4.4 Electronegativity, Bond Polarity, Bond Length and Bond Energy
- Electronegativity is the ability of an atom which is covalently bonded to the other atom to attract the bond pair of electrons towards itself
- The more electronegative an atom is, the higher the tendency of that atom to attract the bond pair of electrons towards itself
- The Pauling scale is commonly used to quantify the value of electronegativity of a particular element
- Fluorine is the most electronegative element because of its small size, followed by oxygen and nitrogen.
Trends of electronegativity values in the Periodic Table
- The value of electronegativity increases across a Period(from left to right).
- This is because the number of proton increases across a Period. Therefore the amount of positive charge in the nucleus also increases
- The shielding effect by inner electrons remains constant
- Therefore the attraction towards the bond pair of electrons increases, making it more electronegative
- The value of electronegativity decreases down a Group.
- This is because the size of the atoms increases down a Group. Therefore the distance between the nucleus and the bond pair of electrons also increases.
- The shielding effect by inner electrons is also greater
- Therefore the attraction towards the bond pair of electrons decreases, making it less electronegative
- When two covalently-bonded atoms have the same electronegativity, the electron cloud is evenly distributed between the two atoms.
- The bond is described as a ‘pure’ covalent bond or non-polar bond.
- Some examples are H2, Cl2 and Br2.
- However, when an atom is more electronegative than the other, the more electronegative atom will attract the bond pair of electrons more towards itself. The electron cloud is not evenly distributed or distorted.
- The more electronegative end acquires a partial negative charge while the less electronegative end acquires a partial positive charge
- The bond is said to be polarised, or, a polar bond
- Some examples of compound which contain polar bond(s) are HCl and CH4.
- The molecule is described as ‘covalent with some ionic character’.
Polar and non-polar molecules
- A molecule is polar, and thus, has a dipole moment(μ ≠ 0) if:
- the bonds are polarised
- the dipole of the bonds do not cancel out each other(in other words, it is asymmetrical)
- The dipole moment, μ is the product of charges and the distance between the centre of the charges
- A liquid containing polar molecules can be deflected by a charged rod brought near to it. This is because there is a positive end and a negative end in polar molecules. So, irregardless of the charge on the rod, one end of the molecule will always be attracted towards the charged rod
- When polar molecules are placed in an electric field, the positive end of the molecule will face towards the negative terminal while the negative end of the molecule will face towards the positive terminal.
- Generally, polar molecules are more reactive than non-polar molecules because many chemical reactions are started by a reagent attacking an electrically- charged end of the polar molecule. An example is CO is more reactive than N2 although both of them have triple bonds because CO is polar while N2 is not
Bond length and bond energy
- Bond length is the distance between two nuclei of two atoms joined by a covalent bond
- Bond energy is the energy needed to break one mole of covalent bonds between two atoms in the gaseous state
A-B(g) → A(g) + B(g) ∆H° = +x J
4) i. The shorter the bond length, the stronger is the bond.
ii. The greater the bond energy, the stronger is the bond
iii. Hence, the shorter and bond length, the greater the bond energy