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(c) doc b(c) doc b

CHEMICAL BONDING Part 2 Ionic bonding, structure and properties of ionic substances

Doc Brown's Science–Chemistry Chemical Bonding GCSE/IGCSE/O/AS/A2 Level  Revision Notes

What is the bonding in sodium chloride? What are ions? How is an ionic bond formed? What are giant ionic lattices? IONIC BONDING DIAGRAMS and PROPERTIES OF IONIC COMPOUNDS 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 are described, discussed and explained. 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 and how to name simple ionic compounds. These notes on ionic bonding and the properties of ionic compounds are designed to meet the highest standards of knowledge and understanding required for students/pupils doing GCSE chemistry, IGCSE chemistry, O Level chemistry, KS4 science courses and a basic primer for AS/A Level chemistry courses.

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.

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(c) doc b2. Ionic Bonding – electron transfer

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 one or more electrons to have a overall net electrical positive charge or negative charge.

In losing or gaining electrons, the atoms try to attain a stable electron arrangement of a noble gas e.g. a full outer shell of electrons.

For a given atom, a nearly full shell will try to gain electrons and a nearly empty shell will tend to lose electrons

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) e.g.

Group 1 alkali metals lose their single outer electron to form single positive ions e.g. Na ==> Na+ + 2e

Group 2 metals lose their two outer electrons to form doubly charged positive ions e.g. Mg ==> Mg2+ + 2e

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 e.g.

Group 7 halogen atoms gain one electron to form a singly charged negative ion e.g. Cl + e ==> Cl

Group 6 non–metals gain two electrons to form a doubly charged negative ion e.g. O + 2e ==> O2–

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!

(c) doc b (c) doc b (c) doc b symbol (atomic number) electron arrangement

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 with a full outer shell of electrons (full highest energy level).

In advanced level chemistry you will encounter examples of electronic structures of ions that are NOT those of a Noble Gas.

Its a good idea to have some idea of where the elements are in the periodic table, and their electronic structure, before looking at the theoretical electronic model for ionic bonding in ionic compounds

The black zig–zag line 'roughly' divides the metals on the left from the non–metals on the right of the elements of the Periodic Table.

(c) doc b

The electronic structures of the first 20 elements of the Periodic Table

You need to know about these to understand the details of ionic chemical bonding

Pd metals Part of the modern Periodic Table

Pd = period, Gp = group

metals => non–metals
Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0
1

1H  Note that H does not readily fit into any group

2He
2 3Li 4Be atomic number Chemical Symbol eg 4Be 5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba Transition Metals 81Tl 82Pb 83Bi 84Po 85At 86Rn
Gp 1 Alkali Metals  Gp 2 Alkaline Earth Metals  Gp 7 Halogens  Gp 0 Noble Gases

Chemical bonding comments about the selected elements highlighted in white

e.g. When the metals on the left combine with the non–metals on the right, an ionic bond is formed e.g. the formation of an ionic compound like sodium chloride NaCl

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) which 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.

eg Na [2.8.1] ==> Na+ [2.8] like neon + e, or Ca [2.8.8.2] ==> Ca2+ [2.8.8] like argon + 2e

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 with a full outer shell of electrons 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.

eg O [2.6] + 2e ==> O2– [2.8] like neon, or Cl [2.8.7] + e ==> Cl [2.8.8] like argon

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.

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Example 1: A Group 1 Alkali Metal + a Group 7 Halogen non–metal

e.g. sodium + chlorine ==> sodium chloride NaCl or ionic formula Na+Cl 

In terms of electron arrangement in the formation of the ionic compound sodium chloride, 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 via electron transfer ...

... sodium becomes like neon (sodium ion, Na+) and chlorine like argon (chloride ion, Cl).

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]

so both the sodium and chloride ions have a full outer shell like a noble gas

ONE (c) doc batom combines withONE(c) doc b atom to form (c) doc b(c) doc b

Note in this electron diagram, only the original outer electrons are shown above.

The outer electron of the sodium atom (2.8.1) is transferred to the outer shell of the chlorine atom (2.8.7) giving it a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the sodium ion also attains a stable noble gas electron structure (2.8).

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.

Only the outer valency electrons of the chloride ion are shown, the 'blob' electron represents the electron from the sodium atom which is accepted by the chlorine atom to form the chloride ion.

The charge on the sodium ion Na+ is +1 units (by convention shown as just +) because there is one more positive proton than there are negative electrons in the sodium ion (11p, 10e).

The charge on the chloride ion Cl is –1 units (by convention shown as just ) because there is one more negative electron than there are positive protons in the chloride ion (17p, 18e).

