1. Standards for Measurement Preparation for College Chemistry Columbia University Department of Chemistry

2. The Scientific Method Observations Theory (Model) Hypothesis Experiment Laws (analysis) (explanation) (measurement) http://antoine.fsu.umd.edu/chem/senese/101/intro/scimethod-quiz.shtml

3. Measurement and Interpretations 12346578 Diameter = 2.5 cm Area =  x r 2 = 8.04 cm 2 Direct Measurement Interpretation Step Art of Chemical Measurement: Recognize what can be measured directly Devise a way to obtain the desired information from measurement data Radius = Diameter/2

4. ms mV xy +++ ++ + --- -- - --- -- - ++ + ++ + Resting Potential (Squid Experiment) -65 40 Na + K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+

5. Resting Potential

6. Experimentation Resting Potential = -65 mV UNIT NUMERICAL VALUE Measured Data Derived Basic Affected by Uncertainty Accuracy Precision Resolution Noise

7. Significant Figures (Sig. Figs.) The mass of an object weighed on a triple beam balance (precision + 0.1g) is found to be 23.6 g. This quantity contains 3 significant figures, i.e., three experimentally meaningful digits. If the same measurement is made with an analytical balance (precision + 0.0001g), the mass might be 23.5820 g (6 sig. fig.)

8. Evaluating Zero Zero is SIGNIFICANT when: Is between nonzero digits: 61.09 has four sig Figs. Appears at the end of a number that includes a decimal point 0.500 has three sig. Figs.; 1000. has four sig. Figs. Zero is NON SIGNIFICANT when: Appears before the first nonzero digit. 0.0025 has two sig. Figs. Leading Zeros are non significant Appears at the end of a number without a decimal point. 1,000 has one sig. Fig.; 590 has two sig. Figs.

9. Exact Numbers Defined numbers, like 12 inches in a foot, 60 minutes in an hour, 1,000mL in one liter. Exact numbers have an infinite number of sig. figs. Exact numbers do not limit the number of sig. figs. in a calculation. Numbers that occur in counting operations.

10. Scientific Notation Number written as a factor between 1 and 10 multiplied by 10 raised to a power. Useful to unequivocally designate the significant figures.

11. Multiplication or Division The answer must contain as many significant figures as in the least precise quantity (measurement with least precision). What is the density of a piece of metal weighing 36.123 g with a volume of 13.4 mL? Round off to 7 Drop these three digits ANSWER:

12. Addition or Subtraction Keep only as many digits after the decimal point as there are in the least precise quantity Ex. Add 1.223 g of sugar to 154.5 g of coffee: Total mass = 1.2 g + 154.5 g = 155.7 g

13. Addition or Subtraction Note that the rule for addition and subtraction does not relate to significant figures. The number of significant figures often decreases upon subtraction. Mass beaker + sample = 52.169 g(5 sig. figs.) - Mass empty beaker = 52.120 g (5 sig. figs.) Mass sample = 0.049 g (2 sig figs)

14. PrefixSymbolValuePower exaE 1,000,000,000,000,000,000 10 18 petaP 1,000,000,000,000,000 10 15 teraT 1,000,000,000,000, 10 12 gigaG 1,000,000,000 10 9 megaM 1,000,000 10 6 kilok 1,000 10 3 hectoh 100 10 2 dekada 10 10 1 SI Units Prefixes (Multiples)

15. PrefixSymbolValuePower attoa 0.000000000000000001 10 -18 femtof 0.000000000000001 10 -15 picop 0.000000000001 10 -12 nanon 0.000000001 10 -9 microµ 0.000001 10 -6 millim 0.001 10 -3 centic 0.01 10 -2 decid 0.1 10 -1 SI Units Prefixes (Submultiples)

16. Système International (SI) The Metric System Physical QuantityNameSymbol Lengthmeterm Masskilogramkg Timeseconds Electric currentampereA TemperaturekelvinK Amount of substancemolemol Luminous intensitycandelacd Based on seven DIMENSIONALLY INDEPENDENT quantities

17. Length 1790s: 10-millionth of the distance from the equator to the North Pole along a meridian. Base unit is the meter (m) 1960-1983: 1,650,763.73 wavelengths of the orange-red emision of 18 Kr at standard conditions Since 1983: 1/299,792,458 of the distance traveled by light in 1 second through vacuum. 1889: Distance between two engraved lines on a Platinum- Iridium alloy bar maintained at 0°C in Sevres-France.

18. Length Atomic dimensions 1 km = 10 3 m1 in. = 2.54cm 1 cm = 10 -2 m1 mile = 1.61km 1 mm = 10 -3 m 1 µm = 10 -6 m 1 Å = 10 -10 m 1 nm = 10 -9 m Engineering dimensions

19. Mass Base unit is the kilogram (kg) 1 kg = 10 3 g; 1 mg = 10 -3 g A mass of 1 kg has a terrestrial weight of 9.8 newtons (2.2 lbs) Depending on the precision required and the amount of material, different balances are used:  The Quadruple Beam Balance ( + 10 mg)  The Top Loading Balance ( + 1 mg).  The Analytical Balance ( + 0.1 mg). International prototype: a platinum-Iridium cylinder maintained in Sevres-France.

