S2.1 The ionic model

1. X²⁺ — Gaining two electrons does give X a complete 3p sublevel, but electrons are negative: 16 protons and 18 electrons give a charge of 16 − 18 = −2, not +2.
2. X⁶⁺ — Losing all six outer electrons would leave the full shell 1s² 2s² 2p⁶, but having a full shell does not make an ion form. Removing six electrons would need a very large amount of energy; a non-metal atom with a nearly complete outer sublevel gains two electrons instead.
3. X²⁻ — X has six outer electrons (3s² 3p⁴). It gains two electrons, giving 1s² 2s² 2p⁶ 3s² 3p⁶, and with 16 protons and 18 electrons its charge is 2−. (X is sulfur, which forms S²⁻, as in sodium sulfide, Na₂S.)
4. X⁶⁻ — The charge is not the number of outer electrons. X needs only two more electrons to complete its 3p sublevel, so it gains two and forms X²⁻ (the argon configuration). Gaining six electrons would carry it four electrons beyond the argon configuration.

Syllabus statement S2.1.1 · Read this in Learn

1. Attraction between oppositely charged ions in all directions — The ionic bond is the electrostatic attraction between oppositely charged ions. It is non-directional: each Na⁺ ion attracts all the Cl⁻ ions around it in the lattice, and each Cl⁻ ion attracts all the Na⁺ ions around it.
2. Transfer of one electron from each sodium atom to a chlorine atom — Electron transfer is one way the ions can be formed, but it is an event, not the bond. The bond is the electrostatic attraction between the Na⁺ and Cl⁻ ions that exist in the lattice.
3. Attraction joining each Na⁺ to one Cl⁻ in an NaCl molecule — Sodium chloride contains no molecules. It is a giant lattice in which each Na⁺ ion is surrounded by, and attracted to, six Cl⁻ ions; NaCl gives the 1 : 1 ratio of the ions.
4. A pair of electrons shared unequally by Na and Cl atoms — A shared electron pair is a covalent bond. In sodium chloride the electron has been transferred, and the particles are Na⁺ and Cl⁻ ions held by electrostatic attraction.

Syllabus statement S2.1.2 · Read this in Learn

1. One molecule of one Mg atom bonded to two Cl atoms — Magnesium chloride contains no molecules. It is a lattice of Mg²⁺ and Cl⁻ ions in which each ion is surrounded by several ions of opposite charge; the formula gives only their ratio.
2. Mg²⁺ ions combined 1 : 1 with chlorine molecules, Cl₂ — The subscript 2 counts chloride ions, not atoms in a Cl₂ molecule. The two Cl⁻ ions are separate, each surrounded by Mg²⁺ ions in the lattice.
3. A Mg atom sharing an electron pair with each Cl atom — Shared electron pairs are covalent bonds. Magnesium chloride is ionic: magnesium has lost two electrons to become Mg²⁺, and the ions are held by electrostatic attraction.
4. The 1 : 2 ratio of Mg²⁺ and Cl⁻ ions in a lattice — Ionic compounds exist as three-dimensional lattices, and their formulas are empirical formulas: MgCl₂ gives the simplest whole-number ratio of the ions, one Mg²⁺ for every two Cl⁻.

Syllabus statement S2.1.3 · Read this in Learn

1. It loses the two electrons from its 4s sublevel. — Copper has only one 4s electron: its configuration is [Ar] 3d¹⁰ 4s¹, one of the exceptions to the Aufbau order. After the 4s electron, the second electron must come from 3d.
2. It loses two 3d electrons and keeps its 4s one. — Electrons are not removed in the reverse of the filling order. The 4s electron of a first-row transition element is removed before any 3d electron, so Cu²⁺ is [Ar] 3d⁹ with no 4s electron.
3. Its nucleus gains two protons; its electrons stay. — Forming an ion never changes the number of protons: two extra protons would make the atom a different element (gallium). Cu²⁺ has 29 protons and 27 electrons.
4. It loses its one 4s electron and one 3d electron. — Copper is [Ar] 3d¹⁰ 4s¹. The 4s electron is removed first, then one 3d electron, giving Cu²⁺ as [Ar] 3d⁹. Losing only the 4s electron gives Cu⁺, [Ar] 3d¹⁰.

