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Chemistry 100(02) Fall 2001

Chemistry 100(02) 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: September 26 Test 2: Chapters 3, 4: October 31 Test 3: Chapters 5, 6: November 14

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Chemistry 100(02) Fall 2001

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  1. Chemistry 100(02) 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: September 26 Test 2: Chapters 3, 4: October 31 Test 3: Chapters 5, 6: November 14 Make-up, Comprehensive, November 15

  2. Chemistry 100(04) 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 30 Test 3: Chapters 5, 6: 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 30 Test 3: Chapters 5, 6: 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

  6. Chemistry • “The study of matter and the changes it undergoes.” • Major divisions • InorganicCompounds of elements other than carbon • OrganicCompounds of carbon • Biochemistry Compounds of living matter • Physical Theory and concepts • Analytical Methods of analysis

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

  8. Matter Pure Substance Mixture Element Compound Homogeneous Heterogeneous Hemoglobin Plasma Blood Iron Classification of matter

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

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

  11. 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

  12. Pure substances • Element • Cannot be converted to a simpler form by a chemical reaction. • Example hydrogen and oxygen • Compound • Combination of two or more elements in a definite, reproducible way. • Example water - H2O • Both elements and compounds have characteristic properties such as color, boiling point and reactivity

  13. Pure substances • The properties of a compound and the elements it is made of can differ greatly. • Formula BP density Other • Hydrogen H2 -253 0.90 Flammable • Oxygen O2 -297 1.14 Supports combustion • Water H2O 100 1000 Not flammable

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

  15. Extensive and intensive properties • Extensive properties • Depend on the quantity of sample measured. • Example - mass and volume of a sample. • Intensive properties • Independent of the sample size. • Properties that are often characteristic of the substance being measured. • Examples - density, melting and boiling points.

  16. Physical properties • Properties that do not involve substances changing into another substance. • Examples • color density • odor melting point • taste boiling point • feel compressibility

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

  18. Example • Which are chemical or physical changes? • Mulching leaves • Milk turning sour • Making wine • Making ice water • Beer going flat • Leaves changing color

  19. Type of Changes • Physical 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 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.

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

  21. 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.

  22. 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

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

  24. 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.

  25. 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

  26. SI units • SI - System International • Systematic subset of the metric system. • Only uses certain metric units. • Mass kilograms • Length meters • Time seconds • Temperature kelvin • Amount mole • Other SI units are derived from SI base units.

  27. Metric prefixes • Changing the prefix alters the size of a unit. Prefix Symbol Factor mega M 106 1 000 000 kilo k 103 1 000 hecto h 102100 deka da 101 10 base - 100 1 deci d 10-1 0.1 centi c 10-2 0.01 milli m 10-3 0.001

  28. Other Units • Derived Units. Units consisting of more than one one base unit. E.g. g/cm3 • English units. • Still commonly used in the United States. Weight ounce, pound, ton Length inch, foot, yard, mile Volume cup, pint, quart, gallon • Not often used in scientific work. • Very confusing and difficult to keep track of the conversions needed.

  29. Converting units • Factor 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 Methoddimensional(Unit) Analysis

  30. mg g g kg 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

  31. 10-3 g 1 mg 1 g 10-6 g ( ) ( ) 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

  32. Common conversion factors • English Factor • 1 gallon = 4 quarts 4 qt/gal • 1 mile = 5280 feet 5280 ft/mile • 1 ton = 2000 pounds 2000 lb/ton • Common English to Metric conversions Factor • 1 liter = 1.057 quarts 1.057 qt/L • 1 kilogram = 2.2 pounds 2.2 lb/kg • 1 meter = 1.094 yards 1.094 yd/m • 1 inch = 2.54 cm 2.54 cm/inch

  33. .Speed of light is 3.00 x 108 m s-1 . Convert the speed of light to miles per year (1 mile = 1.61 km).

  34. Measurement • Number Part Uncertainty in Measurement Significant Figures • Exact Measurements • Extensive and Intensive Properties • Density • Measuring Temperature and Volume

  35. Uncertainty in Measurement • All measurements contain some uncertainty. • We make errors • Tools have limits • Uncertainty is measured with • Accuracy How close to the true value • Precision How close to each other

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

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

  38. 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

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

  40. Certain Digits Uncertain Digit 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

  41. Significant • Non-zero digits are always significant. • Any zeros between two significant digits • Trailing zeros in the decimal portion • 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

  42. Examples • 0.00341........3 sig. digs. • 1.0040.........5 sig. digs. • 0.00005........1 sig. digs. • 65000.......… 2 sig. digs. 6.5 x 104 • 40300..........3 sig. digs. • 200300.........4 sig. digs. 2.003 x 105

  43. Significant figures: Rules for zeros • Leading zeros are notsignificant. 0.421 - three significant figures Leading zero Captive zeros are significant. 4012 - four significant figures Captive zero Trailing zeros are significant. 114.20 - five significant figures Trailing zero

  44. Significant figures • Zeros are what will give you a headache! • They are used/misused all of the time. • Example • The press might report that the federal deficit is three trillion dollars. What did they mean? $3 x 1012 or $3,000,000,000,000.00

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