1. Chemistry 100(02) Fall 2001 Dr. Upali Siriwardane CTH 311 Phone 257-4941 Office Hours: M, Tu, W, Th, F 9:00-11:00 a.m. Test 1 : Chapters 1, 2: September 26 Test 2: Chapters 3, 4: October 24 Test 3: Chapters 4, 5: November 14 Make-up, Comprehensive, November 15

2. Chemistry 100(04) Fall 2001 Dr. Upali Siriwardane CTH 311 Phone 257-4941 Office Hours: 8:00-9:00, 11:00-12:00 M, W Tu, Th, F 10:00-12:00 a.m. Test 1 : Chapters 1, 2: October 2 Test 2: Chapters 3, 4: October 25 Test 3: Chapters 4, 5: November 13 Make-up, Comprehensive, November 15

3. Chemistry 100(05) Fall 2001 Dr. Upali Siriwardane CTH 311 Phone 257-4941 Office Hours: M, Tu, W, Th, F 9:00-11:00 a.m. Test 1 : Chapters 1, 2: October 2. Test 2: Chapters 3, 4: October 25 Test 3: Chapters 4, 5: November 13 Make-up, Comprehensive, November 15

4. KEY CONCEPTS What is chemistry? Physical & chemical changes. Physical & chemical properties. Categories of matter. Separating Mixtures. Scientific Method. Scientific Measurement Observation Uncertainty. Significant figure. Precision. Accuracy. Significant figures in calculations. Unit Conversions. Temperature Conversions. Unit conversion method. Density Calculations.

5. What is chemistry? Chemistry deals with non-reversible changes of matter. Chemistry explains using atoms and molecules. Chemical Concepts and Models improve your problem solving skills Chemistry is a Central Science

7. Chemistry “The study of matter and the changes it undergoes.” Major divisionsMajor divisions InorganicInorganicCompounds of elements other than carbon OrganicOrganicCompounds of carbon BiochemistryBiochemistryCompounds of living matter PhysicalPhysicalTheory and concepts Analytical AnalyticalMethods of analysis

8. What is Matter Matter: Anything that has a mass and volume. Energy: Manifestations of matter. Matter and Energy is intertwined.

9. Classification of matter Matter Pure Substance Mixture ElementCompoundHomogeneous Heterogeneous Iron HemoglobinPlasmaBlood

10. Mixtures Heterogeneous Homogenous Compounds Atoms Electrons Nucleus Protons Neutrons Pure Substances Elements Hierarchy of Matter

11. Mixtures A combination of two or more pure substances. –Homogeneous –Homogeneous - Uniform composition –Heterogeneous –Heterogeneous - Non-uniform composition Which are homogeneous or heterogeneous? –Blood Urine “T-Bone” steak –Gasoline Twinkie Salad Dressing

12. How do you Separate Mixtures? Flotation: based on density Filtration: Solid- liquid Distillation- Liquid-liquid Magnetic Separation- Magnetic- Chromatography: 1) Paper 2) Column 3) Gas

13. Pure substances ElementElement –Cannot be converted to a simpler form by a chemical reaction. –Example –Examplehydrogen and oxygen CompoundCompound –Combination of two or more elements in a definite, reproducible way. –Example –Examplewater - H 2 O Both elements and compounds have characteristic properties such as color, boiling point and reactivity

14. Pure substances The properties of a compound and the elements it is made of can differ greatly. Formula BP density Other HydrogenH 2 -253 0.90 Flammable OxygenO 2 -297 1.14 Supports combustion Water H 2 O 100 1000 Not flammable

15. Properties of Substances Physical propertiesPhysical properties: Physical properties are descriptions of matter such as color, density, viscosity, boiling point, and melting point. Chemical propertiesChemical properties: Chemical properties relates to the changes of substances making up the matter. For example, corrosiveness, Flammability

16. Extensive and intensive properties Extensive propertiesExtensive properties Depend on the quantity of sample measured. ExampleExample - mass and volume of a sample. Intensive propertiesIntensive properties Independent of the sample size. Properties that are often characteristic of the substance being measured. ExamplesExamples - density, melting and boiling points.

17. Physical properties Properties that do not involve substances changing into another substance. ExamplesExamples –colordensity –odormelting point –tasteboiling point –feelcompressibility

18. Chemical properties Properties that involve substances changing into another substance. Chemical ReactionChemical Reaction - one or more substances are changed into other substances. combustionExample A chemical property of wood is it’s ability to burn - combustion. wood + oxygen carbon dioxide + water + heat Reactants Products The reactants and products are very different.

19. Example Which are chemical or physical changes? Mulching leavesMulching leaves Milk turning sourMilk turning sour Making wineMaking wine Making ice waterMaking ice water Beer going flatBeer going flat Leaves changing colorLeaves changing color

20. Type of Changes Physical changePhysical change: A change in the state of matter. It does not involve a change in the substances. E.g. melting of wax and water. Chemical changeChemical change: A change involving at least one of the substances making the matter. E.g. Electrolysis of water, formation of rust: reaction of iron and oxygen to from iron oxide.