Syllabus statement S2.1.1 · Read this in Learn

1. It is bonded only to the Cl⁻ ion that accepted its electron. — Ions carry no record of where their electrons went. The attraction to the Cl⁻ ion that happened to take the electron is exactly the same as the attraction to the other five, so all six are ionic bonds.
2. It forms one bond, as its atom had one electron to give. — An ion does not form a fixed number of bonds set by the electrons its atom lost. Each ion is attracted to every oppositely charged ion around it; the Na⁺ ion here forms six equivalent ionic bonds.
3. It is bonded to all six Cl⁻ ions by shared electron pairs. — Six neighbours is right, but there are no shared electron pairs: shared pairs are covalent bonds. The Na⁺ and Cl⁻ ions are held by electrostatic attraction between opposite charges.
4. It is bonded to all six Cl⁻ ions, attracting each of them equally. — The ionic bond is the electrostatic attraction between oppositely charged ions. It depends only on the charges and the distance between the ions, and all six Cl⁻ ions are at the same distance, so the Na⁺ ion is bonded equally to all six.

Syllabus statement S2.1.2 · Read this in Learn

1. When molten, it conducts because it releases free electrons. — There are no free electrons in molten sodium chloride. The current is carried by ions: Na⁺ moves towards the negative electrode and Cl⁻ towards the positive electrode.
2. When molten, it conducts because its ions are free to move. — A current needs mobile charged particles. In molten (or aqueous) sodium chloride the Na⁺ and Cl⁻ ions are free to move, so it conducts; in the solid the same ions are fixed in the lattice.
3. When solid, it conducts because it is made up of charged ions. — Solid sodium chloride does not conduct. Its ions are held in fixed positions in the lattice; charged particles carry a current only if they can move.
4. When solid, it has no ions; ions form only when it melts. — The ions are already present in the solid lattice. Solid NaCl does not conduct because its ions cannot move, not because it has none; melting only frees the existing ions.

Syllabus statement S2.1.3 · Read this in Learn

1. [Ar] 3d⁶ 4s⁰ — This is Fe²⁺, read from the subscript 2 on Fe. Subscripts count ions: two iron ions must balance three SO₄²⁻ ions (total −6), so each iron ion is Fe³⁺, [Ar] 3d⁵.
2. [Ar] 3d⁵ 4s⁰ — Three sulfate ions carry 3 × (−2) = −6, so the two iron ions carry +6 and each is Fe³⁺. Iron loses its two 4s electrons first and then one 3d electron: [Ar] 3d⁵.
3. [Ar] 3d³ 4s² — The charge, 3+, is right, but the electrons have been taken from 3d while 4s is kept. When iron forms an ion its 4s electrons are removed first, so Fe³⁺ is [Ar] 3d⁵.
4. [Ar] 3d⁶ 4s² — This is the configuration of the iron atom. The ion's charge comes from losing electrons, not from any change in the nucleus, so Fe³⁺ has three fewer electrons than the atom: [Ar] 3d⁵.

Syllabus statement S2.1.1 · Read this in Learn

1. Na₃PO₄, as three Na⁺ ions balance one PO₄³⁻ ion — Sodium forms Na⁺ and phosphate is PO₄³⁻. Three Na⁺ ions (+3) balance one PO₄³⁻ ion (−3). Only one phosphate ion is needed, so no brackets are used: Na₃PO₄.
2. Na₂PO₄, as phosphate has a 2− charge, like sulfate — Phosphate is PO₄³⁻, not PO₄²⁻: the -ate ending does not fix the charge. With PO₄³⁻, three Na⁺ ions are needed: Na₃PO₄.
3. NaPO₄, as one sodium ion pairs with one phosphate ion — One Na⁺ (+1) and one PO₄³⁻ (−3) give a total charge of −2, so NaPO₄ is not neutral. Three Na⁺ ions are needed: Na₃PO₄.
4. Na(PO₄)₃, as each ion's charge is its own subscript — Subscripts count ions; they are not each ion's own charge. Na(PO₄)₃ would carry +1 + 3 × (−3) = −8. Three Na⁺ ions balance one PO₄³⁻ ion: Na₃PO₄.

