WELCOME TO CHEM 1
WITH MRS. KAUR
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Learn about each other
Learn how to be a good student and a
professional employee
Learn general chemistry
General course info in syllabus
My website – use it!
First two chapters will be FAST – they are a
review!
Look at slide: tell me what you can about it.
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Tell me all you can about this picture:
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Chapter 1
Keys to the Study of
Chemistry
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HOW TO LEARN CHEMISTRY?
Think of it as a foreign language
Memorize the new words, names of elements and
compounds, definitions
Use flash cards
Learn how to be a scientist
Also question everything you read or hear - even me
or the textbooks.
Scientific exploration begins with questions, and leads
to more questions.
Read “Study tips and how to take a test” in your packet
of handouts.
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Is the study of matter,
its properties,
the changes that matter undergoes,
and
the energy associated with these changes.
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Definitions
Matter anything that has mass and volume -the “stuff” of the
universe: books, planets, trees, professors, students
Composition the types and amounts of simpler substances that
make up a sample of matter
Properties
the characteristics that give each substance a unique
identity
Physical Properties
Chemical Properties
those which the substance
shows by itself without
interacting with another
substance such as color, melting
point, boiling point, density
those which the substance shows
as it interacts with, or transforms
into, other substances such as
flammability, corrosiveness
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Figure 1.1
The distinction between physical and chemical change.
A Physical change
B Chemical change
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Table 1.1 (4th ed.)
Some Characteristic Properties of Copper
Physical Properties
Chemical Properties
slowly forms a basic blue-green
sulfate in moist air
reddish brown, metallic luster
easily shaped into sheets
(malleable) and wires
(ductile)
reacts with nitric acid
and sulfuric acid
good conductor of heat
and electricity
density = 8.95 g/cm3
slowly form a deep-blue
solution in aqueous ammonia
melting point = 10830C
boiling point = 25700C
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Figure
1.2
The physical states of matter.
What are the phase changes? There are six - see how
many you can remember…
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REVIEW OF DEFINITIONS YOU
SHOULD ALREADY KNOW:
Phases of matter: solid, liquid, and gas, and
phase changes between them
Solutions of matter: s, l or g mixtures. Solute
dissolved in a solvent = a solution.
Physical properties and chemical properties
(see previous slides)
Extensive vs. Intensive properties
Separation processes
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Sample Problem 1.2
Distinguishing Between Physical and
Chemical Change
PROBLEM: Decide whether each of the following process is primarily a
physical or a chemical change, and explain briefly:
(a) Frost forms as the temperature drops on a humid winter night.
(b) A cornstalk grows from a seed that is watered and fertilized.
(c) Dynamite explodes to form a mixture of gases.
(d) Perspiration evaporates when you relax after jogging.
(e) A silver fork tarnishes slowly in air.
SOLUTION:
(a) physical change
(b) chemical change
(d) physical change
(c) chemical change
(e) chemical change
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Energy is the capacity to do work.
Potential Energy
Nonpotential
Energy
Energy due to the position of the object or
energy from a potential chemical reaction
Energy due to the motion of the object E=1/2 mv2
Heat transfer q=mcpDT
Light energy e=hn
Radiant energy, blackbody radiation
Etc.
Potential and nonpotential energy can be interconverted.
(See figure 1.3 in your textbook)
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Fig 1.4
Scientific Approach: Developing a Model
Observations :
Hypothesis:
Natural phenomena and measured events;
universally consistent ones can be stated as a
natural law.
Tentative proposal that explains observations.
revised if
experiments do
not support it
Experiment:
Model (Theory):
Procedure to test hypothesis; measures one
variable at a time.
Set of conceptual assumptions that explains
data from accumulated experiments; predicts
related phenomena.
altered if
predictions do
not support it
Further Experiment: Tests predictions based on model.
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A Systematic Approach to Solving Chemistry Problems
1. Problem statement
Clarify the known and unknown.
2. Plan
Suggest steps from known to unknown.
Prepare a visual summary of steps.
