hem m F ac acts tshe heet et C he Number 71
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Revision Summary: Summary: Electronegativity Electronegativity, Ionisation Energies & Electron Affinities A chemist named Pauling assigned electronegativities to elements on an arbitrary scale from 0 to 4. The following is a list of some such values:
Before reading through this Factsheet you should: • Have a good understanding of atomic structure; • Have a good appreciation of how the Periodic Table is arranged.
Fig 2. Pauling electronegativities electronegativities After working through this Factsheet you will be able to: • Define electronegativity, ionisation energies and electron affinities; • Explain electronegativity, ionisation energies and electron affinities in terms of electron movement; • Recall trends in electronegativity, ionisation energies and electron affinities within the Periodic table. This Factsheet is written as a revision aid, the main objective being to ensure that there are no misconceptions or confusion about e lectronegativity, ionisation energies and electron affinities.
Value
0.9
1 .0
1.2
1.5
1.5
1 .8
2 .0
2.1
Element
Na
Li
Mg
Al
Be
Si
B
H
Value
2.1
2 .5
2.5
3.0
3.0
3 .5
4 .0
Element
P
C
S
Cl
N
O
F
This list simply helps chemists to quantify the power an atom has to attract bond pair electrons.
Electronegativity
Electronegativity can be used to classify the bond type between particles:
Definition: The electronegativity of an atom is the ability of its nucleus to attract electrons in a bond pair.
Ionic bonds: Occur between two particles which differ greatly in electronegativity - commonly bonds between metals and non-metals. The bond pair of electrons are drawn close to the more electronegative particle forming a negative anion, and also a positive cation.
A very electronegative nucleus has the ability to strongly attract bonding electrons, whilst a nucleus which is not very electronegative (is electropositive) does not attract bonding electrons strongly.
Covalent bonds: bonds : Occur between two particles which do not differ significantly in electronegativity - commonly bonds between two nonmetals. The bond pair electrons are shared between the two neutral atoms.
Trends in the electronegativity in the Periodic Table are shown clearly in this diagram, and need to be remembered:
increasing along each period
Pauling estimated the percentage of ionic character of a bond in a binary molecule from the difference in electronegativity of the two atoms in a bond.
H
Fig 3. Estimated % ionic character
Fig 1. Electronegativity trends
Li Be B C N O F Na S Cl K
Br
Rb
I
increasing up each group
Electronegativity increases: Electronegativity • Across each period. • Up each group.
Difference in EN
0.2
0.4
0 .6
0.8
1.0
1.2
1 .4
1.6
% ionic character
1.0
4.0
9.0
15
22
30
39
47
Difference in EN
1.8
2.0
2 .2
2.4
2.6
2.8
3 .0
3.2
% ionic character
55
63
70
76
82
86
89
92
Note that the greater the difference in electronegativity between the two particles, the more ionic the bond between them. Ionic, covalent and intermediate bonding are discussed in detail in Factsheet 5. Examples:
This makes fluorine the most electronegative atom. Electronegativity increases across a period because the effective nuclear charge of the atoms increases, hence increasing the electron-attracting power. Electronegativity decreases down each group because the number of quantum shells of electrons increases, increasing the shielding effect around the nucleus, hence decreasing the e lectron-attracting power.
Carbon has an EN of 2.5 and hydrogen has an EN of 2.1. So the C - H bond has a very small percentage ionic character - 4% only.
•
Aluminium has an EN of 1.5 and chlorine has an EN of 3.0. So we would expect aluminium chloride to b e just over 40% ionic
NB: You do no t hav e to reme mbe r act ual ele ctron ega tiv ity fig ures or the percen tages in the above table !
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71. Revision Summary: Electronegativity, Ionisation Energies & Electron Affinities Affinit ies
Fig 4. EN difference and bonding
Chem Factsheet
Whilst there is a general increase in IE across a period, this is not a smooth trend – as shown in fig 5.
electron distribution
EN Difference
Fig 5. First ionisation energies for the elements of period 2 pure covalent
0
polar bondsome ionic character
δ+
S+
ionic, but polarised, so some covalent character
+
pure ionic spherical ions
+
S-
δ-
_
Ne
y g r e n e n o i t a s i n o i t s r i f
small
filling of 2p sub-shell
F N
filling of 2s sub-shell
C
pairing of 2p electrons O
Be
2p sub-shell half-fills
B Li
large
atomic number
_
very large
As the graph shows, the first ionisation energy decreases from Be to B and from N to O. A similar phenomenon occurs between groups 2 to 3 and 5 to 6 in the other periods. This can be explained by the existence of sub shells:
Ionisation Energy If an atom is supplied with enough energy, energy, it will lose an electron. Additional supplies of energy may cause the loss of a second electron, then a third, and so on.
