Bonding

1. Bonding

2. Why Bonding? • Some atoms are very reluctant to combine with other atoms and exist in the air around us as single atoms. These are the Noble Gases and have very stable electron arrangements i.e. their outer shells are full. • All other atoms bond together to become electronically more stable (they want noble gas configuration). • Bonding produces new substances and usually only the valenceelectrons are involved.

3. Bonding Types • The physical properties of a substance depend on its structure and type of bonding present. Bonding determines the type of structure. • CHEMICALstrong bonds • ionic (or electrovalent) • covalent • dative covalent (or co-ordinate) • metallic • PHYSICAL weak bonds • van der Waals‘ forces (weakest) • dipole-dipole interaction • hydrogen bonds (strongest)

4. Bonding Types • There are 2 major ways in which atoms can bond: • Ionic ally • Covalently • In ionic bonding, electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond. • Covalent bonding occurs when two (or more) elements share electrons.

5. Bonding

6. Ionic Bonding • Ions are electrically charged particles formed when atoms lose or gain electrons. They have the same electronic structures as noble gases. • Metal atoms form positive ions, while non-metal atoms form negative ions. • e.g. Na+, Cl-, Mg2+, O2- etc . • If negative electrons are lost the excess charge from the protons produces an overall positive ion. If negative electrons are gained there is an excess of negative charge, so a negative ion is formed. • The strong electrostatic forces of attraction between oppositely charged ions are called ionic bonds.

7. ONE combines with ONE to form Ionic Bonding • Example 1: Sodium Chloride (NaCl) 2:8:1 2:8:7 2:8 2:8:8 • An electron is transferred from the 3s orbital of sodium to the 3p orbital of chlorine; both species end up with the electronic configuration of the nearest noble gas the resulting ions are held together in a crystal lattice by electrostatic attraction • What are the ionic equations? • What about the bonding in MgCl2? 1s22s22p63s1 1s22s22p63s23p5 1s22s22p6 1s22s22p63s23p6

8. Ionic Bonding • Example 2: Magnesium Chloride (MgCl2) • Because magnesium have two outer shell electrons, they can combine with two chlorine atoms by the transfer of one electron to each atom to form one Mg2+ and two Cl- ions • What are the ionic equations? e¯ Cl Mg Cl e¯

9. Giant Ionic Crystal Lattice • Oppositely charged ions held in a regular 3-dimensional lattice by electrostatic attraction • The Na+ ion is small enough relative to a Cl¯ ion to fit in the spaces so that both ions occur in every plane. Cl- Chloride ion Na+ Sodium ion

10. Giant Ionic Crystal Lattice Each Na+ is surrounded by 6 Cl¯ (co-ordination number = 6) Each Cl¯ is surrounded by 6 Na+ (co-ordination number = 6)

11. - - - - + + + + + + - - + + - - + + + + Physical Properties • Melting pointvery high A large amount of energy must be put in to overcome the strong electrostatic attractions and separate the ions. • StrengthVery brittle Any dislocation leads to the layers moving and similar ions being adjacent. The repulsion splits the crystal.

12. Na+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ Na+ Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- Cl- Physical Properties • Electrical Don’t conduct when solid - ions held strongly in the lattice conduct when molten or in aqueous solution - the ions become mobile and conduction takes place. DISSOLVING AN IONIC COMPOUND IN WATER BREAKS UP THE STRUCTURE SO IONS ARE FREE TO MOVE TO THE ELECTRODES

13. Covalent Bonding • Covalent bonds are formed by atoms sharing electrons to form molecules. • One single covalent bond is a sharing of 1 pair of electrons, two pairs of shared electrons between the same two atoms gives a double bond and it is possible for two atoms to share 3 pairs of electrons and give a triple bond. • This kind of bond or electronic linkage does act in a particular direction i.e. along the 'line' between the two nuclei of the atoms bonded together, this is why molecules have a particular shape. In the case of ionic or metallic bonding, the electrical attractive forces act in all directions around the particles involved.

14. H H H H Covalent Bonding • Example 1: Two hydrogen atoms form the molecule of the element hydrogen (H2) • Example 2: One atom of hydrogen combines with one atom of Chlorine to give you hydrogen chloride (HCl) H H Cl Cl

15. Covalent Bonding • Draw the bonding in ammonia (NH3) H H N H H N H H

16. Covalent Bonding • Draw the bonding in methane (CH4) H H H C H H C H H H

17. Covalent Bonding • What about oxygen? O O O O each atom needs two electrons to complete its outer shell each oxygen shares 2 of its electrons to form a DOUBLE COVALENT BOND

18. Covalent Bonding • Atoms share electrons to get the nearest noble gas electronic configuration • Some don’t achieve an “octet” as they haven’t got enough electrons e.g. Al in AlCl3 • Others share only some - if they share all they will exceed their “octet” e.g. NH3 and H2O • Atoms of elements in the 3rd period onwards can exceed their “octet” if they wish as they are not restricted to eight electrons in their “outer shell” e.g. PCl5 and SF6

19. Physical Properties (Covalent Bonding) For simple covalent molecules (not giant covalent) • Electrical Don’t conduct electricity as they have no mobile ions or electrons • Solubility Tend to be more soluble in organic solvents than in water • Boiling point The forces between molecules (intermolecular forces) are weak and known as van der Waals forces. Attractions between molecules increases as the molecules get more electrons. e.g. CH4 -161°C C2H6 - 88°C C3H8 -42°C as the intermolecular forces are weak, little energy is required to separate molecules from each other so boiling points are low

20. End • Recap

21. Previously... • CHEMICALstrong bonds • Ionic • Covalent Today’s session: • PHYSICAL weak bonds • van der Waals‘ forces • dipole-dipole interactions • hydrogen bonds

