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WHMIS

WHMIS . workplace hazardous materials information system. all chemicals are treated with. respect. WHMIS has been developed to provide guidelines for of reactive materials . handling, storage and disposal.

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WHMIS

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  1. WHMIS • workplace hazardous materials information system • all chemicals are treated with respect • WHMIS has been developed to provide guidelines for of reactive materials handling, storage and disposal material safety data sheet • MSDS is a

  2. compressed gas corrosive poisonous and infectious material causing immediate and serious toxic effects flammable and combustible

  3. oxidizing material poisonous and infectious causing other toxic effects dangerously reactive material biohazardous infectious material

  4. Science 10 Review A. The Atom •  proton p+ positive charge •  neutron n0 zero charge •  electron e– negative charge • electrons are found in a cloud region around the nucleus • nucleus contains the protons and neutrons which make up most of mass of atom

  5. mass number = # protons + # neutrons # neutrons = mass number - # protons • isotope = atoms with the same # of protons but a different # of neutrons (different mass numbers) eg) carbon-12  carbon-14  6 p+, 6 n0 6 p+, 8 n0

  6. B. Periodic Table • arranged in and groups (columns) periods (rows) • group number = number of outer level (valence) e- • period number = number of energy levels occupied by e- Ex. Na Group number = 1 Period number = 3 This info is helpful for drawing Energy Level Diagrams

  7. C. Energy Level Diagrams • atoms are which means that the electrically neutral # of p+ = # of e- • maximum number of e-: 3rd level = 2nd level = 1st level = 8 e- 8 e- 2 e-

  8. Examples 1 e- 1 e- 8 e- 2 e- 2 e- 11 p+ Na 3 p+ Li

  9. D. Ions • ions are particles or groups of particles that have a net charge (either positive or negative) • neutral atoms are unstable if their valence level is not full • atoms will strive to satisfy the in order to become stable…in other words, they strive to have octetrule a full valence level and do so by giving away or taking e-

  10. metals  give away e- and become positive ions eg) Na+, Ca2+, Fe3+ • non-metals  take e- and become negative ions eg) Cl-, P3-, O2-

  11. Examples 1 e- 8 e- 8 e- 2 e- 2 e- 11 p+ Na+ 11 p+ Na sodium atom sodium ion

  12. 7 e- 8 e- 8 e- 8 e- 2 e- 2 e- 17 p+ Cl– 17 p+ Cl chlorine atom chloride ion

  13. Your Assignment: 1. Draw argon and neon and compare to Na+ and Cl– 2. pgs 1,2 in workbook

  14. E. Elements • metals exist as single atoms eg) Li(s), Cu(s), Hg(l) • nonmetals and hydrogen do not exist as single atoms – flagpole! H2 N2 O2 F2 P4 S8 Cl2 Br2 I2

  15. Try These: • 1. Cu(s) = • 2. O2(g) = • Al(s) = • 4. fluorine gas = • 5. barium = • 6. nitrogen gas = copper oxygen gas aluminum F2(g) Ba(s) N2(g)

  16. F. Ionic Compounds metals + nonmetals or polyatomic ions • eg) K+, Be2+ or metals eg) Fe3+, Fe2+ monovalent multivalent • charges on the ions are the result of taking or giving e- • to go from formula to name: name of first ion, then brackets for charge if multivalent, then name for second ion i.e. first element ( ) second element-ide eg) AlCl3 = aluminum chloride Fe2O3= iron (III) oxide

  17. Try These: 1. Zn3P2 = 2. NaNO3 = 3. NiF3 = 4. MnO2 = 5. Cr2(SO4)3 = zinc phosphide sodium nitrate nickel (III) fluoride manganese (IV) oxide chromium (III) sulphate

  18. to go from name to formula: write the symbol for each ion, then add subscripts to balance charges eg) calcium sulphide = iron (II) hydroxide = CaS Fe(OH)2 Try These: 1. lithium bromide = 2. sodium phosphate = 3. magnesium nitride = 4. ammonium sulphate = 5. calcium phosphate = LiBr Na3PO4 Mg3N2 (NH4)2SO4 Ca3(PO4)2

  19. Hydrated Compounds • ionic compounds containing in their structure water “xH2O” • water is represented by in the formula where is the number of water molecules x • prefixes: mono hexa 1 = 6 = 2 = 7 = 3 = 8 = 4 = 9 = 5 = 10 = di hepta octa tri nona tetra penta deca

  20. to go from name to formula: give the for the first part of the compound, then name the part as ionic name “xH2O” prefix +“hydrate” eg) NaF3H2O = CuSO45H2O = sodium fluoride trihydrate copper (II) sulphate pentahydrate

  21. to go from name to formula: first part is the …look up the symbol for each ion then balance the charges using subscripts, then for the hydrate part…add where is the number given in the prefix same as before “xH2O” x eg) iron (III) nitrate nonahydrate = sodium chlorate tetrahydrate = nickel (II) sulphite heptahydrate = Fe(NO3)39H2O NaClO34H2O NiSO37H2O

  22. Your Assignment: pgs 3,4 in workbook

  23. G. Molecular Compounds nonmetals only e- are shared therefore no ions are formed no charges involved use prefixes in naming

