Electronegativity, Ionic and Metallic Bonding

AS · 10 min

Why does sodium chloride form ions while hydrogen chloride forms a molecule? The answer is electronegativity: how strongly each atom pulls on the electrons between them. This note defines electronegativity, explains its trends with the same factors you used for ionisation energy, and uses the difference in electronegativity to decide whether a bond is ionic or covalent. It then sets out ionic and metallic bonding, both of which you must be able to define in the syllabus wording and illustrate.

Electronegativity

When two atoms share a pair of electrons, they do not always share it equally. One atom may pull the pair towards itself more strongly.

Definition

Electronegativity is the power of an atom to attract electrons to itself.

Electronegativity is measured on the Pauling scale, which has no units. Fluorine is the most electronegative element (4.0). Values you will be given in the data booklet include:

H	Li	Be	B	C	N	O	F
2.1	1.0	1.6	2.0	2.5	3.0	3.5	4.0

Na	Mg	Al	Si	P	S	Cl	K
0.9	1.3	1.5	1.9	2.2	2.6	3.0	0.8

Factors affecting electronegativity

An atom attracts the bonding electrons strongly if its nucleus has a strong pull on electrons in its outer shell. The factors are the same as for ionisation energy:

Key result

Nuclear charge: more protons, stronger attraction, higher electronegativity.
Atomic radius: bonding electrons further from the nucleus are attracted less strongly, lower electronegativity.
Shielding by inner shells and sub-shells: more shielding reduces the attraction, lower electronegativity.

Trends

Across a period electronegativity increases. The nuclear charge increases, the atomic radius decreases and the shielding by inner shells stays about the same, so the bonding electrons are attracted more strongly.
Down a group electronegativity decreases. The atomic radius increases and there are more inner shells shielding the bonding pair. These outweigh the increase in nuclear charge, so the bonding electrons are attracted less strongly.

So the most electronegative elements are at the top right of the Periodic Table (F, O, N, Cl), and the least at the bottom left (Cs, K, Na). Noble gases are usually not given values because they rarely form bonds.

Using electronegativity to predict bond type

The difference in electronegativity between two bonded atoms decides how unevenly the bonding electrons are shared.

difference in electronegativity	type of bonding	example
zero	non-polar covalent: equal sharing	$\ce{Cl2}$ ( $3.0 - 3.0 = 0$ )
small	polar covalent: unequal sharing, partial charges $\delta+$ and $\delta-$	$\ce{HCl}$ ( $3.0 - 2.1 = 0.9$ )
large (as a rough guide, more than about 1.7)	ionic: electron transfer, full charges	$\ce{NaCl}$ ( $3.0 - 0.9 = 2.1$ )

Think of it as a continuum. In $\ce{Cl2}$ the electrons sit exactly between the atoms. In $\ce{HCl}$ they are pulled towards chlorine, so $\ce{H}$ carries a small positive charge ( $\delta+$ ) and $\ce{Cl}$ a small negative charge ( $\delta-$ ). In $\ce{NaCl}$ the difference is so large that chlorine effectively takes the electron completely, forming $\ce{Na+}$ and $\ce{Cl-}$ .

Tip

The syllabus does not specify a sharp cut-off; questions give values and expect sensible judgement. "Large difference, so ionic; small difference, so covalent" is the reasoning examiners want. Covalent character in ionic compounds (polarisation of ions) is not assessed at AS.

Predicting bond types

Use Pauling electronegativity values to predict the type of bonding in (a) $\ce{MgO}$ , (b) $\ce{CH4}$ , (c) $\ce{AlCl3}$ , (d) $\ce{CaF2}$ .

Solution

(a) $\ce{MgO}$ : $3.5 - 1.3 = 2.2$ , large: ionic.

(b) $\ce{CH4}$ : $2.5 - 2.1 = 0.4$ , small: covalent (only slightly polar C–H bonds).

(c) $\ce{AlCl3}$ : $3.0 - 1.5 = 1.5$ , intermediate: covalent (polar). This matches its behaviour: aluminium chloride sublimes at a low temperature and exists as $\ce{Al2Cl6}$ molecules in the vapour.

(d) $\ce{CaF2}$ : $4.0 - 1.0 = 3.0$ , very large: ionic.

Ionic bonding

Definition

Ionic bonding is the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions).

Ionic bonds form between metals with low electronegativity and non-metals with high electronegativity. Electrons are transferred from the metal atom to the non-metal atom so that both usually reach a noble-gas configuration. The resulting ions attract each other in all directions, building a giant ionic lattice.

Dot-and-cross diagrams for ionic compounds

A dot-and-cross diagram shows the outer-shell electrons of each ion, with electrons from one atom drawn as dots and from the other as crosses. Each ion is drawn in square brackets with its charge outside.

