measurement in chemistry (and elsewhere). types of observations qualitative properties that can be...
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Measurement in Chemistry (and elsewhere)
Types of observations
QualitativeProperties 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…)
QuantitativeProperties that can be observed and described numerically and which result from measurement.
Commonly measured values in chemistry
Mass (grams) Volume (liters) Length (meters) Temperature (degrees Celsius or
degrees Kelvin) Time (seconds) Pressure (atmospheres) Concentration (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)
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) kilogram (kg)
Volume liter (L) meter3
TemperatureCelsius (C) Kelvin (K)
Common metric prefixes
Table 2.1Table 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) = Desired unit (Given unit)
12.4 kg x (1000 g) = 12,400 g (1kg)
1265 mm x (1 m) = 1.265 m (1 x 103 mm)
Exact and inexact numbers
Exact numbers No uncertainty to their value Value is known exactlyDefinedConversions within a systems
Inexact numbers Uncertainty of their true valueMeasuredConversions between different systems
Expressing numbers in scientific notation
Why do it?
How to enter them into your calculator
1.5 x 1023
1 . 5 EXP (or EE) 2 3
2.67 x 10-16
2 . 6 7 EXP (or EE) +/- 1 6
Making measurements
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
Only applies to measured (inexact) values
Does not apply to defined (exact) values
Measured ValuesMeasured ValuesWhat figures (digits) are significant?What figures (digits) are significant?
(not applied to defined or exact values such as conversions within the same system)
Non zeros are significant
Zeros between non zeros are significant
Zeros at the beginning are not significant
Zeros at end after decimal are significant
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 m
100 kg
1.0300 x 10-4 mg
Rules for working with measured values
Since there is uncertainty in measurement, we risk “amplifying” the uncertainty when we add, subtract, multiply and divide measured values
So…. There are rules for working with measured values
Calculations involving measured values
Multiplying and dividing:Answer can have no more total sig. figs. than the starting value with the fewest total sig. figs.
Adding and Subtracting:Answer can have no more sig. figs. after the decimal than any original number
Dimensional analysis helps solve conversion problems
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
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
calorie: amount of heat 1 cal raises 1 g of water 1° C
60 Calories = 60 kcal = 60,000 calories
Specific heat of any substanceAmount of heat (in calories) required to raise 1 gram of the substance 1° C