Chapters 1 & 2
Matter, Density & Specific
Gravity, Significant Figures et.al.,
Chemical History
1
What is Chemistry?
What does it have to do with
me?
2
What Does It Have To Do With Me?
“What does chemistry have to do with me?” said
Mr. Averageman, as he looked at a page printed with ink
made by a chemical process, and tied his shoes, made of
leather tanned by a chemical process. He glanced through a
pane of glass, made by a chemical process, and saw a bakers
wagon full of bread leavened by a chemical process, and a
draper’s wagon delivering a parcel of silk, dyed by a chemical
process. He put on a hat, shaped by a chemical process, and
stepped out on the asphalt pavement, compounded by a
chemical process, bought a daily paper with a penny refined
by a chemical process. “No,” he added, “of course not,
chemistry has nothing to do with me.”
Herbert Newton Casson
1869-1964
[Merrill Chemistry 1995
3
This is a liquid crystal
molecule.
4
These are liquid crystal fibers.
5
Ever use any of these?
6
What is Chemistry and How
Did We Get to This Point?

Chemistry is called the ‘central science’.

Studies the structure and properties and
changes of matter.
7
8
Scientific Method
9
What is Matter?

Complete ‘Is It Matter?’ probe individually.

When instructed, share your answers and
your reasoning w/ a partner
(5)
(5)
10
Matter - the ‘stuff’ that
everything is made of




Takes up space
Has mass
Has inertia
Has energy
11
Takes up space means that it is
Three dimensional
12
Has mass indicates


the quantity of matter present
Gram units
13
Has inertia refers to

the tendency of a body to maintain its state of
rest or uniform motion unless acted upon by
an external force
inertia
14
Has energy which means

The ability to do work

In science we say that work is done on an object
when you transfer energy to that object.
15
Mass and Inertia
Relationship?
16
Watch the demo!
17
What did you notice about…




Their hands?
Which was more difficult to catch?
Which required greater effort to stop ?
Remember that inertia is the resistance to a
change in motion.
18
What can you conclude about
the relationship between mass
and inertia?

More mass = More inertia
19
Mass and Energy
Relationship?
20
Assignment

Read “Nature of Energy”

p 516 - 518 in text
21
Many different types of energy



Chemical
Nuclear
Radiant




Heat
Light
Electrical
Mechanical


Kinetic
Potential
22
One form of energy can be
transformed into another.
chemical
potential
radiant
electrical
kinetic
23
Significant Figures, Calculations
and Measurements
24
SI system of units

modern metric


Be able to recognize the order of the prefixes.
Be able to convert from one unit to another.
25
Prefix
abbreviatio
n
Power of 10
Mass unit
Length unit
Volume unit
kilo-
k
1000 or 103
kilogram,
kg
kilometer,
km
kiloliter,
kL
kl
hecto-
h
100 or 102
hectogram,
hg
hectometer,
hm
hectoliter,
hlhL
deka-
dk
10
dekagram,
dkg
dekameter,
dkm
dekaliter,
dkL
dkl
1
gram
meter
liter
base unit
deci-
d
.1 or 10-1
decigram,
dg
decimeter,
dm
deciliter,
dl
dL
centi-
c
.01 or 10-2
centigram,
cg
centimeter,
cm
centiliter,
clcL
milli-
m
.001 or 10-3
milligram,
mg
millimeter,
mm
milliliter,
mLmL
26
Kids
much.
kilo
have
hecto
days
deka
they
meter
liter
gram
don’t
deci
care
centi
milli
To change 11.263 cg to deka grams…
To change 1.235 km to millimeters …
27
28
29
Atlantic/Pacific
Rule

To determine the number of sig figs in a given
value, use the Atlantic/Pacific Rule regarding
decimals
Pacific
Ocean
Atlantic
Ocean
30
Significant Figures
(con’t)

If the decimal is Absent, count from the Atlantic side
from the first non-zero digit

If the decimal is Present, count from the Pacific side
from the first non-zero digit

Do NOT stop counting until you’ve run out of digits!
A bar above a zero indicates the marked zero was the
estimated value in the measurement and is to be
included in calculating the accuracy of the instrument
and in determining the number of significant figures.

31
How many sig figs are
in each number?
 408.00 g  5 sig figs
 639.0 g  4 sig figs
 0.0058020 mm  5 sig figs
 5640 m  3 sig figs
 0.002 g  1 sig fig
 3,090, 000 km  3 sig figs
32
Ice Cubes in a Bag - probe

Complete ‘Ice
Cubes in a Bag’
probe individually.
(5)

When instructed,
share your
answers and your
reasoning w/ a
partner
(5)
33
“Matter and energy are two sides
of the same physical entity.”
Albert Einstein

E = mc2



E = Energy
m = Mass
c = speed of light

3 x 108m/s
34
Law of Conservation of Mass and
Energy


matter and energy are interchangeable
the total amount of matter and energy in the
universe is constant
35
Because of this, chemical
rxn…

are always accompanied by changes in
energy (E)



E given off - exothermic
E used – endothermic
Cannot gain or lose

mass of reactants = mass of products
36
A thought experiment…

You have just made lemonade. You taste it
and find that you pucker when you drink it.

