Electronegativity

ELECTRONEGATIVITY
This page explains what electronegativity is, and how and why it varies around the Periodic Table. It looks at the way that electronegativity differences affect bond type and explains what is meant by polar bonds and polar molecules.
If you are interested in electronegativity in an organic chemistry context, you will find a link at the bottom of this page.
What is electronegativity
Definition
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.
The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.
What happens if two atoms of equal electronegativity bond together?
Consider a bond between two atoms, A and B. Each atom may be forming other bonds as well as the one shown - but these are irrelevant to the argument.
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If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so it will be found on average half way between the two atoms. To get a bond like this, A and B would usually have to be the same atom. You will find this sort of bond in, for example, H2 or Cl2 molecules.
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Note: It's important to realise that this is an average picture. The electrons are actually in a molecular orbital, and are moving around all the time within that orbital.
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This sort of bond could be thought of as being a "pure" covalent bond - where the electrons are shared evenly between the two atoms.
What happens if B is slightly more electronegative than A?
B will attract the electron pair rather more than A does.
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That means that the B end of the bond has more than its fair share of electron density and so becomes slightly negative. At the same time, the A end (rather short of electrons) becomes slightly positive. In the diagram, "[pic]" (read as "delta") means "slightly" - so [pic]+ means "slightly positive".
Defining polar bonds
This is described as a polar bond. A polar bond is a covalent bond in which there is a separation of charge between one end and the other - in other words in which one end is slightly positive and the other slightly negative. Examples include most covalent bonds. The hydrogen-chlorine bond in HCl or the hydrogen-oxygen bonds in water are typical.
What happens if B is a lot more electronegative than A?
In this case, the electron pair is dragged right over to B's end of the bond. To all intents and purposes, A has lost control of its electron, and B has complete control over both electrons. Ions have been formed.
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A "spectrum" of bonds
The implication of all this is that there is no clear-cut division between covalent and ionic bonds. In a pure covalent bond, the electrons are held on average exactly half way between the atoms. In a polar bond, the electrons have been dragged slightly towards one end.
How far does this dragging have to go before the bond counts as ionic? There is no real answer to that. You normally think of sodium chloride as being a typically ionic solid, but even here the sodium hasn't completely lost control of its electron. Because of the properties of sodium chloride, however, we tend to count it as if it were purely ionic.
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Note: Don't worry too much about the exact cut-off point between polar covalent bonds and ionic bonds. At A'level, examples will tend to avoid the grey areas - they will be obviously covalent or obviously ionic. You will, however, be expected to realise that those grey areas exist.
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Lithium iodide, on the other hand, would be described as being "ionic with some covalent character". In this case, the pair of electrons hasn't moved entirely over to the iodine end of the bond. Lithium iodide, for example, dissolves in organic solvents like ethanol - not something which ionic substances normally do.
Summary
• No electronegativity difference between two atoms leads to a pure non-polar covalent bond. • A small electronegativity difference leads to a polar covalent bond. • A large electronegativity difference leads to an ionic bond.
Polar bonds and polar molecules
In a simple molecule like HCl, if the bond is polar, so also is the whole molecule. What about more complicated molecules?
In CCl4, each bond is polar.
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Note: Ordinary lines represent bonds in the plane of the screen or paper. Dotted lines represent bonds going away from you into the screen or paper. Wedged lines represent bonds coming out of the screen or paper towards you.
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The molecule as a whole, however, isn't polar - in the sense that it doesn't have an end (or a side) which is slightly negative and one which is slightly positive. The whole of the outside of the molecule is somewhat negative, but there is no overall separation of charge from top to bottom, or from left to right.
By contrast, CHCl3 is polar.
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The hydrogen at the top of the molecule is less electronegative than carbon and so is slightly positive. This means that the molecule now has a slightly positive "top" and a slightly negative "bottom", and so is overall a polar molecule.
A polar molecule will need to be "lop-sided" in some way.
Patterns of electronegativity in the Periodic Table
The most electronegative element is fluorine. If you remember that fact, everything becomes easy, because electronegativity must always increase towards fluorine in the Periodic Table.
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Note: This simplification ignores the noble gases. Historically this is because they were believed not to form bonds - and if they don't form bonds, they can't have an electronegativity value. Even now that we know that some of them do form bonds, data sources still don't quote electronegativity values for them.
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Trends in electronegativity across a period
As you go across a period the electronegativity increases. The chart shows electronegativities from sodium to chlorine - you have to ignore argon. It doesn't have an electronegativity, because it doesn't form bonds.
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Trends in electronegativity down a group
As you go down a group, electronegativity decreases. (If it increases up to fluorine, it must decrease as you go down.) The chart shows the patterns of electronegativity in Groups 1 and 7.