22. Inter/Intra -molecular Forces • Intermolecular attractions are attractions between one molecule and a neighbouring molecule. Weak • All intermolecular attractions are known collectively as van der Waals forces. • There are 2 types: • Dispersion forces (London forces) or instantaneous dipole-induced dipole forces • Dipole-dipole interactions • The forces of attraction which hold an individual molecule together (for example, the covalent bonds) are known as intramolecular attractions. Strong

23. Intermolecular Forces • In the context of this session, the word dipole means an asymmetric distribution of electron electrical charge to give partially positive and partially negative regions in the same molecule. • In a simple sense its a molecule with a partially positive end and a partial negative charge at the other end. • Electric dipoles may be permanent or transient (temporary) and the molecules discussed here are electrically neutral  overall. • These attractive forces can operate between ANY particles whatever their constitution including free atoms in a gas or ions in a crystal etc.

24. van der Waals Forces Instantaneous dipole-induced dipole forces/Dispersion Forces • Because electrons move quickly in orbitals, their position is constantly changing; at any given instant they could be anywhere in an atom. The possibility will exist that one side will have more electrons than the other. This will give rise to a dipole... • The dipole on one atom induces dipoles on nearby atoms • Atoms are now attracted to each other by a weak forces Original temporary dipole Induced dipole

25. van der Waals Forces • There is no reason why this has to be restricted to a few molecules. As long as the molecules are close together this synchronised movement of the electrons can occur over huge numbers of molecules. • Although the bonding within molecules is strong, between molecules it is weak. (Intramolecular vs. intermolecular)

26. Strength of Dispersion Forces • The boiling points of the noble gases are • Helium -269°C • Neon -246°C • Argon -186°C • Krypton -152°C • Xenon -108°C • Radon -62°C • What is happening to the boiling point as you go down the group? • Why?

27. Strength of Dispersion Forces • Group 18 (Noble Gases) • The boiling points increase as you go down the group • All of these elements have monatomic molecules. • The number of electrons increases and hence the radius of the atom. • The more electrons you have, and the more distance over which they can move, the bigger the possible temporary dipoles and therefore the bigger the dispersion forces • Neon molecules will break away from each other at much lower temperatures than argon molecules - hence neon has the lower boiling point.

28. Strength of Dispersion Forces Will molecular shape have an effect on the strength of dispersion forces? Use butane and 2-methyl propane to explain your answer. • Long thin molecules can develop bigger temporary dipoles due to electron movement than short fat ones containing the same numbers of electrons. • Long thin molecules can also lie closer together - these attractions are at their most effective if the molecules are really close. • Butane has a higher boiling point because the dispersion forces are greater. The molecules are longer (and so set up bigger temporary dipoles) and can lie closer together than the shorter, fatter 2-methylpropane molecules.

29. Electronegativity • The ability of an atom to attract the electron pair in a covalent bond to itself • Non-polar bond e.g. Cl2, O2 • similar atoms have the same electronegativity they will both pull on the electrons to the same extent the electrons will be equally shared • Polar bond e.g. HCl • different atoms have different electronegativities one will pull the electron pair closer to its end it will be slightly more negative than average, and the other end slightly more positive. A dipole is formed and the bond is said to be polar. • Greater electronegativity difference = greater polarity

30. Pauling Scale • The Pauling Scale is a scale for measuring electronegativityvalues increase across periods values decrease down groups fluorine has the highest value H 2.1 Li Be B C N O F 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Na Mg Al Si P S Cl 0.9 1.2 1.5 1.8 2.1 2.5 3.0 K Br 0.8 2.8 Predict the polarity for the following: S – Cl ; C – O and C – C

31. Dipole-dipole Interactions • A molecule like HCl has a permanent dipole because chlorine is more electronegative than hydrogen. These permanent, in-built dipoles will cause the molecules to attract each other rather more than they otherwise would if they had to rely only on dispersion forces. • Dipole-dipole interactions occur in addition to Dispersion forces . • Molecules that have permanent dipoles will therefore have higher boiling points than those that have temporary dipoles. • Use ethane and fluoromethaneto explain why a molecule with a permanent dipole has a higher boiling point.

32. Dipole-dipole Interactions Ethane Fluormethane • Both have the same number of electrons so we expect the dispersion forces to be the same. • Fluromethane has a higher boiling point due to the large permanent dipole on the molecule (because of the highly electronegative Fluorine).

33. Hydrogen Bond • the attractive force between hydrogen in a polar bond (particularly H-F, H-O, H-N bond) and an unshared electron pair on a nearby small electronegative atom or ion.

34. Hydrogen Bonding & Water

35. Hydrogen Bond • This other atom may be in the same molecule or in a nearby molecule, but always has to include hydrogen. INTERMOLECULAR HYDROGEN BONDING

36. INTRAMOLECULAR HYDROGEN BONDING • Hydrogen Bonds have about 5% of the strength of an average covalent bond • Hydrogen Bond is the strongest of all intermolecular forces

37. One of the most remarkable consequences of H-bonding is found in the lower density of ice in comparison to liquid water, so ice floats on water. The molecules in the solid are more densely packed than in the liquid. A given mass of ice occupies a greater volume than that of liquid water because of an ordered open H-bonding arrangement in the solid (ice) in comparison to continual forming & breaking H-bonds as a liquid.

38. Intermolecular Forces • generally much weaker than covalent or ionic bonds. Less energy is thus required to vaporize a liquid or melt a solid. Boiling points can be used to reflect the strengths of intermolecular forces (the higher the Bpt, the stronger the forces)