Download CP Chemsitry Unit 6 Notes KEY

CP Chemistry 2016-2017
Unit 6 Notes
Name:
Accuracy
Density
Precision
Newtons
milli
micro
Important Vocabulary
Signif figures
Dimensional Analysis
kilo
Pascals
Mega
Celsius
Fahrenheit
SI
Base unit
Joules
centi
nano
Section 6-1
Systems of measurement:
Name
Where it is used
1. English/Imperial
USA
Percent error
weight
Derived unit
Giga
Kelvin
Advantages
Human appendages
makes estimating simple
Disadvantages
Difficult
conversions
Example units: miles, feet, inches
2. Metric
Rest of the world
Easy conversions
Need to know
prefixes
Easy conversions
Need to know
prefixes
Example units: meter, seconds, kilograms
3. System International (SI)
Science
Example units: meters, seconds, kilograms (note that Kelvin is used and not Celsius)
The 7 SI base units
Name
1 Meter
2 Second
3 Kilogram (not grams)
4 Kelvin
5 Luminous intensity
6 Ampere
7 mole
Symbol
Quantity it measures
distance
Time
Mass
Temperature
candela
Electric current
Amount of substance
M
s
kg
K
cd
A
mol
Examples of standards for the SI base units:
The meter is defined at the distance travelled by light in a vacuum in 1/299,792,458 seconds.
One second is the time that elapses during 9,192,631,770 (9.192631770 x 10 9 ) cycles
of the radiation produced by the transition between two levels of the cesium 133 atom.
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SI derived units
A Derived unit is: a combination of more than 1 base unit
Examples of derived units
Newton = Kgm/s2
Joule = Nm = kgm2/s2
Pascal = N/m2
unit of force
unit of energy
unit of pressure
Density = g/cm3
Volume = m3 (note that ALL volume units are derived and that the Liter is NOT a base unit!)
SI Prefixes to know
Prefix Name Symbol
Meaning
Exponential factor
Example
1 Giga
G
1,000,000,000
1 x109
1 x106 m = 1 Gm
6
2 Mega
M
1,000,000
1 x10
1 x106 m = 1 Mm
3 kilo
K
1000
1 x103
1000m = 1km
4 centi
C
0.01
1 x10-2
1 m = 100 cm
5 milli
m
0.001
1 x10-3
1 m = 1x103 mm
-6
6 micro
µ
0.000001
1 x10
1 m = 1x106 µm
7 nano
N
0.000000001
1x10-9
1 m = 1x109 nm
-12
8 pico
P
0.000000000001 1 x10
1 m = 1x1012 pm
Converting between metric and English units(these equalities given on exam)
1 mile = 1.61 km
2.54 cm = 1 inch
2.2 lbs = 1 kg
Converting between temperature scales
What are the 3 temperature scales?: Celsius, Fahrenheit, Kelvin
What are the equations to convert between them (given on exams)?
K = °C + 273
°C = K -273
°C= (5/9)( °F – 32)
°F = (9/5)( °C) + 32
Examples:
a.
b. T = 70 °F
T = 300 K
°C = (5/9)(70-32) = 21°C
°C = 300 – 273 = 27°C
d. T = -20 °F
°C = (5/9)(-20-32) = -29°C
K = -29 + 273 = 244
e.
2
c. T= 100 °C
8K = 100 + 273 = 373 K
f.
