More on covalent bonding - single, double and triple bond, their length and strength, dative covalent bonds and bond enthalpy trends
Extra Notes on Covalent Bonding and Covalent Compounds
Doc Brown's A level Chemistry Revision
Extra Notes on chemical bonding for advanced A level chemistry students
All the advanced A level 'basics' with lots of examples of dot and cross diagrams of ionic bonding, Lewis diagrams, properties of covalent compounds etc. is on a separate page.
However, advanced electron notation using s, p and d notation was NOT included, only a brief mention of intermolecular forces/intermolecular bonding. Dative covalent bonds and shapes of molecules were not included. These and other covalent molecules are covered on this page.
There are other pages on SHAPES of MOLECULES
The formation of a covalent bond
In covalent bonds there is a balance between the repulsive forces between the positive nuclei and the attractive forces between the nuclei and the negative electrons between them. The covalent bond is mutual attraction between the nuclei of two atoms and the electrons in between them (+) -- (+). The bond length is determined by the two atoms adopting the position of minimum potential energy.
One or more atomic orbitals from each atom overlap so the bonding pairs of electrons are shared between the nuclei.
The 'dot and cross' (Lewis)
electron diagrams only tell part of the story
Dative covalent bond (co-ordinate bond)
A dative covalent is formed when the pair of electrons forming the bond are donated by one atom only. This contrasts with the usual covalent bond by each atom of the bond contributing one electron.
Examples of dative bond formation
1. Formation of the oxonium ion H3O+ (also known as hydroxonium ion, hydronium ion)
H2O(l) + H+(aq) [H3O]+, a pair of electrons (a lone pair) from the oxygen atom of the water is donated to a proton to form an oxygen-hydrogen dative (co-ordinate) covalent bond in the oxonium ion.
This reaction happens whenever you dissolve a soluble acidic substance in water, but the proton can also come from another water molecule in a self-ionisation process 2H2O(l) H3O+ + OH-(aq).
H2O: + H+ [H2OH]+ where the arrow indicates and 'accentuates' the dative (co-ordinate) covalent bond between the oxygen and the hydrogen. BUT, note that all 3 O-H bonds in the oxonium ion are identical.
2. Formation of the ammonium ion NH4+
NH3(aq) + H+(aq) NH4+(aq), the lone pair of electrons on the nitrogen atom is donated to the proton to form a nitrogen-hydrogen dative (co-ordinate) covalent bond in the ammonium ion.
This reaction happens when you dissolve ammonia gas in water (the proton comes from the water) or when you react aqueous ammonia solution with any acid.
H3N: + H+ [H3NH]+ where the arrow indicates and 'accentuates' the dative covalent bond between the nitrogen and the hydrogen. BUT, note that all 4 N-H bonds in the ammonium ion are identical.
3. Transition metal complexes - dative covalent bonds with ligands
The ligands surrounding the central ion of a complex ion donate pairs of electrons to form the ligand-metal ion bond
octahedral complexes with 6 dative (co-ordinate) bonds
tetrahedral complexes with 4 dative (co-ordinate) bonds
Single and multiple covalent bonds - representations
(a) A molecule with all single covalent bonds (known as a σ bond, sigma bond, C-H and C-C in this case)
(b) Double covalent bond = (σ bonds C-H, and a delocalised pi bond, the C=C bond is a σ bond plus a π bond)
(c) The C=C bond is a σ bond plus π bonding)
(d) Double covalent bond = (σ bond and delocalised π bond)
(e) Triple covalent bond ≡ (σ bond and a double π bond)
Alkynes are unsaturated hydrocarbons with a CC carbon-carbon triple bond
Examples: C2H2, ethyne
All the C-H bonds are single σ bonds.
(f) The nitrogen molecule also has a triple bond :NN:
triple bond, all the rest have all single σ bonds C-H, C-C, C-Cl, C-O and O-H.
Relating single, double and triple bonds to average bond enthalpies and bond length
The average bond enthalpy is the 'typical' energy required to break 1 mole of a covalent chemical bond (but only involving gaseous species). Bond enthalpy is a measure of the bond strength.
Bond length is defined as the distance between the two nuclei of the two atoms bonded together.
You find general patterns of decreasing bond length with increasing bond enthalpy - shorter tends to be stronger because the bonding electrons between the nuclei are closer to the nuclei and consequently more strongly attracted.
Some examples and several important patterns to spot:
Some Group VII (Group 7/17) Halogens trends in bond lengths and bond enthalpies
Some general observations, most of which relate to smaller radii giving shorter stronger bonds:
Halogen molecules X2: From fluorine to iodine the bond length increases and, except for fluorine, the bond enthalpy decreases as the radius of the halogen atom increases with increasing number of filled inner electron shells. Fluorine is distinctly anomalous with a much lower than expected bond dissociation energy, though the bond length fits the general trend. This is explained by the close proximity of the small fluorine atoms causing repulsion between them due to the closeness of the outer electron orbitals.
Hydrogen halides HX: From hydrogen fluoride HF(g) to hydrogen iodide HI(g), there is clear trend in increasing bond length and decreasing bond enthalpy. One result is the increasing ease of aqueous ionisation from hydrofluoric acid to hydriodic acid so that the HX(aq) acids become stronger down the group. In fact, hydrofluoric acid HF(aq) is a relatively weak acid but hydrochloric, hydrobromic and hydriodic acids are all very strong. The latter three are so strong in aqueous media you don't really see the difference e.g. from pH readings, but in non–aqueous media the differences can be clearly measured.
Halogenoalkanes R3C–X: Based on polarisation of the bond (Cδ+–Xδ–), you might expect the reactivity order with respect to nucleophiles (electron pair donors) attacking the δ+ carbon bond to be R–F > R–Cl > R–Br > R–I as the electronegativity difference decreases from C–F to C–I. However, it is the decreasing bond enthalpies from C–F to C–I that override this polarisation trend giving the reactivity trend R–I > R–Br > R–Cl > R–F.
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