|IB syllabus > bonding (sl) > 4.2|
4.2 - The covalent bond
4.2.1: Describe the covalent bond as the electrostatic attraction between a pair of electrons and the positively charged nuclei. Single and multiple bonds should be considered. Examples should include O2, N2, CO2, HCN, C2H4 (ethene) and C2H2 (ethyne).
Covalent bonding - overview
4.2.2: Describe how the covalent bond is formed as a result of electron sharing. Dative covalent bonds are required. Examples include CO, NH4+ and H3O+.
4.2.3: Deduce the Lewis electron dot) structure of molecules and ions for up to four electron pairs on each atom. A pair of electrons can be represented by dots and crosses or by a line. For example chlorine can be shown as:
The octet rule
On studying many molecular compounds it appears that a full set of electrons in the outer shell is a particularly stable situation and this gives rise to the octet rule, which states that atoms tend to share electrons so as to obtain eight electrons in their outer shells
To form a single covalent bond one electron is provided by each of the bonding atoms making a shared pair. The simplest situation occurs when two atoms of a non-metal share one pair of electrons.
Here, we can see that although originally each chlorine atom has only seven electrons in its outer shell, after bonding it has eight (an octet), making the molecule more stable than the sum of the original atoms.
It is possible for atoms to share two pairs of electrons to achieve the octet (8 electrons).
When two pairs of electrons are shared between two atoms this is called a double bond. In the diagram above the carbon dioxide molecule has double bonds between the carbon and each of the oxygen atoms. Triple bonds are also possible (three shared pairs of electrons)
4.2.4: State and explain the relationship between the number of bonds, bond length and bond strength. The comparison should include bond lengths and bond strengths of: two carbon atoms joined by single, double and triple bonds, the carbon atom and the two oxygen atoms in the carboxyl group of a carboxylic acid.
A single bond is a shared pair of electrons that is attracted to both
nuclei of the bonded atoms.
The single bonds hold atoms together by the forces of attraction between the electron pair (bonding pair) and the two nuclei. As the nuclei of different atoms are obviously different from one another then this force of attraction and hence the bond strength varies between different pair of atoms.
The strength of bonds can be measured by several techniques, usually by seeing how much energy is needed to break the bond.
Some bond energy values
Although it is difficult to make a direct relationship between bond length
and strength there are some inferences that can be obtained. It can be
seen that as the bond length of the carbon - halogen bonds increases so
the bond energy decreases. As the atoms get larger they are held further
apart by inter-electron repulsions. The attractive force between the bonding
electron pair and the nuclei is consequently weaker.
As the bond strength INcreases, so the bond length DEcreases. This follows from a consideration of the force of attraction between the greater number of pairs of electrons and the two nuclei. Four electrons (two pairs) can pull the two nuclei closer together than two electrons (one pair).
The carboxlate group
4.2.5: Predict whether a compound of two elements would be covalent from the relative electronegativity values or from their positions in the periodic table
Prediction of bonding type
Compounds in which the bonded atoms have a large electronegativity difference for ionic compounds. Where the difference is slight, they are covalent. There is no hard and fast value at which the change occurs. Rather there is a greater and greater degree of covalency as the values become closer together.
Perhaps the closest to a 'cut off' is the compound formed between aluminium and chlorine. In the solid state at 0ºC there is considerable evidence that it is ionic, but at room temperature it seems to be covalent. At higher temperatures it sublimes as a dimer with the formula Al2Cl6.
Aluminium has an electronegativity of 1.5 and chlorine 3.0. That makes the difference in electronegativity = 1.5 units on the Pauling scale. This is a good value to use as a 'rule of thumb'.
Greater than 1.5 units = ionic
Less than 1.5 units = covalent.
Ionic or covalent bond formation
It was stated above that metals reacting with non-metals form ionic bonds. This is the GENERAL rule. It is really the ease with which a metal can lose electrons coupled with the attraction that a non-metal has for more electrons.
When an element loses electrons easily and another element has a high attraction for electrons then this encourages ionic bond formation. As metals lose electrons easily they are said to be electropositive, or to put it another way, to have low electronegativity values.
Non-metals tend to attract electrons and the non-metals closest to group 7 have higher electronegativity values. Electronegativity increases from bottom left to top right in the periodic table.
When the DIFFERENCE between the electronegativity values is large, ionic bonds are formed.
It is usually the case that the difference in electronegativity values between metals and non-metals is large enough to cause ionic bond formation (transfer of electrons etc etc), however this is not the case for all metal:non-metal combinations.
4.2.6: Predict the relative polarity of bonds based on electronegativity values. AIM 7: Simulations may be used here
Comparison of electronegativity difference
From the previous section you can see that a direct comparison of bonds may be done using the electronegativity difference between the atoms in the bonds.
