The Pauling electronegativity values of elements can be used to predict the chemical properties of compounds. Use the information in Table 1.1 to answer the following questions.
(a) (i) Define electronegativity.
(ii) O and S are in Group 16. Explain the difference in the Pauling electronegativity values of O and S.
(b) (i) LiH is an ionic compound. Draw a dot‑and‑cross diagram of LiH. Include all electrons.
(ii) Suggest the shape of a molecule of \(H_2S\).
(c) (i) Write an equation that represents the first ionisation energy of H.
(ii) Explain why there is no information given in Table 1.1 for the second ionisation energy of H.
(iii) Give the full electronic configuration of \(S^{2+}\)(g).
(d) \(CO_2\) and \(SO_2\) are acidic gases.
(i) Write an equation for the reaction of \(SO_2\) with \(H_2O\).
(ii) Write an equation for the reaction of \(SO_2\) with NaOH.
(iii) Construct an equation for the reaction of \(CO_2\) with \(Mg(OH)_2\).
(e) (i) Complete Table 1.2 by placing a tick (✓) to show which of the compounds have molecules with an overall dipole moment.
(ii) At 150°C and 103kPa, all of the compounds listed in Table 1.2 are gases. Under these conditions, 0.284g of one of the compounds occupies a volume of \(127cm^3\). Use this information to calculate the \(M_r\) of the compound. Hence, identify the compound from those given in Table 1.2. Show your working.
▶️ Answer/Explanation
(a)(i) The power of an atom to attract the bonding pair of electrons towards itself in a covalent bond.
(a)(ii) Oxygen has a higher electronegativity than sulfur because its bonding electrons are closer to the nucleus (smaller atomic radius) and experience less shielding from inner shells, leading to a stronger effective nuclear charge and a greater attraction for electrons.
(b)(i)
Explanation: Lithium (Li) has one valence electron, and hydrogen (H) has one. In ionic LiH, lithium transfers its electron to hydrogen, forming \(Li^+\) and \(H^-\) ions.
(b)(ii) Non-linear / Bent / V-shaped
Explanation: \(H_2S\) has two bonding pairs and two lone pairs of electrons around the central sulfur atom, resulting in a bent molecular shape with a bond angle of approximately \(92^\circ\).
(c)(i) \(H(g) \rightarrow H^+(g) + e^-\)
Explanation: The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms.
(c)(ii) A hydrogen atom has only one electron, so it cannot have a second ionisation energy.
Explanation: The second ionisation energy would involve removing an electron from \(H^+\), which has no electrons to remove.
(c)(iii) \(1s^2 2s^2 2p^6 3s^2 3p^2\)
Explanation: A neutral sulfur atom (\(Z=16\)) has the configuration \(1s^2 2s^2 2p^6 3s^2 3p^4\). Removing two electrons gives \(S^{2+}\) with 14 electrons.
(d)(i) \(SO_2(g) + H_2O(l) \rightarrow H_2SO_3(aq)\)
Explanation: Sulfur dioxide dissolves in water to form sulfurous acid.
(d)(ii) \(SO_2(g) + 2NaOH(aq) \rightarrow Na_2SO_3(aq) + H_2O(l)\)
Explanation: Sulfur dioxide acts as an acidic oxide, reacting with a base to form a salt (sodium sulfite) and water.
(d)(iii) \(CO_2(g) + Mg(OH)_2(s) \rightarrow MgCO_3(s) + H_2O(l)\)
Explanation: Carbon dioxide reacts with the insoluble base magnesium hydroxide to form magnesium carbonate and water.
(e)(i)
Explanation: A molecule has an overall dipole moment if it is asymmetric and has polar bonds. \(CO_2\) is linear and symmetric (no dipole), \(SO_2\) is bent (has a dipole), \(H_2S\) is bent (has a dipole), \(CH_4\) is tetrahedral and symmetric (no dipole), \(BF_3\) is trigonal planar and symmetric (no dipole).
(e)(ii) Using the ideal gas equation \(PV = nRT\), converted to find molar mass \(M = \frac{mRT}{PV}\). Mass \(m = 0.284\) g, volume \(V = 0.127\) dm³, pressure \(P = 103\) kPa, temperature \(T = 423\) K, \(R = 8.314\) dm³ kPa mol⁻¹ K⁻¹. Substituting the values gives \(M = \frac{0.284 \times 8.314 \times 423}{103 \times 0.127} = 64.1\) g mol⁻¹. The compound with this \(M_r\) is sulfur dioxide, \(SO_2\).
