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Ionic Bonds

Ionic bonds are formed between metals and non-metals. In order to have a stable outer shell (an octet), a metal 'teams up' with a non-metal. The metal donates its excess electron(s) to the needy non-metal, and everyone is happy.

An ion is an atom which has an electrical charge. The number of protons of an element never changes, so the charge is due to missing or extra electrons. Typically, atoms of elements with high group numbers (6 - 7) gain electrons and become negative when they complete their outer electron shell (8 electrons). Atoms of elements with low group numbers (1 - 3) have these numbers of electrons in their outer shell, so to need to shed these in order to have a complete outer shell, becoming positive as they do so.

Periodic Groups and the Ions they form

Group I elements are alkali metals and form ions with a single positive charge (Li$^{+}$, Na$^{+}$, K$^{+}$, Rb$^{+}$, Cs$^{+}$). Group II elements are alkali-earth metals and form ions with a double positive charge (Be$^{2+}$, Mg$^{2+}$, Ca$^{2+}$, Sr$^{2+}$, Ba$^{2+}$). The transition metals, such as iron, copper, and zinc, have more complex chemical behaviours. For example, there are two forms of iron compounds: ferric (+3) and ferous (+2).

Group IV elements are a mix of non-metals (C and Si) and metals (Ge, Sn, Pb). They can either lose four electrons or gain four electrons to complete their outer shell. Due to the difficulty in losing or gaining so many electrons, these elements are more likely to opt for covalent relationships with other non-metals, and share their electrons to complete their shells.

Group V elements need to gain three electrons to complete their outer shell, so form ions with a charge of negative three (N$^{3-}$, P$^{3-}$, As$^{3-}$, Sb$^{3-}$). Group VI elements need to gain two electrons to complete their outer shell, so form ions with a charge of negative two (O$^{2-}$, S$^{2-}$, Se$^{2-}$, Te$^{2-}$). Group VII elements (the halogens) need to gain only one electron to complete their outer shell, so form ions with a charge of negative one (F$^{-}$, Cl$^{-}$, Br$^{-}$, I$^{-}$).

Group VIII (He, Ne, Ar, Kr, Xe, Rn) are the noble gases. They already have complete outer shells, so do not easily form bonds with other atoms, even of their own kind. They are relatively rare, and are found in nature as single atom gases.

Sodium chloride

Salt
Salt is an ionic compound consisting of sodium and chloride ions

Sodium (a metal) loses one electron, forming Na$^{+}$, and chlorine (a non-metal) gains that electron, forming Cl$^{-}$, to form an ionic compound, NaCl. You know this as table salt.

NaCl
Sodium and chlorine exchange an electron and form attracting positive and negative ions

The changes to the electronic configuration of the outer shell changes dramatically the radius of the atoms. The radius of the chloride ion (Cl$^{-}$) is twice that of the chlorine atom, even though only one electron has been added. Similarly, when a sodium atom loses an electron, it shrinks to half its original radius. These changes have implications for the energy balance of the reaction.

Take for example the reaction of a piece of hot sodium metal with gaseous chlorine (Cl$_2$).

2Na(s) + Cl$_2$(g) → 2 NaCl(s)

The reaction is exothermic, and the product is solid sodium chloride salt. The ionic bonding of Cl and Na is exothermic, giving out net energy, even though energy needed to be put in to break the Cl-Cl bond of the gas. The first ionization energy of sodium is +494 kJ mol$^{-1}$, meaning this not inconsiderable amount of energy is needed to strip off its electron to donate to the chlorine. The conversion of solid sodium to gas also requires an input of energy.

On the other side of the energy balance, energy is returned to the system by the chlorine accepting the electron (electron affinity, -349 kJ mol$^{-1}$). Most of the energy of the exothermic reaction, however, is output by the formation of the very strong ionic crystal lattice (lattice enthalpy). The lattice consists of equal numbers of chloride and sodium ions, arranged so that each chloride is surrounded by 6 equidistant (all as far from each other as possible) positive sodium ions, which in turn are each surrounded by 6 negative chloride ions.

The very high temperature required to melt and gasify salt is testimony to the amount of energy released in the formation of the salt, and the strength of each ionic bond and the lattice as a whole.

Polyatomic Ions

Ions can also be composed of more than one atom, and more than one element. In these, the charge is not due to a single electron displacement from a single atom, but is a 'non-localised' or delocalised charge. This means the charge is a result of all the atoms in the group as a whole.

Examples are: ammonium NH$_4^{+}$, nitrate NO$_3^{-}$ (formed when nitric acid HNO$_3$ loses a proton), hydroxide OH$^{-}$, sulfate SO$_4^{2-}$, carbonate CO$_3^{2-}$, phosphate P$_4^{3-}$, and hydrogen carbonate HC$_3^{-}$.

Electronegativity

The elements of Groups 1, 2 and 3 tend to have low electronegativities. Those of Groups 5, 6 and 7 usually have higher electronegativities. The greater the electronegativity difference between two ions the greater the probability (usually) that the two ions will form an ionic bond rather than a covalent bond.

For example, magnesium has an electronegativity of 1.2. Chlorine and oxygen have electronegativities of 3.0 and 3.5 respectively. The differences are 1.8 and 2.0. Magnesium forms ionic bonds with both (MgCl$_2$ and MgO).

Beryllium and aluminium both have an electronegativity of 1.5. The difference to chlorine (3.0) is 1.5. Neither beryllium not aluminium forms ionic bonds with chlorine. Instead they form covalent molecules (BeCl$_2$ and Al$_2$Cl$_6$).

The rule appears to be: if the difference in electronegativity is less than 1.8, the ions will form covalent bonds in preference to ionic. The rule holds in most cases, but not all, so be careful!

H
2.1
He
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Ne
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
Ar
K
0.8
Ca
1.0
Sc
1.3
Ti
1.5
V
1.6
Cr
1.6
Mn
1.5
Fe
1.8
Co
1.8
Ni
1.8
Cu
1.9
Zn
1.6
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
2.8
Kr
Rb
0.8
I
2.5
Xe
Cs
0.7
At
2.2
Rn

Linus Pauling's scale of electronegativities

There are exceptions to this rule, and some compounds show some degree of ambiguity in their ionic or covalent preference. An experimental test for electronegativity is its conductivity when molten. Ionic bonding is predominant when the compound conducts electricity well when in liquid form of in acqueous solution. Covalent compounds tend not to be very good conductors in either state.

Content © Andrew Bone. All rights reserved. Created : August 21, 2015 Last updated :March 12, 2016

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