21. Chemical verses Physical change Sodium reacting Iodine changing with chlorine. from a solid to a gas

22. Scientific method All scientific studies follow the same approach to examining a problem. The scientific method requires that we: Make observations Apply logical, organized reasoning to observations made. Form a hypothesis. Reject or confirm that hypothesis through experiments.

23. Scientific Method Scientific Method. A method common to all sciences has Four Basic Steps: a) Experiment b) Data or Results c) Hypothesis d) Further experiments to test hypothesis

24. Develop a theory Do experiments Do more experiments Organize Make hypothesis Try new tests Did hypothesis work? Yes Scientific method Make observations No

26. Measurement Measurements or observations are made using our physical senses or using scientific instruments. 1) Qualitative measurements. Changes that cannot be expressed in terms of a number. 2) Quantitative measurements. expressed in terms of a number and an unit.

29. Units are important 45 000 has little meaning, just a number $45,000 has some meaning - money $45,000/yr more meaning - person’s salary

30. SI units System InternationalSI - System International –Systematic subset of the metric system. Only uses certain metric units. Mass kilograms Lengthmeters Timeseconds Temperature kelvin Amountmole Other SI units are derived from SI base units.

31. Metric prefixes Changing the prefix alters the size of a unit. Prefix Symbol Factor megaM10 6 1 000 000 kilok10 3 1 000 hectoh10 2 100 dekada10 1 10 base-10 0 1 decid10 -1 0.1 centic10 -2 0.01 millim10 -3 0.001

33. Other Units Derived UnitsDerived Units. Units consisting of more than one one base unit. E.g. g/cm 3 English units.English units. –Still commonly used in the United States. Weightounce, pound, ton Lengthinch, foot, yard, mile Volumecup, pint, quart, gallon Not often used in scientific work.Not often used in scientific work. – Very confusing and difficult to keep track of the conversions needed.

34. Converting units Factor label methodFactor label method Regardless of conversion, keeping track of units makes thing come out right Must use conversion factors - The relationship between two units Canceling out units is a way of checking that your calculation is set up right! Other names used Unit Conversion Method dimensional(Unit) Analysis

35. Example. Metric conversion How many milligrams are in a kilogram? 1 kg=1000 g 1 g =1000 mg 1 kg x 1000 x 1000 = 1 000 000 mg kg g mg g

36. Example Creatinine is a substance found in blood. If an analysis of blood serum sample detected 0.58 mg of creatinine, how many micrograms were present?  = 10 -6 = micro 0.580 mg = 580  g 10 -3 g 1 mg ( ) 1  g 10 -6 g ( )

37. Common conversion factors FactorEnglish Factor –1 gallon= 4 quarts 4 qt/gal –1 mile= 5280 feet 5280 ft/mile –1 ton= 2000 pounds 2000 lb/ton FactorCommon English to Metric conversions Factor –1 liter= 1.057 quarts1.057 qt/L –1 kilogram= 2.2 pounds2.2 lb/kg –1 meter= 1.094 yards1.094 yd/m –1 inch= 2.54 cm2.54 cm/inch

38. . Speed of light is 3.00 x 10 8 m s -1. Convert the speed of light to miles per year (1 mile = 1.61 km).

39. Measurement Number PartNumber Part Uncertainty in Measurement Significant Figures Exact MeasurementsExact Measurements Extensive and Intensive PropertiesExtensive and Intensive Properties DensityDensity Measuring Temperature and VolumeMeasuring Temperature and Volume

41. Uncertainty in Measurement All measurements contain some uncertainty. We make errors Tools have limits Uncertainty is measured with –Accuracy –AccuracyHow close to the true value –Precision –PrecisionHow close to each other

42. Accuracy Here the average value would give a good number but the numbers don’t agree. Large random error How close our values agree with the true value.

43. Precision Here the numbers are close together so we have good precision. Poor accuracy. Large systematic error. How well our values agree with each other.

44. Accuracy and precision Predict the effect on accuracy and precision.Predict the effect on accuracy and precision. Instrument not ‘zeroed’ properly Reagents made at wrong concentration Temperature in room varies ‘wildly’ Person running test is not properly trained

45. Types of errors Instrument not ‘zeroed’ properly –Reagents made at wrong concentration –Temperature in room varies ‘wildly’ –Person running test is not properly trained Random Systematic

46. Significant figures Method used to express accuracy and precision. You can’t report numbers better than the method used to measure them. 67.2 units = three significant figures Certain Digits Uncertain Digit

47. Leading zeros Trailing zeros in whole numbers (use scientific notion to avoid confusion. Exact numbers: unit definition has an unlimited number of sig. figs. 1 ft = 12 in Not significant Significant Non-zero digits are always significant. Any zeros between two significant digits Trailing zeros in the decimal portion

48. Examples 0.00341........3 sig. digs. 1.0040.........5 sig. digs. 0.00005........1 sig. digs. 65000.......… 2 sig. digs. 6.5 x 10 4 40300..........3 sig. digs. 200300.........4 sig. digs. 2.003 x 10 5

