1. Chapter 1 Matter,Measurement, and Problem Solving

2. Classification of Matter
matter is anything that has mass and occupies space
we can classify matter based on whether it’s solid, liquid, or gas
Fixed = keeps shape when placed in a container
Indefinite = takes the shape of the container

3. Solids
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

4. Crystalline vs. Amorphous solids
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

5. Liquids
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

6. Gases
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

7. Classification of Matter by Composition
made of multiple types of particles
samples may show different intensive properties
made of one type of particle
all samples show the same intensive properties

8. Classification of Pure Substances
made of one type of atom (some elements found as multi-atom molecules in nature)
combine together to make compounds
made of one type of molecule, or array of ions
molecules contain 2 or more different kinds of atoms

9. Classification of Mixtures
made of multiple substances, whose presence can be seen
portions of a sample have different composition and properties
made of multiple substances, but appears to be one substance
all portions of a sample have the same composition and properties

10. Separation of Mixtures
separate mixtures based on different physical properties of the components
Physical change
Centrifugation &
Decanting
Density
Evaporation
Volatility
Chromatography
Adherence to a Surface
Filtration
State of Matter (solid/liquid/gas)
Distillation
Boiling Point
Technique
Different Physical Property

11. Changes in Matter Dissolving of Sugar
C12H22O11(s)
C12H22O11(aq)
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

12. Physical change vs. Chemical change
C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(l)
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
Burning propane gas

13. Energy
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

14. Energy kinetic energy is energy of motion
motion of the atoms, molecules, and subatomic particles
potential 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 Energy
energy is neither created nor destroyed. It is converted from one form to another

15. Spontaneous 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

16. Temperature 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

17. Temperature 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

18. Fahrenheit vs. Celsius TF = 1.8 TC + 32 TC = (TF – 32) 1.8
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
TC = (TF – 32)
TF = 1.8 TC + 32
1.8
K = °C

19. Example
The melting point of gallium is 85.6oF. What is this temperature on
Celsius scale
Kelvin scale

20. The 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
Quantity
Unit
Symbol
length
meter m
mass
kilogram kg
time
second s
temperature
kelvin K

21. Common Prefix Multipliers in the SI System
All units in the SI system are related to the standard unit by a power of 10
The power of 10 is indicated by a prefix multiplier
The prefix multipliers are always the same, regardless of the standard unit
Report measurements with a unit that is close to the size of the quantity being measured
Prefix
Symbol
Decimal
Equivalent
Power 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-
m or mc
Base x 10-6
nano- n
Base x 10-9
pico p
Base x 10-12

22. Common Units and Their Equivalents
Mass
1 kilogram (km) =
2.205 pounds (lb)
1 pound (lb)
grams (g)
1 ounce (oz)
28.35 grams (g)
Volume
1 liter (L) =
1000 milliliters (mL)
1000 cubic centimeters (cm3)
1.057 quarts (qt)
1 U.S. gallon (gal)
3.785 liters (L)
Length
1 kilometer (km) =
mile (mi)
1 meter (m)
39.37 inches (in.)
1.094 yards (yd)
1 foot (ft)
30.48 centimeters (cm)
1 inch (in.)
2.54 centimeters (cm) exactly

23. Volume any length unit cubed Derived unit
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)

24. Mass & 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!
Volume vs. Mass of Brass
y = 8.38x 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
18.0
Volume, cm3
Mass, g

25. Density 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

26. A Measurement
the unit tells you what standard you are comparing your object to
the number tells you
what multiple of the standard the object measures
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.

27. Known & Estimated Digits
For the following volume readings, what would be measured values?
E.g
l l l l l cm
What is the length of the line?
1) cm
2) cm
3) cm

28. Uncertainty in Measured Numbers
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

29. Accuracy, Precision, and Significant 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.

30. Examples
Classify each of the following as (1) exact or (2) measured numbers. 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.

31. Accuracy, Precision, and Significant Figures
Rules for counting significant figures (left-to-right):
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).
g 3 sf or 6.61 x 10-3 g
3. Zeros at the end of a number and after the decimal point are always significant.
K 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

32. Rounding Numbers
If the first digit you remove is 4 or less, it and all following digits are dropped from the number
= (4 s.f)
If the digit you remove is 5 or greater, the last digit of the number is increases by 1
= (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 ÷ = 
3 s.f s.f calculator zeros are needed
result to give 3 s.f
Books sometimes use different rules for this one.

33. Significant figures in calculation
When multiplying or dividing
the final answer must have the same number of significant figures as the measurement with the fewest significant figures.
Example:
x = = (rounded)
4 SF SF calculator SF
When adding or subtracting
the final answer must have the same number of decimal places as the measurement with the fewest decimal places.
one decimal place
two decimal places
calculated answer
final answer with one decimal place

34. Both Multiplication/Division and Addition/Subtraction with Significant Figures
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
× ( – ) =
2 dp dp
× = 12
4 sf dp & 2 sf 2 sf

35. Metric Equalities An equality
states 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 = cm
10-2 m = 1cm
1 m = mm
10-3m = 1 mm

36. Conversion Factors A 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 = cm
2.54 cm in.
Arrange conversion factors so given unit cancels
Arrange conversion factor so given unit is on the bottom of the conversion factor

37. 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 unit
Unit x Unit x Unit = Unit 3
Unit 1 Unit 2

38. Example
If a ski pole is 3.0 feet in length, how long is the ski pole in mm?
Convert cm to yard
Convert 9255 cm3 to gallons

39. 1.8 Density Density compares the mass of an object to its volume.
is the mass of a substance divided by its volume.
Density Expression
Density = 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?

40. Examples
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 mL after the metal is added?
A drop of gasoline has a mass of 22.0 mg and a density of g/cm3. What is its volume in Liters?
25.0 mL
33.0 mL
object