Personal Profile
Name
Birth date
What was your favorite experience this
summer?
How do you like to spend your time?
Who do you like to hang out with?
What activities or hobbies do you do?
What are your parent professions?
What do you hope to learn in Chemistry?
Student Email:
Parent Email:
Vocabulary 1
Scientific method, observation, data, hypothesis,
independent (manipulated) variable, dependent
(responding) variable, controlled variable,
controlled experiment, scientific law, theory,
mass, volume, precision, accuracy, infer,
conclusion, analyze, matter, element, mixture,
atom, molecule, homogeneous, heterogeneous,
solution, colloid, suspension, metals, non-metals,
compound, ionic compound, covalent
compound, metal, non-metal, metalloid, solute,
solvent, significant figure
Vocabulary of the Week

















Matter
Pure substance
Mixture
Element
Atom
Compound
Solution
Homogeneous solution
Heterogeneous solution
Solvent
Solute
Colloid
Suspension
Dilute solution
Concentration solution
Polar solvent
Non-polar solvent

















Metal
Non-metal
Metalloid
Malleable
Luster
Brittle
Ductile
Tensile strength
Decant
Centrifuge
Filter
Chromatography
Distillation
Miscible
Residue
Group
period
Unit 1 The Nature of Chemistry
Chapters 1 and 2
Chemistry and Matter
Science



Science is trying to explain the world
around us
Science is a way of thinking
“Science is a system of knowledge based
on facts or principles”
Pattern Puzzles – Experimental
Method
A)
B)
C)
D)
E)
F)
G)
ottffentettff
mirror numbers
roses & petals
Grandma likes Coffee, but she doesn’t like tea
Days in the months of the year
Next Box?
1st Letters of Months, of 9 Planets
Branches of Science
Science
Biological
Science
Physical
Science
Earth
Science
Branches of Science
Science
Biological
Science
Physical
Science
Zoology
Botany
Ecology
Science of living things
Earth
Science
Branches of Science
Science
Biological
Science
Physical
Science
Earth
Science
Physics
Science of matter and energy
Chemistry
Branches of Science
Science
Biological
Science
Physical
Science
The systems of the earth
Earth
Science
Geology
Meteorology
Astronomy
Science Skills

Observe


gathering of information by using your five senses
Infer

applying reason to explain an observation
Observations

Qualitative – describe with words


Quantitative – describe with numbers




Hot , red, large
100° , 10 meters, 3.46 grams
Scientists prefer quantitative
Easy to agree upon
No personal bias
Scientific Method Notes (1 of 3)
1. Observations – sensed, measured
2. Question – How explain observations?
3. Hypothesis – Proposed explanation that must
be testable, should be able to predict behavior.
 if (change, manipulated variable), then
(responding variable) statement because ….
4. Experiment – test 1 factor at a time, record
procedures used and results observed
Controlled Experiments (2 of 3)






Controlled experiments insure that scientists can make
clear interpretations of results.
Purposefully change 1 factor at a time – the independent
(manipulated) variable.
Measure 1 key effect – the dependent (responding)
variable.
Graphs reflect the experimental variables: independent =
x, dependent = y
All other variables are kept constant (controlled) to
eliminate their possible influence on the results.
Example: Chocolate causes zits? or testing the
effectiveness of a new drug on mice?
Scientific Method (3 of 3)
5. Data – collected measurements & observations
6. Conclusions – interpretation of results to confirm or
reject experimental hypothesis. (Use #s as examples)
7. Theory (mistakenly used in place of hypothesis) – an
explanation of a wide range of observations well
supported by experimental results.
Examples? Evolution, Relativity, Quantum, Germ, Cell
8. Scientific Law – a description of common
predictable behavior, not an explanation.
Examples? Newton’s 3 Laws of Motion, Law of Gravity,
Laws of Thermodynamics
Scientific Method

The scientific method is the systemized testing
of ideas, inferences, predictions, and hypotheses.
A way of thinking about and solving problems
 It is a logical method
 You do it all the time

10 Criteria of Science
1.
2.
3.
4.
5.
Science is logical and rational.
Science makes well-defined claims.
Scientific hypotheses are testable.
Scientific experiments are repeatable.
Science requires that claims be examined
by peers.
10 Criteria of Science (2 of 2)
6. Science views unexplained gaps in theories with
suspicion.
7. Science requires caution in examining evidence.
8. Science requires objectivity.
9. Science does not accept coincidence as proof.
10. Science does not accept anecdotal evidence as
proof.
Scientific Theory