Note:

would represent the full electronic structure diagram of the sodium ion [2.8] and the chloride ion [2.8.8], hence the full electronic structure of sodium chloride. Note that the 'blob' and 'x' electrons are identical, but their use is just a useful visual device to show how the ions are formed. The blue circle represents the nucleus.

See Example 6. aluminium oxide for more highly charged ion analysis.

Li is 2.1, K is 2.8.8.1, F is 2.7, rest of dot and cross diagram is up to you.

Gp1\7 F Cl Br I
Li LiF LiCl LiBr LiI
Na NaF NaCl NaBr NaI
K KF KCl KBr KI
Rb Rb RbCl RbBr RbI
Cs CsF CsCl CsBr CsI

All the formula highlighted in yellow can be described in the same way as sodium chloride

The Group 1 Alkali Metal atom loses one electron to form a singly charged positive ion

The Group 7 Halogen atom gains one electron to form a singly charged negative ion

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 OR there is a section Appendix 1. at the end of this page.

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Example 2: A Group 2 Alkaline Earth Metal + a Group 7 Halogen non–metal

e.g. magnesium + chlorine ==> magnesium chloride MgCl2 or ionic formula Mg2+(Cl)2 

In terms of electron arrangement in the formation of the ionic compound magnesium chloride, the magnesium donates its two outer electrons to two chlorine atoms forming a double positive magnesium ion and two single negative chloride ions via electron transfer.

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  via electron transfer

so both the magnesium and chloride ions have a full outer shell of electrons like a noble gas

ONE (c) doc batom combines withTWO (c) doc b atoms to form (c) doc b(c) doc b

Note in this electron diagram, only the original outer electrons are shown above.

The outer electrons of the magnesium atoms (2.8.2) are transferred to the outer shell of the chlorine atom (2.8.7) giving it a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the magnesium ion also attains a stable noble gas electron structure (2.8).

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.

represents the full electronic structure diagram of the magnesium ion [2.8] and the chloride ion [2.8.8], hence the full electronic structure of magnesium chloride.

Note that the 'blob' and 'x' electrons are identical, but their use is just a useful visual device to show how the ion is formed. The blue circle represents the nucleus.

 

Ca is 2.8.8.2, Cl is 2.8.7, F is 2.7 rest of dot and cross diagrams are up to you, but calcium chloride is shown below.

The calcium atoms transfer their two outer electrons to the outer shell of two chlorine atoms

calcium chloride

The two outer electrons of the calcium atoms (2.8.8.2) are transferred to the outer shell of two chlorine atoms (2.8.7) until it has a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the calcium ion also attains a stable noble gas electron structure (2.8.8). The blue circle represents the nucleus.

Gp2\7 F Cl Br I
Mg MgF2 MgCl2 MgBr2 MgI2
Ca CaF2 CaCl2 CaBr2 CaI2
Sr SrF2 SrCl2 SrBr2 SrI2
Ba BaF2 BaCl2 BaBr2 BaI2

All the formula highlighted in yellow can be described in the same way as magnesium chloride or calcium chloride

The Group 2 Alkaline Earth Metal atom loses two electrons to form a doubly charged positive ion

The Group 7 Halogen atom gains one electron to form a singly charged negative ion

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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 in the formation of the ionic compound aluminium fluoride, 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 via electron transfer.

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

so both the aluminium and fluoride ions have a full outer shell like a noble gas

ONE (c) doc batom combines withTHREE (c) doc batoms to form (c) doc b(c) doc b

Note in this electron diagram, only the original outer electrons are shown above.

The outer electrons of the aluminium atom (2.8.3) is transferred to the outer shell of the fluorine atoms (2.7) giving them a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the aluminium ion also attains a stable noble gas electron structure (2.8).

full electronic structure diagram of aluminium fluoride, the blue circle represents the nucleus.

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!

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Example 4: A Group 1 Alkali Metal + a Group 6 non–metallic element

e.g. sodium/potassium + oxygen ==> sodium/potassium oxide Na2O/K2O or ionic formula (Na+)2O2–/(K+)2O2–

In terms of electron arrangement in the formation of the ionic compound sodium oxide, 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 via electron transfer.

All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like).

Valencies, sodium/potassium 1, oxygen/sulfur 2. giving the following formulae:

Lithium oxide, Li2O, sodium oxide Na2O, sodium sulphide Na2S and potassium K2S etc.

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–

so both the sodium and oxide ions have a full outer shell like a noble gas

TWO (c) doc batoms combine withONE (c) doc batom to form (c) doc b

or

(c) doc b + (c) doc b + (c) doc b ==> (c) doc b(c) doc b(c) doc b

Note in this electron diagram, only the original outer electrons are shown above.