  24. to go from formula to name: name of first element (including prefix if necessary), then name for second element with “ide” ending (including prefix) i.e. ___first element ___second element -ide eg) N2O = CO2 = P4O10 = dinitrogen monoxide carbon dioxide tetraphosphorus decaoxide • to go from name to formula: write the symbol for each element, then use the prefixes to determine the subscripts eg) carbon monoxide = carbon tetrachloride = CO CCl4

  25. remember the memorizers?????? NH3(g)= H2O(l) = H2S(g) = HF, HCl, HBr, HI = CH4(g)= CH3OH(l) = C2H6(g)= C2H5OH(l) = C6H12O6(s)= C12H22O11(s)= O3(g)= H2O2(l)= ammonia water hydrogen sulphide no prefixes methane methanol ethane ethanol glucose sucrose ozone hydrogen peroxide

  26. H. Acids • always have as the state and always have aqueous (aq) hydrogen Rules 1. hydrogen becomes acid 2. hydrogen becomes acid 3. hydrogen becomes acid ____ide hydro___ic ____ate _______ic ____ite ______ous

  27. Try These: 1. hydrogen iodide = 2. hydrogen phosphate = 3. hydrogen nitrite = 4. hydrogen sulphite = hydroiodic acid HI(aq) phosphoric acid H3PO4(aq) nitrous acid HNO2(aq) sulphurous acid H2SO3(aq)

  28. Your Assignment: pgs 5-7 in workbook

  29. I. States • acids – always (aq) • elements – can be (s), (l) or (g)…see periodic table • molecular compounds – can be (s), (l), or (g) - If not in a solution always (s) - If in a solution either (s) or (aq)…look up on the solubility chart • ionic compounds

  30. Try These: 1. NaCH3COO( ) 6. CaCO3( ) 2. BaSO4( ) 7. FeSO4( ) 3. KOH( ) 8. (NH4)2S( ) 4. Pb(NO3)4( ) 9. Pb(SO4)2( ) 5. Hg(CH3COO)2( ) 10. Ca3(PO4)2( ) aq s s aq aq aq aq aq aq s

  31. J. Chemical Reactions • vs. endothermic exothermic • reaction types: 1. hydrocarbon combustion C?H? + O2(g) CO2(g) + H2O(g) eg) CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) 2. simple composition element + element  compound eg) 2 Mg(s) + O2(g) 2 MgO(s)

  32. 3. simple decomposition compound  element + element eg) 2 H2O(l) 2 H2(g) + O2(g) 4. single replacement element + compound  element + compound eg) Cu(s) + 2 AgNO3(aq) 2 Ag(s) + Cu(NO3)2(aq) 5. double replacement compound + compound  compound + compound eg) Pb(NO3)2(aq) + 2 KI(aq) 2 KNO3(aq) + PbI2(s)

  33. Balancing Reactions • law of conservation of matter says that matter cannot be created or destroyed, it can only change forms • we must chemical equations to conserve matter balance CH4 (g) + O2(g)  2 CO2(g) + H2O(g) 2 4 2 C2H6 (g) + O2(g)  7 CO2(g) + H2O(g) 6

  34. Your Assignment: pg 8, 1st half p. 9

  35. Predicting Reactions Try the following: Potassium iodide solution is added to lead (II) nitrate solution. 2 KI(aq) +  + 2 Pb(NO3)2(aq) KNO3(aq) PbI2(s) • NOTE: • SR and DR reactions always happen in solutions so for ionic compounds check solubility table • Composition and decomposition do NOT happen in solutions so ionic compounds are (s)

  36. Predicting: single replacement Copper metal is added to a solution of silver nitrate Cu(s) + 2 AgNO3(aq) 2 Ag(s) + Cu(NO3)2(aq) Chlorine gas is bubbled through a solution of sodium phosphide 6 Cl2(g) + 4 Na3P(aq) P4(s) + 12 NaCl(aq)

  37. Your Assignment: page 11

  38. K. Significant Digits • any digit from is significant 1-9 • are significant eg) trailing zeros 6.3800, 12 000 • are significant eg) “sandwich” zeros 2.04, 1005.002 • are not significant eg) leading zeros 0.0065 • counted objects and constants are not included in sig digs

  39. / : multiply or divide then round answer to the lowest number of sig digs • +/: add or subtract then round answer to the lowest number of decimal places

  40. L. The Mole • it is a = number 6.02 x 1023 “items” 1. Molar Mass • sum of the individual atomic masses for each element in a compound eg) CO2 = Al(OH)3 = Cu(ClO3)2 = 44.01 g/mol 78.01 g/mol 230.45 g/mol

  41. 2. Mole/Mass Calculations n = m M m = nM where: n = m = M= number of moles inmol mass ing molar mass ing/mol

  42. Example 1 How many moles are in 8.06 g of magnesium oxide? n = m M = 8.06 g 40.31 g/mol = 0.1999503 mol = 0.200 mol m = 8.06 g M = 40.31 g/mol n = ?

  43. Example 2 What is the mass of 0.677 mol of potassium sulphide? m = nM = (0.677 mol)(110.27 g/mol) = 74.65…g = 74.7 g n = 0.677 mol M = 110.27 g/mol m = ?

  44. Your Assignment: p. 12 + 13 Your Review Assignment: finish p. 14 - 16

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