A reasoned explanation tested by many
observations and experiments
Tells why things are
Three things
Must explain clearly and simply
 Must be repeatable
 Must be able to make predictions


Theories can be changed or modified by new
evidence
Scientific Laws




Describe what happens
Quantitative – use numbers and equations to
describe
Often equations are part of the law
Mathematics is a universal language
Law vs. Theory
Law
Describes how
Summarizes observations
Usually an equation
Theory
Explains why
Agrees with observations
Predicts new discoveries
Chemistry is…
…the study of the composition,
structure, and properties of matter and
the changes it undergoes
C2H5OH + 3 O2  2 CO2 + 3 H2O + Energy
Reactants

Products
Energy



Energy is the ability to do work.
Work - cause a change or move an object.
Many types- all can be changed into the other.
 Heat- the energy that moves because of a
temperature difference.
 Chemical energy- energy released or
absorbed in a chemical change.
 Electrical energy - energy of moving
charges
Types of energy

Kinetic Energy is energy of motion
KE = ½ m v 2

Potential Energy is stored energy or
energy of position
Batteries (chemical potential energy)
Spring in a watch (mechanical potential energy)
Water trapped above a dam
(gravitational potential energy)
Measuring Energy



Calorie (cal) – one calorie is the amount of heat
needed to raise the temperature of 1 gram of
water by 1 degree Celsius.
Joule (J) – metric system unit for energy
1 cal =4.184
Conservation of Energy

The Law of Conservation of Energy
states that energy can be neither
created or destroyed in ordinary
changes (not nuclear), it can only change
form.
Kinetic Energy
“HOT”
Fast
Kinetic Energy (KE) = ½ m v2
*Vector = gives direction and magnitude
“COLD”
Slow
Temperature Scales
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 136
273 K






Measuring Temperature
Kelvin starts at absolute zero (-273 º C)
degrees are the same size
K -273 = C
C + 273 = K
Kelvin is always bigger.
Kelvin can never be negative.
Wednesday, September 15, 2010
Give three examples of kinetic energy?
Give three examples of potential energy?
What is the metric system unit for temperature?
Energy?
What is absolute zero?
Matter
Chemistry is……the study of the composition, structure, and
properties of matter and the changes it undergoes
Matter is anything that has mass and occupies space
Mass is a measure of the amount of matter
Types of Matter


Pure Substance- a single element or
compound
Mixture- contain more than one kind of
matter
Building Blocks of Matter


Atom
 The smallest unit of an
element that maintains
the properties of that
element
Element
 A pure substance made
of only one kind of
atom
Building Blocks of Matter

Compound
 A substance that is made from the atoms
of two or more elements that are
chemically bonded
Sucrose – C12H22O11
Sucrose is also known
as table sugar, and is
used to make Gummy
Bears!
Physical Properties
A characteristic that can be
observed or measured without
changing the identity of the
substance
Physical Properties

Extensive Property
Dependent upon the amount of matter
present
 Examples: volume, mass, and energy
content


Intensive Property
Independent of the amount of matter
present
 Examples: melting point, boiling point,
density

States of Matter
States of Matter

Solids
 very low kinetic energy particles vibrate but can’t
move around
 Retains size and shape
 Definite shape and volume
States of Matter

Liquids
 low kinetic energy particles can move
around but are still close
together
 Takes the shape of its
container
 definite volume
States of Matter

Gases
 high KE - particles can
separate and move
throughout container
 Easily compressed
 No definite shape
 No definite volume
States of Matter

Plasma
 very high KE - particles collide with
enough energy to break into
charged particles (+/-)
 gas-like, variable
shape & volume
 stars, fluorescent
light bulbs, CRTs
Physical Change
A change in a substance that does not involve a change in the identity
of the substance.
Examples:
Phase Changes – boiling
point, melting point,
freezing point
A substance dissolving
in another substance solubility
Separation Lab Simulation
How do you separate the following mixture based
on the physical properties:
Sand, salt, iron filings, waxy flakes, sawdust
A. List properties of each of the substances
B. Create a flow chart showing how you would
separate these items sequentially & return
them to the dry solid state.
Chemical Properties

