1. Matter and Measurements
Chapter 1
Matter and Measurements

2. How to be Successful in Chemistry
Memorize strategies not equations!
Study a lot!
Work ALL the problems.
Self-evaluate after test results.
Make use of Tutorial.
Start a Study Group
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3. Chapter 1: Matter and Measurement
Overview:
The Study of Chemistry
Classifications of Matter
Properties of Matter
Units of Measurement
Uncertainty in Measurement
Dimensional Analysis
Basic Math Concepts (see Appendix A)
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The study of matter and the changes it undergoes
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Matter
Anything that has mass and occupies space
Characterized by physical and chemical properties
Law of the Conservation of Mass - matter is neither created nor destroyed in chemical reactions
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Element
An element is a pure substance composed of one type of atom.
An atom is the smallest particle of an element that retains the chemical properties of the element.
An element is the most basic form of matter under ordinary circumstances
Simplest chemical substance
Only a few elements are found in their free state (nitrogen, oxygen, gold, etc.)
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7. Elements and the Periodic Table
Each element is represented by a name and a symbol. (Periods/groups - alkali metals, alkaline earth metals, halogens, noble gases)
The first letter is always capitalized the second (and third) are never capitalized.
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Compound
A unique substance composed of two or more elements that are chemically combined (i.e. joined intimately, not just mixed together)
Pure compounds have definite compositions and properties
Require complex chemical procedures to separate into simpler substances (elements)
Compounds include water, table salt, sugar, etc
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9. Properties of Substances
Elements and Compounds are pure substances.
Properties describe the particular characteristics of a substance
Pure substances have definite composition and definite, unchanging properties
Physical properties - can be observed without changing the substance
Chemical properties - require that the substance change into another
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Physical States
The three physical states are solid, liquid and gas
solids - have a definite shape and volume
liquid - have a definite volume but not a definite shape
gas - neither a definite volume or shape
A substance exists in a particular physical state under defined conditions
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Phase Changes
Melting point or freezing point
temperature at which a substance changes from solid to liquid
Boiling point or condensation point
temperature at which a substance changes from liquid to gas
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Density
ratio of the mass of a substance to the volume of that mass
usually measure in g/mL for solids and liquids; g/L for gases
also a conversion factor relating the mass of a substance to it’s volume
Specific gravity is the ratio of the mass of a substance to the mass of an equal volume of water
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What’s happening?
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Density Calculation
Equation
d=m/V
Example:
If an object has a mass of 15.0 g and a volume of 10cm3 what’s the objects density?
d = 15.0 g/ 10.0 cm3 = 1.50 g/cm3
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16. Temperature and its Measurement
Temperature - measure of the intensity of the heat of a substance
Thermometer - device to measure temperature
Kelvin - K - SI unit of temperature
Celsius - °C - commonly used unit
Fahrenheit - °F - only used in USA
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17. Relationships between temperature scales
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The Kelvin scale
The idea of negative temperatures is a problem for any mathematical treatment of temperature dependent properties.
It was found that a practical minimum temperature did exist (absolute zero) which has a value of °C
This is defined as 0 K (no degree sign)
The Kelvin degree is the same size as the Celsius degree (K = °C )
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19. Temperature Scale Comparison
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Chemical Properties
Chemical properties - involve how a substance changes into another
Sometimes quite difficult to determine
Some examples are burning (as opposed to boiling) and color changes
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Mixtures
Combinations of two or more substances
Can be separated by exploiting different physical properties (filtration, distillation, crystallization, chromatography)
Have chemical and physical properties that are different from the substances that make them up
The percentages by mass of the components of a mixture can be varied continuously
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Types of Mixtures
Phase - one physical state with distinct boundaries and uniform properties
Heterogeneous mixture - nonuniform mixture containing two or more phases with definite boundaries between the phases (e.g. ice and water)
Homogeneous mixture - same throughout and contains only one phase (substances are mixed at the atomic or molecular level)
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23. Heterogeneous Vs. Homogeneous Mixtures
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24. Heterogeneous Vs. Homogeneous Mixtures
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Solutions
A type of homogeneous mixture
Usually involves a liquid phase, but can be solid-solid, liquid-liquid, solid-liquid, etc.
The pure substances can be in different phases but form a homogeneous mixture (table salt and water, for example)
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Alloys
important solid solutions of two or more metals
dental fillings (silver and mercury)
stainless steel (iron, chromium and nickel)
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27. Mixtures separated by:
Filtration: Mixture consists of a solid and liquid; liquid distributing components between a mobile separated by filtration.
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28. Separation of Mixtures cont.
Distillation: Liquid mixture is boiled; components in the mixture boil off at different temperatures.
Chromatography: Separates mixtures by and stationary phase.
