1. CHEM110W2 Ms Janine Kasavel kasavel@ukzn.ac.za 031-2607747 Rm: 03-041 1

2. LECTURES MONDAY L22 07h45 THURSDAY L22 13h15 FRIDAY L22 08h40 TUTORIAL WEDNESDAY 10h30 -12h10 2

3. Quantitative Chemistry matter units significant figures atomic structure, isotopes, periodic table basic nomenclature (ions, molecular &inorganic compounds) stoichiometry and balancing equations by inspection moles and Avogadro’s number empirical and molecular formulae limiting reagents 3

4. MATTER “matter is anything that has mass and takes up space” Pure Substance -ELEMENT: can’t be decomposed into simpler substances -COMPOUND: composed of 2/> different elements Mixture -HETEROGENEOUS: visibly different composition, properties or appearance -HOMOGENEOUS: visibly uniform composition, properties & appearance throughout 4

5. Physical properties of matter -measured without changing the identity or composition of the substance Chemical properties of matter -describe the way a substance may change or react to form other substances 5

6. UNITS Système International (SI) MASS kilogram kg LENGTH metre m TIME second s TEMPERATURE Kelvin K G giga 10 9 M mega 10 6 k kilo 10 3 d deci 10 -1 c centi 10 -2 m milli 10 -3 µ micro 10 -6 n nano 10 -9 p pico 10 -12 6 SCIENTIFIC NOTATION

7. SIGNIFICANT FIGURES 1)any figure that is not zero is significant. 2)zeroes between non-zero figures are significant. 3) exact (“counting”) numbers by definition have an ¥ number of s.f., so physical constants defined to be exact numbers do so also. 4) leading zeroes (to the left of the first non-zero figure) are not significant. 5)trailing zeroes (to the right of the last non-zero figure) are significant only if the number has a d.p. 6)in measurements without a d.p., the number of s.f. is ambiguous. 7

8. Using Significant Figures in Calculations multiplication/division Number of s.f. in final answer is the same as the LEAST of numbers of s.f. in each of original measurements. addition/subtraction Number of d.p. in final answer is the same as the LEAST of numbers of d.p. in each of original measurements. 8

9. DENSITY ρ = mass/ volume ρ: gcm -3 mass: g volume: cm 3 9

10. PRACTICE EXAMPLE 10 A nugget of gold with a mass of 521 g is added to 50.0 mL of water. The water level rises to a volume of 77.0 mL. What is the density of the gold?

11. ATOMIC STRUCTURE 11 PROTONS, NEUTRONS in the nucleus surrounded by orbiting ELECTRONS Early Atomic Theory (Dalton 1803 – 1807) Cathode Rays & Particles (Thomson, 1897) Electron Charge & Mass (Millikan, 1909) Nuclear Atom (Rutherford, 1910) Modern Atomic Structure (Rutherford, 1919)

12. 12 ChargeMass Actual/ Coulombs Relative Actual/ g Relative/ u Proton1.602 x 10 -19 + 11.673 x 10 -24 1.00727 Electron1.602 x 10 -19 - 19.109 x 10 -28 0.00054858 Neutron001.675 x 10 -24 1.00866

13. 13 A: mass number = no. protons + no. neutrons Z: atomic number = no. protons /electrons A Z E

14. Isotopes Atoms of the same element with different mass numbers due to: different numbers of neutrons 14

15. Average atomic mass 15 AAM: average atomic mass IM: isotopic mass

16. PRACTICE EXAMPLE 16 Naturally occurring Mg has three isotopes: 24 Mg (78.90 %) 23.9850 u 25 Mg (10.00 % )24.9858 u 26 Mg (11.10 %) 25.9826 u AAM=?

17. 17

18. 18 IONS If electrons are added to or removed from a neutral atom, an ion is formed. When an atom or molecule loses electrons it becomes positively charged  CATION (E + ) 11 p + 11 e - 11 p + 10 e - Na atom Na + ion L.Pillay 2010

19. 19 When an atom or molecule gains electrons it becomes negatively charged  ANION (E - ). Generally, metal atoms tend to lose electrons (forms cations) and non-metal atoms gain electrons (forms anions). 17 p + 18 e - 17 p + 17 e - Cl atomCl - ion L.Pillay 2010

20. 20 COMMON CATIONS

21. 21 COMMON ANIONS

22. Ionic compounds Composed of nonmetal and metal Cations and anions attract each other to form a neutral compound NAMES: Name of metal (cation) written first If metal has more than one common charge, write the charge in roman numerals in brackets Name of nonmetal (anion) written next with –ide ending FORMULAE: compounds are electrically neutral, the formula of a compound can easily be constructed simply by: -writing value of cation charge as subscript on anion -writing value of anion charge as subscript on cation 22

23. PRACTICE EXAMPLE 23 NaCl K 2 SO 4 Ba(OH) 2 cobalt(II) nitrate silver sulfide ferric chloride

24. Oxyanion ClO 4 - perchlorate ion (one more O atom than chlorate) ClO 3 - chlorate ion (one more O atom than chlorite) ClO 2 - chlorite ion (one more O atom than hypochlorite) ClO - hypochlorite ion Acids -acids containing anions whose names end in -ide are named by changing the -ide ending to -ic, adding the prefix hydro- to this anion name, and then following with the word acid - acids containing anions whose names end in -ate/-ite are named by changing the -ate ending to -ic or the -ite ending to -ous and then adding the word acid 24

25. PRACTICE EXAMPLE 25 AnionCorresponding acid Cl - S2-S2- ClO 4 - ClO 3 - ClO 2 - ClO -

26. MOLECULAR COMPOUNDS Generally composed only of nonmetals Diatomic species includes O 2 N 2, F 2, Br 2, I 2 NAMING: name of element furthest left on periodic table generally written first both elements in same group on periodic table, element with higher Z written first name of 2nd element given the ending –ide Greek prefixes used to indicate number of atoms of each element Greek prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca 1 2 3 4 5 6 7 8 9 10 26

27. PRACTICE EXAMPLE 27 SO 2 PCl 5 N 2 O 3 NF 3 P 4 S 10 silicon tetrabromide

28. STOICHIOMETRY “quantities of substances consumed and produced in chemical reactions” Atoms are neither created or destroyed in a chemical reaction. A chemical equation must have equal numbers of atoms of each element on each side of the arrow. The molecular composition of certain ions must remain the same on each side of the arrow. 28

29. PRACTICE EXAMPLE 29 C 2 H 6 + O 2 → CO 2 + H 2 O Al + HCl → AlCl 3 + H 2

30. MOLE & AVOGADRO’S NUMBER Number of atoms/molecules/ions represented as mole amounts Avogadro’s number: N A = 6.022 X 10 23 1 mol 12 C atoms = 6.022 X 10 23 12 C atoms 1 mol H 2 O molecules = 6.022 X 10 23 H 2 O molecules 1 mol NO 3 - ions = 6.022 X 10 23 NO 3 - ions 30

31. Molar mass “Mass in grams of one mole of a substance” Related to mole amount of a substance by the equation: 31 n: number of moles (in mol) m: mass (in grams) MM: molar mass (in grams per mole) m n MM

32. PRACTICE EXAMPLE 32 How many oxygen atoms are in 1.50 mol of sodium carbonate?

33. EMPIRICAL AND MOLECULAR FORMULA “Ratio of atoms of each element in a compound” 33 Mass % elements Grams of each element Moles of each element Empirical formula Assume Use molar Calculate 100g mass mole ratio sample

34. PRACTICE EXAMPLE 34 Determine the empirical formula of a compound with 10.4% C, 27.8% S and 61.8% Cl.

35. 35