Relates to a substances ability to
undergo change that transform it into a
different substance

Ability to : combust, oxidize, neutralize,
etc
Chemical Change
A change in which one or more substances
are converted into different substances.
Evidence of Chemical
Change:
 Heat and light
 Change in color
 Production of gas
 Precipitation of a solid
Physical & Chemical Quiz
Stage
Breathe in O2 &
CO2 out
Shower steam on
mirror
Boiled water at 100C
Stirred cocoa mix
Digested cereal
Brushed teeth
Lit bunsen burner
Marker ink on board
Change (P or C)
Specific property of
substance
Monday, September 20, 2010



Draw pictures of the 3 states of matter?
Identify the process which describe each change
in state of matter?
How is matter separated? What are the key
groupings?
Properties of Matter
Classification of Matter
(elements video)
Matter
Element
Compounds
Atoms
Mixtures
Molecules
Metals
Non-metals
Ionic
Homogeneous
Covalent
Solutions
Suspension
Metalloids
Solute
Heterogeneous
Solvent
Colloid
Classification of Matter


Pure Substance
 Fixed Composition
 A single compound or element
Mixtures
 A blend of two or more kinds of matter,
each with its own properties
 Components of mixtures can be separated
using physical properties
 Can be heterogeneous or homogeneous
Mixtures

Homogeneous Mixture – two or more
substances that are evenly distributed

Solution
Solute – matter which is present in smaller quantity
 Solvent – matter which is present in highest quantity


Heterogeneous solution – unevenly distributed
matter
Colloid – does not settle over time
 Suspension – settles over time

Classification of Matter
Tuesday, September 21, 2010

Classify each of the following kinds of matter:
Obsidian
 Marble
 Granite
 Asprin
 Tungsten

Compound or Mixture
Compound
Mixture
One kind of piece-
More than one kind -
Molecules
Molecule or atoms
Making is a
Making is a
chemical change
physical change
Only one kind
Variable composition

Chromatography Lab (2/2?)
Measure your chromatographs
Distance (start to end of each color)
 Distance (start to end of solvent front)


Calculate your Rf value
Rf value = Distance color moved
Distance to solvent front
Compare your Rf values for each solvent. Do any two
solvents have the same Rf value? Explain.
How could you prove that single colors from different
solvents are the same compound?
Separation of a Compound
The Electrolysis of water
Compounds must be
separated by chemical
means.
With the application of
electricity, water can
be separated into its
elements
Reactant

Water

H2O

Products
Hydrogen + Oxygen
H2
+
O2
Separation of a Mixture
The constituents of the mixture retain their
identity and may be separated by physical
means.
Methods of Separating Mixtures







Magnet
Filter
Decant
Evaporation
Centrifuge
Chromatography
Distillation
Filtration
separates a
liquid
from a
solid
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 40
Setup to heat a solution
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 42
A Distillation Apparatus
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 282
Centrifugation


Spin sample very rapidly: denser
materials go to bottom (outside)
Separate blood into serum and
plasma


Serum (clear)
Plasma (contains red blood cells
‘RBCs’)

Check for anemia (lack of iron)
AFTER
Before
Serum
Blood
RBC’s
A
B
C
The Periodic Table
Group or
family
Period
Periodic Table

Groups or Families
Vertical columns
 These elements have similar chemical
properties


Periods

Horizontal rows
Metals and Nonmetals



A zigzag line separates the metals from
the nonmetals
Metalloids, which straddle the line, are
considered non-metals
Lanthanide and Actinide Series

