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Part 2 Ionic bonding, structure and properties of ionic substances Doc Brown's Science–Chemistry Chemical Bonding GCSE/IGCSE/AS/A2 O Level Revision Notes What is the bonding in sodium chloride? This page describes the formation of an ionic bond by electron transfer, usually from a metal to a non–metal and give detailed annotated dot and cross diagrams of the resulting ionic compounds. ionic bonding i.e. the metal attraction of oppositely charged ions to give ionic bonds and the properties of ionic compounds is described. Some of the ionic compounds covered include sodium chloride NaCl, magnesium chloride MgCl2, aluminium fluoride AlF3, potassium oxide K2O, magnesium oxide MgO, calcium oxide CaO, magnesium sulfide MgS, calcium sulphide CaS and aluminium oxide Al2O3. The 2nd section describes and explains the physical properties that are characteristic of ionic compounds. Also, a section on how to work out the ionic formula of an ionic compound from the charges on the ions. Part 1 Introduction – why do atoms bond together (I suggest you 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 chemistry (in preparation, BUT a lot on intermolecular forces in Equilibria Part 8)
Part 2. IONIC BONDING – compounds and properties Examples of ionic compounds and 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 surrounding negative electrons in their respective energy levels. Charged particles called IONS are atoms, or groups of atoms, which have lost or gained electrons to have a overall net electrical positive or negative charge. The atom losing electrons forms a positive ion (a cation) and is usually a metal.
The atom gaining electrons forms a negative ion (an anion) and is usually a non–metallic element.
Therefore an IONIC BOND IS THE FORCE OF ATTRACTION BETWEEN ADJACENT IONS OF OPPOSITE CHARGE. Which electronic structures are the most stable? because this what atoms will try to get to electronically!
When atoms LOSE OR GAIN ELECTRONS, they try to attain the electron structure (electron configuration) of the electronically very stable atoms of the Group 0 Noble Gases eg helium (2), neon (2.8) or argon (2.8.8) etc. quite simply because these are the most stable electron arrangements and have a full outer shell of electrons (full highest energy level).
Ionically bonded molecules and the Periodic Table
All the atoms of the metallic elements on the left–hand side of the Periodic Table eg (Groups 1/2) have 1/2 electrons in their outer shell (highest energy level) that are readily lost to form a positive ion of charge +1/+2 (cations) eg sodium, potassium, magnesium and calcium etc. The electronic structure of these stable positive ions are those of a Noble Gas with a full outer shell.
The atoms of the non–metallic elements on the right–hand side of the Periodic Table eg (Groups 6/7) have 6/7 electrons in their outer shell and try to gain 2/1 electrons to become electronically stable like a Noble Gas eg oxygen and sulfur in Group 6 and the Group 7 Halogens – fluorine, chlorine, bromine and iodine. The electronic structure of these stable negative ions are those of a Noble Gas with a full outer shell.
Brief summary of the Periodic Table including electronic structure and formula patterns The examples below involve 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 which will alternate between being positive and negative to maximise attraction. Note: In the examples of ionic bonding it is assumed YOU can work out the electron configuration (arrangement in shells or energy levels) given the atomic number from the Periodic Table. 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.
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 (only original outer electrons shown above)
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.
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 (only original outer electrons shown above) NOTE
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.
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 (only original outer electrons shown above)
full electronic structure of aluminium fluoride 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–
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 (only original outer electrons shown above)
full electronic structure of sodium oxide
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 (only original outer electrons shown above)
full electronic structure of potassium oxide
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. electronic structures of sodium sulfide Na2S and potassium sulfide K2S
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–
magnesium oxide ONE (only original outer electrons shown above)
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] + [2,6] ==> [2,8]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 (only original outer electrons shown above)
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). eg electronic structure of magnesium sulfide
electronic structure of calcium sulfide
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
TWO (only original outer electrons shown above)
full electronic structure of aluminium oxide on another web page is how to work out an ionic formula given the ionic charges (combining power)
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