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WELCOME TO SKO16 CHEMISTRY CHEMISTRY CHEMISTRY

WELCOME TO SKO16 CHEMISTRY CHEMISTRY CHEMISTRY. CHEMISTRY SK016. 1.0 Matter 7 2.0 Atomic Structure 7 3.0 Periodic Table 4 4.0 Chemical Bonding 2 5.0 State of Matter 7 6.0 Chemical Equilibrium 5 7.0 Ionic Equilibria 12 Total 54. 08/16/11. matter. 2.

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WELCOME TO SKO16 CHEMISTRY CHEMISTRY CHEMISTRY

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  1. WELCOME TO SKO16 CHEMISTRY CHEMISTRY CHEMISTRY

  2. CHEMISTRY SK016 1.0 Matter 7 2.0 Atomic Structure 7 3.0 Periodic Table 4 4.0 Chemical Bonding 2 5.0 State of Matter 7 6.0 Chemical Equilibrium 5 7.0 Ionic Equilibria 12 Total 54 08/16/11 matter 2

  3. CHEMISTRY SK026 8.0 Thermochemistry 4 9.0 Electrochemistry 6 10.0 Reaction Kinetics 7 11.0 Intro To Organic Chemistry 4 12.0 Hydrocarbons 8 13.0 Aromatic Compounds 3 14.0 Haloalkanes (Alkyl halides) 4 15.0 Hydroxy compounds 3 08/16/11 matter 3

  4. CHEMISTRY SK026 16.0 Carbonyl 4 17.0 Carboxylic acids & Derivatives 4 18.0 Amines 5 19.0 Amino acids and Proteins 2 20.0 Polymers 1 08/16/11 matter 4

  5. ASSESSMENT • 1. COURSEWORK (20%) • Continuous evaluation (tutorial/test/quiz) - 10% • Practical work - 10% 2. MID-SEMESTER EXAMINATION - 10% 3. FINAL EXAMINATION (70%) • Paper 1 (30 multiple choice questions) - 30% • Paper 2 (Part A-structured) (Part B-long structured) -100% 08/16/11 matter 5

  6. REFERENCE BOOKS • CHEMISTRY ,9th Ed. – Raymond Chang, McGraw-Hill • CHEMISTRY –The Molecular Nature of Matter and Change, 3rd Ed.– Martin Silberberg, McGraw Hill • CHEMISTRY – The Central Science, 9th Ed. Theodore L.Brown, H.Eugene LeMay,Jr, Bruce E Bursten, Pearson Education • GENERAL CHEMISTRY – Principle & Structure, 6th Ed. James E Brady, John Wiley and Sons. 08/16/11 matter 6

  7. GENERAL CHEMISTRY – Principle and Modern Applications, 8th Ed. Ralph H. Petrucci, William S. Harwood, Prentice-Hall • ORGANIC CHEMISTRY, 7th Ed – T.W.Graham Solomon,Craig B.Fryhle, John Wiley and Sons • ORGANIC CHEMISTRY, 4th Ed – L.G. Wade, Jr, Prentice Hall • ORGANIC CHEMISTRY, 6th Ed – John McMurry, Thompson – Brooks/Cole 08/16/11 matter 7

  8. Chapter 1 : MATTER 1.1 Atoms and Molecules 1.2 Mole Concept 1.3 Stoichiometry 08/16/11 matter 8

  9. 1.1 Atoms and Molecules 08/16/11 matter 9

  10. Learning Outcome At the end of this topic, students should be able to: (a) Describe proton, electron and neutron in terms of the relative mass and relative charge. (b) Define proton number, Z, nucleon number, A and isotope. (c) Write isotope notation. 08/16/11 matter 10

  11. Introduction • Matter  Anything that occupies space and has mass. e.g: air, water, animals, trees, atoms, etc • Matter may consists of atoms, molecules or ions. 08/16/11 matter 11

  12. Classifying Matter 08/16/11 matter 12

  13. Asubstanceis a form of matter that has a definite or constant composition and distinct properties. Example: H2O, NH3, O2 • Amixtureis a combination of two or more substances in which the substances retain their identity. Example : air, milk, cement 08/16/11 matter 13

  14. Anelementis a substance that cannot be separated into simpler substances by chemical means. Example : Na, K, Al,Fe Acompoundis a substance composed of atoms of two or more elements chemically united in fixed proportion. Example : CO2, H2O, CuO

  15. Three States of Matter SOLID LIQUID GAS 08/16/11 matter 15

  16. 1.1 Atoms and Molecules • a) Atoms • An atom is the smallest unit of a chemical element/compound. • In an atom, there are 3 subatomic particles: - Proton (p) - Neutron (n) - Electron (e) 08/16/11 matter 16

  17. Modern Model of the Atom • Electrons move around the region of the atom. 08/16/11 matter 17

  18. All neutral atoms can be identified by the number of protons and neutrons they contain. Proton number (Z) is the number of protons in the nucleus of the atom of an element (which is equal to the number of electrons). Protons number is also known as atomic number. Nucleon number (A) is the total number of protons and neutrons present in the nucleus of the atom of an element. Also known as mass number.

