2. Ionic Bonding explained What is an ionic bond? Electron transfer How is an ionic bond formed? description of ions crystal lattice oxides chlorides properties of ionic compounds gcse igcse O Level A level GCE AS A2 chemistry







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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 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.

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!

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 and have 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.

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.

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 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^– ==> O^2– [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.

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 (c) doc b atom combines with ONE atom to form

(only original outer electrons shown above)

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 of the sodium ion and the chloride ion, 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 ion is formed.

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.

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 (c) doc b atom combines with TWO atoms to form

(only original outer electrons shown above)

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 of the magnesium ion and the chloride ion, 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.

Ca is 2.8.8.2, F is 2.7 rest of dot and cross diagrams are up to you.

calcium chloride

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 (c) doc b atom combines with THREE atoms to form

(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⁺)₂O^2–/(K⁺)₂O^2–

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) O^2– (2.8)

can be summarised electronically as 2[2,8,1] + [2,6] ==> [2,8]⁺₂ [2,8]^2–

TWO (c) doc b atoms combine with ONE atom to form

(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) O^2– (2.8)

can be summarised electronically as 2[2,8,8,1] + [2,6] ==> [2,8,8]⁺₂ [2,8]^2–

TWO (c) doc b atoms combine with ONE atom to form

(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⁺O^2–/Ca²⁺O^2–

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 (c) doc b atom combines with ONE atom to form

(only original outer electrons shown above)

full electronic structure of magnesium oxide

For magnesium oxide: Mg (2.8.2) + O (2.6) ==> Mg²⁺ (2.8) O^2– (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) O^2– (2.8)

can be summarised electronically as [2,8,8,2] + [2,6] ==> [2,8,8]²⁺ [2,8]^2–

ONE (c) doc b atom combines with ONE atom to form

(only original outer electrons shown above)

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) S^2– (2.8.8)

For calcium sulphide: Ca (2.8.8.2) + S (2.8.6) ==> Ca²⁺ (2.8.8) S^2– (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 MgS

electronic structure of calcium sulfide CaS

Example 6: A Group 3 metal + a Group 6 non–metal

e.g. aluminium + oxygen ==> aluminium oxide Al₂O₃ or ionic formula (Al³⁺)₂(O^2–)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) 3O^2– (2.8)

can be summarised electronically as 2[2,8,3] + 3[2,6] ==> [2,8]³⁺₂ [2,8]^2–₃

TWO (c) doc b atoms combine with THREE atoms to form

(only original outer electrons shown above)

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 O^2– 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)

The diagram on the right is typical of the giant ionic crystal structure of ionic compounds like sodium chloride and magnesium oxide.
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 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 boiling points, so they are not very volatile!
A relatively large amount of energy is needed to melt or boil ionic compounds. 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⁺O^2– has a higher melting point than sodium chloride Na⁺Cl^–.
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 assumptions.
- 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 (see diagrams below).

A '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 an ionic compound formula

selected ions and charges

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 O^2–

number of Al³⁺ x charge on Al³⁺ = number of O^2– x charge on O^2–

the simplest numbers are 2 of Al³⁺ x 3 = 3 of O^2– 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

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 O^2– because 2 x 1 = 1 x 2 gives Na₂O or (Na⁺)₂O^2– 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 SO₄^2– because 2 x 3 = 3 x 2 gives Fe₂(SO₄)₃ or (Fe³⁺)₂(SO₄^2–)₃ iron(III) sulphate

some keywords and formulae: how to work out ionic bonding diagrams for 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)

WHAT NEXT? and other associated Pages

foundation tier (easier) multiple choice quiz on the Structure, Properties and Chemical Bonding of Materials

higher tier (harder) multiple choice quiz on the Structure, Properties and Chemical Bonding of Materials

Atomic Structure notes & diagrams including electron arrangements

Models and properties of Gases, Liquids and Solids

A brief survey–summary of the Periodic Table including electronic structure and formula patterns

How to work out formula of ionic or covalent compounds without going through some demanding electronic thinking is described in the valency section on the Elements, Compounds and Mixtures page and it is followed by a section on naming compounds.

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