Transcript Unit 2 PowerPoint - Mr. Kinton`s Science Classes

Mr. Kinton’s Honors Chemistry
THE MATHEMATICS OF CHEMISTRY
EXACT NUMBERS
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definitive values
Can be counted
Conversion factors
MEASUREMENT
INEXACT NUMBERS
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Measured quantities
Have error
Limitations in equipment
MEASUREMENT
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We are concerned about two things:
 Precision:
how individual measurements agree with
one another
 Accuracy: how individual measurements agree with
the “true” value
SCIENTIFIC UNITS
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In everyday life we use
Standard or Customary
units
In other places in the
world, the metric system
is used.
Science uses SI units
SI UNITS
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There are 7 SI Base Units:
CONVERTING UNITS
Scientists prefer the metric and SI units
because unit conversion is easier
 This is because every unit is some multiple of
10
 Convert to have meaningful measurements
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CONVERTING USING DIMENSIONAL ANALYSIS
Chemists use this as a way of canceling out
units and solving problems involving math.
 This method requires us to know certain
conversion factors.
 What are some examples of conversion factors
that you have used before?
 Let’s look at some examples!
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UNIT CONVERSION
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Here is an easy way to remember conversions:
CONVERSION
Here is how it works: Larger to smaller
 Convert 50 kg to g
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 Start
with the unit given -50 kg
 Determine how many grams are in a kilogram
 Use dimensional analysis or the factor label method
 50 kg
1,000g
50,000 g
1 kg
CONVERSION
Here is how it works: Smaller to larger
 Convert 5000 mL to L
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 Start
with the units given- 5000 mL
 Determine how many L are in a mL
 Use dimensional analysis
 5000 mL
1L
5L
1000 mL
YOU TRY
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Make the following conversions
 252
dam (decameter) to dcm (decimeter)?
 51 cL (centiliter) to hL (hectoliter)?
FINAL NOTE ON CONVERSIONS
Mega (M)- 106, 1Mm= 100,000 m
 Micro (u)- 10-6, 1um= 0.000001 m
 Nano (n)- 10-9, 1nm = 0.000000001 m
 We will use these measurements later in the
course
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SIGNIFICANT FIGURES
Digits of a measured quantity including the
uncertain one
 For Example, let’s examine these 2 balances
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HOW TO COUNT SIGNIFICANT FIGURES
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Zeros within a number are always significant
 4308
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and 40.05 each have 4 significant figures
Zeros at the beginning of a number are not
significant
 0.0026
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has only 2 significant digits
Zeros at the end of a number and after the
decimal are significant
 0.0200
and 3.00 have 3 significant digits
COUNTING SIGNIFICANT FIGURES
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Numbers ending in zero depend on a decimal
point
130 is only 2 significant figures
 130. is 3 significant figures
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How many significant figures are present?
105
 0.005
 40.0
 220
 2220.
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MULTIPLICATION/DIVISION
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Answer must contain a
number with the fewest
significant figures
Ex) Area = (6.221 cm)(5.2
cm) = 32.3492 cm2 = 32
cm2
ADDITION/SUBTRACTION
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Answer must align with the
fewest number of decimal
places
Ex) 20.4 + 1.322 + 83 =
104.722 = 105
SIG FIGS IN CALCULATIONS
MULTIPLICATION/DIVISION
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230 x 12 =?
0.4058/0.003 =?
5482.3/25 =?
74.077 x 2.100 x 16.0037
=?
ADDITION/SUBTRACTION
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230 + 12 =?
0.4058 – 0.003 =?
5482.3 + 25 =?
74,077 + 2,100 + 16,003.7
=?
PRACTICE WITH SIG FIGS
SCIENTIFIC NOTATION
Used to remove ambiguity of zeros at the end
of a number
 Example: 10,300 g has how many significant
figures?
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 Using
Scientific Notation, up to 5 sig figs can be
given
 1.03 x 104
 1.030 x 104
 1.0300 x 104
DENSITY
Amount of mass in a unit volume
 Density = mass/volume
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 Units
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g/cm3 or g/mL
Temperature dependent
SAMPLE PROBLEMS
Calculate the density of mercury if 1.00 x 102 g
occupies a volume of 7.36 cm3
 Calculate the volume of 65.0 g of liquid
methanol (wood alcohol) if its density is 0.791
g/mL
 What is the mass in grams of a cube of gold
(density = 19.32 g/cm3) if the length of the
cube is 2.00 cm.
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YOU TRY!
Calculate the density of a 374.5 g sample of
copper with a volume of 41.8 cm3
 A student needs 15.0 g of ethanol. If the
density of ethanol is 0.789 g/mL, how many
milliliters are needed.
 What is the mass, in grams of 25.0 mL of
mercury (density = 13.6 g/mL)
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