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Bell Work
Prepare for Quiz
Write these in your Bell Work
Composition Book
• Hg
• Br
• Kr
•K
•S
• Atomic #
• Atomic weight
• # of protons
• # of electrons
• # of neutrons
• Group
• Period
Write these in your Bell Work
Composition Book
• Hg
• Br
• Kr
•K
•S
Find the Atomic Weight for MgSO4
• Bell work composition book
Today’s Element
Liquid Metals
Zinc
Chemical Properties
Importance
Physical Properties
Periodic Table
•
•
•
•
Atomic #
Atomic mass
# of Protons
# of Electrons
• # of Neutrons
• Period
• Group
What’s in a Battery?
•
•
•
•
•
Modern batteries use a variety of chemicals to power their reactions. Common
battery chemistries include:
Zinc-carbon battery: The zinc-carbon chemistry is common in many inexpensive
AAA, AA, C and D dry cell batteries. The anode is zinc, the cathode is manganese
dioxide, and the electrolyte is ammonium chloride or zinc chloride.
Alkaline battery: This chemistry is also common in AA, C and D dry cell batteries.
The cathode is composed of a manganese dioxide mixture, while the anode is a
zinc powder. It gets its name from the potassium hydroxide electrolyte, which is an
alkaline substance.
Lithium-ion battery (rechargeable): Lithium chemistry is often used in highperformance devices, such as cell phones, digital cameras and even electric cars. A
variety of substances are used in lithium batteries, but a common combination is a
lithium cobalt oxide cathode and a carbon anode.
Lead-acid battery (rechargeable): This is the chemistry used in a typical car
battery. The electrodes are usually made of lead dioxide and metallic lead, while
the electrolyte is a sulfuric acid solution.
Water: Separation by Electrolysis
Video of Electrolysis: Water to Hydrogen and Oxygen
Atomic Mass of a Compound
• H2O
H2 1.01 x 2 = 2.02
O 16 x 1 = 16.00
1.01 + 1.01 + 16 = 18.02
Try these
• CO2
• C6H12O6
Add totals 2.02
+ 16.00
18.02
Practice – Finding Atomic Mass
• CO2
C 12.01 x 1 = 12.01
O2 16 x 2 =
32.00
Add totals 12.01
+ 32.00
44.01
Practice – Finding Atomic Mass
• C6H12O6
C6 12.01 x 6 = 72.06
H12 1.01 x 12 = 12.12
O2
16 x 6 = 96.00
Add totals 72.06
12.12
+96.00
108.18
Percent Composition of Mass for
Mixtures
• A 6g mixture of sulfur and iron is separated
using a magnet.
Data Sulfur (S)
Iron (Fe)
5g
1g
• Calculate the percent composition of S and Fe.
Percent Composition of Mass for
Mixtures
• A 6g mixture of sulfur and iron is separated
using a magnet.
Data Sulfur (S)
Iron (Fe)
5g
1g
• Calculate the percent composition of S and Fe.
Part / Whole x 100 = % composition
Sulfur: 5g/6g x 100 =
Iron : 1g/6g x 100 =
Percent Composition of Mass for
Mixtures
• A 6g mixture of sulfur and iron is separated
using a magnet.
Data Sulfur (S)
Iron (Fe)
5g
1g
• Calculate the percent composition of S and Fe.
Part / Whole x 100 = % composition
Sulfur: 5g/6g x 100 = 83.33% S
Iron : 1g/6g x 100 = 16.66% Fe
Use Percent Composition to find the
composition of a compound
• Use the periodic table to find the compound’s
percent composition of each element.
