|
|
Doc Brown's Chemistry - Chemical Bonding - Revision Notes Part 2 Ionic bonding, structure and properties of ionic substances Revision KS4 Science IGCSE/O level/GCSE Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel 360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science (revise courses equal to US grades 9-10) Revision notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA OCR Edexcel Salters CIE revising courses for pre-university students (equal to US grade 11 and grade 12 and Honours/honors level courses) Part 1 Introduction - why do atoms bond together? & sub-index for Parts 2-5 (read 1st) Part 2 Ionic Bonding - compounds and properties (this page) Part 3 Covalent Bonding -small simple molecules and properties Part 4 Covalent Bonding - macromolecules and giant covalent structures Part 5 Metallic Bonding - structure and properties of metals Part 6 More advanced concepts for advanced level chemistryPart 6 More advanced concepts for advanced level chemistry (in preparation, BUT a lot on intermolecular forces in Equilibria Part 8)
Part 2. IONIC BONDING - compounds and properties examples of ionic compounds * physical properties of ionic compounds Examples of ionic compounds described: sodium chloride NaCl (exemplar for any Li/Na/K + F/Cl/Br/I combination), magnesium chloride MgCl2 (exemplar for any Mg/Ca + F/Cl/Br combination), aluminium fluoride AlF3, potassium oxide K2O (exemplar for any Li/Na/K + O/S combination), magnesium/calcium oxide MgO/CaO and magnesium/calcium sulphide (MgS/CaS), aluminium oxide Al2O3 (exemplar for Al2S3). If your ionic compound is not listed, look for a compound with a similar formula and you should be able to work it out from the example given. The use of the word exemplar implies you are dealing with the same set of outer electron arrangements (configurations), which is why you can work out lots more dot and cross diagrams of ionic compounds by understanding one example.
Ionic bonds are formed by one atom transferring electrons to another atom to form ions. Elements consist of neutral atoms or molecules, the electrical neutrality is because the number of positive protons equals the number of negative electrons. Ions are atoms, or groups of atoms, which have lost or gained electrons to have a net electrical charge overall . The atom losing electrons forms a positive ion (a cation) and is usually a metal. The overall charge on the ion is positive due to excess positive nuclear charge (protons do NOT change in chemical reactions). The atom gaining electrons forms a negative ion (an anion) and is usually a non-metallic element. The overall charge on the ion is negative because of the gain, and therefore excess, of negative electrons. The examples below combining a metal from Groups 1 (Alkali Metals), 2 or 3, with a non-metal from Group 6 or Group 7 (The Halogens). The electron structures are shown in () or []. Only the outer electrons of the original atoms, and where they end up in the ions, are shown in the dot and cross (ox) diagrams Ionic bonding is not directional like covalent bonding, in the sense that the force of attraction between the positive ions and the negative ions act in every direction around the ions. Example 1: A Group 1 metal + a Group 7 non-metal e.g. sodium + chlorine ==> sodium chloride NaCl or ionic formula Na+Cl- In terms of electron arrangement, the sodium donates its outer electron to a chlorine atom forming a single positive sodium ion and a single negative chloride ion. The atoms have become stable ions, because electronically, sodium becomes like neon and chlorine like argon. Na (2.8.1) + Cl (2.8.7) ==> Na+ (2.8) Cl- (2.8.8) can be summarised electronically to give the stable 'noble gas' structures as [2,8,1] + [2,8,7] ==> [2,8]+ [2,8,8]- ONE The valencies of Na and Cl are both 1, that is, the numerical charge on the ions. sodium fluoride NaF, potassium bromide KBr and lithium iodide LiI etc. will all be electronically similar.
Li is 2.1, K is 2.8.8.1, F is 2.7, rest of dot and cross diagram is up to you. Reminder: How to work out formula of ionic compounds without going through some demanding electronic thinking is described on the "Elements, Compounds and Mixtures" page and it is followed by a section on naming compounds.
Example 2: A Group 2 metal + a Group 7 non-metal e.g. magnesium + chlorine ==> magnesium chloride MgCl2 or ionic formula Mg2+(Cl-)2 In terms of electron arrangement, the magnesium donates its two outer electrons to two chlorine atoms forming a double positive magnesium ion and two single negative chloride ions. The atoms have become stable ions, because electronically, magnesium becomes like neon and chlorine like argon. Mg (2.8.2) + 2Cl (2.8.7) ==> Mg2+ (2.8) 2Cl- (2.8.8) can be summarised electronically as [2,8,2] + 2[2,8,7] ==> [2,8]2+ [2,8,8]-2 ONE * NOTE you can draw two separate chloride ions, but in these examples square brackets and a number subscript have been used, as in ordinary chemical formula. The valency of Mg is 2 and chlorine 1, i.e. the numerical charges of the ions. Beryllium fluoride BeF2, magnesium bromide MgBr2, calcium chloride CaCl2 or barium iodide BaI2 etc. will all be electronically similar. Ca is 2.8.8.2, F is 2.7 rest of dot and cross diagrams are up to you.