The outer electrons of the sodium atoms (2.8.1) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the sodium ion also attains a stable noble gas electron structure (2.8).

full electronic structure diagram of sodium oxide, the blue circle represents the nucleus.

 

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–

so both the potassium and oxide ions have a full outer shell like a noble gas

TWO (c) doc batoms combine withONE (c) doc batom to form (c) doc b(c) doc b

Note in this electron diagram, only the original outer electrons are shown above.

The outer electrons of the potassium atoms (2.8.8.1) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the potassium ion also attains a stable noble gas electron structure (2.8.8).

full electronic structure diagram of potassium oxide,  the blue circle represents the nucleus.

 

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 dot and cross diagrams are up to you.

e.g. electronic structure diagrams for sodium sulfide Na2S and potassium sulfide K2S

sodium sulfide

The outer electrons of the sodium atoms (2.8.1) are transferred to the outer shell of the sulfur atom (2.8.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the sodium ion also attains a stable noble gas electron structure (2.8).

 

potassium sulfide

The outer electrons of the potassium atoms (2.8.8.1) are transferred to the outer shell of the sulfur atom (2.8.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8.8). At the same time, the potassium ion also attains a stable noble gas electron structure (2.8.8).

Gp1\6 O S
Li Li2O Li2S
Na Na2O Na2S
K K2O K2S
Rb Rb2O Rb2S
Cs Cs2O Cs2S

All the formula highlighted in yellow can be described in the same way as sodium oxide, potassium oxide, sodium sulfide or calcium sulphide

The Group 1 Alkali Metal atom loses one electron to form a singly charged positive ion

The Group 6 non–metal atom gains two electrons to form a doubly charged negative ion

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Example 5: A Group 2 Alkaline Earth Metal + a Group 6 non–metallic element

e.g. magnesium/calcium + oxygen ==> magnesium/calcium oxide MgO/CaO or ionic formula Mg2+O2–/Ca2+O2–

In terms of electron arrangement in the formation of the ionic compound magnesium oxide, 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 via electron transfer.

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 (c) doc batom combines withONE (c) doc batom to form (c) doc b(c) doc b

Note in this electron diagram, only the original outer electrons are shown above.

The two outer electrons of the magnesium atoms (2.8.2) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the magnesium ion also attains a stable noble gas electron structure (2.8).

full electronic structure of magnesium oxide

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–

so both the magnesium and oxide ions have a full outer shell like a noble gas

 

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 (c) doc batom combines withONE (c) doc batom to form (c) doc b(c) doc b

Note in this electron diagram, only the original outer electrons are shown above.

The two outer electrons of the calcium atoms (2.8.8.2) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas (2.8). At the same time, the calcium ion also attains a stable noble gas electron structure (2.8.8).

full electronic structure of calcium oxide

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)

so both the magnesium/calcium and sulfide ions have a full outer shell like a noble gas

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 MgS

electronic structure of calcium sulfide CaS

Gp2\6 O S
Mg MgO MgS
Ca CaO CaS
Sr SrO SrS
Ba BaO BaS

All the formula highlighted in yellow can be described in the same way as magnesium oxide, magnesium sulphide, calcium oxide or calcium sulphide

The Group 2 Alkaline Earth Metal atom loses two electrons to form a doubly charged positive ion

The Group 6 non–metal atom gains two electrons to form a doubly charged negative ion

 

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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 in the formation of the ionic compound aluminium oxide, 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 via electron transfer.

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

so both the aluminium and oxide ions have a full outer shell like a noble gas

TWO  (c) doc batoms combine withTHREE (c) doc b atoms to form (c) doc b(c) doc b

Note in this electron diagram, only the original outer electrons are shown above.

The three outer electrons of the aluminium atoms (2.8.3) are transferred to the outer shell of the oxygen atom (2.6) until it has a complete octet shell of outer electrons, just like a noble gas 2.8). At the same time, the aluminium ion also attains a stable noble gas electron structure (2.8).

Note:

The charge on the aluminium ion Al3+ is +3 units (shown as 3+) because there are three more positive protons than there are negative electrons in the aluminium ion.

The charge on the oxide ion O2– is –2 units (shown as 2–) because there are two more negative electrons than there are positive protons in the oxide ion.