  19. Subatomic Particles 08/16/11 matter 19

  20. Isotope • Isotopes are two or more atoms of the same element that have the same proton number in their nucleus but different nucleon number. 08/16/11 matter 20

  21. Examples:

  22. Isotope Notation An atom can be represented by an isotope notation ( atomic symbol ) X = Element symbol Z = Proton number of X (p) A = Nucleon number of X = p + n 08/16/11 matter 22

  23. Nucleon number of mercury, A = 202 Total charge on the ion The number of neutrons = A – Z = 202 – 80 = 122 Proton number of mercury, Z = 80 08/16/11 matter 23

  24. In a neutral atom:  number of protons equals number of electrons In a positive ion:  number of protons is morethan number of electrons In a negative ion:  number of protons is less than number of electrons

  25. Exercise 1 Give the number of protons, neutrons, electrons and charge in each of the following species: 08/16/11 matter 25

  26. Exercise 2 Write the appropriate notation for each of the following nuclide : 08/16/11 matter 26

  27. b) Molecules A molecule consists of a small number of atoms joined together by bonds. 08/16/11 matter 27

  28. A diatomic molecule • Contains only two atoms Ex : H2, N2, O2, Br2, HCl, CO A polyatomic molecule • Contains more than two atoms Ex : O3, H2O, NH3, CH4 08/16/11 matter 28

  29. Learning Outcomes At the end of this topic, student should be able to : (a) Define relative atomic mass, Ar and relative molecular mass, Mr based on the C-12 scale. (b) Calculate the average atomic mass of an element given the relative abundance of isotopes or a mass spectrum. 08/16/11 matter 29

  30. Relative Mass • Relative Atomic Mass, Ar A mass of one atom of an element compared to 1/12 mass of one atom of 12C with the mass 12.000 amu 08/16/11 matter 30

  31. Mass of an atom is often expressed in atomic mass unit, amu (or u). • Atomic mass unit, amu is defined to be one twelfthof the mass of 12C atom • Mass of a 12C atom is given a value of exactly 12 amu 1 u = 1.660538710-24 g • The relative isotopic massis the mass of an atom, scaled with 12C. 08/16/11 31

  32. Example 1 Determine the relative atomic mass of an element Y if the ratio of the atomic mass of Y to carbon-12 atom is 0.45 ANSWER: 08/16/11 matter 32

  33. ii) Relative Molecular Mass, Mr A mass of one molecule of a compound compared to 1/12 mass of one atom of 12C with the mass 12.000amu 08/16/11 matter 33

  34. The relative molecular mass of a compound is the summation of the relative atomic masses of all atoms in a molecular formula.

  35. Example 2 Calculate the relative molecular mass of C5H5N, Ar C = 12.01 Ar H = 1.01 Ar N = 14.01 08/16/11 matter 35

  36. MASS SPECTROMETER • An atom is very light and its mass cannot be measured directly • A mass spectrometer is an instrument used to measure the precise massesand relative quantityof atoms and molecules 08/16/11 36

  37. 08/16/11 matter 37

  38. Mass Spectrum of Monoatomic Elements • Modern mass spectrum converts the abundance into percent abundance 08/16/11 38

  39. 63 9.1 8.1 24 25 26 m/e (amu) Mass Spectrum of Magnesium • The mass spectrum of Mg shows that Mg consists of 3 isotopes: 24Mg, 25Mg and 26Mg. • The height of each lineis proportional to the abundance of each isotope. • 24Mg is the most abundant of the 3 isotopes Relative abundance 08/16/11 matter 39

  40. Learning Outcomes At the end of this topic, student should be able to : (a) Calculate the average atomic mass of an element given the relative abundances of isotopes or a mass spectrum. 08/16/11 matter 40

  41. How to calculate the relative atomic mass, Ar from mass spectrum? • Ar is calculated using data from the mass spectrum. • The average of atomic masses of the entire element’s isotope as found in a particular environment is the relative atomic mass, Ar of the atom. 08/16/11 matter 41

  42. Example 1: Calculate the relative atomic mass of neon from the mass spectrum. 08/16/11 42

  43. Solution: Average atomic = mass of Ne = = 20.2 u Relative atomic mass Ne = 20.2

  44. Example 2: Copper occurs naturally as mixture of 69.09% of 63Cu and 30.91% of 65Cu. The isotopic masses of 63Cu and 65Cu are 62.93 u and 64.93 u respectively. Calculate the relative atomic mass of copper. 08/16/11 44

  45. Solution: Average atomic = mass of Cu = = 63.55 u Relative atomic mass Cu = 63.55

  46. Example 3: Naturally occurring iridium, Ir is composed of two isotopes, 191Ir and 193Ir in the ratio of 5:8. The relative isotopic mass of 191Ir and 193Ir are 191.021 u and 193.025 u respectively. Calculate the relative atomic mass of Iridium 08/16/11 46

  47. Solution: Average atomic = mass of Ir = = 192.254 u Relative atomic mass Ir = 192.254 08/16/11 47

  48. Mass Spectrum of Molecular Elements A sample of chlorine which contains 2 isotopes with nucleon number 35 and 37 is analyzed in a mass spectrometer. How many peaks would be expected in the mass spectrum of chlorine?

  49. MASS SPECTROMETER 35Cl-35Cl 35Cl-37Cl 37Cl-37Cl _ _ + Cl2 35Cl-35Cl+ 35Cl-37Cl+ Cl2 + e  Cl2+ + 2e 37Cl-37Cl+ Cl2 + e  2Cl+ + 2e 35Cl+ 37Cl+

  50. Mass Spectrum of Diatomic Elements

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