CHEMICAL BONDS A chemical bond is formed between two atoms if the resulting arrangement of the atoms and their electrons has a lower energy than the separate atoms themselves. Two main ways of creating a chemical bond: Complete electron transfer IONIC BOND Electron sharing COVALENT BOND
IONIC BONDS : formed between atoms with low ionisation energies and those with high ELECTRON AFFINITIES, e.g. Na+Cl-. You will hear a lot more about ionic compounds and their properties in later modules. We are going to concentrate on COVALENT BONDING Characteristic of elements in the middle of the Periodic Table - complete electron transfer is not practicable.
Nature of covalent bonding first established by G.N.Lewis, in 1916, before modern quantum theory had been thought of. Lewis’ idea was that the formation of a covalent bond occurred when one electron from each atom became shared to form an ELECTRON PAIR. For hydrogen we would have: electrons The two atoms are held together by the attraction between the pair of electrons and both nuclei nuclei
Unequal sharing - polar covalent bonds Homonuclear bond when the two atoms forming a covalent bond arethe same - the electrons will be equally shared. Heteronuclear bond if the atoms are different – the sharing will be unequal! Which atom gets more of the electrons? Depends on the electronegativities of the atoms
Electronegativity Basic definition: The ability of an atom in a molecule to attract electron density to itself How to make this idea quantitative? Simplest (Mulliken) Ionisation energy defined earlier. Electron affinity: see next slide:
Definition of electron affinity: Thus - the harder it is to remove an electron (I), and the more energy you get back by attaching an extra electron (Ea) - the higher the electronegativity. In Periodic Table: high electronegativity at top right, low electronegativity at bottom left. Very large electronegativity difference ionic bond (X+Y-) No electronegativity difference ‘pure’ covalent bond (X-X) Small to moderate electronegativity difference polar covalent bond (Xd+-Yd-, where Y is more electronegative than X)
Lewis Structures We will first concentrate on H and first row elements. Basic rules: 1. Every atom tries to achieve an octet of electrons (2 for H). 2. Each pair of shared electrons one bond - symbolised by a dash (-) How to achieve an octet: Note: if an atom loses an electron (giving a positive charge), then it needs to form more bonds to achieve its octet (see extra sheet)
Note : (1) never more than 4 bonds to central atom (max. 8 electrons) (2) keep formal charges to a minimum In polyatomic molecules - distinguish between central and terminalatoms. Difficult to give general rule to specify which is which. Often - atom with lowest electronegativity is central; also chains of same atoms are very rare. Lewis structures tell us nothing about shapes of molecules.
Some simple examples: O needs 2 electrons to achieve octet Each H needs 1 electron to achieve its full shell H2O (remember - each dash is a pair of shared electrons) NH3 N needs 3 electrons to achieve octet Each H needs 1 electron to achieve octet There will also be a lone pair on N
C needs 4 electrons to achieve octet Each H needs 1 electron to achieve its full shell CH4 No lone pairs - all C electrons involved in bonding. Remember - Lewis structures are 2-dimensional - do not give correct structures. C needs 4 electrons to achieve octet Each Cl needs 1 electron to achieve its octet O needs 2 electrons to achieve octet COCl2 Note double bond between C and O Therefore to fill all octets we have:
CH3CO2H (ethanoic acid, acetic acid) The only way to achieve octets for C, O, and full shells for H is NH4+ Also works for ions N+ needs 4 electrons to achieve octet Each H needs 1 electron to achieve its full shell
Mendeleev vodka: old style bottle
There are many more simple examples of organic and inorganic molecules where the drawing of Lewis structures using the above rules is very straightforward. But - there are exceptions 1. Some elements (towards left hand side of the Periodic Table) cannot achieve a full octet. B only has 3 valence electrons - would need 5 more to fill octet. But can only form 3 bonds. Thus BF3 is BF3
2. Elements in second and subsequent rows of the Periodic Table can accommodate >8 electrons (can ‘expand their octet’) Such elements possess d-orbitals of suitable energy for bonding - so >4 bonds can be formed (not restricted to one s + three p orbitals). Because S can expand its octet, it is quite in order to have the Lewis structure: SF6 Expansion of octet very common for compounds of Si, Ge, P, As, Sb, S, Se, Te, Cl, Br etc.
3. It is often not possible to draw a single, unambiguous, Lewis structure for a given molecule or ion. C is the central atom; O can achieve its octet by forming 2 bonds (=O), or one bond + picking up an electron (-O-) CO32- Also - remember that C cannot form >4 bonds One possible Lewis structure would therefore be: But this is not the only possibility!
The structure could equally be or Each of these structures would predict one double bond and two single bonds, i.e. one shorter and two longer bonds. Experiment shows that all C-O bonds in the carbonate are the same length. The correct structure is a blend of all three forms (NOT hopping between them). This phenomenon is called RESONANCE - symbolised by:
The actual structure is called a resonance hybrid, and the pair of electrons comprising the second component of the double bond is said to be delocalised. Another example of resonance: Acetate (ethanoate) ion, CH3CO2- Here there are TWO possible resonance structures:
Both C-O bond lengths are again the same (and intermediate between the lengths of a single and a double bond) Benzene, C6H6 The bonding in benzene was not properly understood for many years. When the possibility of resonance became recognised, the problem was largely solved.
Simple ideas would give the structure BUT the CC bond lengths are all the same (intermediate between single and double), and benzene does not behave chemically like a molecule containing isolated C=C bonds. Hence we must introduce the second possible form, in resonance with the first:
This explains the observed physical and chemical properties of benzene, and it can also be drawn as with the circle symbolising the delocalised electrons. Resonance is also possible with molecules containing atoms where expansion of the octet has occurred.
oxidizing agent rocket oxidizer (chlorate ion) ClO3- and (perchlorate ion) ClO4- Note that in all oxy-anions, the negative charges are formally situated on theoxygen atoms
Space Shuttle Solid booster Al + NH4ClO4
Resonance involves the blending of two or more Lewis structures, where the atoms have the same relative positions, but with different electronic arrangements. Electrons are delocalised, and the energy of the resonance hybrid is lower than for any single component form. All of the examples of resonance shown so far involve two or more equivalentstructures. It is quite possible to have resonance between non-equivalent forms.
Non-equivalent resonance hybrids e.g. thiocyanate ion, (NCS)- Two forms - the one with the negative charge on N will be more important (electronegativity N > S), although both will contribute to the final hybrid. This introduces the idea of preferred Lewis structures, and this will help us to predict relative stabilities of isomers. Note - do not confuse the resonance forms shown so far with isomers, where the spatial arrangements of the atoms differ
Why is carbon dioxide OCO and not COO? Lewis structures: compared to The latter has (i) large formal charges, and (ii) negative charge on less electronegative atom. Another example: The ion (NCO)- is stable; the ion (CNO)- is known, but very unstable. Why is this?
Lewis structures for the two isomers are: and respectively. This again explains the observed properties. Note that in the former, the negative charge will be predominantly on O rather than N, from their relative electronegativities. Summarise requirements for a preferred Lewis structure:
1. Keep formal charges to a minimum. 2. Where formal charges are necessary, they should be consistent with relative electonegativities (-ve charge on more electronegative atom), and 3. Keep like charges as far apart as possible. Two final examples (more in tutorial work):
H2O2 and F2O2 For H2O2 the only realistic Lewis structure is Because F is so electronegative, in the other case we can have: This accounts for the observation that the O-O distance in H2O2 is much longer than in F2O2