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Explaining the patterns in electronegativity
The attraction that a bonding pair of electrons feels for a particular nucleus depends on: • the number of protons in the nucleus; • the distance from the nucleus; • the amount of screening by inner electrons.
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Note: If you aren't happy about the concept of screening or shielding, it would pay you to read the page on ionisation energies before you go on. The factors influencing ionisation energies are just the same as those influencing electronegativities.
Use the BACK button on your browser to return to this page.
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Why does electronegativity increase across a period?
Consider sodium at the beginning of period 3 and chlorine at the end (ignoring the noble gas, argon). Think of sodium chloride as if it were covalently bonded.
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Both sodium and chlorine have their bonding electrons in the 3-level. The electron pair is screened from both nuclei by the 1s, 2s and 2p electrons, but the chlorine nucleus has 6 more protons in it. It is no wonder the electron pair gets dragged so far towards the chlorine that ions are formed.
Electronegativity increases across a period because the number of charges on the nucleus increases. That attracts the bonding pair of electrons more strongly.
Why does electronegativity fall as you go down a group?
Think of hydrogen fluoride and hydrogen chloride.
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The bonding pair is shielded from the fluorine's nucleus only by the 1s2 electrons. In the chlorine case it is shielded by all the 1s22s22p6 electrons.
In each case there is a net pull from the centre of the fluorine or chlorine of +7. But fluorine has the bonding pair in the 2-level rather than the 3-level as it is in chlorine. If it is closer to the nucleus, the attraction is greater.
As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus.
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Warning! As far as I am aware, none of the UK-based A level (or equivalent) syllabuses any longer want the next bit. It used to be on the AQA syllabus, but has been removed from their new syllabus. At the time of writing, it does, however, still appear on at least one overseas A level syllabus (Malta, but there may be others that I'm not aware of). If in doubt, check your syllabus.
Otherwise, ignore the rest of this page. It is an alternative (and, to my mind, more awkward) way of looking at the formation of a polar bond. Reading it unnecessarily just risks confusing you.
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The polarising ability of positive ions
What do we mean by "polarising ability"?
In the discussion so far, we've looked at the formation of polar bonds from the point of view of the distortions which occur in a covalent bond if one atom is more electronegative than the other. But you can also look at the formation of polar covalent bonds by imagining that you start from ions.
Solid aluminium chloride is covalent. Imagine instead that it was ionic. It would contain Al3+ and Cl- ions.
The aluminium ion is very small and is packed with three positive charges - the "charge density" is therefore very high. That will have a considerable effect on any nearby electrons.
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We say that the aluminium ions polarise the chloride ions.
In the case of aluminium chloride, the electron pairs are dragged back towards the aluminium to such an extent that the bonds become covalent. But because the chlorine is more electronegative than aluminium, the electron pairs won't be pulled half way between the two atoms, and so the bond formed will be polar.
Factors affecting polarising ability
Positive ions can have the effect of polarising (electrically distorting) nearby negative ions. The polarising ability depends on the charge density in the positive ion.
Polarising ability increases as the positive ion gets smaller and the number of charges gets larger.
As a negative ion gets bigger, it becomes easier to polarise. For example, in an iodide ion, I-, the outer electrons are in the 5-level - relatively distant from the nucleus.
A positive ion would be more effective in attracting a pair of electrons from an iodide ion than the corresponding electrons in, say, a fluoride ion where they are much closer to the nucleus.
Aluminium iodide is covalent because the electron pair is easily dragged away from the iodide ion. On the other hand, aluminium fluoride is ionic because the aluminium ion can't polarise the small fluoride ion sufficiently to form a covalent bond.
ELECTRONEGATIVITY
This page deals with electronegativity in an organic chemistry context. If you want a wider view of electronegativity, there is a link at the bottom of the page.
What is electronegativity?
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is given a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.
What happens if two atoms of equal electronegativity bond together?
The most obvious example of this is the bond between two carbon atoms. Both atoms will attract the bonding pair to exactly the same extent. That means that on average the electron pair will be found half way between the two nuclei, and you could draw a picture of the bond like this:
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It is important to realise that this is an average picture. The electrons are actually in a sigma orbital, and are moving constantly within that orbital.
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Help! A sigma orbital is a molecular orbital formed by end-to-end overlap between two atomic orbitals. If you aren't happy about this, read the articles on orbitals and the bonding in methane and ethane.
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The carbon-fluorine bond
Fluorine is much more electronegative than carbon. The actual values on the Pauling scale are
|carbon |2.5 |
|fluorine |4.0 |

That means that fluorine attracts the bonding pair much more strongly than carbon does. The bond - on average - will look like this:
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Why is fluorine more electronegative than carbon?
A simple dots-and-crosses diagram of a C-F bond is perfectly adequate to explain it.