•
group 2 elements have full subshell stability, stability, hence are more difficult to ionise than group 3 elements.
•
group 5 elements have half subshell stability, stability, hence are more difficult to ionise than group 6 elements.
If a neutral atom loses an electron, it becomes a positive ion (a cation).
Electronic Structure from Ionisation Energies Examining successive ionisation energies for an element gives us an insight into the electronic structure of that element.
Definition First ionisation energy – the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms.
Fig 6. Successive ionisation energies for sodium
X(g) → X +(g) + e−
1000000
Second ionisation energy – the energy required to remove 1 mole of electrons from 1 mole of gaseous 1+ charged cations.
100000
X +(g) → X 2+(g) + e−
1 -
l o 10000 m J k / 1000 y g r e n e 100 n o i t a s i n 10 o i
Trends in ionisation energies in the Periodic Table • •
Ionisation energies decrease down down a group. Ionisation energies generally increase across a period.
Going down a group group,, we are adding more electron shells. So in each successive period, the outer electrons are:
• •
More shielded from the nucleus by the inner electrons Accordingly, ionisation energies decrease down a group.
this electron is in the n=3 shell, furthest from the nucleus. This electron experiences comparatively weak attractive forces from the nucleus.
The pattern shows us that sodium has 11 electrons arranged in three quantum shells, or energy levels. The first electron is relatively easy to remove as:
Moving across a period the nuclear charge increases, but we are adding electrons in the same shell, so they add little to the shielding. This means that the outer shell electrons are attracted increasingly strongly to the nucleus – thus more energy is required to remove them. So generally we would expect ionisation energy to increase across the period.
• •
these 8 electrons are in the n=2 shell, between the n=1 and n=3 shells. These electrons experience strong attractive forces from the nucleus
0
Further away from the nucleus
Exam Hint: - Many
these 2 electrons are in the n=1 shell, closest to the nucleus. These experience very strong attractive forces from the nucleus.
• •
It exists further from the nucleus The electrons orbiting closer to the nucleus shield the outer electron from the positive centre so attractive forces are comparatively weak.
The next 8 electrons have similar ionisation energies (as they are all a similar distance from the nucleus) but do get successively more difficult to remove as the relative positive charge within the ion is increasing.
candidates lose marks by:
when explaining the group trends, not referring to shielding, when explaining trends across the period, just referring to "increased nuclear charge" without explaining why it is not counterbalanced by increasing numbers of electrons
The last 2 electrons are very difficult to remove as they exist very close to the unshielded nucleus.
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71. Revision Summary: Electronegativity, Ionisation Energies & Electron Affinities Affinit ies
Electron Affinities
Chem Factsheet
Practice Questions
It is possible to add electrons to an atom, forming a negative anion. The energy change associated with this is the electron affinity.
1. Exp Explai lain n the term ‘elec ‘electro tronega negativ tivity ity’. ’. 2. (a) Use the the data data below below to plot plot a graph of ‘electr ‘electronegat onegativity ivity diffe difference’ rence’ (y-axis) against ‘% ionic character’ (x-axis).
Definition First electron affinity – the energy change when 1 mole of electrons is added to 1 mole of gaseous atoms.
Difference in EN
0.2
0.4
0 .6
0.8
1.0
1.2
1 .4
1.6
% ionic character
1.0
4.0
9.0
15
22
30
39
47
o f electrons Second electron affinity – the energy required to add 1 mole of to 1 mole of gaseous 1- charged anions.
Difference in EN
1.8
2.0
2 .2
2.4
2.6
2.8
3 .0
3.2
O− (g) + e− → O2− (g)
% ionic character
55
63
70
76
82
86
89
92
−
O(g) + e
−
→ O
(g)
ê ∆ H ê
−
= -141 kJmol
1
EA1
∆ H ê ê
= +798 kJmol− 1
EA2
(b) Use the data below to calculate the electronegati electronegativity vity difference difference between the following bonded elements: The first EA is always exothermic (energy is released) because the electron goes into a vacancy in the outer energy level. This is ‘bond-making’ so so energy is released. However, this creates a 1- charged anion, so to add a second electron requires energy to overcome the repulsion (-ve to –ve ) between the electron and the anion. Consequently the second EA is always endothermic (energy is absorbed).