Metals
Properties of Metals
 Metals are good
conductors of heat and
electricity
 Metals are malleable
 Metals are ductile
 Metals have high
tensile strength
 Metals have luster
Examples of Metals
Potassium, K
reacts with
water and
must be
stored in
kerosene
Copper, Cu, is a relatively soft
metal, and a very good electrical
conductor.
Zinc, Zn, is
more stable
than potassium
Mercury, Hg, is the only
metal that exists as a
liquid at room temperature
Properties of Nonmetals
Carbon, the graphite in “pencil lead” is a great example of a
nonmetallic element.
 Nonmetals are poor conductors of
heat and electricity
 Nonmetals tend to be brittle
 Many nonmetals are gases at room
temperature
Examples of Nonmetals
Sulfur, S, was
once known as
“brimstone”
Graphite is not the only
pure form of carbon, C.
Diamond is also carbon;
the color comes from
impurities caught within
the crystal structure
Microspheres
of phosphorus,
P, a reactive
nonmetal
Properties of Metalloids
Metalloids straddle
the border between
metals and nonmetals
on the periodic
table.
 They have properties of both metals and
nonmetals.
Metalloids are more brittle than metals, less
brittle than most nonmetallic solids
 Metalloids are semiconductors of electricity
 Some metalloids possess metallic luster
Silicon, Si – A Metalloid
 Silicon has metallic luster
 Silicon is brittle like a nonmetal
 Silicon is a semiconductor of
electricity
Other metalloids include:





Boron, B
Germanium, Ge
Arsenic, As
Antimony, Sb
Tellurium, Te
Which is it?
Mixture
Element
Compound
Measuring Skills



Interpreting Scales – May be 1/10ths, 1/5ths,
1/4ths, 1/2s, 1s, 2s, 5s, 10s
Estimating Between the Lines – The last digit of
your measurement is estimated 1 digit beyond
what the hashmarks give.
Parallax & Meniscus
Estimating Between the Lines




You are expected to estimate 1 digit beyond what is
given by the lines.
The estimate will have an understood uncertainty of
+/-1 of the estimated last digit.
A final zero is used to indicate that the measurement is
right on the line.
Odd scales are estimated to round #s in between lines
with adjusted uncertainty.
Parallax & Meniscus





The meniscus is the curvature of water as it creeps up
the sides.
Always look for the bottom of the meniscus.
Parallax demo with fingers vs stars.
Parallax causes error when viewing is not straight on.
(Fuel gauge)
Graduated cylinders have hexagonal bases, plastic
bumpers & circles (for avoiding parallax errors)
Parallax & Meniscus

http://www.wisc-online.com/objects/index_tj.asp?objID=GCH302
The Metric System



A metric unit has two parts.
A prefix and a base unit.
prefix tells you how many times to divide or
multiply by 10.
Nature of Measurement
Measurement - quantitative observation
consisting of 2 parts

Part 1 - number
 Part 2 - scale (unit)


Examples:
20 grams
 6.63 x 10-34 Joule seconds

The Fundamental SI Units
(le Système International, SI)
Physical Quantity
Mass
Name
kilogram
Abbreviation
kg
Length
meter
m
Time
second
s
Temperature
Kelvin
K
Electric Current
Ampere
A
mole
mol
candela
cd
Amount of Substance
Luminous Intensity
Metric Units
UNIT
meter
liter
second
joule
gram
celsius
Kelvin
Mole
Molarity
(Abbreviation)
(m)
(L)
(s)
(J)
(g)
(oC)
(K)
(mol)
(M)
QUANTITY
distance
volume
time
heat
mass
temperature
absolute temperature
substance
concentration
Metric Prefixes
Prefix
nano
micro
milli
centi
deci
Deca
Hecto
Kilo
Mega
Giga
Sym
N
m
m
c
d
D
H
K
M
G
Size
Billionth
Millionth
Thousandth
Hundredth
Tenth
Ten x
Hundred x
Thousand x
Million x
Billion x
Decimal
.000000001
.000001
.001
.01
.1
10
100
1000
1000000
1000000000
Power
10-9
10-6
10-3
10-2
10-1
101
102
103
106
109
Examples of Common Metric Units
Metric Unit
 Liter
 milliliter
 Millimeter
 centimeter
 Meter
 kilometer
 Gram
 kilogram
Common Example
 1 L bottle of pop
 half teaspoon
 Thickness of dime
 Width of pinky nail
 Little more than a yard
 2.5 x around track
 Paperclip, 2 aspirin (500mg)
 Baseball bat (2.2 pounds)
Quiz – Units & Prefixes
1.
nanosecond
Billionth of a time unit
2.
kiloJoule
1000 units of heat
3.
Gigameter
Billion units of distance
4.
Microliter
Millionth of a volume unit
5.
Centimole
Hundredth of a unit of
substance
Converting
k h


D
d
c
m
how far you have to move on this chart, tells
you how far, and which direction to move the
decimal place.
The box is the base unit, meters, Liters,
grams, etc.
Conversions
k h

