2. 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

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   combines with ONE   to form 

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.

Note:

would represent the full electronic structure of the sodium ion.

would represent the full electronic structure of the chloride ion and 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.

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 (shown as just +) because there is one more positive proton than there are negative electrons in

the sodium ion.

The charge on the chloride ion Cl- is -1 units (shown as just -) because there is one more negative electron than there are positive protons in

the chloride ion.

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   combines with TWO   to form   see *

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.

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   combines with THREE  to form 

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   combine with ONE   to form 

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   combine with ONE   to form 

 

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   combines with ONE  to form 

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   combines with ONE  to form 

 

 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    combine with THREE   to form 

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

theoxide ion.

on another web page is how to work out an ionic formula given the ionic charges (combining power)

The properties of Ionic Compounds

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 way in a giant ionic

lattice structure shown on the left.

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+O2- 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 (see below).

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. 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.