1. Octet rule / Chemical bonding The octet rule is a simple chemical rule of thumb that states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. The rule is applicable to the main- group elements, especially carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium or magnesium. In simple terms, molecules or ions tend to be most stable when the outermost electron shells of their atoms contain eight electronschemicalrule of thumbatomselectronsvalence shellselectronic configurationnoble gascarbonnitrogenoxygenhalogenssodiummagnesiumelectron shells

2. Octet rule The octet rule states that atoms are most stable when they have a full shell of electrons in the outside electron shell. The first shell has only two electrons in a single s subshell. Helium has a full shell, so it is stable, an inert element. All the other shells have an s and a p subshell, giving them at least eight electrons on the outside. The s and p subshells often are the only valence electrons, thus the octet rule is named for the eight s and p electrons

3. Chemical reactions Chemical reactions take place because atoms need to achieve a more stable electronic configuration. Chemical reactions involve the re-distribution of the electrons in the outter shell. Outter shell electrons are called valence electrons

4. Octet rule / Periodic table All atoms,except the group 8 atoms need to achieve 8 electrons in the outter shell. In order to do these atoms have two choices, ie to loose electrons or to gain electrons. Atoms in groups 1-3 will loose,ie donate their, electrons when they react. Atoms in groups 5-7 will accept, ie gain electrons, when they react.

5. Valence shells The valence shell is the outermost shell of an atom. It is usually said that the electrons in this shell make up its valence electrons, that is, the electrons that determine how the atom behaves in chemical reactions. atomvalence electrons Atoms with complete valence shells,Group 8 or 0 = noble gases) are inert (unreactive). Those with only one electron in their valence shells (alkalis) or just missing one electron from having a complete shell (halogens) are the most reactive.alkalishalogens

6. Periodic table and valence On the Periodic Table with shell totals you can easily see the octet rule. A valence is a likely charge on an element ion. All of the Group 1 elements have one electron in the outside shell and they all have a valence of plus one. Group 1 elements will lose one and only one electron, to become a single positive ion with a full electron shell of eight electrons (an octet) in the s and p subshells under it. Example Na ---  Na+ + e- – 2,8,1 2,8

7. Group 2 Group 2 elements all have two electrons in the outer shell and all have a valence of plus two. Beryllium can be a bit different about this, but all other Group 2 elements can lose two electrons to become +2 ions. They do not lose only one electron, but two. Example Ca ---  Ca 2+ + 2e- 2,8,8,2 2, 8, 8

8. Group 3 Group 3 elements have a valence of plus three. Boron is a bit of an exception to this because it is so small it tends to bond covalently. Aluminum has a valence of +3. Example Al ---  Al 3+ + 3e-

9. Group 4 The smallest Group 4 elements, carbon and silicon, are non-metals because it is difficult to lose the entire four electrons in the outer shell. Small Group 4 elements tend to make only covalent bonds, sharing electrons. Larger Group 4 elements have more than one valence, usually including +4. We will explain this later

10. Group 5 Small Group 5 elements, nitrogen and phosphorus, are non-metals. They tend to either gain three electrons to make an octet or bond covalently. The larger Group 5 elements have more metallic character. We will explain these later.

11. Group 6 Small Group 6 elements, oxygen and sulfur, tend to either gain two electrons or bond covalently. The larger Group 6 elements have more metallic character. Example O + 2e- ---  O 2- 2,6, 2, 8,

12. Group 7 Group 7 elements all have seven electrons in the outer shell and either gain one electron to become a -1 ion or they make one covalent bond. Example F + 1e- --  F - – 2,7, 2, 8, The Group 7 elements are diatomic gases due to the strong tendency to bond to each other with a covalent bond.

13. Group 0/8 = Inert All of the inert elements, the noble gases, have a full octet in the outside shell (or two in the first shell). They do not naturally combine chemically with other elements. Inert = chemically un-reactive.

14. Chemical bonding: ionic Two basic types: ionic or covalent Ionic bonding: When atoms react they loose or gain electrons. Atoms from groups 1-3 all react by loosing/or donating their electrons. All metals behave this way, including Fe,Cu,Zn. On the other hand atoms in groups 6 and 7 react by accepting /or gaining electrons. These are the non-metals such as F, Cl, I, O, and S.

15. Ionic bonding/ ions All metals loose electrons to form positive ions. These are called cations, i.e Na +, Ca 2+, Fe 2+, Al 3+ All non-metals gain electrons to form negatively charged anions, i.e. Cl -,S 2- and O 2-

16. Ionic bonds

17. Ionic bonding Ionic Bonding – Some atoms gain electrons to become anions – Others lose electrons to become cations – Ions are attracted by their opposing charges – Electrical Neutrality Maintained – Most Important Bonding in Rocks and Minerals

18. Ionic bonds The Ionic Bond: Ionic bonds are formed when there is a complete transfer of electrons from one atom to another, resulting in two ions, one positively charged and the other negatively charged. For example, when a sodium atom (Na) donates the one electron in its outer valence shell to a chlorine (Cl) atom, which needs one electron to fill its outer valence shell, NaCl (table salt) results. Ionic bonds are often 4-7 kcal/mol in strength

19. Chemical bonding: covalent Covalent bonding is when atoms share their valence electrons in order to have 8 in their outer shells. This kind of bonding takes place between non-metals groups 4, 5 and 6. This bonding does not form ions. Such bonds lead to stable molecules if they share electrons in such a way as to create a noble gas configuration for each atom. Each atom donates one electron to form a single covalent bond

20. Lewis structure The Lewis structures are just an attempt to show these valence electrons in a graphic manner as they are used to combine with other elements. The element symbol is in the center and as many as four groups of two electrons are shown as dots above, below, to the right and left of the element symbol to show the valence electrons. All of the inert gases (noble gases) have all eight of the electrons around the element symbol, except for helium, which has only two electrons even with a full shell. Below is a demonstration of the noble gases written in Lewis structure. Notice the electrons are in red just to emphasize them.