49. Significant figures: Rules for zeros are notLeading zeros are not significant. 0.421 - three significant figures Leading zero are Captive zeros are significant. 4012 - four significant figures are Trailing zeros are significant. 114.20 - five significant figures Captive zero Trailing zero

50. Significant figures Zeros are what will give you a headache!Zeros are what will give you a headache! They are used/misused all of the time. ExampleExample The press might report that the federal deficit is three trillion dollars. What did they mean? $3 x 10 12 or $3,000,000,000,000.00

51. Significant figures In science, all of our numbers are either measured or exact. Exact Exact - Infinite number of significant figures. Measured Measured - the tool used will tell you the level of significance. Varies based on the tool. ExampleExample Ruler with lines at 1/16” intervals. A balance might be able to measure to the nearest 0.1 grams.

52. Most calculators use scientific notation when the numbers get very large or small. How scientific notation is displayed can vary. It may use x10 n or may be displayed using an E. They usually have an Exp or EE –This is to enter in the exponent. Scientific notation 1.44939 E-2

53. Examples 378 000 3.78 x 10 5 8931.5 8.9315 x 10 3 0.000 593 5.93 x 10 - 4 0.000 000 4 4 x 10 - 7

54. Significant figures and calculations An answer can’t have more significant figures than the quantities used to produce it. –Example How fast did you run if you went 1.0 km in 3.0 minutes? speed = 1.0 km / 3.0 min = 0.33 km / min 0.333333333

55. Significant figure in Calculations Different rules apply in each case Addition and subtraction In multiplication and division The root or power of a measurement Exact Numbers: Numbers coming from definitions such as 12 in = 1 foot. They are not considered in Sig. Fig. Calculations.

56. Significant figures and calculations Addition and subtractionAddition and subtraction Report your answer with the same number of digits to the right of the decimal point as the number having the fewest to start with. 123.45987 g + 234.11 g 357.57 g 805.4 g - 721.67912 g 83.7 g

57. Significant figures and calculations Multiplication and division.Multiplication and division. Report your answer with the same number of digits as the quantity have the smallest number of significant figures. Example. Density of a rectangular solid.Example. Density of a rectangular solid. 25.12 kg / [ (18.5 m) ( 0.2351 m) (2.1m) ] = 2.8 kg / m 3 (2.1 m - only has two significant figures)

58. Rounding off numbers After calculations, you may need to round off. If the first insignificant digit is 5 or more, - you round up If the first insignificant digit is 4 or less, - you round down. If the digit 5 exactly rounded off to a even

60. Temperature Conversions o F -- > o C ; C = 5/9 (F - 32) o C -- > o F ; F =9/5 C + 32 o C -- > K ; K = C + 273.15 Human body temperature is 98.6 o F. Convert this temperature to o C and K scale o C = 5/9 (98.6 - 32) = 5/9 (66.6) = 37.0 o C--> K = 37.0 o C +273.15 = 310.2 K

61. Example. o F to K If the temperature is 75.0 o F, what is it in K? First convert to o C Then convert to K o C = (75.0 o F - 32) 5 9 = 23.9 K= 23.9 o C + 273 = 297

62. Measuring volume VolumeVolume - the amount of space that an object occupies. liter (L)The base metric unit is the liter (L). milliliter (mL)The common unit used in the lab is the milliliter (mL). cm 3One milliliter is exactly equal to one cm 3. SIm 3The derived SI unit for volume is the m 3 which is too large for convenient use.

63. Density Density is an intensive property of a substance based on two extensive properties. Common units are g / cm 3 or g / mL. »g / cm 3 g / cm 3 Air 0.0013Bone1.7 - 2.0 Water 1.0Urine1.01 - 1.03 Gold19.3Gasoline0.66 - 0.69 Density = Mass Volume cm 3 = mL

64. Example. Density calculation What is the density of 5.00 mL of a fluid if it has a mass of 5.23 grams? d = mass / volume d = 5.23 g / 5.00 mL d = 1.05 g / mL What would be the mass of 1.00 liters of this sample?

65. Density Calculations Equation method: Density = mass ÷ volume; d = m/v Factor Label method:14.2 g -- > ? cm3 conversion factor? 2.70 g 1 cm 3 -------- or ------ 1 cm 3 2.70 g 14.2 gx 1 cm 3 --------------------- = 5.26 cm 3 2.70 g

66. Example. Density calculation What would be the mass of 1.00 liters of the fluid sample? The density was 1.05 g/mL. »density = mass / volume –somass = volume x density »mass = 1.00 L x 1000 x 1.05 » = 1.05 x 10 3 g ml L g mL

67. Specific gravity The density of a substance compared to a reference substance. –Specific Gravity = Specific Gravity is unitless. Reference is commonly water at 4 o C. At 4 o C, density = specific gravity. Commonly used to test urine. density of substance density of reference

68. Specific gravity measurement Hydrometer Float height will be based on Specific Gravity.