It is safe to say that the C-F bond (electronegativity difference = 1.5) has a greater dipole than the C-Br bond (electronegativity difference = 0.3)
4.2.7: Predict the shape and bond angles for species wth four, three, and two negative charge centres on the central atom using the valence shell electron pair repulsion theory VSEPR. Example should include CH4, NH3, H2O, NH4+, H3O+, BF3, C2H4, SO2, C2H2, and CO2. Simulations are available to study the three dimensional structure of these and the structures in 4.2.9 and 4.2.10
The central oxygen atom has its electron arranged into four pairs in four distinct regions (the orbitals are sp3 hybridised)
These repel one another and adopt a tetrahedral arrangement.
However only two of the electron pairs are used in bonding and the other two pair are 'lone' (ie cannot be seen). The shape of the molecule is therefore 'angular' or 'bent'
Tetrahedral electronically but the molecular shape is angular
The valence shell electron pair repulsion theory now looks at the relative strength of the repulsions between the lone pairs and the bonding pairs. As the lone pairs are not drawn further away from the central atom by another atomic nucleus then they exert a greater repulsion on each other than the bonding pairs do on each other. Intermediate is the repulsion felt between a bonding pair and a one pair.
Order of repulsion strength:
lone pair- lone pair >> lone pair - bonding pair >> bonding pair - bonding pair
This causes the tetrahedral electronic shape to distort and squeezes the bonding pairs together. The bond angle then closes sightly from 109,5º to 104,5º
(click to see in 3D - if it doesn't work you have to download chime)
4.2.8: Predict whether a molecule is polar from its molecular shape and bond polarities.
This is caused by a difference in electronegativity between two bonded atoms. Most bonds are polar, but in reality only those with an electronegativity difference of at least 1 unit on the Pauling scale shows the effect.
For example carbon has an electronegativity of 2.5 and hydrogen 2.1. In principle they are polarised and the bond has a dipole, but the two values are close enough together as to be insignificant. Carbon - hydrogen bonds are not said to be polar.
Resolution of dipoles
Dipoles are vector quantities in that they have both magnitude and direction. They can be resolved in the same way as forces.
If the resolution of all vector dipoles in a molecue leaves an overall dipole in any specific direction then the molecule is polar.
Water has two dipoles on the oxygen - hydrogen bonds. If the molecule is orientated with the oxygen at the top these cancel out in a horizontal direction but add up in a vertical direction.
Water is consequently a polar molecule.
Example: Carbon dioxide
In this case the carbon - oxygen bonds are also polarised, but they cancel out in the horizontal axis, while there is no component in the vertical axis. Carbon dioxide is therefore non-polar.
Polar molecules attract one another increasing the degree of intermolecular force experienced. This means that they have higher boiling ponts than expected from purely van der Waaal's forces.
4.2.9: Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C60 fullerene)
Allotropes are different physical forms of the same element. All elements are made up uniquely of their own atoms and therefore any physical differences must be a consequence of how the atoms are joined together - their arrangement within the bulk structure.
Many elements exhibit allotropy as there are often varous ways in which the atoms can be linked together into molecules and also different ways in which the molecules can be arranged to make larger structures.
In the case of carbon, the atoms form either giant macromolecular structures (diamond and graphite) in which all of the atoms in the bulk structure are joined together by covalent bonds making giant molecules, or smaller molecules (buckminster fullerene) in which there are only discrete molecules made up of 60 carbons in a structure resembling a football (hence the nickname 'bucky balls')
Each carbon in a diamond crystal is bonded to four other carbon atoms making a giant macromolecular array (lattice). As each carbon has four single bonds it is sp3 hybridised and has tetrahedral bond angles of 109º 28'
Properties of diamond
Physical properties of diamond explained by considering the structure and bonding
Again the carbon atoms are bonded together to make a giant structure but in this case all of the carbons are bonded to only three neighbour and are sp2 hybridised. As the sp2 hybridisation results in planar structures, there are giant 2 dimensional layers of carbon atoms and each layer is only weakly linked to the next layer by Van der Waal's forces.
Physical properties of graphite explained by considering the structure and bonding
These are small molecules of carbon in which the giant structure is closed over into spheres of atoms (bucky balls) or tubes (sometimes caled nano-tubes). The smallest fullerene has 60 carbon atoms arranged in pentagons and hexagons like a football. This is called Buckminsterfullerene.
The name 'buckminster fullerene' comes from the inventor of the geodhesic dome (Richard Buckminster Fuller) which has a similar structure to a fullerene. Fullerenes were first isolated from the soot of chimineys and extracted from solvents as red crystals.
The bonding has delocalised pi molecular orbitals extending throughout the structure and the carbon atoms are a mixture of sp2 and sp3 hybridised systems.
Fullerenes are insoluble in water but soluble in methyl benzene. They are non- conductors as the individual molecules are only held to each other by weak van der Waal's forces.
Physical properties of fullerenes explained by considering the structure and bonding
4.2.10: Describe and structure and bonding in silicon and silicon dioxide
These are giant covalent structures, with the bonding covalent from atom to atom in a never ending array. The bond angles at each silicon atom is 109º 28'.
The oxygen atoms act as bridges between silicon atoms in silicon dioxide.
|Copyright: 2013- IsisSoft Interactive|