CHEMISTRY – Chapter 1 & 2
Matter, Measurements, and
Calculations
Chapter 1 – Section 1
Objectives:
1. Define chemistry
2. List examples of branches of
chemistry
3. Compare and contrast basic research,
applied research, and technological
development
What objects in this room are
related to chemistry?









Plastics
Fabrics
Clothes
Cooking oil
Motor oil
Make-up
Radio
Batteries
Computers
Chemistry in our daily lives.







Antibiotics
Food
Transportation
Sports
Farming
Military
Industry
Chemistry
 Study of the composition and
properties of matter and the changes
that matter undergoes
- What something is made of
- What is the internal arrangement
Chemical
 Any substance that has a definite
composition
6 Main Branches of Chemistry
1.
2.
3.
4.
5.
Organic – substances containing C
Inorganic – substances other than organic
Biochemistry – living things
Physical chemistry – changes of matter
Analytical chemistry – id components of
materials
6. Theoretical chemistry – use math and
computers to understand chemical behavior
All branches involve some type of
research.
Basic research – to increase knowledge
- how and why
Applied research – to solve problems
Technological development – production
and use of products
- lags behind discoveries
- application of knowledge
Review and Assignment
1. Define chemistry
2. List examples of branches of
chemistry
3. Compare and contrast basic research,
applied research, and technological
development
Assignment: WS 1-1
Quiz
1. Name two branches of chemistry.
2. List two ways that chemistry affects
our daily lives.
3. Definition of chemistry.
Chapter 1 - Matter
Chapter 1 – Section 2
Objectives:
1. Distinguish between a mixture and a pure
substance.
2. Define what matter is.
Matter
-
anything that has mass and occupies space
includes almost everything
exceptions are light, heat, and sound
properties are used to measure matter
ex. mass
Mass – measure of quantity of matter
- not affected by temp, location, or any
other factor
Demo.
 Mass vs. matter
 What caused the change in mass?
 Is air matter?
Matter (cont.)
Classified into 2 groups:
1. pure substances
2. mixtures
Pure substance – matter that has the same
properties throughout
ex. element or compound
Pure Substances
Element – substance that cannot be broken down by
ordinary chemical change
- only 1 type of atom
- symbols abbreviated w/1 or 2 letters
- can be an allotrope
allotrope – one of a number of different molecular
forms of an element in the same state
Compound – substance made up of 2 or more elements
chemically combined
- can be broken down by chemical change
- more than 1 type of atom
Compounds
1. Elements that make up a compound are
combined in definite proportion by mass
ex. 100 g water has 11.2 g H and 88.8 g of O
2. Chemical and physical properties of compound
differ from those of its parts
ex. water is liquid, H and O are gases
3. Compounds can be formed from simpler
substances by chem change and can be
broken down into simpler substances
example
100 of water has 11.2 g H and 88.8 g O
How many g of H is in a 120g sample of
water?
120 g water| 11.2 g H = 13.4 g H
| 100 g water
Mixtures
- contain 2 or more substances that have
-
different properties
- vary in composition and properties from
sample to sample
ex. rock, wood, salt water
Not chemically combined
-
Can be separated by simple physical means
-
ie. filtration, evaporation, distillation
Formation of Mixtures
A mixture can be formed 3 ways:
1. Element mixed w/1 or more other
elements
ex. carbon w/sulfur
2. Compound mixed w/ 1 or more other
compounds
ex. salt w/sugar
3. 1 or more elements mixed w/1 or more
compounds
ex. sulfur w/sugar
Characteristics of Mixtures
- retain properties of each of its parts
ex. iron and sulfur
- iron remains magnetic
- composition can vary widely
- can be homogeneous or heterogeneous
Types of mixtures
Homogeneous – uniform composition
throughout
- called solutions
ex. alloys, pop, air, coffee
Heterogeneous – not uniform
throughout
ex. concrete, soil, dry soup, spaghetti
and meat balls
Matter
Pure substance
Element
Compound
Mixture
Homogeneous
Heterogeneous
Review and Assignment
1.
Distinguish between a mixture and a pure
substance.
2. Define what matter is.
Assignment: WS
Chapter 1 – Section 2
Objectives:
1. Distinguish between the physical properties
and chemical properties of matter.
2. Classify changes of matter as physical or
chemical.
3. Explain the gas, liquid, and solid states in
terms of particles.