3. Solution
4. Check
5. Comment and 6. Follow-up Problem
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SI Base Units
Table 1. 1
Physical Quantity
(Dimension)
Unit Name
Unit
Abbreviation
mass
kilogram
kg
length
meter
m
time
second
s
temperature
kelvin
K
electric current
ampere
A
amount of substance
mole
mol
luminous intensity
candela
cd
volume
liter
Know the ones outlined in red.1-
L
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Table 1.2
Common Decimal Prefixes Used with SI Units
Prefix
Prefix
Symbol
Word
tera
giga
mega
kilo
hecto
deka
----deci
centi
milli
micro
nano
pico
femto
T
G
M
k
h
da
---d
c
m

n
p
f
trillion
billion
million
thousand
hundred
ten
one
tenth
hundredth
thousandth
millionth
billionth
trillionth
quadrillionth
Conventional
Notation
1,000,000,000,000
1,000,000,000
1,000,000
1,000
100
10
1
0.1
0.01
0.001
0.000001
0.000000001
0.000000000001
0.000000000000001
Memorize these
prefixes!
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Exponential
Notation
1x1012
1x109
1x106
1x103
1x102
1x101
1x100
1x10-1
1x10-2
1x10-3
1x10-6
1x10-9
1x10-12
1x10-15
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Just for fun!
Try to find out the distance from the earth to the
sun in meters:_____________
Also the diameter of the helium
atom:_________________
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Common SI-English Equivalent Quantities
Quantity
SI Unit
SI Equivalent
English Equivalent
Length
English to
1 kilometer(km)
1000(103)m
SI Equivalent
Table 1.3
0.62miles(mi)
1 mi = 1.61km
1 meter(m)
1 mi = 5280 ft
100(102)cm
1000(103)mm
1.094yards(yd)
39.37inches(in)
1 yd = 3 ft, etc.
1 yd = 0.9144m
1 foot (ft) = 0.3048m
1 centimeter(cm)
0.01(10-2)m
0.3937in
1 in = 2.54cm
(exactly!)
1 kilometer(km)
1000(103)m
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0.62mi
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Common SI-English Equivalent Quantities
Quantity
SI Unit
SI Equivalent
English Equivalent
Volume
English to
SI Equivalent
1 cubic meter(m3)
Table 1.3
1,000,000(106)
cubic centimeters
35.2cubic feet (ft3)
1 ft3 = 0.0283m3
1 cubic decimeter(dm3)
1000cm3
0.2642 gallon (gal)
1.057 quarts (qt)
1 cubic centimeter (cm3)
0.001
1 gal = 4 qt
1 gal = 3.785 dm3
1 qt = 0.9464 dm3
1 L = 1.0567 qt
dm3
0.0338 fluid ounce
1 qt = 946.4 cm3
1 fluid ounce = 29.6 cm3
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Common SI-English Equivalent Quantities
Quantity
SI Unit
SI Equivalent
English Equivalent
Mass
English to
1 kilogram (kg)
SI Equivalent
1000 grams
1 lb = 453.6 g
2,205 pounds (lb)
Table 1.3
1 (lb) = 0.4536 kg
1 gram (g)
1 lb = 16 oz
1000 milligrams
0.03527 ounce(oz)
1 lb = 453.6 g
1 ounce = 28.35 g
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English-Metric Conversions to
Memorize
1" = 2.54 cm exactly
1 lb = 453.6 g
1 cal = 4.184 J
1.0 L = 1.0567 qt
1.000 atm = 14.70 psi
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Practice:
The O-H bond length is 95.7 pm. Convert it to nm.
Convert 750 mL to L:
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Weight = mass x acceleration
Metric:
Newton = mass in grams x accel of gravity
English:
Pound = mass in slugs x accel of gravity
On earth we use 1 lb = 453.6 g, even though we
should use slugs!
Practice:
Convert 2.65 g to kg:
Convert 2.65 g to mg:
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The three dimensions
Length is one dimensional.
Area is two-dimensional: area = length x width
Units will be ft2, m2, cm2, etc.
Volume is three-dimensional
V = length x width x height
Units will be L, mL, cm3,
MEMORIZE: 1 L = 103 mL, 1 mL = 1 cm3
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Equipment measuring liquid volume:
From previous lab experience, describe a
graduated cylinder, a buret, a volumetric
flask, and a pipette.
Decide which of them would provide the most
accurate volume reading.
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DENSITY
Density is a physical property that represents
the amount of volume a certain mass
displaces
D = m/V
Find the mass of 250.0mL of ethanol if D =
0.789g/mL at 20C.