Examples of dimensional analysis problems
a. convert 5.0 km into m
5 km *
1000 m
1 km
= 5000 m
b. Convert 200. lbs into kg
200 lbs *
1 kg
2.2 lbs
= 91 kg
c. Convert 2.5 hrs into seconds
2.5 hrs *
60 min
1 hr
*
60 sec
1 min
= 9000 s
d. Convert 4.5 mm into m
4.5 mm *
1m
= 0.0045 m
1000 mm
e. Convert 0.500 kg into mg
0.500 kg *
1000g
1 kg
*
f. Convert 5.0 miles into inches
5.0 miles *
5280 ft
1 mi
g. Convert 1 yr into seconds
1 yr * 365.25 days
1 yr
*
*
h. Convert 2.45 nJ into MJ
2.45 nJ *
1J
*
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1 x10 nJ
i. Convert 75 MW into cW
75 MW *
1 x106 W
1 MW
j. Convert 0.0050 µs into ks
0.0050 µs * 1 s
*
1 x106 µs
1000 mg
1g
12 in
1 ft
= 5.0 x105 mg
= 316,800 in
24 hrs
1 day
*
1 MJ
1 x 106 J
= 2.45 x10-15 MJ
*
100 cW
1W
1 ks
= 5.0 x10-12 ks
1000 s
3
3600 s
1 hr
= 75 x108 cW
= 3.16 x107s
Section 6-2
Density equation:
density = mass/volume
How to rearrange the equation: One can switch the variables V and D and cross multiply to get M = DV
V = M/D
Why is density important? (examples):
How do metal ships float in water? Isn’t the metal more dense than the water?
Why do regular sodas sink to the bottom of a cooler when diet sodas float?
Density example problems (copy from board):
1. Mass of sample of zinc = 20.0 grams
If the density of zinc is 7.14 g/mL, how much
Volume does the zinc sample occupy?
2. A sample of gold has a volume of 40 cm3.
If the density of gold is 19.3 g/mL, what is the
mass of the gold sample?
V = M/D = 20 g/(7.14 g/mL) = 2.80 mL
M = DV = 19.3 g/mL (40 cm3) = 772 g
3.A sample of metal displaces 30.0 mL of
water. If the mass of the sample is 210. grams,
is the sample zinc?
4. A cube of aluminum has a length of one side
of 4.0 cm. What is the mass of the cube?
The density of aluminum is 2.7 g/cm3.
D = m/v = 210/30 = 7.0 g/cm3; yes
V cube = (4.0 cm)3 = 64 cm3
M = DV = (2.7 g/cm3)(64 cm3) = 173 g
5.
A gold bar is a rectangular shape and has the
dimensions of 12 inches long by 4 inches high
by 5 inches wide. What is the mass of this bar?
V = 240 in3
240 in3 *
2.54 cm
1 in
*
2.54 cm
1 in
*
2.54 cm
1 in
= 3933 cm3
M = DV = (19.3 g/cm3) (3933 cm3) = 75, 905 g = 75.9 kg = 167 lbs
Scientific Notation Review
Why we use scientific notation: to show the proper amount of significant figures. Yes, it also helps with
writing large and small numbers in manageable form.
Writing large numbers in scientific notation: The decimal point is moved to the LEFT until there is one
non zero integer. The number of places moved is equal to the exponent.
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Examples: 5000. = 5.0 x103
12,000,000. = 1.2 x107
Writing small numbers in scientific notation: The decimal point is moved to the RIGHT until there is one
non zero integer. The number of places moved is equal to the negative exponent.
Examples: 0.002 = 2.0 x10-3
0.000 000 005
= 5.0 x10-9
Adding and subtracting in scientific notation examples:
The exponents must be the same to add or subtract
2.0 x105 + 1.0 x103 = 200 x103 + 1.0 x103 = 201 x103 OR 2.01 x105
Multiplying and dividing in scientific notation examples:
Multiplying = add exponents; division = subtract
(5.0 x102)(2.0 x103) = 5.0 x108
2.0 x10-2
Significant Figures
Definition: all certain(known) digits and 1 estimated digit
Reasons why we use them:
1. To relay the proper number of measured digits from a measurement
2. To convey the type of measuring instrument used
What do significant figures tell us?:
Examples
a. Ruler
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b. Balances
0.01 g precision = about $10.0
0.001 g precision = about $200.0
0.0001 g precision = about $2500.0
Rules for Significant Figures (SF)
1. All non zero integers are significant
Examples
123 = 3 SF
512.5 = 4 SF
2. “trapped” zeros are significant
505 = 3 SF
2002.1 = 5 SF
3. “trailing” zeroes are significant
UNLESS they are before a non-existent decimal point
500.000 = 6 SF
0.5050 = 4 SF
1000 = 1 SF
20 = 1 SF
5000. = 4 SF
4. “leading” zeroes are NOT significant
0.00005 = 1 SF
0.00009090 = 4 SF
* note that writing these in scientific notation tells exactly how many SF there are
5. Some numbers have an infinite number of SF
Pi = 3.1415……..