• List the atomic weight of each element in the
compound
• Note how many of each type of atom is in the
compound
• Add it all up to get the atomic weight of the
whole compound
Atomic Mass of a Compound
• H2O
H2 1.01 x 2 = 2.02
O 16 x 1 = 16.00
1.01 + 1.01 + 16 = 18.02
Try these
• CO2
• C6H12O6
Add totals 2.02
+ 16.00
18.02
Practice – Percent Composition
H2O
part / whole x 100 = % composition
% composition of H
% composition of O
Practice – Percent Composition
CO2
part / whole x 100 = % composition
% composition of C
% composition of O
Practice – Percent Composition
C6H12O6
part / whole x 100 = % composition
% composition of C
% composition of H
% composition of O
Law of Conservation of Mass
• Mass is neither created nor destroyed in any
process. It is conserved.
Mass reactants = Mass products
2H2O + electricity yields 2H2 + O2
Isotopes
• The atomic weight found on the periodic table
is based on the average weight of all the
isotopes of the element
• Isotope – atoms of the same element with the
same number of protons but different
numbers of neutrons
• M&M activity
Writing Isotopes
Reading Isotopes
Mass number - the sum of
the protons and neutrons
Isotopes of Hydrogen
Write Isotopes for Iron
More isotopes
Argon 36,
Argon 37…
M&Mium Isotope Activity
http://www.chem.memphis.edu/bridson/Fund
Chem/T07a1100.htm
S.I.Units
• http://2012books.lardbucket.org/books/gener
al-chemistry-principles-patterns-andapplications-v1.0/section_05.html#averill_1.0ch01_s09_s01_s02_t02
http://chemwiki.ucdavis.edu/Physical_Chemist
ry/Atomic_Theory/The_Mole_and_Avogadro's
_Constant
Measurements
and
Calculations
Where to Round Song
Steps in the Scientific Method
• 1.
Observations
•
quantitative
•
qualitative
• 2.
Formulating hypotheses
•
possible explanation for the
observation
• 3.
Performing experiments
•
gathering new information to decide
whether the hypothesis is valid
Outcomes Over the Long-Term
•
•
Theory (Model)
A set of tested hypotheses that give
an
overall explanation of some natural phenomenon.
•
•
Natural Law
The same observation applies to
many
different systems
•
-
Mass
Example - Law of Conservation of
Law vs. Theory
 A law summarizes what happens
 A theory (model) is an attempt to explain why
it happens.
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)
International System of Units
a system of measurement units agreed on by scientists to aid in
the comparison of results worldwide.
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
SI Units
Metric Prefixes
Common to Chemistry
Prefix
Unit Abbr.
Exponent
Kilo
k
103
Deci
d
10-1
Centi
c
10-2
Milli
m
10-3
Micro

10-6
Nano
n
10-9
Metric Prefixes and Conversion Examples
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?
Precision and Accuracy
•
Accuracy refers to the agreement of a
particular value with the true value.
•
Precision refers to the degree of
agreement among several measurements
made in the same manner.
Neither
accurate nor
precise
Precise but not
accurate
Precise AND
accurate
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.
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 inch = 2.54 cm, exactly
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
Rules for Significant Figures in Mathematical
Operations
•
Multiplication and Division: # sig figs in the
result equals the number in the least precise
measurement used in the calculation.
• 6.38 x 2.0 =
• 12.76  13 (2 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 4.22 g/cm3
23 m2
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
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 #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
Scientific Notation
In science, we deal with some very LARGE
numbers:
1 mole = 602000000000000000000000
In science, we deal with some very SMALL
numbers:
Mass of an electron =
0.000000000000000000000000000000091 kg
Imagine the difficulty of calculating the
mass of 1 mole of electrons!
0.000000000000000000000000000000091 kg
x 602000000000000000000000
???????????????????????????????????
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
7 x 106
IF the exponents are the
same, we simply add or
subtract the numbers in
front and bring the
exponent down
unchanged.
-
4 x 106
3 x 106
1 x 106
The same holds true for
subtraction in scientific
notation.
106
4x
+ 3 x 105
If the exponents are NOT
the same, we must move
a decimal to make them
the same.
5
6
40.0
4.00 xx 10
10
+ 3.00
x
43.00
= 4.300 x
5
10
x
5
10
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
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