Chapter 1Chapter 1Matter,Measurement, Matter,Measurement,
and Problem and Problem SolvingSolving
Classification of MatterClassification of Matter
matter is anything that has mass and occupies space we can classify matter based on whether it’s solid,
liquid, or gas
State Shape Volume Compress Flow
Solid Fixed Fixed No No
Liquid Indef. Fixed No Yes
Gas Indef. Indef. Yes Yes
• Fixed = keeps shape when placed in a container • Indefinite = takes the shape of the container
SolidsSolids the particles in a solid are
packed close together and are fixed in position◦ though they may vibrate
incompressible the inability of the particles
to move around results in solids retaining their shape and volume when placed in a new container, and prevents the particles from flowing
Crystalline vs. Amorphous Crystalline vs. Amorphous solidssolids
some solids have their particles arranged in an orderly geometric pattern – we call these crystalline solids◦ salt and diamonds
some solids have their particles randomly distributed without any long-range pattern – we call these amorphous solids◦ plastic◦ glass◦ charcoal
LiquidsLiquids
the particles in a liquid are closely packed, but they have some ability to move around
Incompressible take the shape of their container
and to flow however, they don’t have
enough freedom to escape and expand to fill the container
GasesGases in the gas state, the
particles have complete freedom from each other
the particles are constantly flying around, bumping into each other and the container
Compressible
because there is a lot of empty space, the particles can be squeezed closer together – therefore gases are compressible
because the particles are not held in close contact and are moving freely, gases expand to fill and take the shape of their container, and will flow
Classification of MatterClassification of Matterby Compositionby Composition
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1) made of one type of particle
2) all samples show the same intensive properties
1) made of multiple types of particles
2) samples may show different intensive properties
Classification of Pure Classification of Pure SubstancesSubstances
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1) made of one type of atom (some elements found as multi-atom molecules in nature)
2) combine together to make compounds
1) made of one type of molecule, or array of ions
2) molecules contain 2 or more different kinds of atoms
Classification of MixturesClassification of Mixtures
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1) made of multiple substances, but appears to be one substance
2) all portions of a sample have the same composition and properties
1) made of multiple substances, whose presence can be seen
2) portions of a sample have different composition and properties
Separation of MixturesSeparation of Mixtures separate mixtures based on different physical
properties of the components◦ Physical change
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Centrifugation &Decanting
Density
EvaporationVolatility
ChromatographyAdherence to a Surface
FiltrationState of Matter (solid/liquid/gas)
DistillationBoiling Point
TechniqueDifferent Physical Property
Changes in MatterChanges in Matter changes that alter the
state or appearance of the matter without altering the composition are called physical changes
state changes◦ boiling / condensing◦ melting / freezing◦ subliming
Dissolving of Sugar
C12H22O11(s)
C12H22O11(aq)
Physical change vs. Physical change vs. Chemical changeChemical change changes that alter the
composition of the matter are called chemical changes◦ during the chemical
change, the atoms that are present rearrange into new molecules, but all of the original atoms are still present
rusting processes that release
lots of energy burning
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C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(l)
Burning propane gas
EnergyEnergy changes in matter, both physical and chemical, result
in the matter either gaining or releasing energy energy is the capacity to do work work is the action of a force applied across a distance
◦ a force is a push or a pull on an object◦ electrostatic force is the push or pull on objects that
have an electrical charge
EnergyEnergykinetic energy is energy of motion
motion of the atoms, molecules, and subatomic particlespotential energy is energy that is stored in the matter
due to the composition of the matter and its position in the universe
chemical potential energy arises from electrostatic forces between atoms, molecules, and subatomic particles
Law of Conservation of Energyenergy is neither created nor destroyed. It is converted from one form to another
Spontaneous ProcessesSpontaneous Processes
processes in nature tend to occur on their own when the result is material(s) with lower total potential energy
◦ processes that result in materials with higher total potential energy can occur, but generally will not happen without input of energy from an outside source
when a process results in materials with less potential energy at the end than there was at the beginning, the difference in energy is released into the environment
TemperatureTemperature measure of the average amount of kinetic energy
◦ higher temperature = larger average kinetic energy
heat flows from the matter that has high thermal energy into matter that has low thermal energy◦ until they reach the same temperature◦ heat is exchanged through molecular collisions
between the two materials
Temperature ScalesTemperature Scales
Fahrenheit Scale, °F◦ used in the U.S.