20. Freezing point of water Boiling point of water 100 180 32 212 Fahrenheit 100 0 Celsius 273.15 373.15 Kelvin Comparison of Temperature Scales K = °C + 273.15 °F = (1.8 x °C ) + 32 °C = (°F - 32) / 1.8

21. Temperature Conversion K = °C + 273.15 °F = (1.8 x °C ) + 32 °C = (°F - 32) / 1.8 Ex. 2.20 Convert 110°F to °C and K °C = (68° – 32°) / 1.8 = 20°C K = 20 + 273 = 293 K

22. Derived Units Physical quantity Name Symbol Definition Frequency Hertz Hz s -1 Force newton N m.kg.s -2 Pressure, stress pascal Pa N.m -2 = m -1.kg.s -2 Energy work, heat joule J N.m = m 2.kg.s -2 Electric charge coulomb C A.s Electromotive Force volt V J.C -1 = m 2.kg.s -3. A -1 Electric Resistance ohm  V.A -1 = m 2.kg.s -3. A -2

23. Measurement of Volume

24. Conversion Factors Two conversion factors

25. Dimensional Analysis (pp. 23 - 24) Read Problem. What needs to be solved for? Write it down Tabulate data given. Label all factors with proper units Determine principles involved and unit relationships Set up the problem deciding for the proper conversion factor Perform mathematical operations Check if the answer is reasonable

26. Simple, One Step Conversions CBS News reported the barometric pressure to be 99.6 kPa. Express this in mm Hg. 760 mm Hg 101.3 kPa Unit given Unit needed Unit given = 747mmHg x Conversion factor : 101.3 kPa = 760 mm Hg pressure (mmHg) = 99.6kPa

27. Simple, One Step Conversions A rainbow trout is measured to be 16.2 in. long. What is the length in cm? length in cm = 16.2 in 2.54 cm 1 in = 41.1 cm x Note the cancellation of units. To convert from centimeters to inches, the conversion factor would be 1 in / 2.54 cm.

28. Multiple Conversion Factors Three sig. figs. A baseball is thrown at 89.6 miles per hour. What is the speed in meters per second? m/s 1 mile = 1.609 km = 1.609 x 10 -3 m; 1 h = 3600 s speed = 89.6 = 40.0 m / s miles hour 1.609 x 10 – 3 m 1 mile 1 h 3600 s Mile/hour m/hour x x

29. Properties of Substances Intensive vs. Extensive: density vs. mass Chemical vs. Physical Density and Solubility

30. Extensive Properties Vary with the amount of material Mass Volume Internal Energy Enthalpy Entropy

31. Intensive Properties Independent of the amount of material Density (mass per unit volume) Temperature (average energy per particle)

32. Chemical vs. Physical Properties Chemical Properties Physical Properties Molecules or ions undergo a change in structure or composition Can be studied without a change in structure or composition

33. Density: Conversion factor mass volume An empty flask weighs 22.138 g. You pipet 5.00 mL of octane into the flask. The total mass is 25.598 g. What is the density? What is the volume occupied by ten grams of octane? 1 mL 0.692g V = 10.00g x = 14.5 mL Octane amount in g: 25.598 -22.138 3.460g 5.00 mL d = = 0.692g / mL

34. Solubility Expressed as grams of solute per 100 g of solvent in the CRC (Chemical Rubber Company) Chemistry and Physics Handbook. For lead nitrate in aqueous solution: Solubility (g/100g water)T (°C) 5010 140100

35. Solubility Conversion factor (from table) How much water is required to dissolve 80 g of lead nitrate at 100°C? Mass water = 80g lead nitrate 100g water 140g lead nitrate = 57g water x

36. Cool the solution to 10 ° C. How much lead nitrate remains in solution? Mass of lead nitrate = 57g water 50g lead nitrate 100g water = 28g lead nitrate 28g lead nitrate remain in solution 80g – 28g = 52g lead nitrate crystallizes out of solution x 80g lead nitrate were initially in solution

37. Classification of Matter Elements Compounds Mixtures

38. States of Matter Solid Liquid Gas Plasma

39. Matter Classification Matter Mixtures Pure Substances Heterogeneous More than one phase Homogeneous One Phase (Solutions) Elements Compounds

40. Elements’ Distribution (earth, sea, atmosphere) ElementMass %ElementMass % Oxygen49.20Chlorine0.19 Silicon25.67Phosphorus0.11 Aluminum7.50Manganese0.09 Iron4.71Carbon0.08 Calcium3.39Sulfur0.06 Sodium2.63Barium0.04 Potassium2.40Nitrogen0.03 Magnesium1.93Fluorine0.03 Hydrogen0.87 Titanium0.58All others0.47

41. Average Elemental Composition of Human Body ElementMass % Oxygen65.0 Carbon18.0 Hydrogen10.0 Nitrogen3.0 Calcium2.0 Phosphorus1.0 Traces of other elements 1.0

