2. 4. Electronegativity – bond polarity in covalent bonds. 3. Bonding Learning Objectives:  State what is meant by the term electronegativity.  State what makes one atom more electronegative than another.  State what the symbols δ + and δ - mean when placed above atoms in a covalent bond.

3. 4. Electronegativity – bond polarity in covalent bonds. Forces that hold atoms together are all to do with the attraction between positive and negative. In ionic compounds this is caused by electrons being completely transferred between the atoms. In a covalent bond, the electrons are shared between the two atoms, however, sometimes one atoms is better at pulling electrons towards it than another. This results in said atom having a greater share of the electrons in the bond. That atom would be said to be more electronegative.

4. Electronegativity Electronegativity is a chemical property of an atom Electronegativity is the tendency of an atom to attract (the bonding pair of) electrons towards itself When those electrons are part of a chemical bond we can begin to explain why covalent bonds involves ‘sharing’ and ionic bonding involves ‘donating’ electrons

5. 4. Electronegativity – bond polarity in covalent bonds. Electronegativity is the power of an atom to attract the electron density of a covalent bond towards itself. When considering electrons as clouds of charge, the term electron density is often used to describe the way the charge is distributed. What factors do you think affect how electronegative an atom is?

6. Just a reminder about electron density…

7. Electron density The electrons move very quickly throughout space This creates areas of electron density we call orbitals

8. Electron density Imagine a p-orbital which contains two electrons However at any given moment you might find both of the electrons in the same ‘lobe’ of the orbital

9. 4. Electronegativity – bond polarity in covalent bonds. Electronegativity depends on: 1.The nuclear charge of the atom. 2.The distance between the nucleus and the outer electrons. 3.The shielding of nuclear charge by electrons in inner shells. Use these statements to describe and explain the trend in electronegativity across a period and down a group. What is the most electronegative atom?

10. Electronegativity Electronegativity increases across the period because the number of protons and hence nuclear charge increases. A greater charge creates a greater attraction to the pair of electrons.

11. Electronegativity Electronegativity increases up the group because the bonding pair of electrons are closer to the nucleus

12. 4. Electronegativity – bond polarity in covalent bonds. Electronegativity is measured on the Pauli scale where 4.0 is the largest electronegativity possible. Noble gases do not have a value as it is unusual for them to form covalent bonds. F>O>N>Cl

13. 4. Electronegativity – bond polarity in covalent bonds. Polarity is the unequal sharing of electrons in a bond, if the electron sharing is equal, the bond is described as non- polar. Both chlorine atoms have the same electronegativity (3.0). The electron cloud is shared evenly between the 2 atoms. So the bond is described as non-polar. Chlorine has a higher electronegativity (3.0) than hydrogen (2.1). The chlorine is able to pull more of the electron cloud towards itself and so the bond is described as polar.

14. 4. Electronegativity – bond polarity in covalent bonds. If two atoms that are bonded together do have different electronegativity's, this is represented using partial charges. They are described as ‘delta positive’ or ’delta negative.’

15. 4. Electronegativity – bond polarity in covalent bonds. A molecule can be made up of several polar bonds, but be non-polar overall. Because the molecules are symmetrical, there is no overall polarity, and the electron cloud is shared evenly over the whole molecule.

16. 5. Forces acting between molecules. 3. Bonding Learning Objectives:  State the three types of intermolecular force.  Describe how dipole-dipole and van der Waals forces arise.  State what is needed for hydrogen bonding to occur.  Explain why NH 3, H 2 O and HF have higher boiling points than might be expected.

17. 5. Forces acting between molecules. Types of intermolecular force: Van der Waals force – also known as temporary dipole – induced dipole. Dipole – dipole forces. Hydrogen bonding. Act between ALL atoms and molecules. Act between certain types of molecule. } Increasing strength Intermolecular forces are much weaker than covalent bonds – do not get them confused!

18. 5. Forces acting between molecules. Dipole – Dipole forces: Occur between molecules that possess an overall dipole. For example: HCl. The δ + hydrogen part of one molecule is attracted to the δ - chlorine part of another molecule, as opposite charges attract. Lots of molecules can the join together is this manner. Remember, these forces are weak! This term dipole refers to a molecules where the charge is shared unevenly

19. 5. Forces acting between molecules. Van der Waals forces (instantaneous dipole – induced dipole, AKA London forces): These occur between all molecules, even if there are other intermolecular acting on that molecule. These also occur between atoms. Helium has 2 electrons in its neutral atom, but these electrons could be anywhere at anytime – they are constantly on the move!. If they happen to both be on one side, a temporary dipole is created. This can the induce a dipole in a neighbouring atom, causing the two atoms to be temporarily attracted to each other – until the electrons move again…..

20. 5. Forces acting between molecules. Explain the trend in boiling point of the noble gases… As you go down the group, there are more electrons. As the temporary dipole is caused by electron distortion (moving), the more electrons, the greater the distortion and the stronger the van der Waals forces.