Value
0.9
1 .0
1.2
1.5
1.5
1 .8
2 .0
2.1
Element
Na
Li
Mg
Al
Be
Si
B
H
Value
2.1
2 .5
2.5
3.0
3.0
3 .5
4 .0
Element
P
C
S
Cl
N
O
F
Trends in Electron Affinities in the Periodic Table Bond Electron Elect ron Affin Affinity ity • Increases across a period (to group 7) • Decreases down a group.
Electron affinity follows the same trends as electronegativity for similar reasons.
EN difference
Bond
LiF
HO
NO
BeCl
CO
N aF
PO
SCl
EN difference
(c) Use the graph to estimate estimate the ionic character character,, and then place in order of increasing ionic character, the above bonds.
As we move across a period there is increased effective nuclear charge, so an additional electron would be more strongly bound to the nucleus and more stable, releasing more energy on the formation of this bond. Note that this is not a uniform increase - full sub-shell stability and half sub-shell stability result in a lower than expected first electron affinity in groups 2 and 5.
3. Define the fourth fourth ionisa ionisation tion energy energy of an element element.. 4. Use the idea idea of ionisatio ionisation n energies energies to explain explain why group group 2 metals metals are less reactive than group 1 metals.
As we move down a group, the available site for an additional electron would be further away from the positive nucleus, in a position where the re is greater electron shielding from the nucleus. Hence there is a weaker force of attraction between the additional electron and the nucleus, so EA is less.
5. Define the fourth fourth electr electron on affinity affinity of an element element.. 6. State whether whether the fourth electro electron n affinity affinity of an element element is likely to to be exothermic or endothermic.
. y g r e n e e l l i w n o i s l u p e r s i h t g n i m o c r e v o r i u q e r d n a , n o i n a - 3 d e g r a h c y l e v i t a g e n a o t n o r t c e l e d e g r a h c y l e v i t a g e n a g n i d d a s e v l o v n i t i s a , c i m r e h t o d n e e b o t y l e k i l s i y t i n i f f a n o r t c e l e h t r u o f e h T . 6 → e ) g ( 4 O + ) g ( - 3 O −
−
. s n o i n a d e g r a h c - 3 s u o e s a g f o e l o m 1 o t s n o r t c e l e f o e l o m 1 d d a o t d e r i u q e r y g r e n e e h t – y t i n i f f a n o r t c e l e h t r u o F . 5 . s l a t e m 1 p u o r g e h t f o s n o i t c a e r r o f r e w o l e r a s e i g r e n e n o i t a v i t c a e c n e H . s n o i t a c + 1 e l b a t s m r o f 1 p u o r g s a e r e h w ) d e r i u q e r s e i g r e n e n o i t a s n o i 2 e c n e h ( s n o r t c e l e 2 f o n o i t a n o d e h t g n i v l o v n i , s n o i t a c + 2 e l b a t s m r o f s l a t e m 2 p u o r g o s l A . 1 p u o r g n a h t s n o r t c e l e e v o m e r o t s l a t e m 2 p u o r g r o f d e r i u q e r e b d l u o w y g r e n e e r o m g n i n a e m ) 2 p u o r g o t 1 p u o r g m o r f ( d o i r e p a s s o r c a s e s a e r c n i y g r e n e n o i t a s i n o I . ) s t n e g a g n i c u d e r ( s n o r t c e l e e t a n o d e h t t c a e r s l a t e m n e h W . 4 e + ) g ( + 4 X → ) g ( + 3 X
−
. s n o i t a c d e g r a h c + 3 s u o e s a g f o e l o m 1 m o r f s n o r t c e l e f o e l o m 1 e v o m e r o t d e r i u q e r y g r e n e e h t – y g r e n e n o i t a s i n o i h t r u o F . 3 . l C a N F , i L , l C e B , O P , O H , O C , l C S , O N : r e d r O ) c ( 5 . 0 l C S ; 4 . 1 O P ; 1 . 3 F a N ; 0 . 1 O C ; 5 . 1 l C e B ; 5 . 0 O N ; 4 . 1 O H ; 0 . 3 F i L ) b ( . 2 . r i a p d n o b a n i s n o r t c e l e t c a r t t a o t s u e l c u n s t i f o y t i l i b a e h t s i m o t a n a f o y t i v i t a g e n o r t c e l e e h T . 1
s r e w s n A Acknowledgements: This Factsheet was researched and written by Kieron Heath. Curriculum Press, Bank House, 105 King Street, Wellington, Shropshire, TF1 Acknowledgements: 1NU. ChemistryFactsheets may be copied free of charge by teaching staff or students, provided that their school is a registered subscriber. No part of these Factsheets may be reproduced, stored in a retrieval system, or transmitted, in any other form or by any other means, without the prior permission of the pub lish er. ISSN 1351- 5136
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