1. Section 1- Matter and Change
Unit 1 – Matter and Energy

2. What is chemistry?
 Chemistry:
the study of the composition, structure, and
properties of matter, the process that matter
undergoes, and the energy changes that
accompany these processes.
 Chemicals:
any substance with a definite composition.
Examples: H2O and H2O2

3. Matter: anything with mass and volume.
Weight: Mass:
Force of
Gravity on
an Object
(Measured with a
scale)
Amount of
Matter in an
Object
(Measured with a
balance)

4. Matter and Its Properties
 Atom: Smallest unit of an element…
 Element: A pure substance that CANNOT be
broken down into a simpler substance (one type
of atom)
Example: Cu or He
 Compound: A substance that can be broken
down into a simpler substance (2 or more
elements chemically bonded together)
Example: H2O or H2SO4

5. Physical
Property
Measured without changing the
substance
Intensive
DOES NOT
depend on the
amount
Extensive
DEPENDS
on the
amount
Melting
Point, Density, Freezing
point
Volume, Mass, Energy

6. Chemical Property
Behavior of a substance interacting with
another
Coal Burning Iron Rusting

7. Matter and Its Properties
Physical
Change:
Chemical
Change:
Same substance
(chemically) present
BEFORE and AFTER
the change
Produces
ONE or MORE
new
substances.

8. Matter and Its Properties
Indicators of Chemical Change
1. Precipitate (solid formed)
2. Gas is given off
3. Color change
4. Change in energy (Hot OR
Cold)
5. Light given off (emission)

9. Classification of Matter
Homogeneous
Materials
Elements
Is it the same
throughout?
Is it a
combination of
things?
Can it be
separated in to
smaller
substances?

10. Groups = COLUMNS Periods = ROWS

11. Unit 1: Matter & Energy
Section 2:
Measurements and
Calculations

12. Scientific Method
Theory Publish Results

13. Units of Measure
Qualitative
Descriptive
(Non-numeric)
The fact that
the sky is blue
Quantitative
Measurement
(Numeric)
A sample of
copper ore has
a mass of 25.7
grams

14. Units of Measure
The International System (SI)--1960s
 Commonly used SI units in CHEMISTRY.
1. Length meters (m)
2. Mass kilograms (kg)
3. Time seconds (s)
4. Quantity moles (mol)

15. Units of Measure
 SI units PREFIXES
Prefi
x
Symbo
l
How many in a base unit?
Nano- n 1,000,000,000
Micro- μ 1,000,000
Milli- m 1,000
Centi- c 100
Deci- d 10
Prefix Symbo
l
How many base units?
Kilo- k 1,000
Mega- M 1,000,000
Giga- G 1,000,000,000
Example: 1 m = 100 cm
1 g = 1,000 mg
Example:
1 km = 1000 m
1 mg = 1,000,000 g

16. Units of Measure
 Derived Units: A combination of SI base units.
Examples: Velocity, Density
 Conversion Factors
Math used to relate 2 units that measure
the same quantity (written as a fraction);
Equal to 1
Example 1 m = 1000 mm

17. Units of Measure
 Conversion Factor Practice
Example 1 – Convert 22,000 g to kg.
Example 2 – Convert 0.0290 m to
mm.
Example 3 – How many seconds are
in 3.11 hours?

18. Units of Measure
 Example 1 – Convert 22 000 g to kg
 22000 g x 1 kg = 22000 kg =
1 1000 g 1000
 Example 2 – Convert 0.0290 m to millimeters
 0.0290 m x 1000 mm = 29 mm =
1 1 m 1
 Example 3 – How many seconds in 3.11 hours?
 3.11 hours x 60 min x 60 sec = 11196 sec =
1 1 hr 1 min 1
22 kg
29 mm
11196
sec

19. Using Scientific Measurement
Accuracy
Precision
Closeness of a measurement to
the true or CORRECT value.
How close a set of measurements
are to one another, regardless of
correctness

20. Using Scientific Measurement
 Experiments will always have errors.
(human, mechanical, environmental)
 PERCENT ERROR determines the
accuracy of the experiment. (the lower the
percent, the better)

21. Using Scientific Measurement
 Significant Figures
All digits that occupy places for which
ACTUAL measurement was made, including
the last estimated digit.
See Significant Figure/Scientific Notation Rules Reference Sheet
When taking
measurements,
you ALWAYS read the
instrument to one more
place than is marked. You
estimate that digit.

22. Significant Figures Practice
1) 30 504
 5 sig. figs.
2) 32.001 20
 7 sig. figs.
3) 0.000 123 0
 4 sig figs.
4) 560 000
 2 sig. figs.
5) 200.003
 6 sig. figs.

23. Math w/ Significant Figures
 Put answers in correct Sig. Figs.
1) 2.301 + 6.12 + 4.1158
 12.5368 = 12.54
2) 1500 – 301.05 – 251.223
 947.727 = 948
3) (9.554)(5.10)(1.5)
 73.0881 = 73
4) 44581.2/235.2
 189.545918367 = 189.5

24. Using Scientific Measurement
 Scientific Notation
 How scientists show either BIG or SMALL numbers.
 602,200,000,000,000,000,000,000 = 6.022 x 1023
Hint:
 Positive exponents – Move decimal to the right
 Negative exponents – Move decimal to the left.
 Note: Only the numbers before the scientific notation are
significant.
(ex: 6.022 x 1023 has 4 sig figs)

25. Scientific Notation
1) Put 325 000 000 in scientific notation
 3.25 x 108
2) Put 0.002405 in scientific notation
 2.405 x 10-3
3) Put 7.54 x 104 in regular notation
 75 400
4) Put 5.001 x 10-5 in regular notation
 0.00005001

26. Using Scientific Measurement
Direct proportion Inverse Proportion
As one variable
increase, so
does the other.
As one
variable
increases, the
other
decreases

27. Unit 1: Matter & Energy
Section 3: Energy
Transfer

28. What is Energy?
Energy
Chemical
Energy due to
chemical bonds
Kinetic
Energy due to
motion
Heat (q) is the
total KE in a
sample
Potential
Energy due to
position

29. Energy Units
Name Symbol Info
Joule J SI Unit for energy
calorie c
(lowercase)
Old energy unit
1 calorie = 4.184 J
Calorie C 1000 c = 1 Calorie
Food Calorie

30. Energy Transfer
 Energy in the form of HEAT moves
back and forth between a system and
its surroundings.
 Temperature: Average KE of the random
motion of particles in a substance.

31. Energy Transfer
 Specific Heat Capacity
Amount of HEAT needed to raise the
temperature of 1 gram of a substance 1
Kelvin (K).
 Calorimeter
 A device used to measure energy absorbed
or released during chemical or physical
change.

32. Energy Transfer
 How to calculate specific heat capacity:
q = (m)(ΔT)(Cp)
q = Heat (J)
m= Mass (g)
ΔT = Change in Temperature (K = °C + 273)
Cp = Specific Heat Capacity (Given)

33. Example Problem #1
 40.00 g of water is heated from 10.00°C to 30.00°C. The
specific heat of water is 4.184 J/gK.
q = (m)(ΔT)(Cp)
q=Heat (J)
m=Mass (g)
ΔT = Change in Temperature
(K = °C + 273)
Cp = Specific Heat Capacity (Given)
q = ?
m = 40.00 g
ΔT = (30 + 273) - (10 + 273) = 20.00 K
Cp = 4.184 J/gK
q = (40.00 g) ( 20.00 K) ( 4.184 J/gK)
q = 3347.2 J
q = 3350 J

34. Example Problem #2
 135.6 g of water is cooled from 95.80°C to 21.60°C. The
specific heat of water is 4.184 J/gK.
q = (m)(ΔT)(Cp)
q=Heat (J)
m=Mass (g)
ΔT = Change in Temperature
(K = °C + 273)
Cp = Specific Heat Capacity (Given)
q = ?
m = 135.6 g
ΔT = (21.6 + 273) - (95.8 + 273) =
-74.20 K
Cp = 4.184 J/gK
q = (135.6 g) ( -74.20 K) ( 4.184 J/gK)
q = -42097.39 J
q = -4.210 x 104 J

35. Energy Transfer Practice
Determine the amount of heat absorbed or released in each
of the following changes. Use the table of specific heat
capacity on page 533 of your text. The specific heat of
water is 4.184 J/gK.
1. 40.00 g of water is heated from 10.00°C to 30.00°C.
2. 135.6 g of water is cooled from 95.80°C to 21.60°C.
3. 30.00 g of aluminum is heated from 15.00°C to 35.00°C.
4. 450.0 g of iron is cooled from 125.0°C to 45.00°C.
5. 62.30 g of lead is heated from 21.70°C to 136.4°C.

36. Energy Transfer
Answers to previous slide
1. +3347.2 J = + 3347 J
2. -42,097.4 J = -4.210 x 104 J
3. +538.2 J
4. -16,164.0 J = -16160 J
5. +921.8 J