Periodic trends in electronegativity

1. Periodic trends in electronegativity Learning intention Learn the definition of electronegativity, and how to explain periodic trends in terms of nuclear charge, covalent radius and the screening effect of inner electrons.

2. Electronegativity Electronegativity is a numerical measure of the relative ability of an atom in a molecule to attract the bonding electrons towards itself.

3. e e C H Electronegativity Electronegativity is a measure of an atom’s attraction for the shared pair of electrons in a bond Which atom would have a greater attraction for the electrons in this bond and why?

4. Linus Pauling Linus Pauling, an American chemist (and winner of two Nobel prizes!) came up with the concept of electronegativity in 1932 to help explain the nature of chemical bonds. Today we still measure electronegativities of elements using the Pauling scale. Since fluorine is the most electronegative element (has the greatest attraction for the bonding electrons) he assigned it a value and compared all other elements to fluorine. Values for electronegativity can be found on page 10 of the data book

5. Electronegativities Looking across a row or down a group of the periodic table we can see a trend in values. We can explain these trends by applying the same reasoning used for ionisation energies.

6. IncreasingElectronegativity Looking across a period F C B Li Be N O 2.0 2.5 3.0 3.5 4.0 1.0 1.5 What are the electronegativities of these elements? Across a period electronegativity increases The charge in the nucleus increases across a period. Greater number of protons = Greater attraction for bonding electrons

7. F DecreasingElectronegativity Cl Br I Looking down a group 4.0 3.0 What are the electronegativities of these halogens? 2.8 2.6 Down a group electronegativity decreases Atoms have a bigger radius (more electron shells) The positive charge of the nucleus is further away from the bonding electrons and is shielded by the extra electron shells.

8. Trends in electronegativity Electronegativity increases across a period. Electronegativity decreases down a group Going across the period, the nuclear charge increases. This pulls the electron shells closer to the nucleus. As a results, the electronegativity increases. Going down the group, the nuclear charge increases but the number of electron shells also increases. As a result of ‘shielding’ and an increase distance the outer shell is from the nucleus, electronegativity decreases.

9. Chemical bonds: types of bonds Explores how different types of bonds are formed due to variations in the electronegativity of the bonded atoms. The distortion of the orbitals and the polarity of the bond is also displayed. Linus Pauling (1901-1994) An account of the life and work of the Nobel Prize-winning chemist, Linus Pauling. Periodic Table of Data Visual database of the physical and thermochemical properties of the chemical elements which allows the user to plot graphs and tables, play games and view diagrams.

10. Polar covalent bonds Learning intention Learn how differences in electronegativity between bonding atoms lead to the formation of polar covalent bonds.

11. Polar Covalent Bonds Non-polar covalent bond – electrons shared equally between atoms (same electronegativity) Polar covalent bond – electrons shared unequally between atoms (atom B is more electronegative)

12. - - - - + + + + Covalent Bonding A covalent bond is a shared pair of electrons electrostatically attracted to the positive nuclei of two atoms. Both nuclei try to pull the electrons towards themselves The atoms achieve a stable outer electron arrangement (a noble gas arrangement) by sharing electrons. This is like a tug-of-war where both sides are pulling on the same object. It creates a strong bond between the two atoms.

13. Covalent Bonding Picture a tug-of-war: If both teams pull with the same force the mid-point of the rope will not move.

14. H H e e Pure Covalent Bond This even sharing of the rope can be compared to a pure covalent bond, where the bonding pair of electrons are held at the mid-point between the nuclei of the bonding atoms.

15. CovalentBonding What if it was an uneven tug-of-war? The team on the right are far stronger, so will pull the rope harder and the mid-point of the rope will move to the right.

16. Polar Covalent Bond A polar covalent bond is a bond formed when the shared pair of electrons in a covalent bond are not shared equally. This is due to different elements having different electronegativities.

17. δ- δ+ H e e I Polar Covalent Bond e.g. Hydrogen Iodide If hydrogen iodide contained a pure covalent bond, the electrons would be shared equally as shown above. This makes iodine slightly negative and hydrogen slightly positive. This is known as a dipole. However, iodine has a higher electronegativity and pulls the bonding electrons towards itself (winning the tug-of-war)

18. δ- δ+ C Cl Electronegativities Polar Covalent Bond In general, the electrons in a covalent bond are not equally shared. e.g. 3.0 2.5 δ- indicates where the bonding electrons are most likely to be found.

19. C Cl 2.5 3.0 P H 2.2 2.2 O H C Cl O H 3.5 2.2 P H δ- δ+ δ- δ+ Polar Covalent Bond Consider the polarities of the following bonds: Difference Electronegativities Bond 0.5 0 1.3 Increasing Polarity Complete a similar table for C-N, C-O and P-F bonds.

20. Polar Covalent Bonds In the covalent bond between fluorine and hydrogen. The bonding electrons are not shared equally between the two atoms. Hydrogen The fluorine nucleus has more protons and has a stronger pull on the electrons than the hydrogen nucleus.. Fluorine

21. Thus the fluorine atom has a greater share of the bonding electrons and acquires a slight negative charge. - + F H The hydrogen atom is then made slightly positive. The bond is a polar covalent bond and we use the symbols + and - to show this. The dipole produced is permanent. Fluorine is the most electronegative element. It is small atom compared to others and its nucleus is massive for its atomic size. Some other polar covalent bonds are O-H and N-H