Syllabus statement S2.1.2 · Read this in Learn

1. Cu₂O, copper(II) oxide — The subscript 2 counts copper ions; it is not their charge. Two copper ions balance one O²⁻ ion (−2), so each is Cu⁺: Cu₂O is copper(I) oxide.
2. Fe₂O₃, iron trioxide — 'Iron trioxide' would mean one iron to three oxygens, which is not the ratio in Fe₂O₃. Number prefixes are the convention for covalent compounds such as sulfur trioxide; an ionic compound is named from its ions, with a Roman numeral for the charge of the iron ion.
3. CuO, copper(II) oxide — One O²⁻ ion (−2) is balanced by one copper ion, so the copper ion is Cu²⁺ and the Roman numeral is (II).
4. CaS, calcium sulfur — The anion is the sulfide ion, S²⁻. The anion of a binary ionic compound takes the suffix -ide: CaS is calcium sulfide.

Syllabus statement S2.1.2 · Read this in Learn

1. NH₄⁻(aq) — Ammonium is the one positive polyatomic ion: NH₄⁺, formed when ammonia accepts H⁺. In its compounds it takes the place of a metal cation, as in NH₄Cl.
2. OH⁻(aq) — The hydroxide ion is OH⁻. The correct formulas of the other three ions are NH₄⁺ (ammonium), NO₃⁻ (nitrate) and HCO₃⁻ (hydrogencarbonate).
3. NO₃²⁻(aq) — Not every -ate ion is 2−. Sulfate and carbonate are 2−, but nitrate is NO₃⁻, as in sodium nitrate NaNO₃.
4. HCO₃²⁻(aq) — The hydrogen in hydrogencarbonate is present as H⁺ added to CO₃²⁻, which reduces the charge by one: HCO₃⁻, as in sodium hydrogencarbonate NaHCO₃.

Syllabus statement S2.1.2 · Read this in Learn

1. Al₂O₃, as two Al³⁺ ions cancel the charge of three O²⁻ — Aluminium loses three electrons to form Al³⁺ and oxygen gains two to form O²⁻. The lowest ratio with zero total charge is 2 × (+3) + 3 × (−2) = 0, so the formula is Al₂O₃.
2. Al₃O₂, as each ion's charge is written as its own subscript — Al₃O₂ would have a total charge of 3 × (+3) + 2 × (−2) = +5. The charge of each ion becomes the subscript of the OTHER ion, which gives Al₂O₃.
3. AlO, as each aluminium ion is paired with one oxide ion — Al³⁺ + O²⁻ has a total charge of +1, so AlO is not neutral. The ions combine in the ratio that cancels the charges: two Al³⁺ to three O²⁻.
4. Al₂O, as oxygen has six outer electrons and forms O⁶⁻ — Oxygen does not form O⁶⁻. It needs two electrons to complete its 2p sublevel, so it gains two and forms O²⁻; with Al³⁺ this gives Al₂O₃.

Syllabus statement S2.1.2 · Read this in Learn

1. Water lowers the melting point of NaCl so that it melts at room temperature. — Dissolving is not melting: NaCl melts at 801 °C. It dissolves at room temperature because water molecules are attracted to the ions and pull them into solution.
2. The ions fit into the spaces between water molecules but not hexane ones. — Dissolving does not depend on gaps between solvent molecules. It depends on attraction: polar water molecules are attracted to the ions, and non-polar hexane molecules are not.
3. Polar water molecules are attracted to the ions and separate them from the lattice. — Water molecules are polar: their partially negative O atoms are attracted to Na⁺ and their partially positive H atoms to Cl⁻, so the ions leave the lattice and are surrounded by water. Hexane is non-polar and is not attracted strongly enough to the ions.
4. Water turns the neutral NaCl units into ions, which then spread out. — The ions already exist in the solid lattice; water does not create them. Water dissolves NaCl because its polar molecules are attracted to the existing Na⁺ and Cl⁻ ions.