What should you do to ‘fix’ the lemonade?
37


“Add sugar”
... qualitative solution
“Add 1 cup of sugar” ... quantitative solution


*has Numeral
*has UNIT
38
Mass v. Weight


Mass – amount of matter present
Weight – measure of the pull of gravity on the
mass.
80 kg
39
Mass v. Weight

The gravitational force on the moon is 1/6th
that of the earth.

weight is less on the moon.
13.3 kg
13.3 kg
40
Basic unit

original measurement
41
Derived unit

- result of a mathematical function performed on
basic units

volume of a box [l x w x h],

speed [distance/time]
42
43
Measuring

Compares a feature of an object to a
standard

For example:



the length of a book is compared to the standard length of a
ruler
the mass of a jar of pickles is compared to the standard mass
of a gram
Measurement tools we will often use in the
lab:



Balance (measures mass)
Graduated cylinder (measures volume)
Meter stick (measures length)
44
When reading an
instrument…

Consider all forms of error that could occur



Human error (we all make mistakes)
Equipment error
Changing environmental conditions


Ex: Pressure
No instrument is 100% accurate
45
All measurements MUST have…



Number value
Unit
Indicator of the degree of accuracy

Significant figures indicate the degree of accuracy

The number of digits in the measurement shows the
degree of precision of the measuring instrument,
which shows how accurate the measurement is
46
When taking a measurement,
remember:


Read all numbers from the instrument PLUS
one additional ESTIMATE
When reading a measurement from a digital
readout, treat the last number on a digital
readout as an estimate
47
Significant Figures

All of the numbers that can be directly read
from an instrument PLUS one additional
estimated number


For digital instruments such as the electronic
balance, the last digit IS the mechanically
estimated value
Significant figures reflect the precision of an
instrument, which shows the accuracy of a
measurement
48
Accuracy




Nearness to an accepted value
Measured in ERROR
ABSOLUTE ERROR
EA = O – A
 EA – absolute error
 O – observed; may be individual or experimental average
 A – accepted value
RELATIVE ERROR (Percent Error)
ER = (EA  A)  100
or
Percent error = experimental value – accepted value X 100
accepted value
49
Precision



Repeatability; consistency; closeness of multiple
measurements to each other
Measured in DEVIATION
ABSOLUTE DEVIATION
DA = OI - M


OI – individual observed trial
 M – average of all trials
RELATIVE DEVIATION
DR = (DA  M)  100
50
Accuracy vs. Precision
(con’t)

Two archers are taking target practice. Here
are their shots…

Which of these archers is accurate and which is precise? 51
How do you know??
Measuring & Calculating Activity
52
Measuring the Mass
~Digital v. Triple Beam Balances~

Always use weighing paper or a glass container – never put
Digitalor
Balances
Triple Beam Balances
chemicals
hot objects directly on the balance
 Turn the scale on and be sure it reads 0.0 • Use the zero adjust knob until the pointer
g before placing anything on it to be
rests at zero if it is not already. (The
massed.
pointer must be at zero before use)
Mass the container / paper to be used in
the weighing.
 Press the Tare button to bring the scale
back to 0.0g.
Add the substance to be weighed to the
container / paper.
No estimation is necessary: the last
•Pre-weigh the paper or glassware that the
solid or liquid will be in. After recording
that mass, put the solid or liquid into the
glassware and remass the glassware.
•Mass of Glassware & Substance – Mass
of
Glassware = Mass of Substance
•Adjust the sliding weights on the various
53
scales as necessary
Measuring the Volume
~Graduated Cylinders~


The lowest curvature of the liquid level in a graduated cylinder is
the meniscus
To read the volume of liquid in a graduated cylinder, you need to
carefully read the BOTTOM of the meniscus at eye level
 If you read the markings on the top of the meniscus, your
reading will be incorrect!
 If you read the meniscus from an angle and not at eye level,
your reading will either be too high or too low
 Don’t forget, when taking a measurement, you want to read
all of the numbers from the instrument PLUS one estimated
number (can be a zero)
54
Evaluation

Measurement Using Significant Figures
55
DENSITY



derived unit
mass per volume
D = m/V
Density


= mass of orange
volume of orange
= __x_grams
4/3 π r3
56
Units of Density


g/cm3 - unit of density for solids and liquids
g/L - unit of density for gases
57
Density Example





