Matter and
Measurement
Chapter 1
What is Chemistry?
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The science that deals with matter and the
changes that matter undergoes
Sometimes called the central science
because so many things involve chemistry
Scientific Problem Solving
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Involves three basic steps
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State the problem and make observations
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Observations can be quantitative (involving
numbers) or qualitative (not involving numbers)
Formulate a possible solution (hypothesis)
Perform experiments to test the hypothesis

Results and observations from experiments lead to
modification of the hypothesis and more
experiments
Scientific Problem Solving

After several experiments the hypothesis may
become a theory

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Theory – gives a universally accepted explanation of
the problem
Theories should be challenged and reviewed as and
when new data and scientific evidence comes to light
Theories are different than laws

Laws state what general behavior is observed to
occur naturally

Example – law of conservation of mass exists since it has
been consistently observed that during all chemical changes
mass remains unchanged
Measurement


Allow the determination of some of the
quantitative properties of a substance
Example – mass and density
Scientific Notation

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Measurements and calculations in
chemistry often require the use of very
large numbers or very small numbers
To make handling them easier we express
them in scientific notation
Scientific Notation

All numbers expressed in this notation are
represented by a number between 1 and
10 multiplied by 10 raised to a power
Scientific Notation

The number of places the decimal point has moved
determines the power of 10

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If a decimal point moved to the left then the power is positive
If a decimal point moved to the right then the power is
negative
Examples
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42000.0
24500
0.222
0.000985
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4.2 x 104
2.45 x 104
2.22 x 10-1
9.85 x 10-4
SI Units

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Units tell the scale that is being used for
measurement
Common units
Base Quantity
Name of Unit
Mass
Kilogram
Length
Meter
Time
Second
Amount of a substance Mole
Temperature
Kelvin
Symbol
kg
m
s
mol
K
SI Units

Common prefixes
Prefix
Symbol
Giga
G
Mega
M
Kilo
k
Deci
d
Centi
c
Milli
m
Micro
m
Nano
n
Meaning
109
106
103
10-1
10-2
10-3
10-6
10-9
Uncertainty, Significant Figures and
Rounding

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When reading the scale on a piece of lab
equipment there is always some degree of
uncertainty in recording the measurement
The reading will often fall between two
divisions on the scale and you have to
estimate to record the final digit
The estimated final digit is said to be
uncertain and is indicated by using a +/-
Uncertainty, Significant Figures and
Rounding

Determining the number of significant figures
present in a number

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Any non-zero integers are always counted as sig figs
Leading zeros are those that precede all of the nonzero digits and are never counted as sig figs
Captive zeros are those that fall between non-zero
digits and are always counted as sig figs
Trailing zeros are those at the end of a number and
are only counted if the number is written with a
decimal point
Exact numbers have an unlimited number of sig figs
In scientific notation, the 10x part of the number is
never counted as significant
Uncertainty, Significant Figures and
Rounding

Examples
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250.7
0.00077
1024
4.7 x 10-5
34000000
4
2
4
2
2
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500.0
0.230970
0.03400
0.34030
26
3
5
4
5
2
Uncertainty, Significant Figures and
Rounding

Determining the correct number of
significant figures to be shown as the
result of a calculation

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When multiplying or dividing – limit the
answer to the same number of sig figs that
appear in the original data with the fewest
number of sig figs
When adding or subtracting – limit the answer
to the same number of decimal places that
appear in the original data with the fewest
number of decimal places
Uncertainty, Significant Figures and
Rounding

Examples
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34.5 x 23.46
123/3
2.61 x 10-1 x 356
21.78 + 45.86
23.888897 – 11.2
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Answers
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809
40
92.9
67.64
12.7
Temperature

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There are 3 scales of temperature that you may
come across in chemistry – Celsius, Fahrenheit
and Kelvin
The following conversion factors will be useful
Celsius to Kelvin K = oC + 273
Kelvin to Celcius oC = K – 273
Celcius to Fahrenheit oF = 1.8(oC) + 32
Fahrenheit to Celcius oC = oF -32
1.8
Derived Units

Density – the ratio of the mass of an
object to its volume

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Density = mass
volume
Particularly useful when dealing with liquids
and need to find a mass
Converting Units

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One unit can be converted into another by
using a conversion factor (fraction)
Basic process

(Unit a)(conversion factor) = Unit b
Converting Units

A conversion factor is derived from the
equivalence statement of two units

Example – 1.00 inch = 2.54 cm can be written in two
ways

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Or
1.00 inch
2.54 cm
The correct choice is the one that allows the
cancellation of the unwanted units

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2.54 cm
1.00 inch
Example – convert 9.00 inches to cm
Example – convert 5.00 cm to inches
This is called the factor-labeling method or
dimensional analysis
Examples
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Convert the following
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30 m to miles (1 mile = 1760 yd)
206 miles to m (1 m = 1.094 yd)
34 lb to kg (1 kg = 2.205 lb)
38 K to oF
1390 oC to K
3000 oC to oF
2500 m to km
500mg to g
States of Matter

All matter has 2 distinct characteristics

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Has mass
Occupies space
Solids
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Definite shape and volume
Particles are packed tightly together and
vibrate gently around fixed positions
States of Matter
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Liquids
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No shape of their own, take the shape of their
container
Definite volume
Particles are free to move
Gases
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No definite shape or volume
Particles spread apart filling all the space of
the container available to them
Physical and Chemical Properties
and Changes

All matter exhibits physical and chemical
properties to classify it

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Physical property examples – color, odor,
density, hardness, solubility, melting point and
boiling point
Chemical properties – those exhibited when a
substance reacts with other substances

Examples – reactions with acids and bases,
oxidation/reduction reactions, etc.
Physical and Chemical Properties
and Changes

Changes in which physical or chemical
properties are altered are considered
physical or chemical changes
Physical Change

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When some aspect of the physical state of
matter is altered, but the chemical
composition stays the same
Most common physical changes

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Solid  liquid (melting)
Liquid  gas (boiling or evaporation)
Gas  liquid (condensing)
Gas  solid (deposition)
Liquid  solid (freezing, solidifying or
crystallizing)
Chemical Change
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Often called a chemical reaction
Atoms of a substance are rearranged to
form new substances
Require that the new substance or
substances formed have a different
chemical composition than the original
substance or substances
Often accompanied by color changes
and/or heat
Elements, Mixtures and
Compounds

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Element – a substance that cannot be
broken down into other substances by
chemical means; found on the periodic
table
Compound – formed when a number of
elements bond together; always have a
fixed composition; example – water
Elements, Mixtures and
Compounds
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Mixture – has varying composition and is made
up of a number of pure substances
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Homogeneous – uniform in composition throughout a
given sample; example – salt water
Heterogeneous – have separate, distinct regions
within the sample; composition and properties vary
from one part of the mixture to another; example –
chocolate chip cookie
All pure substances are either elements or
compounds
Classification of Matter
Matter
Homogeneous
Mixtures
Heterogeneous
Mixtures
Physical
methods
Physical methods
Pure Substances
Compound
Chemical
methods
Elements
Atoms
Nucleus
Protons
Electrons
Neutrons