What is Chemistry

1. REVIEW CHAPTERS 1, 2, 3 and 10 (part) 0 What is Chemistry • Chapter Overview •   An understanding of the history of chemical investigation. • The history of experimentation and scientific inquiry. • 1.1 Science and Technology • 1.2 Matter • What Is Matter? • A. Occupies space and has mass • B. Atom – smallest unit of matter • C. Molecule – atoms joined together • See next slide for classification of matter into Pure substances and mixtures • Classifying Matter According to Its State: Solid, Liquid, and Gas • A. Solid (fixed volume, incompressible) • 1. Crystalline • 2. Amorphous • B. Liquid • 1. Fixed volume • 2. Fluid • C. Gas (lot if empty space) • 1. Compressible • 2. Fluid

2. Classifying Matter

3. 1.3, 1.4 and Inserts section 3.5 and 3.6 • How We Tell Matter Apart: Physical and Chemical Properties • A. Physical property • 1. Observable without changing the identity • 2. Melting point, odor, color • B. Chemical property • 1. Observable only by changing the identity-Chemical reactions • 2. Flammability • How Matter Changes: Physical and Chemical Changes • A. Physical change • 1. Appearance and properties can change • 2. Composition does not change • B. Chemical change • 1. Appearance and properties can change • 2. Composition changes • C. Separation of mixtures through physical changes • 1. Decanting • 2. Distillation • 3. Filtration • 1.5 The Scientific Method: How Chemists Think • Observation – hypothesis – law – theory - experiment • Scientific law (e.g. law of conservation of mass) • Dalton's atomic theory

4. Numerical Side of Chemistry 0 Chapter Overview   A cornerstone of the chemical sciences, the manipulation of numbers and their associated units. Measurement accuracies, significant figures, rounding and scientific notation. • 2.1 and 2.2 Numbers in chemistry–Units and precision and accuracy in reporting it • 2.3 Significant Figures: Writing Numbers to Reflect Position • A. How many digits can I report? How many should I report? • B. Certain digits and estimated digits • C. Counting significant figures • 1. All nonzero digits are significant 1234 = 4 Sig fig • 2. Interior zeros are significant 505 = 3 sig fig • 3. Trailing zeros after a decimal are significant 55.00 = 4 sig fig • 4. Leading zeros are not significant 0.012 = 2 sig fig • 5. Zeros at the end of a number, without a decimal point, are ambiguous 150 = ambiguous • D. Exact numbers • 2.4 Scientific Notation: Writing Big and Small Numbers • A. Shorthand notation for numbers • B. Two main pieces: decimal and power-of-10 exponent • C. Measured value does not change, just how you report it (550.6 to 1 sig fig?)

5. 0 • 2.5 Significant Figures in Calculations • Multiplication and division: • Result carries as many significant digits as the factor with the fewest significant digits • B. Rounding • 1. If leftmost dropped digit is 4 or less, round down (leave it same) • 2. If leftmost dropped digit is 5 or higher, round up (increment it by 1) • C. Addition and Subtraction • Result carries as many decimal places as the quantity with the fewestdecimal places • D. Calculations Involving Both Multiplication/Division and Addition/Subtraction • Do steps in parentheses first • Determine the number of significant figures in intermediate answer • Do remaining steps • 2.6 The Basic Units of Measurement • A. English, metric, SI • B. SI Units (Mass – kg; Length – m; Time – sec) • C. Prefix Multipliers • milli (m) 0.001 • centi (c) 0.01 • kilo (k) 1000 • Mega (M) 1,000,000 • D. Derived Units • 1. Area – cm2 • 2. Volume – cm3 or L

6. Common Prefixes in the SI System MEMORIZE 0

7. 0 • 2.8 Converting from One Unit to Another (UNIT 1 to UNIT 2) • A. Units are important, most numbers get one • B. Include units in all calculations • C. Conversion factors (Unit you have comes in the bottom, unit you want comes in the TOP) • Unit 1 X Unit 2 = UNIT 2 • Unit 1 • D. Significant figure of the final answer depends on UNIT 1 given in problem NOT the sig. fig of the conversion factor • Solving Multistep Conversion Problems • A. Understand where you are going first • B. Not all calculations can be done in one step • 2.8 Units Raised to a Power • A. 1 inch = 2.54 cm so 1 inch3 = 2.54 cm3 = 16.4 cm3 • 2.7 Density (D= Mass/Volume) • A. Mass per unit volume (D= Mass/Volume) • B. Derived unit (Volume = Mass/Density) (Mass = Density X Volume) • C. Can be used as a conversion factor between mass and volume • Numerical Problem Solving Strategies and the Solution Map • A. Come up with a plan before you pull out your calculator • B. Use the units to guide your plan

8. 0 2.10 Energy A. Energy cannot be created or destroyed B. Units of energy 1. Joule (J) 2. calorie (cal) (1 cal = 4.184J) 3. Calorie (Cal) (1Cal = 1000cal = 1kcal) 4. Kilowatt-hour (kWh)- - Will not be used in CH19 Temperature: Random Molecular and Atomic Motion A. Fahrenheit (F) B. Celsius (C) C. Kelvin (K) Conversions And Calorimetry: Measuring Quantities of Heat Read definitions of Specific heat (cal/gºC) meaning of it Specific of heat of water = 1cal/g ºC

9. 0 EXAM 1- 100 POINTS Feb. 21, 2012, 6:00 – 7:30 pm, Room T-109 Multiple Choice Fill in the Blanks scientific notation, chemical change, physical change, chemical and physical properties, Pure substances and mixtures Show calculations for partial/full credit Simple Conversions –show all work round to correct to correct Sig. Fig, scientific notation when indicated Density, calories, temperature Show all calculations/work Simple Calorimetry problems for the first test