D
Change 5.6 m to millimeters
starts
at the base unit and move three to the right.
move
the decimal point three to the right
56 00
d
c
m
Conversions
k h




D
d
c
convert 25 mg to grams
convert 0.45 km to mm
convert 35 mL to liters
It works because the math works, we are
dividing or multiplying by 10 the correct
number of times.
m
Conversion Practice
Kilo1000
units
To convert to a smaller unit, move the
decimal point to the right or multiply.
Hecto
100
Dekaunits 10
units Base
Unit
Deci0.1
units
To convert to a larger unit, move
the decimal point to the left or
divide.
Centi0.01
Milliunits
0.001
units
Conversions
k h

D
Change 5.6 km to millimeters
d
c
m
Metric Units
•Convert 98 milligrams to grams.
98
mg
10-3 g
1 mg
=
0.098 g
025
Metric Units
•Convert 1.025 kilometers to meters.
km
103 m
1 km
= 1.025 x 103 m
.00
Metric Units
•Convert 52.00 microliters to liters.
µl
10-6 l
1 µl
= 5.200 x 10-5
l
Metric Units
•Convert 0.250 milliliters to microliters.
0.250
ml
10-3 l
1 ml
1 µl
10-6 l
=
2.50 x 102 µl
Multistep & Dimension Conversions
1.
2.
3.
4.
869ft = ?m
1.
2.5 yd2 = ?m2
0.63tons = ?oz
2.
47.2L = ?cm3
7960oz = ?kg
3.
3.6ft3 = ?m3
427,000in = ?km
4.
83,000ml = ?in3
Unit Conversions Quiz
1.
64 oz = ?g
2.
4,693mm = ?ft
3.
327in3 = ?L
Quiz - October 05, 2010

Convert the following metric units.
5.79 centimeters to meters
 7.5 Megaliters to liters
 8.94 centigrams to kilograms
 34.98 milligrams to kilograms
 0.093 micrograms to milligrams

Uncertainty in Measurement

A digit that must be estimated is
called uncertain. A measurement always
has some degree of uncertainty.
Why Is there Uncertainty?
 Measurements are performed with instruments
 No instrument can read to an infinite number of decimal
places
Which of these balances has the greatest uncertainty in
measurement?
Significant Figures (1 of 6)
A. Uncertainty


We will assume an uncertainty of +/- 1 of the
last significant figure.
Industry operates on an assumption of +/- 0.5
of the last significant figure unless otherwise
noted.
Significant Figures (2 of 6)
A. Uncertainty
Value
Dodge 3.8L
F100 351 cu. in.
Mazda B2200
(Manual) 2159 cu.in.
21.59 cm
2.159 cm
21,590 cm
215.90 cm
21,590,000 m
Uncertainty
+/- .1
+/- 1
+/- 100
+/- 1
+/- .01
+/- .001
+/- 10
+/- .01
+/- 10,000
Differences



Accuracy can be true of an individual
measurement or the average of several.
Precision requires several measurements before
anything can be said about it.
examples
Precision and Accuracy


Accuracy refers to the agreement of a particular
value with the true value.
Precision refers to

A) the degree of agreement among several
measurements made in the same manner.

B) the level of detail in a measurement.
Neither accurate
nor precise
Precise but not
accurate
Precise AND
accurate
Let’s use a golf anaolgy
Accurate?
Precise?
No
Yes
Accurate?
Yes
Precise?
Yes
Precise?
Accurate?
No
Maybe?
Accurate?
Precise?
Yes
We cant say!
Types of Error

Random Error (Indeterminate Error) measurement has an equal probability of being
high or low.