Sodium chloride. Sodium ( $2,8,1$ ) transfers its single outer electron to chlorine ( $2,8,7$ ), giving $\ce{Na+}$ ( $2,8$ ) and $\ce{Cl-}$ ( $2,8,8$ ).

Dot-and-cross diagram for sodium chloride. The cross in the chloride ion is the electron transferred from sodium. The sodium ion's full outer shell (its second shell) is shown with crosses.

Magnesium oxide. Magnesium ( $2,8,2$ ) transfers two electrons to oxygen ( $2,6$ ), forming $\ce{Mg^2+}$ and $\ce{O^2-}$ .

Dot-and-cross diagram for magnesium oxide. The two crosses in the oxide ion came from magnesium.

Calcium fluoride. Calcium ( $2,8,8,2$ ) transfers one electron to each of two fluorine atoms, forming $\ce{Ca^2+}$ and two $\ce{F-}$ ions. That is why the formula is $\ce{CaF2}$ .

Dot-and-cross diagram for calcium fluoride: one calcium ion and two separate fluoride ions, each fluoride ion containing one electron (cross) from calcium.

Watch out

In an ionic dot-and-cross diagram, the ions are drawn separately, each in its own brackets with its charge. Do not draw overlapping circles or a shared pair: that is a covalent diagram. For $\ce{CaF2}$ , draw two separate $\ce{F-}$ ions (or one with " $\times 2$ " written beside it).

What makes an ionic bond strong

The electrostatic attraction between ions is stronger when the ions have higher charges and smaller radii (so their centres are closer). That is why magnesium oxide ( $\ce{Mg^2+}$ , $\ce{O^2-}$ ) has a far higher melting point (2852 °C) than sodium chloride ( $\ce{Na+}$ , $\ce{Cl-}$ ; 801 °C).

Metallic bonding

Definition

Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons.

In a metal, each atom releases its outer-shell electrons into a "sea" of delocalised electrons that are free to move throughout the structure. The remaining positive ions are arranged in a regular giant metallic lattice. The bonding is the attraction between the lattice of positive ions and the delocalised electrons, and it acts in all directions.

The strength of metallic bonding increases with:

the charge on the metal ions (and so the number of delocalised electrons per atom);
smaller ionic radius, so the delocalised electrons are closer to the nuclei.

metal	ion	delocalised electrons per atom	melting point / °C
sodium	$\ce{Na+}$	1	98
magnesium	$\ce{Mg^2+}$	2	650
aluminium	$\ce{Al^3+}$	3	660

Explaining melting points of metals

Explain why magnesium has a higher melting point than sodium.

Solution

Both have giant metallic lattices. Magnesium forms $\ce{Mg^2+}$ ions and contributes two delocalised electrons per atom; sodium forms $\ce{Na+}$ and contributes one. The $\ce{Mg^2+}$ ion also has a smaller radius than $\ce{Na+}$ . So there is a stronger electrostatic attraction between the positive ions and the delocalised electrons in magnesium, and more energy is needed to overcome it.

Metallic bonding explains the properties of metals: delocalised electrons can move and carry charge, so metals conduct electricity in the solid and liquid states; layers of ions can slide over each other without breaking the non-directional bonding, so metals are malleable and ductile. You will meet these again in the note on structure and properties.

Explaining an electronegativity trend

Explain why fluorine is more electronegative than chlorine.

Solution

A fluorine atom is smaller than a chlorine atom (two occupied shells instead of three), so the bonding pair of electrons is closer to the nucleus. It is also shielded by fewer inner electrons. Although chlorine has the greater nuclear charge, these effects outweigh it, so fluorine's nucleus attracts the bonding pair more strongly.

Exam-style: from data to bonding

Element X has electronegativity 0.8 and element Y has electronegativity 3.5. X is in Period 4 and forms a $1+$ ion; Y is in Period 2.

(a) Identify X and Y. (b) Predict the type of bonding in the compound they form, give its formula and draw its dot-and-cross diagram in words. (c) Explain why the compound has a high melting point.

Solution

(a) An element in Period 4 forming $1+$ ions with electronegativity 0.8 is potassium. Electronegativity 3.5 in Period 2 is oxygen.

(b) Difference $= 3.5 - 0.8 = 2.7$ : ionic. Potassium forms $\ce{K+}$ , oxygen forms $\ce{O^2-}$ , so the formula is $\ce{K2O}$ . Diagram: two separate $[\ce{K}]^+$ ions, each with a full outer shell of eight (crosses), and one $[\ce{O}]^{2-}$ ion with six dots and two crosses (one cross from each potassium atom), each ion in brackets with its charge.

(c) $\ce{K2O}$ has a giant ionic lattice with strong electrostatic attractions between oppositely charged ions acting in all directions. A large amount of energy is needed to overcome these many strong attractions.