: A brick [ingot ] of silver is 3 cm x 4.5 cm x
7.5 cm. What is the density of silver if the
brick is massed at 1062 g?
m = 1062 g
v = 3 cm x 4.5 cm x 7.5 cm
= 101.25 cm3
D = 1062 g
101.25 cm3
= 10.49 g/cm3
58
SPECIFIC GRAVITY

comparative value of a substance to a
standard
S.G. = density / standard

Standard Values
solids/liquids - 1.0 g/cm3 [density of water]
gases - 1.29 g/L [density of air]


59
What is the density and specific gravity of mercury
(Hg) with a mass of 300.0 g and a volume of 22.1 mL?





D=m
v
= 300.0 g
22.1 mL
D = 13.6 g/mL





SG = Density
Std
SG = 13.6 g/mL
1.0 g/cm3
SG = 13.6
60
Very Important Equality to
Remember !!!
1 gram = 1 ml =1 cm3
Water @ 4oC and standard pressure.
61
62
When you solve a math
calculation, is there a ‘right’
answer?
63
Calculate the mass of a silver ingot, which measures
1.14 cm x 2.35 cm x 1.88 cm and has a density of 10.49
g/cm3.


Group A
Solve this calculation
rounding the answers
to 3 numbers after
each step.


Group B
Solve this calculation
saving all numbers
and rounding only
once at the very end.
Ans.
1.14 cm x 2.35 cm x 1.88 cm = 5.04
cm3
Ans.
1.14 cm x 2.35 cm x 1.88 cm = 5.03652
cm3
10.49 g = mass
5.04 cm3
52.9 g = mass
10.49 g = mass
5.03652 cm3
52.8 g = mass
64
What would you predict will happen
if…

The calculation had
more steps in it?

The calculations had
much larger or much
smaller numbers?
65
Now we see the need for having
procedures for getting the same
answer(s).
66
Sig Fig Rules for Addition/Subtraction



Step #1: Perform the math
Step #2: Determine which number you are
adding or subtracting has the least number of
decimal places
Step #3: Round your answer to show the same
number of decimal places as the term in the
problem with the fewest decimal places
67
Addition/Subtraction Example
Step 1
5.6079 m
3.14 m
+ 6.704 m
15.4519 m
Step 2  3.14 has the fewest decimal places of any of the numbers in the calculation
Step 3  Final answer should be rounded to match this number of decimal places
[in this case, there should be two decimal places]
Answer: 15.45 m
68
Sig Fig Rules for
Multiplication/Division

Step #1: Perform the math

Step #2: Determine which term in the
calculation has the fewest number of
significant figures

Step #3: Round your answer to match that
number of significant figures. Double check
your units.
69
Multiplication/Division
Example
8.563cm x 9.23cm x 3.487cm = ?
4
3
Step
1 (Perform
the 4math):
8.563cm x 9.23cm x 3.487cm = 275.60024 cm3
Step 2 (Round answer to least number of sig figs):
Answer: 276 cm3
Step 3 (Double check to make sure your units make
sense)
70
Scientific Notation Review


If a number is larger than the thousands place, or
if the number is smaller than the thousandth place put into scientific
notation

Numbers, when written in scientific notation, have only
ONE DIGIT to the left of the decimal place. This is the
acceptable and correct way of writing numbers in scientific
notation.

Examples: How would you write the following expressions in scientific
notation?
Standard Form
Scientific Notation Form
14, 890
0.456
71
0.007532
Scientific Notation Review


If a number is larger than the thousands place, or
if the number is smaller than the thousandth place put into scientific
notation

Numbers, when written in scientific notation, have only
ONE DIGIT to the left of the decimal place. This is the
acceptable and correct way of writing numbers in scientific
notation.

Examples: How would you write the following expressions in scientific
notation?
Standard Form
Scientific Notation Form
14, 890
1.489 x 104
0.456
0.456 or 4.56 x 10-1
72
0.007532
7.532 x
10-3
Changing a Number into
Scientific Notation



If the exponent increases, the number before the
x10 (known as the mantissa) decreases by the
number of decimal places the power is changed
If the exponent decreases, the mantissa increases
by the number of decimal places the power is
changed.
Example: 2.71 x 102 to 103


The exponent is increased by 1 power of 10, so the
mantissa is decreased by 1 decimal place
0.271 x 103
73
Try these…

