Transcript Measurement in Chemistry (and elsewhere)

Measurement in Chemistry
(and elsewhere)
Types of observations

Qualitative
Properties that can be observed and
described that do not involve measurement.
If they do refer to quantities, they are vague
(ie fast, hot, large etc…)
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Quantitative
Properties that can be observed and
described numerically and which result from
measurement.
Commonly measured values in
chemistry
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Mass
Volume
Length
Temperature
Time
Pressure
Concentration
(grams)
(liters)
(meters)
(degrees Celsius or
degrees Kelvin)
(seconds)
(atmospheres)
(percent, molar)
Length, Mass and Volume

Length - distance between two points

Mass - amount of matter in an object
(weight is dependent upon the force of
gravity on the object)
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Volume - the amount of space an object
occupies
Length, Mass and Volume
Fig 2.2
Volume is derived from length
Some units of measurement
Metric Base Unit
SI Unit
Length
meter (m)
meter
Mass
gram (g)
Volume
liter (L)
meter3
Temperature
Celsius (C)
Kelvin (K)
kilogram (kg)
Common metric prefixes
Table 2.1
Common metric prefixes
1000 base / 1 kilo
1000 g / 1 kilogram
1 x 103
100 centi / 1 base
100 centimeters/ 1 meter
1 x 102
1000 milli / 1 base
1000 millimeters/ 1 meter 1 x 103
1,000,000 micro / 1 base
1,000,000 micrometers/1 meter 1 x 106
Converting between units
(Dimensional analysis – factor label method)
Given unit
x
(Desired unit)
(Given unit)
=
Desired unit
12.4 kg
x
(1000 g)
(1kg)
=
12,400 g
1265 mm
x
(1 m)
(1 x 103 mm)
=
1.265 m
Exact and inexact numbers

Exact numbers
No uncertainty to their value
Value is known exactly
Defined
Conversions within a systems
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Inexact numbers
Uncertainty of their true value
Measured
Conversions between different
systems
Expressing numbers in scientific
notation
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Why do it?
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How to enter them into your calculator
1.5 x 1023
1 .
5
2.67 x 10-16
2 . 6 7
EXP (or EE)
EXP (or EE)
2
3
+/-
1
6
Making measurements
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Accuracy: How close a measured value is to
the true value
Precision: How close multiple measured
values are to each other
There is estimation
(and therefore uncertainty)
in all measurements
Significant figures

The digits in a measurement that are known
with certainty, plus the single estimated digit
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Only applies to measured (inexact) values
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Does not apply to defined (exact) values
Measured Values
What figures (digits) are significant?
(not applied to defined or exact values such as conversions within the
same system)
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Non zeros are significant
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Zeros between non zeros are significant
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Zeros at the beginning are not significant
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Zeros at end after decimal are significant
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Zeros at end before the decimal  depend
Three ways to represent these zeros
How many significant figures are in
these measured values?
0.2304 cm
30.030 L
0.0034 mm
100 kg
1.0300 x 10-4 mg
Rules for working with measured values
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Since there is uncertainty in measurement,
we risk “amplifying” the uncertainty when we
add, subtract, multiply and divide measured
values
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So…. There are rules for working with
measured values
Calculations involving measured values
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Multiplying and dividing:
Answer can have no more total sig.
figs. than the starting value with the
fewest total sig. figs.
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Adding and Subtracting:
Answer can have no more sig. figs.
after the decimal than any original
number
Dimensional analysis helps solve
conversion problems
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What are you starting with?
What do you need to convert it into?
What conversion factor(s) do you need?
Must know conversions within the metric
system.
Must know other conversions we will
identify.
Do not have to memorize conversions
between systems.
English/Metric conversions (Table 2.2)
Density
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Mass of material per given volume
Commonly: grams/mL
SI: kg/m3
Density is a conversion factor for
converting between mass and volume
grams
(mL/g)
=
milliliter
milliliter (g/mL)
=
grams
Temperature scales
K = C + 273
C = K - 273
Calories and specific heat
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calorie: amount of heat
1 cal raises 1 g of water 1° C
60 Calories = 60 kcal = 60,000 calories
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Specific heat of any substance
Amount of heat (in calories) required to
raise 1 gram of the substance 1° C