Syllabus statement S2.1.3 · Read this in Learn

1. NaF — NaF does have the smallest sum of ionic radii (235 pm), but its ions are only singly charged. CaO's 2+ and 2− ions attract far more strongly, and a 5 pm difference in radii cannot make up for that.
2. CsF — CsF has the largest electronegativity difference, but lattice enthalpy depends on the charges and radii of the ions. Cs⁺ is the largest cation here (167 pm) and both ions are singly charged, so CsF's lattice enthalpy is much smaller than CaO's.
3. BaS — BaS has the largest ions here (sum of radii 319 pm). Larger ions attract each other less strongly because their centres are further apart, so BaS has a smaller lattice enthalpy than CaO, whose ions carry the same charges.
4. CaO — Lattice enthalpy increases with ion charge and decreases with ion radius. CaO and BaS contain 2+ and 2− ions, which attract far more strongly than the 1+ and 1− ions of NaF and CsF. Of the two, CaO has the smaller ions (sum of radii 240 pm against 319 pm), so its lattice enthalpy is the largest.

Syllabus statement S2.1.3 · Read this in Learn

1. KCl > NaCl > NaF — This is the reverse order. Larger ions have their centres further apart, so they attract each other less strongly: KCl, with the largest ions, has the smallest lattice enthalpy.
2. NaF > NaCl > KCl — All three compounds contain 1+ and 1− ions, so the radii decide. The sums of the radii are NaF 235 pm, NaCl 283 pm and KCl 319 pm; the closer the ion centres, the stronger the attraction, so NaF > NaCl > KCl.
3. NaF > KCl > NaCl — This follows the electronegativity difference, which is larger for K–Cl than for Na–Cl. Lattice enthalpy depends on the charges and radii of the ions; KCl has larger ions than NaCl, so its lattice enthalpy is smaller.
4. NaF = NaCl = KCl — The charges are the same, but lattice enthalpy also depends on ion radius. The sums of the radii differ (235, 283 and 319 pm), so the lattice enthalpies differ, decreasing as the ions get larger.

Syllabus statement S2.1.3 · Read this in Learn

1. Every ion is strongly attracted to many ions of opposite charge. — In the lattice every ion is surrounded by ions of opposite charge and attracted to all of them. Separating the ions means overcoming many strong electrostatic attractions, which needs a large amount of energy, so sodium chloride has a low volatility.
2. Its particles are so heavy that they cannot move fast enough to escape. — Sodium chloride's relative formula mass, M_r, is only 58.44, lower than that of many volatile liquids such as octane (114.26). Its low volatility comes from the strong attractions between its ions, not from mass.
3. Evaporating it would break the strong bond in every NaCl molecule. — Solid sodium chloride contains no molecules. It is a lattice in which each ion is attracted to all the oppositely charged ions around it, and these attractions must be overcome to melt or vaporize it.
4. Sodium is a metal with a high boiling point, and NaCl keeps this property. — A compound does not keep the properties of its elements: sodium chloride is not a metal. Its low volatility comes from its own structure, a lattice of ions held by strong electrostatic attractions.

Syllabus statement S2.1.3 · Read this in Learn

1. Ammonium hydrogencarbonate has the chemical formula (NH₄)₂HCO₃ — Hydrogencarbonate is HCO₃⁻: adding H⁺ to CO₃²⁻ changes the charge from 2− to 1−. One NH₄⁺ balances one HCO₃⁻, so ammonium hydrogencarbonate is NH₄HCO₃.
2. Aluminium phosphate has the chemical formula Al₂(PO₄)₃ — Phosphate is PO₄³⁻, not PO₄²⁻: the -ate ending does not fix the charge. Al³⁺ and PO₄³⁻ combine 1 : 1, so aluminium phosphate is AlPO₄.
3. Ammonium carbonate has the chemical formula (NH₄)₂CO₃ — Ammonium is NH₄⁺ and carbonate is CO₃²⁻. Two NH₄⁺ ions (+2) balance one CO₃²⁻ ion (−2), and the ammonium ion is kept intact and bracketed.
4. Potassium phosphate has the chemical formula KPO₄ — One K⁺ (+1) and one PO₄³⁻ (−3) give a total charge of −2, so KPO₄ is not neutral. Three K⁺ ions balance one PO₄³⁻ ion: K₃PO₄.

Syllabus statement S2.1.2 · Read this in Learn