2/3 = 0.66666666666……
1 m = 100 cm is infinitely true
5 people = 5.000000000……
Determine the number of SF in the following (copy from board)
a. 5.000
b. 50.00090
d. 0.005
e. 20.0 grams
g. 10.0
h. 32.0050
j. 100
k. 100.00050
m. 5 people
n. 90.0 cars
p. 0.000006
q. 6/7
6
c. 1000 mm = 1 m
f. 25.00000
i. 0.909090
l. 50000
o. 2 cars
r. 20.0 L
Practice determining the number of SF in the following
0.50
2
0.070
2
2.000 4
90
1
1.010 4
20 apples
inf
0.90901 5
0.0000910
3
100 cm = 1m inf
50 nickels
inf
18.00
4
0.100000
5
0.0010010
5
800
1
1.0
0.00009
1000
50.
12
5.0100
55.0010
2
1
1
2
2
5
6
Guidelines for rounding numbers
If the last number removed is a 5 or above, then round the last number kept up one digit.
If the last number removed is a 4 or below, leave the last number kept alone.
Rule for adding and subtracting SF: round to the least number of decimal places in the original numbers
Examples:
a.
5.000 + 4.0 - 1.00
= 8.0
b.
5.0 + 9.0000 + 1.
10.
c.
d.
16.0 + 6. - 10.000
12.
10.00 - 6.0 + 15.00
= 19.0
Rule for multiplying and dividing SF: round to the least number of SF in the original numbers
Examples:
a.
10.0(5.0)(1.)
c.
= 5 x101
50.0(5.0)(1.00) = 2.5 x102
b.
(15.00)(1.) = 3
(5.00)
d.
(1.000)(5.) = 5
(1.0)
What happens when you have different operations in the same problem?
One must separate the rules from each other and perform them separately
Examples:
1. 1.000 + 2.0(5) = 1.000 + 1 x101
= 1 x101
2. 10.00 - 2.0 = 8
1.0
7
3. 50.0 – 2.0 + (40.0)(2)
= 50.0 – 2.0 + 8 x101
= 128 = 1.3 x102
4. 100.0 (2.00) + 5.0
2
2
= 1 x10 + 5.0
= 1 x102
Section 6-3
Accuracy: the agreement between the observed value and the accepted/true value
Precision: the agreement between different experimental values
Examples of the difference between accuracy and precision:
1.Dartboards
2. A ruler: the actual length of an object is 10.0 cm
a. 12.0 cm is recorded
b. 12.0 cm is recorded
These observations are not very accurate, but they are precise. Perhaps the person systematically read the
ruler incorrectly?
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Percent Error (know this equation – not given on exams)
Why we use percent error: the determine accuracy
The equation is: % error = (observed value – true/accepted value) X 100
True value
Examples:
1.In an experiment, the density of water
was determined to be 1.2 g/mL
% error = (1.0 – 1.2)
1.0
x 100 = 20%
2. The mass of a sample of gold was recorded as 20.0 grams. When placed into a graduated
cylinder with 50.0 mL of water, the volume changed to 51.5 mL. What is the percent error for
the density determination?
V = 51.5 – 50.0 = 1.5 mL
D = m/V = 20.0g/1.5 mL = 13.3 g/mL Should be 2 SF = 13 g/mL
TV = 19.3 g/mL (from earlier in the notes)
% error = (13 – 19.3) x100 = - 6 x 100 = 31.09 % = 30%
19.3
19.3
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