Example 3: A Group 3 metal + a Group 7 non-metal e.g. aluminium + fluorine ==> aluminium fluoride AlF3 or ionic formula Al3+(F-)3 In terms of electron arrangement, the aluminium donates its three outer electrons to three fluorine atoms forming a triple positive aluminium ion and three single negative fluoride ions. The atoms have become stable ions, because aluminium and fluorine becomes electronically like neon. Valency of Al is 3 and F is 1, i.e. equal to the charges on the ions. Al (2.8.3) + 3F (2.7) ==> Al3+ (2.8) 3F- (2.8) can be summarised electronically as [2,8,3] + 3[2,7] ==> [2,8]3+ [2,8]-3 ONE Solid aluminium chloride/bromide/iodide have similar formula but are covalent when vapourised into Al2X6 dimer molecules, but AlCl3 has an ionic lattice in the solid, not sure on solid AlBr3 and AlI3, but these points are best left for an advanced AS-A2 chemistry discussion, not for GCSE students! Example 4: A Group 1 metal + a Group 6 non-metal e.g. sodium/potassium + oxygen ==> sodium/potassium oxide Na2O/K2O or ionic formula (Na+)2O2-/(K+)2O2- In terms of electron arrangement, the two sodium/potassium atoms donate their outer electron to one oxygen atom. This results in two single positive potassium ions to one double negative oxide ion. All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like). Valencies, K 1, oxygen 2. Lithium oxide, Li2O, sodium oxide Na2O, sodium sulphide Na2S and potassium K2S etc. will be similar. sodium oxide 2Na (2.8.1) + O (2.6) ==> 2Na+ (2.8.8) O2- (2.8) can be summarised electronically as 2[2,8,1] + [2,6] ==> [2,8]+2 [2,8]2- TWO or
potassium oxide 2K (2.8.8.1) + O (2.6) ==> 2K+ (2.8.8) O2- (2.8) can be summarised electronically as 2[2,8,8,1] + [2,6] ==> [2,8,8]+2 [2,8]2- TWO
The electronic similarities between the two examples are very obvious. Li is 2.1, Na is 2.8.1, S is 2.8.6 (for group 1 sulphide compound), rest of dots and crosses diagrams are up to you.
Example 5: A Group 2 metal + a Group 6 non-metal e.g. magnesium/calcium + oxygen ==> magnesium/calcium oxide MgO/CaO or ionic formula Mg2+O2-/Ca2+O2- In terms of electron arrangement, one magnesium/calcium atom donates its two outer electrons to one oxygen atom. This results in a double positive calcium ion to one double negative oxide ion. All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like). the valency of both calcium and oxygen is 2. magnesium oxide ONE For magnesium oxide: Mg (2.8.2) + O (2.6) ==> Mg2+ (2.8) O2- (2.8) the stable 'noble gas' structures can be summarised electronically as [2,8,2] + [2,6] ==> [2,8,8]2+ [2,8]2-
calcium oxide Ca (2.8.8.2) + O (2.6) ==> Ca2+ (2.8.8) O2- (2.8) can be summarised electronically as [2,8,8,2] + [2,6] ==> [2,8,8]2+ [2,8]2- ONE
Magnesium oxide MgO, magnesium sulphide MgS and calcium sulphide CaS will be similar electronically and give identical giant ionic lattice structures. Group 2 metals lose the two outer electrons to give the stable 2+ positive ion (cation) and S and O, both non-metals in Group 6, have 6 outer electrons and gain 2 electrons to form 2- negative ion (anion). For magnesium sulphide: Mg (2.8.2) + S (2.8.6) ==> Mg2+ (2.8) S2- (2.8.8) For calcium sulphide: Ca (2.8.8.2) + S (2.8.6) ==> Ca2+ (2.8.8) S2- (2.8.8) The dot and cross (ox) diagrams will be identical to that for calcium oxide above, except Mg instead of Ca (same group) and S instead of O (same group of Periodic Table). Example 6: A Group 3 metal + a Group 6 non-metal e.g. aluminium + oxygen ==> aluminium oxide Al2O3 or ionic formula (Al3+)2(O2-)3 In terms of electron arrangement, two aluminium atoms donate their three outer electrons to three oxygen atoms. This results in two triple positive aluminium ions to three double negative oxide ions. All the ions have the stable electronic structure of neon 2.8. Valencies, Al 3 and O 2. 2Al (2.8.3) + 3O (2.6) ==> 2Al3+ (2.8) 3O2- (2.8) can be summarised electronically as 2[2,8,3] + 3[2,6] ==> [2,8]3+2 [2,8]2-3 TWO
on another web page is how to work out an ionic formula given the ionic charges (combining power)
|
combines with ONE 
to form 
combines with TWO 
see *
combines with THREE 
to form 
to form
combine with ONE 
combines with ONE 
The
diagram on the right is typical of the giant ionic crystal structure of ionic
compounds like sodium chloride and magnesium oxide.
|
|
|
|
|||||||||
  |