21. 5. Forces acting between molecules. Hydrogen bonding: It is the strongest type of IM force, as it has some character of dipole – dipole interactions, and some of a dative covalent bond, though it is still around 10 times weaker than a covalent bond. It can occur when a molecules has: A very electronegative atom (F, O, N) that has a lone pair of electrons. A hydrogen atom covalently bonded to this electronegative atom. Remember the term lone pair refers to a pair of outer electrons that are not involved in a bond. An example of such a molecule is water. Lone pair Strong positive charge

22. Where do they occur Hydrogen bonds form between a δ + Hydrogen atom and a δ - atom on a SMALL atom which has a lone pair in the 2p orbital Chlorine and other similar electronegative atoms are too big to hydrogen bond This means only N, O and F form hydrogen bonds

23. 5. Forces acting between molecules. Drawing hydrogen bonds: H H O H H O δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ-δ- δ-δ-.. 1. Draw your molecules. ------- 2. Show any lone pairs and partial charges. 3. Draw a straight dotted line between the lone pair on one atom and the hydrogen atom. Note: The bond angle around the hydrogen atom must be shown as 180 o, i.e. in a perfectly straight line!

24. Ammonia and HF Ammonia – 1 lone pair – 3 hydrogens Hydrogen fluoride – 3 lone pairs – 1 hydrogen

25. H-bonding in Ammonia

26. 5. Forces acting between molecules. Quick check: 1. What intermolecular forces are present between water molecules? 2. Show how a hydrogen bond might occur between water and ammonia! Water molecules have van der Waals forces (present between all molecules), dipole – dipole forces (water is a polar molecules) and hydrogen bonds! H H O H H N δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ-δ- δ-δ-.. ------- H δ+δ+

27. 5. Forces acting between molecules. Describe and explain the trends in boiling point of the hydrides. These 3 hydrides have higher than expected boiling points as they are all able to form hydrogen bonds. Carbon is not electronegative enough to form H-bonds, and therefore the boiling point of methane is as expected. Remember: the general trend is an increase in in boiling point as the size of the molecule increases, due to greater electron density – leading to stronger intermolecular forces that require more energy to break.

28. H 2 O HF H 2 Te NH 3 SbH 2 H 2 SeHI H 2 S HBr HCl SnH 4 PH 3 GeH 4 SiH 4 CH 4 Water Water forms the most ‘effective’ hydrogen bonding hence it has the highest boiling point

29. 5. Forces acting between molecules. Look at the information above : describe which interesting and very important phenomenon that it is explaining!

30. Hydrogen bonding in water Each molecule has two hydrogens and two lone pairs allowing two hydrogen bonds per molecule

31. Water Hydrogen bonding is one of the reasons water is so special! Forming four bonds the water molecules can take a tetrahedral shape when it becomes a solid. This is why ice is strong, uniform and less dense than liquid water

32. What are these? C H H H H C H H H H C H H H H C H H H H Molecule – the atoms are bonded Compound – It contains different elements bonded together

33. Let’s look at the properties of molecules C H H H H C H H H H C H H H H C H H H H When a substance is a liquid, weak intermolecular forces hold the molecules together: An intermolecular force is a force that acts between nearby molecules Weak Intermolecular forces hold molecules together

34. The larger the molecule, the stronger the intermolecular forces: Larger molecules have more places for intermolecular forces to happen! C H H H H C H H C H H C H H C H H C H H C H H C H H H H C H H C H H C H H C H H C H H C H H More Intermolecular forces hold molecules together

35. Which do you think would have a higher boiling point? C H H H H C H H H H C H H C H H C H H C H H C H H C H H VS

36. The larger the molecule, the more intermolecular forces that hold its molecules together. This means more energy is required to break the intermolecular forces of larger molecules. This means a higher boiling point (the temperature when it evaporates) C H H H H C H H C H H C H H C H H C H H C H H C H H H H C H H C H H C H H C H H C H H C H H C H H H H C H H H H

37. So… The boiling point is the __________ at which a substance __________. Shorter molecules have _______ boiling points. Larger molecules have ________ boiling points. This is because, larger molecules have more ____________ forces to break so the molecules can be separated. This takes more ________.

38. Dry Ice Dry ice has the formula CO 2 Use your codes, what is it? M = Molecule E = Element C = Compound Compound – It contains different elements bonded together Molecule – the atoms are bonded O C O

39. Dry Ice Dry Ice – CO 2 – Is more commonly known as carbon dioxide. Dry ice, is the frozen solid version of carbon dioxide. What holds the CO 2 molecules together in the solid? O C O O C O O C O O C O O C O Intermolecular Forces

40. Are the intermolecular forces in dry ice strong? NO! They are weak! It is a very small molecule. Only three atoms. O C O O C O What do weak intermolecular forces mean about the boiling point of dry ice?

41. CO 2 is a small molecule This means it has weak intermolecular forces. Very little energy is needed to break these forces. So dry ice has a low boiling point. So low that it turns straight from a solid to a gas! Sublimation O C O