1. Chapter 2
Metric System

2. The Metric System

4. Measuring The numbers are only half of a measurement. It is 10 long.
10 what?
Numbers without units are meaningless.
How many feet in a yard?
A mile?
A rod?

5. The Metric System Easier to use because it is a decimal system.
Every conversion is by some power of 10.
A metric unit has two parts.
A prefix and a base unit.
prefix tells you how many times to divide or multiply by 10.

6. Length - straight distance between two points -Meters (m)
Mass - how much matter in an object
-grams (g)
Volume - amount of space taken up by an object
Cubic meters (m3) or
-Liters (L)
Length
meter m
Mass
gram g
Time
Second s
Temperature
Kelvin K
Amount of a substance
Mole
Mol
Volume
Liter L

7. Base Units Length - meter - more than a yard - m
Mass - grams - about a raisin - g
Time - second - s
Temperature - Kelvin or ºCelsius K or ºC
Energy - Joules- J
Volume - Liter - half of a two liter bottle- L
Amount of substance - mole - mol

8. Metric System
Prefixes convert the base units into units that are appropriate for the item being measured.
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9. Prefix
Kilo-
Hecta-
Deka-
UNITS
Deci-
Centi-
Mili- k h da d c m

10. Prefixes giga- G 1,000,000,000 109 mega - M 1,000,000 106
kilo - k 1,
deci- d
centi- c
milli- m
micro- m
nano- n

11. Prefixes kilo k 1000 times deci d 1/10 centi c 1/100 milli m 1/1000
kilometer - about 0.6 miles
centimeter - less than half an inch
millimeter - the width of a paper clip wire

12. Using the units to solve problems
Dimensional Analysis
Using the units to solve problems

13. A Problem-Solving Method
Chapter 1: Chemistry: Matter and Measurement
A Problem-Solving Method
Chemistry problems usually require calculations, and yield quantitative (numerical) answers
The unit-conversion method is useful for solving most chemistry problems – the focus here is on “unit equivalents”
For example,
1 inch = 2.54 cm
EOS

14. 2km = ________________mm .34km = ________ cm 5km = ____________m
40mm = __________m
67m = __________ hm
135cm = _____________km
0.1km = _________dm
2km = ________________mm
.34km = ________ cm
5km = ____________m
19cm = __________mm
98m = __________km
135m = _____________km
33 km = _________dm
87m = ________________mm
For the first problem which color do you start with? Which color do you end with? How may spaces did you move your pencil? How many spaces should you move your decimal??

15. Converting k h D d c m
how far you have to move on this chart, tells you how far, and which direction to move the decimal place.
The box is the base unit, meters, Liters, grams, etc.

16. k h D d c m Conversions convert 25 mg to grams convert 0.45 km to mm
convert 35 mL to liters
It works because the math works, we are dividing or multiplying by 10 the correct number of times.

17. Other Equivalents and Conversion Factors
Chapter 1: Chemistry: Matter and Measurement
Other Equivalents and Conversion Factors
A conversion factor is the fractional expression of the equivalents
EOS

18. Dimensional Analysis
We use dimensional analysis to convert one quantity to another.
Most commonly dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm)
1 in.
2.54 cm
2.54 cm
1 in. or
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19. © 2009, Prentice-Hall, Inc.
Dimensional Analysis
Use the form of the conversion factor that puts the sought-for unit in the numerator.
Given unit 
 desired unit
desired unit
given unit
Conversion factor

20. Dimensional Analysis For example, to convert 8.00 m to inches, 100 cm
© 2009, Prentice-Hall, Inc.
Dimensional Analysis
For example, to convert 8.00 m to inches,
convert m to cm
convert cm to in.
100 cm 
1 m 
1 in.
2.54 cm 
315 in.
8.00 m

21. Dimensional Analysis Use conversion factors to change the units
1 foot = 12 inches (equivalence statement)
12 in = = 1 ft ft in
2 conversion factors
multiply by the one that will give you the correct units in your answer.

22. Chapter 1: Chemistry: Matter and Measurement
Two Examples
How many cm are in 26 inches?
2.54 1 cm
26 in ×
= 66 cm in
EOS

23. Examples 11 yards = 2 rod 40 rods = 1 furlong 8 furlongs = 1 mile
The Kentucky Derby race is 1.25 miles. How long is the race in rods, furlongs, meters, and kilometers?
A marathon race is 26 miles, 385 yards. What is this distance in rods and kilometers?

24. Units to a Power 3 1 m 100 cm 1 m 100 cm 1 m 100 cm 1500 cm3 1 m
How many m3 is 1500 cm3?
1 m
100 cm
1 m
100 cm
1 m
100 cm
1500 cm3
1 m
100 cm 3
1500 cm3

25. Units to a Power
How many cm2 is 15 m2?
36 cm3 is how many mm3?

26. Multiple units 65 mi hr 1760 yd 1 m 1 hr 1 min 1 mi 1.094 yd 60 min
The speed limit is 65 mi/hr. What is this in m/s?
1 mile = 1760 yds
1 meter = yds
65 mi hr
1760 yd
1 m
1 hr
1 min
1 mi
1.094 yd
60 min
60 s

27. Multiple units
Lead has a density of 11.4 g/cm3. What is this in pounds per quart?
454 g = 1 lb
1 L = qt

28. Uncertainy in Measurement
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Uncertainy in Measurement

29. © 2009, Prentice-Hall, Inc.
Significant Figures
The term significant figures refers to digits that were measured.
When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

30. Chapter 1: Chemistry: Matter and Measurement
Significant Figures
All digits in a number that are known with certainty plus the first uncertain digit
The more significant digits obtained, the better the precision of a measurement
The concept of significant figures applies only to measurements
Exact values have an unlimited number of significant figures
EOS

31. Significant Figures All nonzero digits are significant.
© 2009, Prentice-Hall, Inc.
Significant Figures
All nonzero digits are significant.
Zeroes between two significant figures are themselves significant.
Zeroes at the beginning of a number are never significant.
Zeroes at the end of a number are significant if a decimal point is written in the number.