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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(c) doc bThe crystal structure and properties of Ionic Compounds

  • (c) doc bA GIANT IONIC LATTICE – explaining its properties
  • The diagram on the right is typical of the giant ionic crystal structure of ionic compounds like sodium chloride and magnesium oxide.
  • Solid ionic compounds consist of a giant lattice of closely packed ions which are all combine together to form a crystal.
  • The alternate positive and negative ions in an ionic solid are arranged in an orderly/regular way in a giant ionic lattice structure shown on the right.
  • The ionic bond is the strong electrical attraction between the oppositely charged positive and negative ions next to each other in the lattice.
  • The bonding extends throughout the crystal in all directions.
  • Salts and metal oxides are typical ionic compounds.
  • This strong bonding force makes the structure hard (if brittle) and have high melting and very high boiling points, so they are not very volatile!
  • A relatively large amount of energy is needed to melt or boil ionic compounds to reduce/overcome the strong bonding forces.
    • Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.
  • The bigger the charges on the ions the stronger the bonding attraction e.g. magnesium oxide Mg2+O2– has a much higher melting point than sodium chloride Na+Cl.
    • The ions of magnesium oxide are both doubly charged so the electrostatic attraction is much greater (its actually about 4x as strong attractive force).
      • As it happens in this case, the ions in magnesium oxide are smaller than the ions in sodium chloride, so the ions in magnesium oxide can pack closer together and this also increase the attractive bonding force.
    • This double effect results in a much stronger ionic bond in magnesium oxide, so a much greater thermal kinetic energy i.e. a much greater temperature, is required to weaken the giant ionic lattice and melt the crystals of magnesium oxide compared to sodium chloride.
    • Simple experimental evidence – sodium chloride melts at 801oC, whereas magnesium oxide melts much higher at 2852oC.
  • Unlike covalent molecules, ALL ionic compounds are crystalline solids at room temperature.
  • They are hard but brittle, when stressed the bonds are broken along planes of ions which shear away.
    • They are NOT malleable like metals.
  • Many ionic compounds are soluble in water, but not all, so don't make this assumption.
    • Salts can dissolve in water because the ions can separate and become surrounded by water molecules which weakly bond to the ions (see diagrams below).
    • This reduces the attractive forces between the ions, preventing the crystal structure to exist.
    • Evaporating the water from a salt solution will eventually allow the ionic crystal lattice to reform.
  • The solid crystals DO NOT conduct electricity because the ions are not free to move to carry an electric current.
    • However, if the ionic compound is melted or dissolved in water, the liquid will now conduct electricity, as the ion particles are now free to move and carry the electric current in the molten salt or the solution of the salt in aqueous solution (see diagrams below).
    • This electrical conduction under these conditions is evidence for the existence of ions in this type of compound.

 

An 'advanced' molecular particle picture of sodium chloride dissolving in water

(the partial electrical charges δ+ and δ– are for advanced level students only)

 

==>

solid sodium chloride ==> molten sodium chloride (fixed ions to free moving ions)


Appendix 1. How to work out the formula for an ionic compound

Table 1a.table of ions, names and symbols (c) doc bselected ions and charges

Table 1b. The periodic table pattern of charges on ions

The charge is based on the number of electrons lost (giving positive ions) or gained (giving negative ions) forming a noble gas electron structure, i.e. to make a full outer shell of electrons

CATIONS from metals ANIONS from non–metals
Group 1 Group 2 Group 3 Group 6 Group 7
lithium ion Li+ beryllium ion Be2+   oxide ion O2– fluoride ion F
sodium ion Na+ magnesium ion Mg2+ aluminium ion Al3+ sulfide ion S2– chloride Cl
potassium ion K+ calcium ion Ca2+     bromide ion Br
        iodide ion I

In the electrically balanced stable formula, the total positive ionic charge must equal the total negative ionic charge.

To work out an ionic formula by combining ion 'A' with ion 'B' the rule is:

number of ion 'A' x charge of ion 'A' = number of ion 'B' x charge of ion 'B' (you ignore charge sign)

Example: How do we work out that the formula of aluminium oxide is Al2O3?

As difficult an example as any you will have to work out!

