1. Measurement in Chemistry (and elsewhere)

2. 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…)
Quantitative
Properties that can be observed and described numerically and which result from measurement.

3. Commonly measured values in chemistry
Mass (grams)
Volume (liters)
Length (meters)
Temperature (degrees Celsius or degrees Kelvin)
Time (seconds)
Pressure (atmospheres)
Concentration (percent, molar)

4. 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

5. Length, Mass and Volume
Fig 2.2

6. Volume is derived from length

7. Some units of measurement
Metric Base Unit SI Unit
Length meter (m) meter
Mass gram (g) kilogram (kg)
Volume liter (L) meter3
Temperature Celsius (C) Kelvin (K)

8. Common metric prefixes
Table 2.1

9. Common metric prefixes
1000 base / 1 kilo g / 1 kilogram x 103
100 centi / 1 base 100 centimeters/ 1 meter 1 x 102
1000 milli / 1 base millimeters/ 1 meter 1 x 103
1,000,000 micro / 1 base 1,000,000 micrometers/1 meter 1 x 106

10. 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) = m
(1 x 103 mm)

11. Exact and inexact numbers
Exact numbers No uncertainty to their value
Value is known exactly
Defined
Conversions within a systems
Inexact numbers Uncertainty of their true value
Measured
Conversions between different
systems

12. Expressing numbers in scientific notation
Why do it?
How to enter them into your calculator
1.5 x 1023
EXP (or EE)
2.67 x 10-16
EXP (or EE) /

13. Making measurements
Accuracy: How close a measured value is to the true value
Precision: How close multiple measured values are to each other

14. There is estimation (and therefore uncertainty) in all measurements

15. 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

16. Measured Values 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

17. How many significant figures are in these measured values?cm L mm
100 kg
x mg

18. 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

19. 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

20. 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.

21. English/Metric conversions (Table 2.2)

22. 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

23. Temperature scales
K = C
C = K

24. 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 substance
Amount of heat (in calories) required to raise 1 gram of the substance 1° C