1. What is Chemistry REVIEW CHAPTERS 1, 2, 3 and 10 (part)
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. 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
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
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 = 4 Sig fig
2. Interior zeros are significant 505 = 3 sig fig
3. Trailing zeros after a decimal are significant = 4 sig fig
4. Leading zeros are not significant = 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. 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 fewest decimal 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
Prefix
Symbol
Decimal
Equivalent
Power of 10
mega- M
1,000,000
Base x 106
kilo- k
1,000
Base x 103
deci- d
0.1
Base x 10-1
centi- c
0.01
Base x 10-2
milli- m
0.001
Base x 10-3
micro-
m or mc
Base x 10-6
nano- n
Base x 10-9

7. 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
Units Raised to a Power
A. 1 inch = 2.54 cm so 1 inch3 = 2.54 cm3 = 16.4 cm3
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. And Calorimetry: Measuring Quantities of Heat
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. EXAM 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