Aluminium oxide consists of aluminium ions Al3+ and oxide ions O2– 

number of Al3+ ions x charge on Al3+ balances the number of  O2– ions x charge on O2– 

the simplest numbers are 2 of Al3+ x 3 balances 3 of  O2– x 2 (total 6+ balances total 6–)

so the simplest whole number formula for aluminium oxide is Al2O3 

6 more examples of working out an ionic formula, starting with some easy ones!

numerically ion charge = valency of A or B to deduce the formula

valency or ionic charge = the combining power of the ion

'molecular' or ionic style of formula and compound name

1 of K+ balances 1 of Br because 1 x 1 = 1 x 1 gives KBr or K+Br  potassium bromide

2 of Na+ balances 1 of O2– because 2 x 1 = 1 x 2 gives Na2O or (Na+)2O2–  sodium oxide

1 of Mg2+ balances 2 of Cl because 1 x 2 = 2 x 1 gives MgCl2 or Mg2+(Cl)2  magnesium chloride

1 of Fe3+ balances 3 of F because 1 x 3 = 3 x 1 gives FeF3 or   Fe3+(F)3  iron(III) fluoride

1 of Ca2+ balances 2 of NO3 because 1 x 2 = 2 x 1 gives Ca(NO3)2 or Ca2+(NO3)2  calcium nitrate

2 of Fe3+ balances 3 of SO42– because 2 x 3 = 3 x 2 gives Fe2(SO4)3 or (Fe3+)2(SO42–)3  iron(III) sulphate

 

By applying similar logic you can work out the charge on one ion, knowing the formula and charge on the other ion

e.g. supposing a metal M forms an ionic chloride compound MCl2, Cl will be the chloride ion (charge single –), so to balance the two chloride ions, the metal ion must carry a charge of 2+ i.e. the M2+ ion.

If a metal that forms a singly charged positive ion M+, forms an ionic sulfide compound M2S, the charge on the sulfur ion must be 2– i.e. S2–, to balance the two + charges of the metal ion.

 

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Appendix 2 How do you name simple ionic compounds?

How to name ionic compounds: Naming simple ionic compounds isn't difficult and through your course you get used to the names as they crop up.

The name usually consists of two parts, first the name of the positive metal ion, and secondly, the name of the negative ion derived from the non–metal.


For the metal ions (cations) it couldn't be easier, its just the metal name itself e.g. sodium, magnesium, aluminium etc.

For the Na+ ion e.g. sodium chloride

For the Mg2+ ion e.g. magnesium bromide

For the Al3+ ion eg. aluminium oxide etc.

However, there is one complication when a metal can form two different ions like copper or iron.

In these cases the two ions are distinguished in the name by a Roman numerals number in brackets after the name of the metal which corresponds to the numerical value of the positive charge on the metal ion.

e.g. Cu2O copper(I) oxide contains the Cu+ copper(I0 ion, and CuO copper(II) oxide contains the Cu2+ copper(II) ion

and FeCl2 iron(II) chloride contains the Fe2+ iron(II) ion, and FeCl3 contains the Fe3+ iron(III) ion.


However, things are a bit more complicated for the negative ions (anions) because although the name of the ion is derived from the name of the non–metallic element, it is a bit different.

Oxygen forms the oxide ion, sulfur the sulfide ion etc. the names of these anions from group 6 end in ..ide,

e.g. sodium oxide, magnesium oxide, aluminum oxide

Fluorine forms the fluoride ion, chlorine the chloride ion, bromine the bromide ion and iodine the iodide ion.

The names of the anions from group 7 halogens (naming ending ...ine) end in ...ide, the halide ions),

e.g. potassium fluoride, sodium chloride, calcium bromide etc.

For anions where two or more non–metallic atoms are combined in a single ion, and one of them is oxygen, the name often ends in ..ate

e.g. carbonate (C + O), sulfate (S + O), nitrate (N + O), chlorate (Cl + O) etc. see their formula in the table 1a. of Appendix 1.


Appendix 3. Tests for Cations and Anions

These are all written up in detail on other pages, so see ...

TESTS for Metal cations (positive ions) Group 1 The Alkali Metals  and  Transition Metals

TESTS for Anions (negative ions)  and  Group 7 The Halogens


some keywords and formulae: how to work out ionic bonding diagrams for the ionic compounds Li2O Na2O K2O Li2S Na2S K2S LiF LiCl LiBr LiI NaF NaCl NaBr NaI KF KCl KBr KI MgF2 MgCl2 MgBr2 MgI2 CaF2 CaCl2 CaBr2 CaI2 MgO CaO AlF3 Al2O3 Al2S3 worked out


Revision notes information to help revise KS4 Science Additional Science Triple Award Separate Sciences GCSE/IGCSE/O level Chemistry Revision–Information Study Notes for revising for AQA GCSE Science, Edexcel GCSE Science/IGCSE Chemistry & OCR 21st Century Science, OCR Gateway Science WJEC/CBAC GCSE science–chemistry CCEA/CEA GCSE science–chemistry (and courses equal to US grades 8, 9, 10) basic aid notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA OCR Edexcel Salters CIE, CCEA/CEA & WJEC advanced level courses for pre–university students (equal to US grade 11 and grade 12 and Honours/honors level courses)


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