Properties of Matter
- allow us to distinguish btwn
substances
- characteristics of a substance
- what can be observed
- way that a substance behaves
ex. color, taste, odor, gas, liquid, solid
Properties (cont.)
- can be extensive or intensive
Extensive – d/o amount of matter
ex. volume, weight, mass, and E
Intensive – does not d/o amount of matter
ex. melting point, boiling point, density,
and conductivity
Demonstration
 Properties
- water and glycerin
How do they compare?
- look, feel, weight, flow
- water and salt water
How do they compare?
- conductivity
Physical Properties
 Can be observed or measured w/out
changing the substance
 Can describe the substance
 Odor, taste, hardness, density, melting
point, and boiling point
 Metals – ductile (pulled into wire),
malleable (hammered into sheets), luster
(shine), good conductors
Chemical Properties
 A transformation of a substance into
a different one
 rusting, flammability, tarnishing, new
substance formed
Physical Change
 No new substance is formed
 CHANGE IN PHASE, pounding, grinding,
cutting
 Changes of phase
 When a substance changes phase there
is no change in composition
 Physically different, chemically the same
 Solid, liquid, or gas are the three states
of matter
States of Matter
 Solid – definite volume and shape
 Particles are in fixed positions
 Held w/strong attractive forces
 Liquid – definite volume and no
definite shape
 Takes shape of container
 Particles can move past each other
States of Matter (cont.)
 Gas – neither definite volume nor
definite shape
 Particles move easily and are very far
apart
 Plasma – high temperature state in
which atoms lose their electrons
Chemical Change
 One or more substance is changed to
something new
 Rusting, burning, gas formed, digestion,
heat or light added, explosion, color
change, odor change, water formed
Review and Assignment
1. Distinguish between the physical properties
and chemical properties of matter.
2. Classify changes of matter as physical or
chemical.
3. Explain the gas, liquid, and solid states in
terms of particles.
Assignment: p. 18 and WS
CHEMISTRY –
Chapter 1 – Section 3
Objectives:
1. Perform density calculations.
2. Describe conservation of mass.
Properties of Matter
- E is always involved in both
physical and chemical changes
- Physical are not at noticable
- Chemical are more noticable
- Heat and light are given off
Density
 is a physical property
 is always the same for
a solid substance
 in gases and some
liquids a change in
temperature will
change the density
 increase in
temperature will
decrease density
 D = m/V
Density problem
Use the 5 steps in problem solving to solve
the following problem.
Lead has a mass of 22.7 g and its volume
is 2.00 cm3. What is its density?
m = 22.7 g
V = 2.00 cm3
D = m/V = 22.7 g/2.00 cm3 = 11.4 g/ cm3
Examples
Conservation of Mass
 In reactions matter cannot be created
or destroyed by a chemical change
- mass stays the same, it may just
change form
Density Lab Results
Group 1 –
Group 2 –
Group 3 –
Group 4 –
Group 5 -
Review and Assignment
1. Perform density calculations.
2. Describe conservation of mass.
Assignment: WS and Density lab
Chapter 2 - Sec.1
Objectives:
1. Describe the purpose of the scientific
method.
2. Distinguish between qualitative and
quantitative observations.
3. Describe the steps to making a graph.
4. Distinguish between inversely and
directly proportional relationships.
Scientific Method
- a logical approach to solving problems
1. Make observations
-
observe your surroundings
2. State the problem
- stated as a question
3. Collect data
4. Form hypothesis
- testable statement
5. Test hypothesis
6. Conclusion
7. Modify hypothesis and retest
Observing
 Involves making measurements and
collecting data
 Data can be qualitative or quantitative
Qualitative – non-numerical information
- descriptive (the sky is blue)
Quantitative – numerical information
- the mass is 25.7 grams
Conclusion
 Can be explained by using models
Model – explanation of how phenomena
occur or how things are related
- visual
- verbal
- mathmatical
Theory
-
models may become part a theory
Theory – broad generalization that
explains facts or phenomena
- must be able to predict results
ex. kinetic-molecular theory
collision theory
Controlled Experiments




Use manipulated variable (independent)
Use responding variable (dependent)
One variable manipulated at a time
Measurements are called data
Making a Graph
 Shows results of an experiment in a
meaningful pattern
 Dependent variable is on the vertical axis
1. Always include a title
2. Determine variables
3. Set up scale
4. Plot points
5. Draw best-fit line
Oxygen obtained from
electrolysis of water
3
10
18
23
28
oxygen (g)
Oxygen Water
2.7
8.9
16
20.4
24.9
Electrolysis of water
25
20
15
10
5
0
28
23
18
10
3
0
5
10
15
water (g)
20
25
Relationships in graphs
 Directly proportional – if dividing one by
the other gives you a constant value
 If one increases so does the other
 If started at point (0,0)
 Inversely proportional – if their product is
constant
 If one increases the other decreases
 Produce a curve
Review and Assignment
1. Describe the purpose of the scientific
method.
