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. Fluid
Classifying Matter
•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
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 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?)
2.5 Significant Figures in Calculations
A.
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
1.
Do steps in parentheses first
2.
Determine the number of significant figures in intermediate answer
3.
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
Common Prefixes in the
SI System MEMORIZE
Prefix
Symbol
Decimal
Equivalent
Power of 10
1,000,000
Base x 106
1,000
Base x 103
mega-
M
kilo-
k
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
0.000 001
Base x 10-6
nano-
n
0.000 000 001 Base x 10-9
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
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

F - 32
C 
1.8
And
1.8  C  32  F
K  C  273
Calorimetry: Measuring Quantities of Heat
Read definitions of Specific heat (cal/gºC) meaning of it
Specific of heat of water = 1cal/g ºC
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