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Table 1.5 (4th ed.) Densities of Some Common Substances*
Substance
Physical State
Density (g/cm3)
Hydrogen
Gas
0.0000899
Oxygen
Gas
0.00133
Grain alcohol
Liquid
0. 789
Water
Liquid
0.998
Table salt
Solid
2.16
Aluminum
Solid
2.70
Lead
Solid
11.3
Gold
Solid
19.3
*At
room temperature(200C) and normal
atmospheric pressure(1atm). 27
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Figure 1.6
The freezing and boiling points of water.
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Temperature Scales and Interconversions
Kelvin ( K ) - The “Absolute temperature scale” begins at
absolute zero and only has positive values.
Celsius ( oC ) - The temperature scale used by science,
formally called centigrade, most commonly used scale around the
world; water freezes at 0oC, and boils at 100oC.
Fahrenheit ( oF ) - Commonly used scale in the U.S. for our
weather reports; water freezes at 32oF and boils at 212oF.
K = oC + 273.15
oC = K - 273.15
oF
= 1.8 oC + 32
oC = [oF - 32 ]/ 1.8
Memorize these!!!
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TEMPERATURE CALCULATIONS
Convert room temperature of 72.0oF to oC and
to Kelvin. Apply sig fig rules of math. (Do
you remember them?)
In the northeast, we say “below zero” for very
cold temperatures, but this is referring to
Fahrenheit. What is “five below zero” in
Celsius?
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PERCENT CALCULATIONS
Percent: can be by mass or by some other
measurement.
Percent = part/whole * 100
If a 17.6 g silver bracelet is only 14.1 g Ag,
because it is in an alloy with copper for
strengthening. What is the percent-mass Ag?
And the percent-mass Cu?
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Figure 1.7
The number of significant figures in a measurement depends
upon the measuring device.
32.330C
32.30C
The last digit, shaded gray, is the estimated digit
that you would read. It is the digit with some
uncertainty.
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Rules for Determining Which Digits are Significant:*
All digits are significant except zeros that are used only to position the
decimal point.
•Make sure that the measured quantity has a decimal point.
•Start at the left of the number and move right until you reach the first
nonzero digit.
•Count that digit and every digit to its right as significant.
Zeros that end a number and lie either after or before the decimal
point are significant; thus 1.030 ml has four significant figures,
and 5300. L has four significant figures also.
Numbers such as 5300 L are assumed to only have 2 significant
figures. A terminal decimal point is often used to clarify the
situation, but scientific notation is the best!
*See informational handout “The Rules of
Significant Figures”
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Rules for Significant Figures in Calculations
1. For addition and subtraction. The answer has the
same number of decimal places as there are in the
measurement with the fewest decimal places.
Example: adding two volumes
83.5 mL
+ 23.28 mL
106.78 mL = 106.8 mL
Example: subtracting two volumes
865.9
mL
- 2.8121 mL
863.0879 mL = 863.1 mL
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Rules for Significant Figures in Answers
2. For multiplication and division. The number with the least
certainty limits the certainty of the result. Therefore, the answer
contains the same number of significant figures as there are in the
measurement with the fewest significant figures.
Multiply the following numbers:
9.2 cm x 6.8 cm x 0.3744 cm = 23.4225 cm3 = 23 cm3
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Rules for Rounding Off Numbers
1. If the digit removed is greater than or equal to 5, the
preceding number increases by 1.
5.379 rounds to 5.38 if three significant figures are retained
and to 5.4 if two significant figures are retained.
2. If the digit removed is less than 5, the preceding number is
unchanged.
0.2413 rounds to 0.241 if three significant figures are retained
and to 0.24 if two significant figures are retained.
Be sure to carry two or more additional significant figures
through a multi-step calculation and round off only the final
answer.
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Practice
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Pair up with your lab partner or someone in
your study group. Look at sample problem
1.9, sig figs and rounding. Then do its followup problem. Be prepared to show me your
work.
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Precision and Accuracy
Errors in Scientific Measurements
Precision Refers to reproducibility or how close the measurements are to each
other.
Accuracy Refers to how close a measurement is to the real value.
Systematic error Values that are either all higher or all lower than the actual value.
Random Error In the absence of systematic error, some values that are higher and
some that are lower than the actual value.
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Figure
1.8
Precision and accuracy in the laboratory.
precise and accurate
precise but not accurate
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Figure 1.8
Precision and accuracy in the laboratory.
continued
random error
systematic error
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