42. Metalloids Ag Na Li Fr Cs Rb K Be Ca Mg Ra Ba Sr Sc La Y Ac Ti Hf Zr Rf V Ta Nb Ha Cr W Mo Sg Mn Re Tc Bh Fe Os Ru Hs Co Ir Rh Mt Ni Pt Pd Cu Au Zn Hg Cd Ga Al Tl In Pb Sn BiPo B Ge Si As SbTe At H CN P O Se S F Br Cl I Ne Kr Ar Rn Xe HeCe Th Pr Pa Nd U Pm Np Sm Pu Eu Am Gd Cm Tb Bk Dy Cf Ho Es Er Fm Tmi Md Yb No Lu Lr Inner-Transition Metals Transition Metals Main-Group Elements Lantanides Actinides Metals, Nonmetals, 47 11 3 87 55 37 19 4 20 12 88 56 38 21 57 39 89 22 72 40 104 23 73 41 105 24 74 42 106 25 75 43 107 26 76 44 108 27 77 45 109 28 78 46 29 79 30 80 48 31 13 8l 49 82 50 8384 5 32 14 33 5152 85 1 67 15 8 34 16 9 35 17 53 10 36 18 86 54 2 58 90 59 91 60 92 61 93 62 94 63 95 64 96 65 97 66 98 67 99 68 100 69 101 70 102 71 103

43. Elements that Exist as Diatomic Molecules ElementSymbol Molecular Formula Normal State HydrogenHH2H2 Colorless gas NitrogenNN2N2 Colorless gas OxygenOO2O2 Colorless gas FluorineFF2F2 Pale yellow gas ChlorineClCl 2 Yellow-green gas BromineBrBr 2 Reddish-brown liquid IodineII2I2 Bluish-black solid

44. Depending upon Bonding type Compounds Ionic (Coulombic forces) Molecular (Covalent bonds) Molecules Cations Anions

45. Compounds Contain two or more elements with fixed mass percents Sodium chloride: 39.34% Na 60.66% Cl Glucose: 40.00% C 6.71% H 53.29% O Covalent: Ionic:

46. Information in a Chemical Formula Ca(NO 3 ) 2 Nitrate group: two nitrate groups Per each calcium atom Calcium atom Total elements: 1 Ca 2 N 6 O

47. Allotropic Forms (Allotropes) Carbon Graphite Diamond Nanotubes Buckyballs (C 60 )

49. Energy Heat: Quantitative Measurement Energy in Chemical Changes

50. Radiant (light) Thermal (heat) Chemical Electrical Mechanical In any chemical or physical change, energy can be converted from one form to another, but it is neither created nor destroyed Law of Conservation of Energy: Is the energy available but not being used or is it in use? Forms of Energy Types of Energy Kinetic Energy (Motion Energy) Energy (Capacity to do work) Potential Energy (Stored Energy) o Position, o Composition o Condition

51. Heat Energy and Specific Heat Units of Energy: Amount of kinetic energy possessed by a 2kg object moving at a speed of 1m/s. Substituting these values in the equation that defines kinetic energy: Joule : Equivalent to the amount of energy you will feel if you drop 4.4 lb from about 4 in. onto your foot. calorie (cal) : Amount of heat energy needed to raise the temperature of one gram of water by one degree Celsius measured between 14.5 and 15.5°C.

52. 1kcal = 4.3184kJ 1C = 1kcal = 10 3 cal Sprite™ contains 140 C: 1 BTU (British Thermal Unit): 140,000 cal of energy is released when the soft drink is metabolized within the body. Amount of heat needed to raise the temperature of a lb of water one °F 1 cal = 4.3184 J The joule and calorie are rather small units. The large calorie (Cal, C) is used to express the energy content of foods. Units of Energy 1BTU =.818 kcal

53. Specific Heat Amount of heat needed to raise the temperature of 1 g of a substance in a specific physical state by 1°C Units: cal /g °C or J/g °C “Amount of heat needed to raise the temperature of a substance by the same amount depends on the substance” Amount of heat needed to raise the temperature of a given quantity of substance in a specific physical state. Joseph Black (~1750): Heat Capacity and Specific Heat Heat Capacity

54. The specific heat of a substance changes when the physical state of the substance changes 2.1 J / g °C 2. 0 J / g °C 4. 18 J / g °C The higher the specific heat of a substance, the less its temperature will change when it absorbs a given amount of heat. At the beach, sand has a lower specific heat than water, so it heats up while water stays cool. metals heat up quickly, but cool quickly Water (ice) Water (steam) Water (liquid) Ex.

55. Solving problems Heat transferred = mass x Specific heat x ∆T q = m x C s x ∆T 1. Amount of heat energy needed to cause a fixed amount of a substance to undergo a specific temperature change without causing a change of state. 2.Transfer of heat from one body to another. I.Heat always flows from the warmer body to the colder body. II.The heat loss by the warmer body is equal to the heat gained by the colder body. Generalizations:

56. Heat in Chemical Change Potential Energy Diagrams H2OH2O Electrolysis Direct synthesis H2H2 O2O2 + H2OH2O time Potential energy time H2H2 O2O2 +