2. Distinguish between qualitative and
quantitative observations.
3. Describe the steps to making a graph.
4. Distinguish between inversely and
directly proportional relationships.
Assignment: graphing WS
Quiz
1. List three steps of the scientific
method.
2. List two steps in making a graph.
Chapter 2 Sec.2
Objectives:
1. Distinguish between a quantity, a unit, and
a measurement standard.
2. Name SI units for length, mass, time,
volume, and density.
3. Distinguish between mass and weight.
Measurements
 Basic part of science
 Make observations more meaningful
 Needs to be more than just a number
or quantity
 Need a common system of units
 For consistency
 Measure your desk w/anything you have
available
SI System
-
The International System of Units
Used in all science
A standard
Based on 10
-
Makes it easier to convert from one unit
to another
SI System (continued)
- 7 base units
1. Length – meter (m)
2. Mass – kilogram (kg)
3. Time – second (s)
4. Amount – mole (mol)
5. Temperature – Kelvin (K)
6. Electric Current – ampere (amp)
7. Luminous intensity – candela (cd)
Weight vs. mass
Mass – quantity of matter
- how much space it takes up
- measured w/a balance
- unit kg
Weight – F gravity pulls on matter with
- measured w/spring scale
- unit Newton
On the moon will our weight or mass stay the
same?
SI Prefixes
You must know these.
Kilo- 1000
Deca – 10
Base unit (m, s, L)
Centi – 1/100 or 0.01
Milli – 1/1000 0r 0.001
Derived Units
- combination of base units
Examples
- Area = m2
- Volume = m3
- Density = kg/m3
- Newton = m٠kg/s2
Derived Units (cont.)
Area – determined by multiplying 2
lengths
Volume – determined by multiplying 3
lengths for a solid
- for liquids unit is cm3 or mL
** 1 mL = 1 cm3
Review and Assignment
1. Distinguish between a quantity, a
unit, and a measurement standard.
2. Name SI units for length, mass, time,
volume, and density.
3. Distinguish between mass and
weight.
Assignment: WS 2-2 and p. 42 ~1-3
Quiz
1.
2.
3.
4.
5.
What is the base SI unit for mass?
Kilo = ______
Centi = _____
What is a derived unit?
1 cm3 = _____ mL
Chapter 2 - Sec.3
Objectives:
1. Distinguish between accuracy and
precision.
2. Determine the number of significant
figures in measurements.
3. Perform mathematical operations involving
significant figures.
Accuracy and Precision
 Accuracy – closeness of a
measurement to correct value
 Precision – closeness of a set of
measurements to each other
 Consistency
 Do not have to be correct
 d/o measuring instrument
Bullseyes
Significant Figures
- digits in a measurement that are
know with certainty and one digit that
is estimated
- CALCULATORS DO NOT KEEP TRACK
OF SIGNIFICANT FIGURES
Significant Figure Rules
1. Digits other than zero are ALWAYS significant
2.
3.
4.
5.
ex. 61.4
3 sig. fig.
All zeros at the end of a number and to the right of the
decimal with a # preceding the decimal are ALWAYS sig
ex. 4.7200 km
5 sig. fig.
Zeros used only for spacing are NOT significant
ex. 7000
1 sig. fig.
20
1 sig. fig.
100.0
4 sig. fig.
Zeros between sig. fig are significant
Zeros in front of a non-zero are NOT sig.
- don’t count until you get to 1st non-zero from lf to rt
0.004
1 sig. fig.
0.0009
1 sig. fig.
Significant Figures
1,000 = _____ sig figs
100.0 = _____ sig figs
0.00012340 = _____ sig fig
10.0340 = _____ sig fig
Calculating w/Significant Figures
Addition and Subtraction
- use same # of decimal places as the
measurement w/the least decimal places
ex.
2.098
3 DECIMAL places
+6.2
1 DECIMAL place
8.298
round to 1 Decimal
8.3 is the final answer
Adding and Subtracting
10.0 + 123 = _____
23.456 – 23.0 = _____
100.12 + 56.45 = _____
1,000 + 12.234 = _____
Calculating w/sig. figs (cont.)