Celsius Scale, °C◦ used in all other
countries Kelvin Scale, K
◦ absolute scale no negative numbers
◦ directly proportional to average amount of kinetic energy
◦ 0 K = absolute zero
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Fahrenheit vs. CelsiusFahrenheit vs. Celsius a Celsius degree is 1.8 times larger than a Fahrenheit
degree the standard used for 0° on the Fahrenheit scale is a
lower temperature than the standard used for 0° on the Celsius scale
the size of a “degree” on the Kelvin scale is the same as on the Celsius scale◦ so 1 kelvin is 1.8 times larger than 1°F
the 0 standard on the Kelvin scale is a much lower temperature than on the Celsius scale
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TF = 1.8 TC + 32
TC = (TF – 32)
1.8
K = °C + 273.15
ExampleExample
The melting point of gallium is 85.6oF. What is this temperature on ◦ Celsius scale
◦ Kelvin scale
The Standard UnitsThe Standard Units
Scientists have agreed on a set of international standard units for comparing all our measurements called the SI units◦ Système International = International System
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Quantity Unit Symbol
length meter m
mass kilogram kg
time second s
temperature kelvin K
Common Prefix Multipliers in the Common Prefix Multipliers in the SI SystemSI System
Prefix SymbolDecimal
EquivalentPower of 10
mega- M 1,000,000 Base x 106
kilo- k 1,000 Base x 103
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- or mc 0.000 001 Base x 10-6
nano- n 0.000 000 001 Base x 10-9
pico p 0.000 000 000 001 Base x 10-12
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All units in the SI system are related to the standard unit by a power of 10The power of 10 is indicated by a prefix multiplierThe prefix multipliers are always the same, regardless of the standard unitReport measurements with a unit that is close to the size of the quantity being measured
Common Units and Their Common Units and Their EquivalentsEquivalents
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Volume
1 liter (L) = 1000 milliliters (mL)
1 liter (L) = 1000 cubic centimeters (cm3)
1 liter (L) = 1.057 quarts (qt)
1 U.S. gallon (gal) = 3.785 liters (L)
Mass
1 kilogram (km) = 2.205 pounds (lb)
1 pound (lb) = 453.59 grams (g)
1 ounce (oz) = 28.35 grams (g)
Length
1 kilometer (km) = 0.6214 mile (mi)
1 meter (m) = 39.37 inches (in.)
1 meter (m) = 1.094 yards (yd)
1 foot (ft) = 30.48 centimeters (cm)
1 inch (in.) = 2.54 centimeters (cm) exactly
VolumeVolume Derived unit◦ any length unit cubed
Measure of the amount of space occupied
SI unit = cubic meter (m3)
Commonly measure solid volume in cubic centimeters (cm3)
Commonly measure liquid or gas volume in milliliters (mL)
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Mass & VolumeMass & Volume two main physical properties of matter mass and volume are extensive properties
◦ the value depends on the quantity of matter◦ extensive properties cannot be used to identify what
type of matter something is Large iceberg and small ice cube
even though mass and volume are individual properties, for a given type of matter they are related to each other!
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Volume vs. Mass of Brass y = 8.38x
0
20
40
60
80
100
120
140
160
0.0 2.0 4.0 6.0 8.0 10.0 12.0 14.0 16.0 18.0Volume, cm3
Mas
s, g
DensityDensity◦ Ratio of mass:volume is
an intensive property value independent of
the quantity of matter Solids = g/cm3
◦ 1 cm3 = 1 mL Liquids = g/mL Gases = g/L Volume of a solid can be
determined by water displacement – Archimedes Principle
Density : solids > liquids >>> gases◦ except ice is less dense
than liquid water!
For equal volumes, denser object has larger mass
For equal masses, denser object has smaller volume
Heating an object generally causes it to expand, therefore the density changes with temperature
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Volume
MassDensity
A MeasurementA Measurement the unit tells you what standard you are comparing
your object to the number tells you
1. what multiple of the standard the object measures2. the uncertainty in the measurement
scientific measurements are reported so that every digit written is certain, except the last one which is estimated
If the length is reported as 3.26 cm,
• the digits 3 and 2 are certain (known).
• the final digit, 6, is estimated (uncertain).
• all three digits (2, 7, and 6) are significant, including the estimated digit.
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Known & Estimated DigitsKnown & Estimated Digits
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E.g
l8. . . . l . . . . l9. . . . l . . . . l10. . cm
What is the length of the line?
1) 9.2 cm
2) 9.13 cm
3) 9.19 cm
For the following volume readings, what would be measured values?
Uncertainty in Measured Uncertainty in Measured NumbersNumbers
accuracy is an indication of how close a measurement comes to the actual value of the quantity
precision is an indication of how reproducible a measurement is
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Accuracy, Precision, and Accuracy, Precision, and Significant FiguresSignificant Figures
Significant figures: The number of meaningful digits in a measured or calculated quantity. They come from uncertainty in any measurement.
Generally the last digit in a reported measurement is uncertain (estimated).