Systematic Error (Determinate Error) Occurs in the same direction each time (high
or low), often resulting from poor technique or
incorrect calibration.
Rules for Counting Significant
Figures - Details

Nonzero integers always count as
significant figures.
 3456
has
 4 sig figs.
Significant Figures (3 of 6)
B. Atlantic – Pacific Rule (page 26)
Pacific Rule
If the decimal is present,
then count from the
Pacific (left) side.
Start with the 1st non-zero
& every number after is
also significant.
Atlantic Rule
If the decimal is absent,
then count from the
Atlantic (right) side.
Start with the 1st non-zero
& every number after is
also significant.
Sig Fig Practice #1
How many significant figures in each of the following?
1.0070 m 
5 sig figs
17.10 kg 
4 sig figs
100,890 L 
5 sig figs
3.29 x 103 s 
3 sig figs
0.0054 cm 
2 sig figs
3,200,000 
2 sig figs
Significant Figures (3 of 4)
C. Maximum & Minimum Areas
7.6 cm+/-.1
3.4 cm+/-.1
Area
A(max)
A(mean)
A(min)
Area
= length x width
= 7.7cm x 3.5cm
= 26.95cm2
= 7.6cm x 3.4cm
= 25.84cm2
= 7.5cm x 3.3cm
= 24.75cm2
= 25.84cm2+/-1.11
= 26cm2
Significant Figures (6 of 6)
D. Rounding
Adding & Subtracting
Round to the greatest
uncertainty (least precision)
used in the problem.
Ex1:
29,000 gal
+ 1.4 gal
= 29,000 gal
Ex2:
44.76ml – 31.4ml
= 13.4ml
Multiplying & Dividing
Round to the least sig figs used
in the problem.
Ex1:
2.24g/7ml
= 0.3 g/ml
Ex2:
(4.6m)(3.25m)
= 15 m2
Mathematical Operations

Multiplication and Division: # sig
figs in the result equals the number in
the least precise measurement used in
the calculation.
 6.38m
x 2.0m =
 12.76m2  13m2 (2 sig figs)
Rules for Significant Figures in
Mathematical Operations

Addition and Subtraction: The
number of decimal places in the result
equals the number of decimal places in
the least precise measurement.
 6.8
+ 11.934 =
 18.734  18.7 (3 sig figs)
Sig Fig Practice #2
Calculation
Calculator says:
Answer
3.24 m x 7.0 m
22.68 m2
100.0 g ÷ 23.7 cm3
4.219409283 g/cm3
0.02 cm x 2.371 cm
0.04742 cm2
0.05 cm2
710 m ÷ 3.0 s
236.6666667 m/s
240 m/s
1818.2 lb x 3.23 ft
5872.786 lb·ft
5870 lb·ft
1.030 g ÷ 2.87 mL
2.9561 g/mL
2.96 g/mL
23 m2
4.22 g/cm3
Sig Fig Practice #3
Calculation
Calculator says:
Answer
3.24 m + 7.0 m
10.24 m
10.2 m
100.0 g - 23.73 g
76.27 g
76.3 g
0.02 cm + 2.371 cm
2.391 cm
2.39 cm
713.1 L - 3.872 L
709.228 L
709.2 L
1818.2 lb + 3.37 lb
1821.57 lb
1821.6 lb
2.030 mL - 1.870 mL
0.16 mL
0.160 mL
Significant Figures Quiz
How many sig figs?
1) 720,900kg
Round appropriately
5) 3.65g + 3.964g
2)
629.0m
6) 7180m – 62m
3)
.00076mg
7) 8.269m/240s
4)
.390s
8) (.039m)(4.260m)
Significant Figures Quiz

1.
2.
3.
4.
How many sig figs?
29.2 m
0.0036g
36,000cm
.00740L

1.
2.
Solve & round:
720m -84.23m
8.63m/800s
Significant Figures Quiz
How many sig figs?
1) 720,900kg
4 sf
2) 629.0m
4 sf
3) .00076mg
2 sf
4) .390s
3 sf
Round appropriately
5) 3.65g + 3.964g
7.61g
6) 7180m – 62m
7120m
7) 8.269m/240s
0.034m/s
8) (.039m)(4.260m)
.17m2
Rules for Counting Significant
Figures - Details

Zeros

Leading zeros do not count
as significant figures.
-
0.0486 has
 3 sig figs.

Rules for Counting Significant
Figures - Details


Zeros
- Captive zeros always count as
significant figures.


16.07 has
4 sig figs.
Rules for Counting Significant
Figures - Details


Zeros
Trailing zeros are significant only if
the number contains a decimal
point.