Multiplication and Division
- use same # sig. fig. as the
measurement w/the least sig. fig.
ex.
2.38
3 sig. fig
x 9.0
2 sig. fig
21.42 round to 2 sig. Fig
21 is the final answer
Multiplying and Dividing
100.0 x 10 = _____
34.56 x 23.45 = _____
12.045 x 34.008 = _____
50.04 x 23 = _____
Review and Assignment
1. Distinguish between accuracy and
precision.
2. Determine the number of significant
figures in measurements.
3. Perform mathematical operations
involving significant figures.
Assignment: WS 2-6 and sig fig WS
Quiz
How many significant figures are in the
following numbers?
1. 8,000 _____
2. 100.01 _____
3. 0.00056_____
4. 4500.10 _____
5. What is precision?
Chapter 2 - Sec.3 Day 2
Objectives:
1. Perform mathematical operations involving
percent error.
Percent Error
 Observed value – based on lab
measurements
 True value – based on generally accepted
references

Error exists in any measurement
 d/o measurer, instrument, conditions
Percent Error
% error = true value – obs. value x 100
true value
Example
atomic mass of Al = 28.9 g
measured mass = 27.0 g
What is the % error?
28.9 g – 27.0 g x 100 = 7.00 %
28.9 g
Review and Assignment
1. Perform mathematical operations
involving percent error.
Assignment: WS 2-5 and % error WS
Quiz
How many significant figures are in the
following numbers?
1. 8,104 _____
2. 100.01 _____
3. What does % error tell us?
4. What is accuracy?
5. What is precision?
CHEMISTRY –
Chapter 2 Sec.3 Day 3
Objectives:
1. Use dimensional analysis to convert
measurements.
2. Convert measurements into scientific
notation.
3. Perform mathematical operations using
exponents.
Problem Solving Rules
Write down what is known.
- mass = 346 g
volume = 34.6 cm3
2. Write down unknown.
- density = ?
3. Write the equation to use.
D = m/V
4. Fill in knowns.
D = 346 g/34.6 cm3
5. Solve for unknown and label.
D = 200 g/cm3
6. Check your work.
Dimensional Analysis
- use with conversion factors to change from
one unit to another
Steps:
convert 2550 m to km
1. Determine conversion factor
- 1000 m to 1 km
2. Set up T-bars
3. Write given # in first box
Dimensional Analysis (cont.)
4. Write conversion factor in 2nd box
- unit on bottom matches unit of given #
5. Matching labels cancel
- if 1 from conversion factor is on top divide
- if 1 from conversion factor is on bottom multiple
Scientific Notation
- Used to represent very large or very
small numbers
- There are two parts
- Basic form is M x 10n
- M is a number
- n is a number representing how many
places to move the decimal
Scientific Notation (cont.)
 If n is negative, your number is a decimal
 If n is positive, your number is a large
number
Examples:
60,000,000 = 6 x 107
0.000005 = 5 x 10-6
125,000 = 1.25 x 105
Scientific Notation (cont.)
Write the following in scientific notation.
 1,000,000,000
 23,456
 0.0005678
 0.034
 14,239.1
Scientific Notation (cont.)
Write the following in long hand.
1. 1 x 10-9
2. 3.5 x 105
3. 7.123 x 10-3
4. 5 x 102
5. 4.56 x 10-2
Multiplication w/exponents
 Step 1
 Multiply coefficients
 Step 2
 Add exponents
ex. (2 x 102) (2.5 x 105) = 5 x 107
Division w/exponents
 Step 1
 Divide coefficients
 Step 2
 Subtract exponents
ex. (5 x 10-2) (1.0 x 107) = 5 x 10-9
Addition & Subtraction
w/exponents
 All numbers must be written in the
same power of 10
ex. 5.8 x 103 + 2.16 x 104
- change to
0.58 x 104 + 2.16 x 104 = 2.74 x 104
Scientific Notation & sig figs
 All numbers in front of the x 10 are
significant
ex. 2.00 x 102 = 3 sig fig
2 x 102 = 1 sig fig
Scientific Notation & calculators
 5.44 x 107/8.1 x 104
 5.44 (EE or exp) 7 / 8.1 (EE or exp) 4
= 6.7 x 102
Review and Assignment
1. Use dimensional analysis to convert
measurements.
2. Convert measurements into scientific
notation.
3. Perform mathematical operations using
exponents.
Assignment: p. 57 ~ 1-7 and WS