Exact numbers and relationships (7 days in a week, 30 students in a class, etc.) effectively have an infinite number of significant figures.
ExamplesExamplesClassify each of the following as (1) exact or (2)
measurednumbers.
A.__Gold melts at 1064 °C.
B.__1 yard = 3 feet
C.__The diameter of a red blood cell is 6 x 10-4 cm.
D.__There are 6 hats on the shelf.
E.__A can of soda contains 355 mL of soda.
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Accuracy, Precision, and Significant Accuracy, Precision, and Significant FiguresFiguresRules for counting significant figures (left-to-right):
1. Zeros in the middle of a number are like any other digit; they are always significant.
◦ 4.803 cm 4 sf
2. Rules for counting significant figures (left-to-right):◦ Zero at the beginning of a number are not significant
(placeholders).
0.00661 g 3 sf or 6.61 x 10-3 g3. Zeros at the end of a number and after the decimal point are
always significant.55.220 K 5 sf
4. Zeros at the end of a number without a written decimal point are ambiguous and should be avoided by using scientific notation
if 150 has 2 sig. figs. then 1.5 x 102
but if 150 has 3 sig. figs. then 1.50 x 102
Rounding NumbersRounding NumbersIf the first digit you remove is 4 or less, it and all following
digits are dropped from the number 5.664 425 = 5.664 (4 s.f)
If the digit you remove is 5 or greater, the last digit of the number is increases by 1 5.664 525 = 5.665 (4 s.f)
Sometimes, a calculator displays a small whole number. To give an answer with the correct number of significant figures, significant zeros may need to be written after the calculator result.
E.g 8.00 ÷ 2.00 = 4 4.00
3 s.f 3 s.f calculator 2 zeros are needed
result to give 3 s.f
Significant figures in Significant figures in calculationcalculation
When multiplying or dividing
• the final answer must have the same number of significant figures as the measurement with the fewest significant figures.
Example:
110.5 x 0.048 = 5.304 = 5.3 (rounded)
4 SF 2 SF calculator 2 SF
When adding or subtracting
• the final answer must have the same number of decimal places as the measurement with the fewest decimal places.
25.2 one decimal place
+ 1.34 two decimal places
26.54 calculated answer
26.5 final answer with one decimal place
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Both Multiplication/Division and Both Multiplication/Division and Addition/Subtraction with Significant Addition/Subtraction with Significant FiguresFigures
when doing different kinds of operations with measurements with significant figures, do whatever is in parentheses first, evaluate the significant figures in the intermediate answer, then do the remaining steps
3.489 × (5.67 – 2.3) =
2 dp 1 dp
3.489 × 3.37 = 12
4 sf 1 dp & 2 sf 2 sf
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Metric EqualitiesMetric Equalities
An equalitystates the same measurement in two different units.can be written using the relationships between two
metric units.
Example: 1 meter is the same as 100 cm and 1000 mm.
1 m = 100 cm10-2 m= 1cm1 m = 1000 mm10-3m = 1 mm
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Conversion FactorsConversion FactorsA conversion factor is
• obtained from an equality.
• E.g Metric – U.S system
Equality: 1 in. = 2.54 cm
• written as a fraction (ratio) with a numerator and denominator.
• inverted to give two conversion factors for every equality.
1 in. = 1 = 2.54 cm 2.54 cm 1 in.
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Arrange conversion factors so given unit cancelsArrange conversion factor so given unit is on the bottom of the conversion factor
unit desiredunitgiven
unit desiredunitgiven
Using Two or More Factors Using Two or More Factors
• Often, two or more conversion factors are required to obtain the unit needed for the answer.
Unit 1 Unit 2 Unit 3
• Additional conversion factors are placed in the setup problem to cancel each preceding unit.
Given unit x factor 1 x factor 2 = needed unitUnit 1 x Unit 2 x Unit 3 = Unit 3
Unit 1 Unit 2
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ExampleExample If a ski pole is 3.0 feet in length, how long is the ski pole
in mm?
Convert 288.0 cm to yard
Convert 9255 cm3 to gallons
1.8 Density1.8 DensityDensity
• compares the mass of an object to its volume.
• is the mass of a substance divided by its volume.
Density ExpressionDensity = mass = g or g = g/cm3 volume mL cm3
Note: 1 mL = 1 cm3
Can we use density as a conversion factor to calculate mass or volume?
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ExamplesExamples What is the density (g/cm3) of 48.0 g of a
metal if the level of water in a graduated cylinder rises from 25.0 mL to 33.0 mL after the metal is added?
A drop of gasoline has a mass of 22.0 mg and a density of 0.754 g/cm3. What is its volume in Liters?
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object
33.0 mL 25.0 mL