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The bonding pair is in the second energy level of both carbon and fluorine, so in the absence of any other effect, the distance of the pair from both nuclei would be the same.
The electron pair is shielded from the full force of both nuclei by the 1s electrons - again there is nothing to pull it closer to one atom than the other.
BUT, the fluorine nucleus has 9 protons whereas the carbon nucleus has only 6.
Allowing for the shielding effect of the 1s electrons, the bonding pair feels a net pull of about 4+ from the carbon, but about 7+ from the fluorine. It is this extra nuclear charge which pulls the bonding pair (on average) closer to the fluorine than the carbon.
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Help! You have to imagine what the bonding pair "sees" if it looks in towards the nucleus. In the carbon case, it sees 6 positive protons, and 2 negative electrons. That means that there will be a net pull from the carbon of about 4+. The shielding wouldn't actually be quite as high as 2-, because the 1s electrons spend some of their time on the far side of the carbon nucleus - and so aren't always between the bonding pair and the nucleus.
Incidentally, thinking about electrons looking towards the nucleus may be helpful in picturing what is going on, but avoid using terms like this in exams.
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The carbon-chlorine bond
The electronegativities are:
|carbon |2.5 |
|chlorine |3.0 |

The bonding pair of electrons will be dragged towards the chlorine but not as much as in the fluorine case. Chlorine isn't as electronegative as fluorine.
Why isn't chlorine as electronegative as fluorine?
Chlorine is a bigger atom than fluorine. fluorine: 1s22s22px22py22pz1 chlorine: 1s22s22px22py22pz23s23px23py23pz1
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Help! If you aren't happy about this, read the article on orbitals. Use the BACK button on your browser to get back to here again.
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In the chlorine case, the bonding pair will be shielded by all the 1-level and 2-level electrons. The 17 protons on the nucleus will be shielded by a total of 10 electrons, giving a net pull from the chlorine of about 7+.
That is the same as the pull from the fluorine, but with chlorine the bonding pair starts off further away from the nucleus because it is in the 3-level. Since it is further away, it feels the pull from the nucleus less strongly.
Bond polarity and inductive effects
Polar bonds
Think about the carbon-fluorine bond again. Because the bonding pair is pulled towards the fluorine end of the bond, that end is left rather more negative than it would otherwise be. The carbon end is left rather short of electrons and so becomes slightly positive.
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The symbols [pic]+ and [pic]- mean "slightly positive" and "slightly negative". You read [pic]+ as "delta plus" or "delta positive".
We describe a bond having one end slightly positive and the other end slightly negative as being polar.
Inductive effects
An atom like fluorine which can pull the bonding pair away from the atom it is attached to is said to have a negative inductive effect.
Most atoms that you will come across have a negative inductive effect when they are attached to a carbon atom, because they are mostly more electronegative than carbon.
You will come across some groups of atoms which have a slight positive inductive effect - they "push" electrons towards the carbon they are attached to, making it slightly negative.
Inductive effects are sometimes given symbols: -I (a negative inductive effect) and +I (a positive inductive effect).
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Note: You should be aware of terms like "negative inductive effect", but don't get bogged down in them. Provided that you understand what happens when electronegative atoms like fluorine or chlorine are attached to carbon atoms in terms of the polarity of the bonds, that's really all you need for most purposes.
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Some important examples of polar bonds
Hydrogen bromide (and other hydrogen halides)
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Bromine (and the other halogens) are all more electronegative than hydrogen, and so all the hydrogen halides have polar bonds with the hydrogen end slightly positive and the halogen end slightly negative.
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Help! Halogen: a member of group VII of the Periodic Table - fluorine, chlorine, bromine and iodine.
Halide: a compound of one of these - e.g. hydrogen chloride, hydrogen bromide, etc.
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The polarity of these molecules is important in their reactions with alkenes.
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Note: These reactions are explored in the section dealing with the addition of hydrogen halides to alkenes.
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The carbon-bromine bond in halogenoalkanes
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Note: You may come across halogenoalkanes under the names "haloalkanes" or "alkyl halides".
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Bromine is more electronegative than carbon and so the bond is polarised in the way that we have already described with C-F and C-Cl.
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The polarity of the carbon-halogen bonds is important in the reactions of the halogenoalkanes.
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Note: This link will take you to the nucleophilic substitution reactions of the halogenoalkanes in which this polarity is important.
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The carbon-oxygen double bond
An orbital model of the C=O bond in methanal, HCHO, looks like this:
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Note: If you aren't sure about this, read the article on bonding in the carbonyl group (C=O).
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The very electronegative oxygen atom pulls both bonding pairs towards itself - in the sigma bond and the pi bond. That leaves the oxygen fairly negative and the carbon fairly positive.
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