REVIEW CHAPTERS 1, 2, 3 and 4
Introduction to Chemistry
1
Chapter 1 Overview
An understanding of the history of chemical investigation.
• The history of experimentation and scientific inquiry.
1.1-1.2 The Scientific Method: How Chemists Think
Observation – hypothesis – law – theory – experiment
Scientific law (e.g. law of conservation of mass
Dalton's atomic theory)
1.3 Matter
What Is Matter?
A. Occupies space and has mass
B. Atom – smallest unit of matter
C. Molecule – atoms joined together
1.4 Classifying Matter According to Its State: Solid, Liquid, and Gas
A. Solid (fixed volume, incompressible) strongest attractive forces
1. Crystalline
2. Amorphous
B. Liquid- Medium inter molecular forces of attraction
1. Fixed volume
2. Fluid
C. Gas (lot if empty space)
1. Compressible
2. Fluid
3. No attractive forces
Classifying Matter Section 1.4
Numerical Side of Chemistry
2
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 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?)
2.2, 2.4 Numbers in chemistry–Units and precision and accuracy in reporting it; Uncertainty in
measurement, last digit on the right is uncertain
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 = 2 SF
D. Exact numbers/definition, conversion factors have infinite sig fig.
2.4 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
2.5 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
2.6
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 cm)3 = 16.4 cm3
2.7 Temperature: Random Molecular and Atomic Motion
A. Fahrenheit (F)
B. Celsius (C)
C. Kelvin (K)
Conversions

F - 32
C 
K  C  273
2.8
(1.8  C)  32  F
1.8
Density = Mass/Volume; Remember Mass is in grams, volume in mL or cm3
Unit of density = g/mL
or g/cm3
Atoms and Elements
3.1
3
ELEMENTS A. Atoms make up all matter
B. ~91 different naturally occurring elements
1. Each element is composed of tiny indestructible
particles called atoms
2. All atoms of a given element have the same
mass and other properties that distinguish them from atoms
of other elements
3. Atoms combine in simple, whole-number ratios
to form compounds
3.3, 3.4 Names vs symbols of elements- Memorize List
from web
Metal, non-metals, metalloids
-
3.5 and 3.6 Looking for Patterns: The Periodic Law and the Periodic
Table
A. Mendeleev (1834 - 1907)
B. Periodic law
C. Metals (usually loose electrons to form cations)
D. Nonmetals (usually gains electrons to form anions)
E. Metalloids, also known as semiconductors
F. Individual group names
1. Group 1 – alkali metals
+1 ion only
2. Group 2 – alkali earth metals +2 ion only
3. Group 7 – halogens usually –1 ion
4. Group 8 – noble gases
3.7 Elements which are diatomics N2 O2 F2 Cl2 Br2 I2 and H2
3.8 Ionic compound – It must conatin a metal, Meatl forms + ion
CATIONS
Nonmatels form – charge ions ANIONS
3.9 Number of atoms in a chemical formula
4 Properties of Matter
Section 4.1-4.2
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
Section 4.3
Separation of mixtures through physical changes
1. Decanting
2. Distillation
3. Filtration
Section 4.4 Energy
A. Energy cannot be created or destroyed
B. Units of energy and heat
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 CHE100
Exothermic Process: Heat is Released by a system or process
Heat q = negative; Here products are lower energy than reactants
Endothermic Process: Heat is Absorbed by a system or process
Heat q = positive;
Here products are HIGHER energy than reactants
Section 4.5 Heat, Specific heat
The specific heat of a substance is the quantity of heat required to change the temperature of 1 g of
that substance by 1ºC units J/gºC
Specific heat of water very high
Specific heat of metals are low, hence are good conductors of heat
Heat = mass of substance X Sp. Heat X Change in temp
= M S DT
2 body problem Heat lost by hot object = heat gained by cold object
Read definitions of Specific heat (cal/gºC) meaning of it
Specific of heat of water = 1cal/g ºC = 4.184J/g ºC
EXAM 1- 100 POINTS
10 points Bonus question!!
Part 1 Multiple Choice –Show calculations for partial/full credit
20 questions- chapters (1-4)
Part 2 Fill in the Blanks scientific notation, chemical change,
physical change, chemical and physical properties, Pure substances
and mixtures, Simple Conversions –show all work and Density,
calories problem
Bonus question: