Chapter 2
Matter and Energy
UNIT ESSENTIAL
QUESTIONS:
1) WHAT IS THE
RELATIONSHIP BETWEEN
MATTER AND ENERGY?
2) HOW IS MATTER
STUDIED AND WHAT IS
NECESSARY TO PERFORM
THESE STUDIES?
Lesson Essential Question:
WHAT IS ENERGY AND
WHAT FORMS DOES IT
TAKE?
Section 1: Energy

Energy: the capacity to do work.
◦ Whenever matter changes, energy is involved!
 Can be endothermic or exothermic.
 Endothermic – energy is absorbed
 Exothermic – energy is released
Law of Conservation of
Mass(Matter)/Energy

Energy (E) cannot be created or destroyed,
only transferred. (The same is true for
matter!)
*This is what happens in a chemical or
physical change.
◦ System – all components being studied
◦ Surroundings – everything outside the system
◦ Energy is transferred between system and
surroundings
◦ E can be changed into other forms of E.
 Ex: Light, heat, chemical, mechanical, electrical,
sound
Energy as Heat

Heat – energy transferred between two
objects at different temperatures
◦ Always transferred from high E (hotter object)
to low E (cooler object)
Kinetic energy – energy of motion
 Temperature – measure of average
kinetic energy of particles in the object

◦ Kelvin scale – SI unit
◦ Absolute zero = no kinetic energy
◦ K = oC + 273.15
Heat vs. Temperature

Addition of heat does not always change
temperature
◦ Example: boiling water
◦ Adding more heat at the boiling point does not
cause it to change temperature
◦ So what is happening to the energy being
transferred to the water at the boiling point?
Think about what happens to water
molecules at the boiling point.
Phase change!
Heating/Cooling
Curve
Temperature change =
change in molecular
motion (kinetic
energy)
No temperature
change = state change
Specific Heat
Quantity of heat required to raise one
gram of a material 1 K (or 1 oC)
 SI unit for energy = Joule (J)

◦ Units for specific heat = J/(g∙oC) or J/(g∙K)
◦ Metals = low specific heats- they heat up/cool
down easily!
 Aluminum: 0.897J/g∙K
 Copper: 0.385J/g∙K
 Gold: 0.129J/g∙K
◦ Water = high specific heat- does not heat
up/cool down as easily: 4.18J/g∙K
Calculating Specific Heat
What do we need to know to calculate
specific heat? E added as heat, mass, & T
 Formula: Cp = q / (m x T)

◦ Note: T = Tf – Ti (change in anything is always
final minus initial)
Lesson Essential Question:
HOW ARE IDEAS AND
QUESTIONS
APPROACHED IN
SCIENCE?
Scientific Method
Revise
hypothesis
Form
Hypothesis
Construct a
Theory
Ask
Questions
Test
Hypothesis
Make
Observations
Publish
Results
Analyze
Results
Draw
Conclusions
Experiments

Hypothesis – a prediction or educated
guess as to what will happen.
◦ Represents cause and effect- ‘if, then’ statement

Testing
◦ Variable – factor that could effect results
 Change only 1 at a time
◦ Control – variable that is kept constant
 Many of these in experiment.

Theory – explains why things happen.
◦ Repeated testing needed
◦ Based on lots of data and observations
Laws

Law – a summary or description of events
◦ Tells how things work, not why
◦ Helps predict events/behavior (because they
follow a pattern according to the law)

Law of conservation of mass – mass cannot
be created or destroyed in ordinary physical
or chemical changes
◦ Same as law of conservation of energy

Model – represents an object, a system, a
process, or an idea.
◦ Computer generated, 3D, drawing, etc.
Theories vs. Laws

Planets move in an ellipse with a star at a focus.
Kepler’s 1st Law- describes motion of planets.


The amount of disorder in an isolated system never
decreases. 2nd Law of Thermodynamics- describes chaos.
The universe was created when a massive explosion
occurred. Big Bang Theory- explains where
universe & planets came from.


As the pressure of a gas increases, the volume of the
gas decreases. Boyle’s Law- describes P & V effect on gases.
Continents developed from one massive continent
(Pangaea) where they broke apart and moved due to
tectonic plates in the Earth’s lithosphere.
Plate Tectonics Theory – explains where continents came from.
Lesson Essential Questions:
HOW DO WE OBTAIN THE
CORRECT NUMBER OF
DIGITS IN
CALCULATIONS?
HOW ARE VERY SMALL OR
LARGE NUMBERS
REPRESENTED?
Section 3: Measurements &
Calculations in Chemistry

Accuracy vs. Precision
◦ Accuracy – how close a measurement is to
the true/correct value
◦ Precision – how close several measurements
are to each other
Introduction to Sig Figs
Use the ‘ruler’ to measure the width of your
table. Use each ‘side’ of the ruler to make the
measurements. You should have a total of
four measurements.
 Record these on a piece of paper. Include
units!
 Each side should have the following number
of decimal places:
 #1: 1
#2: 1
#3: 1
#4: 2

Significant Figures (significant digits)

D = 3.421g/5.957mL = 0.5742823568…g/mL
◦ How do we know where to round?

Significant Figures are all digits known with
certainty plus one more uncertain/estimated
digit.
◦ Rules that govern how you determine where to “cut
off” a number
◦ Calculators do not “know” these rules, so it’s up to
YOU to know where to round!

Also helps to show degrees of accuracy and
precision- more sig figs = better accuracy and
also helps multiple measurements be precise!
Rules for determining significant digits
Rule #1: Nonzero digits are always significant.
◦
◦
46.3 m
6.295 g
3 sig figs
4 sig figs
Rule #2: Zeros between significant digits
(typically nonzero digits) are significant.
◦
◦
40.7 L
3 sig figs
87,009 km
5 sig figs
Rule #3: Zeros in front of nonzero digits are not
significant.
◦
◦
0.009 587 m
0.000 09 kg
4 sig figs
1 sig fig
Rules for sig figs continued…
Rule #4: Zeros both at the end of a number AND to
the right of the decimal are significant.
◦
◦
85.00 g
4 sig figs
10 sig figs
9.070 000 000 cm
Rule #5: Zeros at the end of a number but to the left
of a decimal point may or may not be significant.
*If a zero has not been measured or estimated, it is not
significant.
*A decimal point placed after zeros indicates that the
zeros are significant.
 2000 m
1 sig fig
4 sig figs
 2000. m
Rules for sig figs continued…
Sig figs & scientific notation
 If a number is written in scientific notation, only
look at the first number for sig figs!
 The x10Y does not impact sig figs- it only
changes size!


2.0 x 103m
3.041 x 10-2g
2 sig figs
4 sig figs
Rules for Using Significant Figures in
Calculations
1) In multiplication and division problems,
the answer cannot have more sig figs
than there are in the measurement with
the least sig figs.
*Look for the # with the least sig figs!
Ex: 12.2257 m
6 sig figs
4 sig figs
x 1.162 m
14.2062634 m2
round off to 4 sig figs
= 14.21 m2
Rules for Calculating continued…
2) In addition and subtraction, the result can be
no more certain than the least certain number
in the calculation.
* Look for the # with the least decimal places!
Ex:
3.95 g
2.879 g
+ 213.6 g
220.429 g
= 220.4 g
2 decimal places
3 decimal places
1 decimal place
round to 1 decimal place
Finally…
3) If a calculation has addition/subtraction
and multiplication/division, round after each
operation.
Ex: 7.92g – 8.5g2 = 7.92g – 3.5g = 4.4g
2.46g
4) In chemistry you will follow sig fig rules
to know where to round off all of your
calculations.
Unlimited Significant Figures
Numbers that are exact or counted have
infinite sig figs.
 Have no impact in determining sig figs in
an answer from a calculation.

◦ Examples:
 35 cars = infinite sf
 1 m = 1000 mm
counted!
exact!
 Conversion factors often have infinite sig figs!
Warm-Up!
Is there an easier way to
write such large and small
numbers ?? YES!
Average distance between sun and earth:
93,000,000 miles
Diameter of an atom:
0.000 000 000 062 m
Imagine you wanted to measure the distance
in between planets of our solar system and
the diameter of an atom. What would the
size of your measurements look like?
Scientific Notation
Very large or very small numbers are easier to write
using scientific notation.
 Form = M x 10y
◦ M = number between 1 and 10 (not including 10!)
◦ y = integer (can be positive or negative)
 Examples:
◦ 299 800 000 m/s = 2.998 x 108 m/s


◦ 0.000 001 23 cm3 = 1.23 x 10-6 cm3
◦ 4500. g = 4.500 x 103 g
◦ 6.79 x 10-7m = 0.000 000 679m
◦ 5.307 x 105L = 530,700L
Follow sig figs when calculating!