21. Lewis structure

22. Covalent bonding = sharing of electrons

23. Dot/ cross diagram

24. Water: bond pair and lone pair

25. Diatomic molecules/elements Why are some elements diatomic? Hydrogen gas forms the simplest covalent bond in the diatomic hydrogen molecule.hydrogen molecule The halogens such as chlorine also exist as diatomic gases by forming covalent bonds. The nitrogen and oxygen which makes up the bulk of the atmosphere also exhibits covalent bonding in forming diatomic molecules.

26. Covalent bonds

27. Covalent bonding Covalent Bonding – Electrons share electrons to fill incomplete shells – Most Important Bonding in Organic Materials (and Organisms)

28. Covalent bonds The Covalent Bond: Covalent Bonds are the strongest chemical bonds, and are formed by the sharing of a pair of electrons. The energy of a typical single covalent bond is ~80 kilocalories per mole (kcal/mol). However, this bond energy can vary from ~50 kcal/mol to ~110 kcal/mol depending on the elements involved. Once formed, covalent bonds rarely break spontaneously.

29. Octet rule

30. Ionisation energy

31. Ionization energy Definition: The first ionisation energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. This is more easily seen in symbol terms. It is the energy needed to carry out this change per mole of Na. Na (g) ---  Na + (g) + e-

32. Ionisation energy The state symbols - (g) - are essential. When you are talking about ionisation energies, everything must be present in the gas state. Ionisation energies are measured in kJ mol -1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high). All elements have a first ionisation.

33. Ionization energy It is the amount of energy required to remove one electron form the outter shell of an atom. An atom can have as many ionization energies as it has electrons. Sodium has 11 electrons so it can have 11 ionization energies. First,second third and so on… till eleven.

34. Ionisation energies for Na

35. Patterns The first 20 elements First ionisation energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar. These variations in first ionisation energy can all be explained in terms of the structures of the atoms involved.

36. Trends in ionisation energy

37. Ionization energy Notice that the highest ionization energies are seen in the Group O elements, known more commonly as the inert, or noble gases. This is because the atoms of these elements have p-orbitals that are full of electrons. This electronic configuration is particularly stable, so that the atoms of these elements are extremely reluctant to lose any electrons. as Na + and K +.

38. Ionization energy At the other extreme, the lowest ionization energies are seen in the Group 1 elements, known as the alkali elements. These elements only have one electron in their outermost orbital, and loss of this one electron will give an ion with a stable, noble gas-like, electronic configuration. Hence, they readily give up this outer electron. This means they readily form singly charged cations.

39. Explaining the pattern in the first few elements Hydrogen has an electronic structure of 1s 1. It is a very small atom, and the single electron is close to the nucleus and therefore strongly attracted. There are no electrons screening it from the nucleus and so the ionisation energy is high (1310 kJ mol -1 ). Helium has a structure 1s 2. The electron is being removed from the same orbital as in hydrogen's case. It is close to the nucleus and unscreened. The value of the ionisation energy (2370 kJ mol -1 ) is much higher than hydrogen, because the nucleus now has 2 protons attracting the electrons instead of 1.

40. Ionization energy It is influenced: – By distance of outer electron from nucleus – The further away the less energy. – The balance between remaining electrons and protons in the nucleus.

41. Factors affecting the size of ionisation energy A high value of ionisation energy shows a high attraction between the electron and the nucleus.The size of that attraction depends: The charge on the nucleus.The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it. The distance of the electron from the nucleus. Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

42. Trends in ionisation energy down a group As you go down a group in the Periodic Table ionisation energies generally fall. Taking Group 1 as a typical example: Why is the sodium value less than that of lithium?

43. Ionisation energy of Group 1

44. Trends in Group 1 There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is much greater. You might have expected a much larger ionisation energy in sodium, but offsetting the nuclear charge is a greater distance from the nucleus and more screening. Li 1s 2 2s 1 1st I.E. = 519 kJ mol -1 Na 1s 2 2s 2 2p 6 3s 1 1st I.E. = 494 kJ mol -1

45. Explaining pattern of ionisation energy

47. Ionisation energy and reactivity

48. Ionisation energies and reactivity The lower the ionisation energy, the more reactive is the atom/element. You can explain the increase in reactivity of the Group 1 metals (Li, Na, K, Rb, Cs) as you go down the group in terms of the fall in ionisation energy. Whatever these metals react with, they have to form positive ions in the process, and so the lower the ionisation energy, the more easily those ions will form.

49. Sodium: very reactive Sodium ((Latin natrium), is a soft, silvery white, highly reactive element and is a member of the alkali metals within "group 1“ It has only one stable isotope, 23 Na.Latinsilvery whitealkali metalsisotope Sodium quickly oxidizes in air and is violently reactive with water, so it must be stored in an inert medium, such as kerosene. Sodium is present in great quantities in the earth's oceans as sodium chloride (common salt).airwaterkerosenesodium chloride

50. Sodium: ionization energy Na (1s 2 2s 2 2p 6 3s 1 ) ----> Na + (1s 2 2s 2 2p 6 ) + e - First ionization energy for sodium Can have a maximum of 11 ionization energies.

51. Sodium Na

52. Chemical bonding

53. Ionic bonding

54. C-H Covalent Bond

55. Carbon-Hydrogen Covalent Bonding Bond NumberExample Energy (kcal/mol) single H | H--C--H | H ~80 double H H | | H--C==C--H | | H H ~150 triple H | C ||| C | H ~200

56. Covalent bonds