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Welcome to GCSE Edexcel Science revision.

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Unit S C 5: Ionic bonding.

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An ion is an atom or group of atoms with an overall charge because its numbers of protons and electrons do not balance.

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Ordinary ion formation changes electrons, not protons.

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Losing negatively charged electrons produces a positive ion, a cation.

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Gaining electrons produces a negative ion, an anion.

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Metals commonly lose outer electrons.

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Sodium, two, eight, one, loses one to form N A, charge plus, with configuration two, eight; magnesium, two, eight, two, loses two to form M G, charge 2 plus, with two, eight.

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Non-metals commonly gain electrons.

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Chlorine, two, eight, seven, gains one to form C L, charge minus, with two, eight, eight; oxygen, two, six, gains two to form O, charge 2 minus, with two, eight.

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Common simple ion charges follow main-group patterns: group 1 plus 1, group 2 plus 2, group 3 plus 3, group 6 minus 2 and group 7 minus 1.

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Transition metals may have different charges, specified when needed.

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Common main-group patterns; transition-metal charges can vary.

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These ions often have a full outer shell, a more stable electronic arrangement.

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The proton number and element identity remain unchanged.

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An ionic bond is the strong attraction between positively and negatively charged ions.

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This attraction between opposite charges is called electrostatic attraction.

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Ionic compounds commonly form between metals and non-metals.

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In sodium chloride formation, a sodium atom transfers one outer electron to a chlorine atom.

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N A, charge plus, and C L, charge minus, form, and their opposite charges attract.

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A transferred electron forms ions; attraction between their charges is the ionic bond.

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In magnesium oxide, magnesium transfers two electrons to oxygen.

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M G, charge 2 plus, and O, charge 2 minus, attract; each ion has a complete outer shell.

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In magnesium chloride, one magnesium atom loses two electrons, one to each of two chlorine atoms.

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The ions are M G, charge 2 plus, and two C L, charge minus,.

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Dot-and-cross diagrams distinguish electrons originally from different atoms.

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Draw brackets and charges for final ions and full outer shells; dots and crosses are not different kinds of electron.

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Outer electrons involved in transfer are shown; N A, charge plus, also has a full inner shell (two, eight).

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Electron transfer explains formation of ions.

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The bond itself is the attraction between ions, not an electron travelling continuously between two atoms.

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An ionic compound is electrically neutral overall.

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Its formula gives the simplest whole-number ratio of ions, not a separate molecule.

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N A, charge plus, and C L, charge minus, balance 1 to 1, giving N A C L.

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M G, charge 2 plus, and O, charge 2 minus, also balance 1 to 1, giving M G O.

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M G, charge 2 plus, needs two C L, charge minus, ions to balance charge, giving M G C L 2.

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Two N A, charge plus, balance one O, charge 2 minus, giving N A 2 O.

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Two A L, charge 3 plus, give plus 6 and three O, charge 2 minus, give minus 6, so aluminium oxide is A L 2 O 3.

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Always simplify a ratio to its lowest whole numbers.

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Balance total positive and negative charge, then simplify the ratio.

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For groups of atoms,

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preserve the group when using brackets: calcium nitrate from C A,

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charge two plus,

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and N O three,

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charge minus,

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is C A,

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open bracket N O three close bracket,

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subscript two.

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Ammonium is N H four, charge plus, and sulfate is S O four, charge two minus.

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A small number below the line (a subscript) counts atoms or groups.

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A charge above the line (a superscript) gives an ion’s charge.

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N A C L is neutral overall, so its formula has no overall plus or minus charge.

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Ionic solids form giant regular three-dimensional lattices of alternating positive and negative ions.

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Each ion attracts oppositely charged neighbours in many directions.

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Ions attract opposite charges throughout the lattice; this is not a collection of separate molecules.

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Strong electrostatic attractions need much energy to overcome, so ionic compounds generally have high melting and boiling points.

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A solid ionic compound does not conduct electricity because its ions are fixed in position, even though they are charged.

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When molten, ions can move and carry charge.

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When an ionic compound dissolves in water, mobile dissolved ions can also conduct.

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Mobile charged ions carry current in molten and aqueous ionic compounds.

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Not all ionic compounds dissolve in water.

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An insoluble compound cannot make a useful conducting aqueous solution simply by adding water.

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Ionic crystals are brittle: they can break when layers shift and bring ions with the same charge next to each other.

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These ions repel, splitting the crystal.

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Metals are different: their layers can slide without this splitting.

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A high-melting solid that conducts when molten but not solid is consistent with an ionic structure.

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Combine evidence rather than deciding from appearance alone.

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The charge carriers in molten or aqueous ionic compounds are ions.

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They are not a sea of delocalised electrons as in a metal.

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In electrolysis, cations move to the negative cathode and anions to the positive anode.

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The electrode signs describe an electrolytic cell connected to a power supply.

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A flat lattice diagram shows only a slice of a three-dimensional arrangement and is not to scale.

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Ion sizes and spacing are simplified.

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Use particle-level explanations: strong attractions explain high melting point, while ability of ions to move explains conduction. “It has bonds” alone is not sufficient.

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That completes Ionic bonding.

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Revisit the notes and test yourself on the revision website.