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29. Putting it All Together
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30. Measurements and Units
Measurement - determines the quantity, dimensions or extent of something
1.Consist of two parts
a. a numerical quantity (1.23)
b. a specific unit (meters)
Unit - a definite quantity adapted to as a standard of measurement
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31. Features of Measured Quantities
When we measure a number, there are physical constraints to the measurement
Instruments and scientists are not perfect, so the measurement is not perfect (i. e., it has error)
The error in the measurement is related to the accuracy and the precision of the measurement
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32. Accuracy and Precision
Accuracy – how close the measurement is to the “true” value (of course we have to know what the “true” value is)
Precision - the degree to which the measurement is reproducible
1. expressed through how we write the number – significant figures
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33. Example: Accuracy and Precision
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34. Equations for Precision and Accuracy
Percent Relative error
% RE = Avg. Dev / Avg Value X 100
2. Accuracy
Absolute Error
% AE = (True value-Avg Value) X 100
True Value
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Significant figures
Significant figure – a digit that is either reliably known or estimated
Assume that the last digit is uncertain
ex: in the number 1.23, there are three significant figures and assume that the last digit, the 3, is uncertain
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Reading Scales
Estimate measurement to one digit beyond the smallest scale division
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37. Counting Significant Figures
Atlantic / Pacific Method
a. Absent Decimal- Start on “atlantic” side of number & cross out all zeroes until 1st nonzero digit is reached, remaining digits are significant
b. Present decimal- start on the “pacific” side of the number & cross out all zeros until the 1st nonzero digit Is reached, remaining digits are significant
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2. Examples:
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Practice
Count the significant figures in each of the following measurements:
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Examples
Round each of the following measurements to 3 significant figures
16.892
14670
1.9220
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Practice
Round each of the following measurements to 2 significant figures
7.711
5.01
13500
0.0985
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Practice
Round each of the following measurements to 2 significant figures
7.711
5.01
13500
0.0985
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43. Significant digits in calculations
Addition and subtraction – answer is expressed to the same number of significant figures as the number in the calculation with the fewest digits to the right of the decimal point
ex = 12.11
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2. In multiplication or division - the answer is expressed to the same number of significant digits as the number with the fewest significant digits ex  = = 7.55
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Calculation Practice =
x 3.10 x = = =
16.2
500,000 / =
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46. Scientific notation -Expressing very large and small numbers
Using scientific Notation - Numbers are expressed with one nonzero digit to the left of the decimal point multiplied by 10 raised to an appropriate power
1. The base is the number with all of the appropriate significant digits
2. The exponent is the power of ten the base is multiplied by
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47. Scientific notation format
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48. Scientific notation and calculators
1. Calculators handle scientific notation by only inputting the exponent, using an EXP or EE key
a. enter the base as you would a regular number, then press EXP or EE, then enter the exponent
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49. Scientific notation and significant figures
1. When using scientific notation the base must be written with the correct number of significant digits
2. All zeroes are significant when using scientific notation
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50. Measurement of mass, length and volume
In the United States, we use a fairly awkward system of measurement for most things - the English system Scientists use the metric and SI systems of units for the measurement of physical quantities
This system using standard units based on very precisely known properties of matter and light
Prefixes are used in from of the units to indicate powers of ten
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SI Units
Measurement
Unit
Symbol
Mass
Kilogram kg
Length
Meter M
Time
Second s
Temperature
Kelvin K
Quantity
Mole
mol
Energy
Joule J
Pressure
Pascal Pa
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SI Prefixes
Prefix
Symbol
Power
tera- T
1012
deci- d
10-1
giga- G
109
centi- c
10-2
mega- M
106
milli- m
10-3
kilo- k
103
micro- 
10-6
hecto- h
102
nano- n
10-9
deca- da
101
pico- p
10-12
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. Base Units
Mass - the quantity of matter that a sample contains
Note that weight is a measure of the attraction of gravity for a sample and it varies depending on the distance of the mass to a planet or moon
Scientists often speak imprecisely of the “weight” of an amount of substance. They really mean mass.
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54. Basic SI units/Derived units
Used to generate new Units
Volume - space a given quantity of matter occupies
Volume - expressed in terms of length - m3
m3 - an inconveniently large volume, so we use liter (L; one cubic decimeter)
We often use a mL (1 cubic centimeter) for more manageable amounts of matter
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55. Converting between units
The standard method to convert between two different units is the factor-label or dimensional analysis method
Dimensional analysis converts a measurement in one unit to another by the use of a conversion factor
Conversion factors are developed from relationships between units
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Conversion factors
Unit factors - factors that relate a quantity in a certain unit to 1 of another unit e.g. 103 m = 1 km
The conversion factor is created by divided both sides by the same quantity
103 m = 1 = 1 km 103 m m
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Dimensional Analysis
Multiplying a quantity in one unit by an appropriate conversion factor converts the number into the new unit
Note that conversion factors are exact relationships
Exact relationships have unlimited precision, so they can be ignored for the purposes of decided the number of significant digits in a calculation
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