32. Rules for Zeros in Significant Figures
Chapter 1: Chemistry: Matter and Measurement
Rules for Zeros in Significant Figures
Zeros between two other significant digits ARE significant
e.g., 10023
A zero preceding a decimal point is not significant e.g.,
EOS
Zeros between the decimal point and the first nonzero digit are not significant
e.g.,

33. Rules for Zeros in Significant Figures
Chapter 1: Chemistry: Matter and Measurement
Rules for Zeros in Significant Figures
Zeros at the end of a number are significant if they are to the right of the decimal point
e.g.,
EOS
Zeros at the end of a number may or may not be significant if the number is written without a decimal point
e.g., compared to 1000

34. Rules for Significant Figures in Calculations
Chapter 1: Chemistry: Matter and Measurement
Rules for Significant Figures in Calculations
KEY POINT: A calculated quantity can be no more precise than the least precise data used in the calculation
… and the reported result should reflect this fact
EOS
Analogy: a chain is only as strong as its weakest link

35. © 2009, Prentice-Hall, Inc.
Significant Figures
When addition or subtraction is performed, answers are rounded to the least significant decimal place.
When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation.

36. Significant Figures in Calculations
Chapter 1: Chemistry: Matter and Measurement
Significant Figures in Calculations
EOS
0.762 has 3 sigfigs so the reported answer is 1.39 m2
Multiplication and Division: the reported results should have no more significant figures than the factor with the fewest significant figures
1.827 m × m = ?

37. Significant Figures in Calculations
Chapter 1: Chemistry: Matter and Measurement
Significant Figures in Calculations
Addition and Subtraction: the reported results should have the same number of decimal places as the number with the fewest decimal places
EOS
NOTE - Be cautious of round-off errors in multi-step problems. Wait until calculating the final answer before rounding.

38. Conservation of Mass
Law of Conservation of Mass- in a physical or chemical reaction, mass is neither created nor destroyed; it is conserved.
All mass can be accounted for.
Mass of the Reactants = Mass of Products

39. Weight vs. Mass Measures the force of gravity on an object
Weight can change if the force of gravity acting on the object changes
How much matter is in an object
Remains constant (the same) no matter where it is

40. Mass and Weight Mass is measure of resistance to change in motion
Weight is force of gravity.
Sometimes used interchangeably
Mass can’t change, weight can

41. Mass Weight is a force. Mass is the amount of matter.
1 gram is defined as the mass of 1 cm3 of water at 4 ºC.
1000 g = 1000 cm3 of water
1 kg = 1 L of water

42. Mass
1 kg = 2.5 lbs
1 g = 1 paper clip
1 mg = 10 grains of salt

43. Volume calculated by multiplying L x W x H
Liter the volume of a cube 1 dm (10 cm) on a side
1L = 1 dm3
so 1 L = 10 cm x 10 cm x 10 cm
1 L = 1000 cm3
1/1000 L = 1 cm3
1 mL = 1 cm3

44. Volume 1 L about 1/4 of a gallon - a quart
1 mL is about 20 drops of water or 1 sugar cube

45. Uncertainty Basis for significant figures
All measurements are uncertain to some degree
Precision- how repeatable
Accuracy- how correct - closeness to true value.
Random error - equal chance of being high or low- addressed by averaging measurements - expected

46. Accuracy versus Precision
Accuracy refers to the proximity of a measurement to the true value of a quantity.
Precision refers to the proximity of several measurements to each other.
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47. Uncertainty Systematic error- same direction each time
Want to avoid this
Bad equipment or bad technique.
Better precision implies better accuracy
You can have precision without accuracy
You can’t have accuracy without precision (unless you’re really lucky).

48. Measuring Temperature
0ºC
Measuring Temperature
Celsius scale.
water freezes at 0ºC
water boils at 100ºC
body temperature 37ºC
room temperature ºC

49. Measuring Temperature
273 K
Measuring Temperature
Kelvin starts at absolute zero (-273 º C)
degrees are the same size
C = K -273
K = C + 273
Kelvin is always bigger.
Kelvin can never be negative.

50. Temperature is different
from heat.
Temperature is which way heat will flow. (from hot to cold)
Heat is energy, ability to do work.
A drop of boiling water hurts,
kilogram of boiling water kills.

51. Units of heat are calories or Joules
1 calorie is the amount of heat needed to raise the temperature of 1 gram of water by 1ºC.
A food Calorie is really a kilocalorie.
How much energy is absorbed to heat 15 grams of water by 25ºC.
1 calorie = 4.18 J

52. Conservation of Mass
Law of Conservation of Mass- in a physical or chemical reaction, mass is neither created nor destroyed; it is conserved.
All mass can be accounted for.
Mass of the Reactants = Mass of Products

53. Energy The ability to do work.
Work - cause a change or move an object.
Many types- all can be changed into the other.

54. Types of energy Potential- stored energy
Kinetic Energy- energy something has because its moving
Heat- the energy that moves because of a temperature difference.
Chemical energy- energy released or absorbed in a chemical change.
Electrical energy - energy of moving charges

55. Types of Energy
Radiant Energy- energy that can travel through empty space (light, UV, infrared, radio)
All types of energy can be converted into others.
If you trace the source far enough back, you will end up at nuclear energy.

56. Conservation of Energy
Energy can be neither created or destroyed in ordinary changes (not nuclear), it can only change form.
Its not just a good idea, its the law.