3.561 x 106 to 104

3.0100 x 10-2 to 10-5

6.211 x 1022 to 1025
74
Try these…

3.561 x 106 to 104


3.0100 x 10-2 to 10-5


Answer: 356.1 x 104
Answer: 3010.0 x 10-5
6.211 x 1022 to 1025

Answer: 0.006211 x 1025
75
Adding/Subtracting Numbers in
Scientific Notation


When adding or subtracting numbers in scientific
notation, the numbers must be in the same power
of 10 to round answer properly, so the form may
need to be TEMPORARILY manipulated to solve
the calculation, but the answer must be restored to
the accepted form.
Use rules for significant figures for your final
answer
76
Adding/Subtracting Numbers in
Scientific Notation
Ex: [2.71 x 102 L] + [7.10 x 103 L] + [1.2 x 103 L] = ?
0.271 x 103 L
7.10 x 103 L
+ 1.2 x 103 L
8.571 x 103 L


1.2 has the fewest number of decimal places in the
calculation
Round the answer to match this number of decimal
points
Answer: 8.6 x 103
77
Multiplying/Dividing Numbers in
Scientific Notation


Numbers can be to any power of 10 and still
be able to be multiplied or divided without first
manipulating the number
Example:
[2.15 x 105 cm] x [5.030 x 10-2] = ?
 First, input the numbers into your calculator
 Calculator shows either 10814.5cm2 or 1.08145 x 104 cm2
 2.15 has the fewest significant figures, so your answer should
be rounded to match that same number of sig figs
78
Answer: 10800cm2 or 1.08 x 104 cm2


ASSIGN: Density & Specific Gravity
ASSIGN: Density & Specific Gravity II
79
Accuracy & Precision
Revisited
80
Accuracy vs. Precision (con’t)

Remember the two archers taking target
practice.

Which of these archers is accurate and which is precise?
Accuracy



Nearness to an accepted value
Measured in ERROR
ABSOLUTE ERROR


EA – absolute error
O – observed; may be individual or experimental average


EA = O – A
A – accepted value
RELATIVE ERROR (Percent Error)
ER = (EA  A)  100
or
Percent error = experimental value – accepted value X 100
accepted value
Precision



Repeatability; consistency; closeness of multiple
measurements to each other
Measured in DEVIATION
ABSOLUTE DEVIATION
DA = OI - M


OI – individual observed trial
 M – average of all trials
RELATIVE DEVIATION
DR = (DA  M)  100
Determine the accuracy and
precision for the EXPERIMENT
Trial
1
Amount of
Product
(g/mL)
3.92
2
3.97
3
3.95
Accepted
Value
3.96
84

O = 3.95 g/mL


Ea= 3.95 g/mL-3.96 g/mL = -.01 g/mL
Er = -.01 g/mL x 100 = -.25%
3.96 g/mL

M = 3.95 g/mL

Da = +/-.03 + +/-.02 + 0 = +/- .017 g/mL
3
Dr = +/-.017 g/mL x 100 = +/- .42%
3.95 g/mL

[3.92 g/mL + 3.97 g/mL + 3.95 g/mL ] ÷ 3
[3.92 g/mL + 3.97 g/mL + 3.95 g/mL ] ÷ 3
85
In the Lab, you should be able
to…




convert between units of measure using the SI
System,
determine whether your data is accurate,
precise, both, or neither,
report data in the correct number of significant
figures, and
work with numbers in scientific notation.


ASSIGN: Accuracy and Precision w.s.
ACTIVITY: Accuracy & Percision in
Measurements Lab
87
What is Chemistry and How Did
We Get to This Point?
CHEMICAL HISTORY
An Internet Activity

Research four eras that made a significant contribution to
chemistry.
Egyptian Era
Dark Ages
Pilgrim Era
Am. Revolution

Content will be included on the Ch. 1-2 test
88
INTERNET SITES

Egyptian Era
http://www.zompist.com/ver
sci.htm

Scroll down to ‘Science’


Choose ‘Chemical History’

Scroll down to the ‘Egyptian,
Mesopotamian’ section

What contribution did chemistry
make to the Egyptians?
What are some products they
produced?

89
INTERNET SITES



Dark Ages
http://www.rsc.org/Publishing
/CurrentAwareness/index.as
p
In the ‘site search’ box type
– alchemical symbols



Choose first selection
What is the purpose(s) of
alchemy?
With what other practices
was alchemy associated?
90
INTERNET SITES



Pilgrim Era
http://www.jimloy.com/p
hysics/phlogstn.htmCho
ose #5 Phlogiston
.


What was the phlogiston
theory?
What were it’s beliefs?
91
INTERNET SITES
American Revolution


library.thinkquest.org/C006439/scientists/
Choose Robert Boyle


Also choose Antoine Lavoisier



What did Robert Boyle contribute to our modern
concept of "element".
Why was Lavoisier's weighing of substances
important?
Why is Lavoisier remembered as the "father" of
chemistry
http://www.ehow.com/info_8233221_disco
veries-inventions-advances-1700s.html

What was were the major advances in chemistry
during this time period?
92