium has one more proton, so the higher effective nuclear charge draws the electron cloud in more tightly, making Mg2+ smaller than Na+.

Question 41Inorganic chemistry: s-, p-, d-block elements and their compounds

Which of the following is the correct electron configuration of a neutral chlorine atom in its ground state?

  1. [Ne] 3s¹ 3p⁶
  2. [Ne] 3s² 3p⁵ Correct answer.
  3. [Ne] 3s² 3p⁶
  4. [Ar] 3s² 3p⁵

Explanation

Chlorine has 17 electrons, so after the neon core (10 electrons) it places two in 3s and five in 3p, giving [Ne] 3s² 3p⁵. The p subshell holds a maximum of six, so one vacancy remains, which is why chlorine commonly forms a single bond or the chloride ion.

Question 42Inorganic chemistry: s-, p-, d-block elements and their compounds

Which group 2 element has the highest first ionisation energy?

  1. Barium
  2. Strontium
  3. Calcium
  4. Beryllium Correct answer.

Explanation

First ionisation energy falls down group 2 because the outer electron occupies a higher shell, is further from the nucleus, and is shielded by more inner electrons. Beryllium, at the top of the group, therefore has the highest value.

Question 43Inorganic chemistry: s-, p-, d-block elements and their compounds

What is the shape of the PCl₃ molecule?

  1. Trigonal planar
  2. T-shaped
  3. Tetrahedral
  4. Trigonal pyramidal Correct answer.

Explanation

Phosphorus in PCl₃ has three bonding pairs and one lone pair, giving four electron domains arranged tetrahedrally. The lone pair compresses the bond angles to about 100°, so the molecular shape is trigonal pyramidal.

Question 44Inorganic chemistry: s-, p-, d-block elements and their compounds

Which oxide of a period 3 element is amphoteric?

  1. MgO
  2. Al₂O₃ Correct answer.
  3. Na₂O
  4. SO₃

Explanation

Aluminium oxide reacts with both acids to give aluminium salts and with alkalis to give aluminates, so it is amphoteric. The oxides at the extremes of period 3 are either basic (Na₂O, MgO) or acidic (SO₃).

Question 45Inorganic chemistry: s-, p-, d-block elements and their compounds

Which property of transition metals arises because they have partially filled d orbitals?

  1. Low electrical conductivity
  2. Formation of only one oxidation state
  3. Low melting points
  4. Formation of coloured ions Correct answer.

Explanation

Partially filled d orbitals allow d–d electronic transitions in the visible region, so transition metal ions absorb some wavelengths and appear coloured. The same partially filled d set also underlies variable oxidation states and catalytic behaviour.

Question 46Inorganic chemistry: s-, p-, d-block elements and their compounds

Which reagent is used in the test for a sulfate ion in aqueous solution?

  1. Sodium hydroxide solution
  2. Silver nitrate solution
  3. Potassium iodide solution
  4. Acidified barium chloride solution Correct answer.

Explanation

Barium ions in acidified solution precipitate sulfate as barium sulfate, a white solid that is insoluble in dilute acid. Acidification with dilute hydrochloric acid prevents false positives from carbonate or sulfite ions, which would also give white precipitates.

Question 47Inorganic chemistry: s-, p-, d-block elements and their compounds

Which of the following is the correct formula of sodium peroxide?

  1. Na₂O
  2. NaO₂
  3. NaO
  4. Na₂O₂ Correct answer.

Explanation

Sodium peroxide contains the peroxide ion O₂²⁻, so two sodium ions are needed to balance the charge, giving Na₂O₂. Sodium superoxide, by contrast, contains the O₂⁻ ion and has the formula NaO₂.

Question 48Inorganic chemistry: s-, p-, d-block elements and their compounds

Which halogen has the lowest bond dissociation enthalpy in the gas phase?

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

Explanation

Bond dissociation enthalpy decreases from chlorine to iodine as the bonding orbitals become larger and overlap less effectively. Fluorine is an exception to the smooth trend because its very short F–F bond suffers strong lone-pair repulsion, making it weaker than chlorine but still stronger than iodine.

Question 49Inorganic chemistry: s-, p-, d-block elements and their compounds

A solution of a transition metal ion gives a pale green precipitate with dilute sodium hydroxide that turns brown on standing in air. Which ion is present?

  1. Fe²⁺ Correct answer.
  2. Fe³⁺
  3. Cu²⁺
  4. Zn²⁺

Explanation

Iron(II) ions give a pale green hydroxide precipitate, Fe(OH)₂, which is readily oxidised by air to brown iron(III) hydroxide. This colour change on standing distinguishes Fe²⁺ from Fe³⁺, which precipitates directly as a brown solid.

Question 50Inorganic chemistry: s-, p-, d-block elements and their compounds

Which explanation best accounts for the increase in thermal stability of the group 2 carbonates down the group?

  1. The decomposition temperature decreases as the cation radius increases
  2. The cation becomes more polarising as its charge density increases
  3. The carbonate ion becomes more polarisable as the cation grows Correct answer.
  4. The lattice enthalpy decreases more than the decomposition enthalpy

Explanation

As the group 2 cation becomes larger, its charge density falls and it distorts the carbonate ion less. A less polarised carbonate ion is harder to decompose, so thermal stability increases down the group.