9.300 has
4 sig figs.
Rules for Counting Significant
Figures - Details

Exact numbers have an infinite
number of significant figures.
1
foot = 12 inches, exactly by
definition.
Wednesday, September 29, 2010

DMA
1.
2.
3.
4.
5.
6.
7.
Convert 5.7 cm to kilometers.
Convert 986 µg to grams.
Convert 893 kilograms to milligrams.
Convert 68.58 dg to megagrams.
Convert 986 µl to liters
Convert 83 grams to nanograms.
Convert 5 gigabytes to kilobytes
Thursday, September 30, 2010

DMA
1.
2.
3.
4.
5.
Convert 6.902 milligrams to grams
Convert 890,029 seconds to nanoseconds
Convert 9.30 joules to kilojoules.
Convert 495 kilopascals to micropascals.
Convert 293 Kelvin to Celcius
Scientific Notation:
A method of representing very large or very
small numbers in the form:
M x 10n
 M is a number between 1 and 10
 n is an integer
2 500 000 000
.
9 8 7 6 5 4 3 2 1
Step #1: Insert an understood decimal point
Step #2: Decide where the decimal must end
up so that one number is to its left
Step #3: Count how many places you bounce
the decimal point
Step #4: Re-write in the form M x 10n
2.5 x
9
10
The exponent is the
number of places we
moved the decimal.
0.0000579
1
2 3
4
5
Step #2: Decide where the decimal must end
up so that one number is to its left
Step #3: Count how many places you bounce
the decimal point
Step #4: Re-write in the form M x 10n
5.79 x
-5
10
The exponent is negative
because the number we
started with was less than 1.
PERFORMING
CALCULATIONS
IN SCIENTIFIC
NOTATION
ADDITION AND SUBTRACTION
Review:
Scientific notation expresses a
number in the form:
Mx
1  M < 10
n
10
n is an
integer
4 x 106
+ 3 x 106
6
7 x 10
IF the exponents are the
same, we simply add or
subtract the numbers in
front and bring the
exponent down
unchanged.
-
6
4 x 10
3 x 106
1 x 106
The same holds true for
subtraction in scientific
notation.
6
10
4x
5
+ 3 x 10
If the exponents are
NOT the same, we must
move a decimal to make
them the same.
5
6
40.0
4.00 x 10
+ 3.00
x
5
10
5
10
43.00
x
= 4.300 x
6
10
Student A
To avoid this
NO!
problem, move
 Is this good
the decimal on
scientific
the smaller
notation?
number!
6
10
4.00 x
6
5
.30xx10
10
+ 3.00
4.30 x
6
10
Student B
YES!
 Is this good
scientific
notation?
A Problem for you…
-6
10
2.37 x
-4
+ 3.48 x 10
Solution…
-6
-6
-4
002.37
2.37xx10
0.0237
10
x 10
-4
+ 3.48 x 10
-4
3.5037 x 10
Which is heavier?
it depends
Volume and Density
Relationship Between Volume and Density for Identical Masses of Common Substances
Substance
Cube of substance
(face shown actual size)
Mass
(g)
Volume
(cm3)
Density
(g.cm3)
Lithium
10
19
0.53
Water
Aluminum
10
10
1.0
10
3.7
2.7
10
0.58
Lead
11.4
Density of Some
Common Substance
Density of Some Common Substances
Substance
Density
(g / cm3)
Air
Lithium
Ice
Water
Aluminum
Iron
Lead
Gold
*at 0oC and 1 atm pressure
0.0013*
0.53
0.917
1.00
2.70
7.86
11.4
19.3
Density
D = M
V
M
M = DxV
ass
D
ensity
V
olume
V = M
D
Calculating


A piece of wood has a mass of 11.2 g and a
volume of 23 mL what is the density?
A piece of wood has a density of 0.93 g/mL and
a volume of 23 mL what is the mass?
Calculating

A piece of wood has a density of 0.93 g/mL
and a mass of 23 g what is the volume?
Floating





Lower density floats on higher density.
Ice is less dense than water.
Most wood is less dense than water.
Helium is less dense than air.
A ship is less dense than water.
Density of water



