Covalent Bonding

1. Covalent Bonding Chapter 7:

2. What is covalent bonding? • Covalent bonding is the force that holds two or more atoms together when electrons are shared between atoms to achieve a stable number of electrons in the outer shell. • The shared electrons continually move between the two atoms acting to fill two outer shell orbitals at once. • Covalent molecules usually have a zero charge and individual molecules are held together by intermolecular forces.

3. Electron shell diagrams Both Chlorine atoms want one more electron in their outer shell. By sharing their unpaired electrons both atoms achieve a full outer shell.

4. Electron dot diagrams – Lewis diagrams • Electrons occur in pairs in their shells. In the outer shell the first four electrons will fill singularly creating four single electrons and the opportunity for four bonds. • After this, the electrons start to pair up, so the number of bonds that can be formed decreases until the outer shell is full.

5. Single, double and triple bonds • Some covalent molecules can form double bonds. These are stronger than a single bond. • Triple bonds may also be formed.

6. Drawing covalent molecules • The symbol for each atom is represented with lines branching from it to the atom(s) it is bonded to. This should be done so as to represent the 3D shape of the molecule. Benzene

7. Shapes of common molecules

8. Naming covalent molecules SO3 would be called Sulphur trioxide N2O5 would be called dinitrogen pentoxide Cl2O would be called oxygen dichloride

9. Naming covalent molecules Have a go at naming the following: NCl3 SF6 CO PH3 SiH4 N2O

10. Polar and non polar molecules • To add to the confusion many covalent molecules are what we call polar. • This means that one end of the molecule has a slightly positive ( + ) charge and the other end a slightly negative ( - ) charge. • This is because some atoms attract electrons more than others. This is called electronegativity.

11. Draw a diagram of the following polar molecules: • HCl • H2O • CH3Cl

12. Identifying polarity • Does the molecule have a dipole covalent bond? • Can dipole charges be assigned to the ends of the molecule? • Do these dipole charges cancel each other out? • Is there 0 or 1 line of symmetry? • Does the dipole charges have a direction?

13. Intramolecular forces So far we have been talking about forces within a molecule. Is the molecule polar? Or non-polar? However, most covalent substances are made up of more than one singular molecule so we also need to talk about forces that hold these individual molecules together. Eg. Water is made up of many H2O molecules, how are these molecules held together??

14. Intermolecular attractions • Covalent molecules can form liquids and solids which means there must be some kind of forces holding these molecules together. • Dispersion forces are very weak forces of attraction between molecules.

15. Dispersion Forces (Van der Waals)Weakest Intermolecular force In a sample of molecules, the nuclei of atoms in one molecule are able to attract the electrons of atoms in neighbouring molecules. All electrons are attracted by neighbouring nuclei, thus, dispersion forces are present in all molecules. • Instantaneous dipoles can also contain dispersion forces

16. Dispersion Forces (Van der Waals)Weakest Intermolecular force Two factors influence the strength of dispersion forces: • The number of electrons in the molecules. In general, the more electrons the molecules of a substance have, the stronger the dispersion forces between them. • The shapes of the molecules. Shape affects how closely the molecules may approach each other. The closer the molecules can get, the stronger the attraction will be.

17. Dipole – Dipole interactions • Polar molecules attract one another as the positive dipole will be attracted to the negative dipole of another molecule. • This will result in compounds with higher boiling points. • Polarity also determines which covalently bonded liquids will mix and which will not. • Water is a polar molecule and oil is non polar. These liquids will not mix – immiscible.

18. Hydrogen Bonding • Hydrogen bonding is a special case of Dipole-dipole attraction • When H bonds to a more electronegative atom, its sole electron is drawn to the other atom, leaving its nucleus exposed. This forms a dipole and as the H nucleus in unshielded, other molecules can approach far more closely. • This is especially the case when hydrogen is bonded to Nitrogen, Fluorine and Oxygen. • It creates a much stronger dipole bond, known as a Hydrogen bond. • High melting and boiling points • Hydrogen bonds are vitally important for life as they keep water as a liquid at room temperature, give proteins their unique shape and hold the two DNA strands together.

19. Properties of Covalent structures A great variety of physical properties occurs due largely to the differences in the strength of the intermolecular attractions. • Don’t conduct electricity in sold/molten form as molecules are electrically neutral • Some molecules dissolve in water to form ions and thus can conduct electricity • Like dissolves like (polar molecules dissolve in polar solvents and non-polar dissolves in non-polar solvents) • Low melting and boiling points due to weak forces • Substances are soft and easily scratched • Mostly gas or liquid at room temperature

20. Covalent lattices • Carbon and Silicon can form covalent lattices. • These are often extremely hard and durable substances such as diamond, graphite and quartz.

21. Covalent lattices • Covalent Network Lattice -Some non-metals form giant structures in which no individual molecules exist. They consist of countless numbers of atoms covalently bonded to each other, forming a 3D lattice. Eg Diamond, Silicon dioxide • Covalent Layer Lattice - Each layer is held together covalently (strong bonds) and the layers are held together by dispersion forces (weak bonds) Eg Graphite

22. Properties of Covalent lattices Network Lattice: • Large amount energy to break lattice so high boiling and melting points • No free moving electrons so no conductivity • Hard, as lattice holds atoms in fixed positions • Unreactive Layer Lattice: • Large amount energy to break lattice so high boiling and melting points • Graphite has a single free electron so does conduct electricity • Forces between layers weak so layers slide over each other, graphite appears soft. • Unreactive