Exam tip

Learn the three definitions exactly: electronegativity "the power of an atom to attract electrons to itself"; ionic bonding "electrostatic attraction between oppositely charged ions"; metallic bonding "electrostatic attraction between positive metal ions and delocalised electrons". The word electrostatic is part of the mark.
In metallic bonding, write "positive ions", never "positive nuclei" or "atoms", and "delocalised electrons", not "free electrons" alone.
For a trend explanation, give the factor, then its effect on attraction for the bonding electrons.

Summary

Electronegativity: the power of an atom to attract electrons to itself. Pauling scale, F = 4.0 highest.
Increases across a period (nuclear charge up, radius down, similar shielding); decreases down a group (radius and shielding up).
Electronegativity difference: zero, non-polar covalent; small, polar covalent; large, ionic.
Ionic bonding: electrostatic attraction between oppositely charged ions; stronger for higher charge and smaller ions.
Dot-and-cross for ionic compounds: separate ions, square brackets, charges.
Metallic bonding: electrostatic attraction between positive metal ions and delocalised electrons; stronger with more delocalised electrons and smaller ions.

Practice

Question

Define electronegativity.
State and explain the trend in electronegativity across Period 3 from sodium to chlorine.
Using the data booklet values, classify the bonds in $\ce{KCl}$ , $\ce{NH3}$ , $\ce{SiCl4}$ , $\ce{O2}$ and $\ce{LiF}$ as ionic, polar covalent or non-polar covalent.
Define ionic bonding and describe the formation of magnesium oxide from its atoms in terms of electron transfer.
Describe the dot-and-cross diagram for calcium fluoride, explaining why the formula is $\ce{CaF2}$ .
Define metallic bonding and explain why aluminium has a higher melting point than sodium.
In the bond $\ce{O-H}$ , state which atom carries the $\delta-$ charge and explain why.
Draw (or describe) the dot-and-cross diagram of lithium oxide, $\ce{Li2O}$ .
Explain why the electronegativity of nitrogen (3.0) is greater than that of phosphorus (2.2) but less than that of oxygen (3.5).
Element Q is in Group 2 and Period 3; element R has the highest electronegativity in Period 3. Predict the formula of the compound of Q and R, the type of bonding and one physical property, explaining each prediction.

Answers

The power of an atom to attract electrons to itself.
Electronegativity increases from Na (0.9) to Cl (3.0). The nuclear charge increases and the atomic radius decreases, while the shielding by inner shells (the first and second shells) stays about the same, so the bonding electrons are attracted more strongly.
$\ce{KCl}$ : $3.0 - 0.8 = 2.2$ , ionic. $\ce{NH3}$ : $3.0 - 2.1 = 0.9$ , polar covalent. $\ce{SiCl4}$ : $3.0 - 1.9 = 1.1$ , polar covalent. $\ce{O2}$ : 0, non-polar covalent. $\ce{LiF}$ : $4.0 - 1.0 = 3.0$ , ionic.
Electrostatic attraction between oppositely charged ions. Each magnesium atom ( $2,8,2$ ) transfers its two outer electrons to an oxygen atom ( $2,6$ ), forming $\ce{Mg^2+}$ ( $2,8$ ) and $\ce{O^2-}$ ( $2,8$ ); the ions attract each other in a giant lattice.
$[\ce{Ca}]^{2+}$ with a full outer shell (eight crosses) between two $[\ce{F}]^-$ ions, each with seven dots and one cross. Calcium has two outer electrons to lose but each fluorine atom can accept only one, so two fluorine atoms are needed for each calcium atom.
Electrostatic attraction between positive metal ions and delocalised electrons. Aluminium forms $\ce{Al^3+}$ ions, which are smaller than $\ce{Na+}$ ions, and provides three delocalised electrons per atom compared with one for sodium. The attraction between ions and delocalised electrons is much stronger, so more energy is needed to overcome it.
Oxygen is $\delta-$ because it is more electronegative (3.5) than hydrogen (2.1): it attracts the bonding pair more strongly, so the electron density is shifted towards O.
Two separate $[\ce{Li}]^+$ ions (each with the full first shell, two crosses, or shown empty of outer electrons) and one $[\ce{O}]^{2-}$ ion with six dots and two crosses, all in brackets with charges.
N vs P: nitrogen is smaller with fewer inner shells shielding, so its nucleus attracts the bonding pair more strongly, outweighing phosphorus's greater nuclear charge. N vs O: oxygen has a greater nuclear charge with the electrons in the same shell and the same shielding, and a slightly smaller radius, so oxygen attracts bonding electrons more strongly.
Q is magnesium; R is chlorine. Formula $\ce{MgCl2}$ (Mg loses two electrons, each Cl gains one). Difference $3.0 - 1.3 = 1.7$ , at the ionic end of the scale, and Mg is a reactive metal: ionic bonding (giant ionic lattice). Property: high melting point (strong electrostatic attractions between many ions), or conducts electricity when molten or in aqueous solution (ions free to move), or soluble in water.