1 g of water is 1 mL of water.
density of water is 1 g/mL at 4ºC
otherwise it is less
Temperature vs Density (1 of 4)
Factor
Effect
Temperat Increasing
ure
temperature
decreases
density
Volume
expands &
contracts
Application
Hot air & water
rise while cold
stuff sinks.
Convection
currents form
in the earth,
oceans &
atmosphere
Materials vs Density (2 of 4)
Material Each
material
has its own
distinct
density
Density can be
used to ID
materials.
Materials can be
separated by
density as they
float or sink.
Concentration vs Density (3 of 4)
Factor
Effect
Concent Increasing
ration
concentration
increases
density
Application
Salt water sinks
below fresh water
in oceans &
estuaries. Melted
ice floats on top
of beverages.
Polluted waters
may sink to
bottom.
Size vs Density (4 of 4)
Size
Size has no
effect on
density
Small
samples can
be used to
characterize
large
quantities.
Heat
a form of energy
Definitions



Temperature is a measure of the average kinetic energy
of the particles in a sample of matter.
A joule is the SI unit of heat as well as all other forms
of energy.
Heat can be thought of as the energy transferred
between sample of matter because of a difference in
their temperatures.
Units for Measuring Heat
The Joule is the SI system unit for measuring
heat:
kg  m
1 J  N m 
2
s
2
The calorie is the heat required to raise the
temperature of 1 gram of water by 1 Celsius degree
1 calorie  4.18 Joules
A calorimeter is a
instrument that
measures the energy
absorbed or released as
heat in a chemical or
physical reaction.
Specific Heat
The amount of heat required to raise
the temperature of one gram of
substance by one degree Celsius.
Calculations involving Specific Heat
cp = ___q__
m x T
OR
q = cp x m x T
cp = Specific Heat
q = Heat lost or gained
T = Temperature change
Table of Specific Heats
Simple Problems

What is the specific heat of a substance that
absorbs 2.5 x 103 joules of heat when a sample
of 1.0 x 104 g of the substance increases in
temperature from 10.0C to 70.0C?
Problems


It takes 24.3 calories to heat 15.4 g of a metal
from 22 ºC to 33ºC. What is the specific heat of
the metal?
Iron has a specific heat of 0.11 cal/gºC. How
much heat will it take to change the temperature
of 48.3 g of iron by 32.4ºC?


A 1.0 kg sample of metal with a specific heat of
0.50 KJ/KgC is heated to 100.0C and then
placed in a 50.0 g sample of water at
20.0C. What is the final temperature of the
metal and the water?
3) A 2.8 kg sample of a metal with a specific
heat of 0.43KJ/KgC is heated to 100.0C then
placed in a 50.0 g sample of water at
30.0C. What is the final temperature of the
metal and the water?



Graphs
Circle graphs – compare
parts in a whole
Bar graphs – compare
quantities
Line graphs – compare
sets of data, show change
and patterns over time.
100
80
East
60
West
40
North
20
0
1st Qtr 2nd Qtr 3rd Qtr 4th Qtr
Proportions

Direct Proportion
y x

y
x
Inverse Proportion
1
y
x
y
x
Direct Proportions
 The quotient of two variables is a
constant
 As the value of one variable increases,
the other must also increase
 As the value of one variable decreases,
the other must also decrease
 The graph of a direct proportion is a
straight line
Inverse Proportions
 The product of two variables is
a constant
 As the value of one variable
increases, the other must
decrease
 As the value of one variable
decreases, the other must
increase
 The graph of an inverse
proportion is a hyperbola
High Quality Graphing
Make each graph at least ½ page large.
Use a ruler or straight edge to make it neat.
Place the manipulated variable on the x-axis.
Spread the range of data as far as possible.
Use logical spacing between numbers on axis.
Use hashmarks to clearly locate numbers.
Title the axis with quantity measured and (units in parentheses)
Accurately place each data point.
Use a best fit line or curve to show trend.
Title the graph with Responding vs Manipulated in Context
Features of Quality Graphs









Fills at least half of a page.
Manipulated variable on X & responding on Y axis.
Neatly constructed with straight edges.
Hash marks are evenly spaced.
Hash marks use round numbers and are spaced logically for
easy estimation in between.
X & Y Axes labeled with quantities and units in parentheses.
Title written as RV(Y) vs. MV(X) in the appropriate context.
Data points are accurately plotted.
Best fit line or curve is used to characterize pattern of data.
Joke of the Day

A group of organic molecules were having a
party, when a group of robbers broke into the
room and stole all of the guest's joules. A tall,
strong man, armed with a machine gun came
into the room and killed the robbers one by one.
The guests were very grateful to this man, and
they wanted to know who he was. He replied:

My name is BOND, Covalent Bond.