1. Chapter 3 “Scientific Measurement”
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Mt. Hebron High School
Chemistry

2. Units of Measurement
When you measure things, what units do you typically use?:
Feet, inches, miles, pounds, gallons..
This is referred to as the “English System.”
In science, the metric system is used because it is a universal language that anyone can understand regardless of what country they are in. 2

3. The Fundamental SI Units (Le Système International, SI)3

4. Other units and Variables used in Chemistry.
Area
Volume
Force
Pressure
Energy
Power
Voltage
Frequency
Electric Charge
Meter squared
Cubic meter
Newton
Pascal
Joule
Watt
Volt
Hertz
Coulomb 4

5. The Metric System
The units in the metric system are known as SI units (International System of Units).
Sometimes there are variables that are larger or smaller than the base unit, so we use metric prefixes to make measurements more convenient.
Examples of Base units… meter, liter, gram, pascal,… 5

6. Scientific Notation
In science it is very common to work with really large or really small numbers.
602,000,000,000,000,000,000,000
6.02 x 1023 atoms = 1 mole of a substance

7. Scientific Notation
In science it is very common to work with really large or really small numbers.
602,000,000,000,000,000,000,000
6.02 x 1023 atoms = 1 mole of a substance
_____ x (# of decimal places)
(+….moved left)
Number between (-…..moved right)
1 and 9

8. Rules for Scientific Notation
If the original number is larger than 1, then you move the decimal to the left and your exponent will be positive.
= x 103
* Moved the decimal point 3 places to the left.
If the original number is less than 1, then you move the decimal point to the right and your exponent will be negative.
= x 10-3
*Moved decimal point 3 places to the right.

9. If you take a number out of scientific notation look at the exponent first.
Negative Exponent = move to the left
Positive Exponent = move to the right

10. Try These: A. 43,812g = B. .000000943m = C. 5.66 x 10-3 L =
D x 106 g =
x 104
9.43 x 10-7
.00566
7,254,000

11. Uncertainity in Measurements
Measurements are an important aspect in Chemistry.
Three common measurements that we typically make are for Mass, Volume, and Temperature.

12. When you make a measurement, write down all the numbers you can. WHY??
Measuring instruments are never completely free from flaws.
Measuring always involves some estimation.
(Digital-Balance or Scale-Graduated Cylinder)

13. How do you know when a measurement is reliable?

14. Accuracy, Precision, and Error
It is necessary to make good, reliable measurements in the lab
Accuracy – how close a measurement is to the true value
Precision – how close the measurements are to each other (reproducibility)

15. Precision and Accuracy
Precise, but not accurate
Neither accurate nor precise
Precise AND accurate

16. Why Is there Uncertainty?
Measurements are performed with instruments, and no instrument can read to an infinite number of decimal places
Which of the balances shown has the greatest uncertainty in measurement?

17. Percents and Percent Error
How do you calculate a percentage?
Many times in Chemistry your results will need to be written as a percent.
Percent = (Actual / Total) x 100

18. Percent Error Percent Error measures how accurate you data really is.
measured value – accepted value x 100
accepted value
This answer should be really low.

19. Measurement Activity

20. Percent Error Calculate the percent error for each measurement in the activity.
Length of Clothespin cm
Length of Pen cm
Volume of cyl. 1
Volume of cyl. 2
Volume of cyl. 3
Volume of cyl. 4
Block 1 volume
Block 2 volume
Volume of rock
Volume of stopper
Mass of cup
Mass of scissors
Mass of water
Density of stopper
Temperature of liquid
8.2cm
15.2cm
54mL
10.2mL
83mL
35.5mL
27cm3
54cm3
2.5mL
10.1mL
354.56g
32.53g
10.0g
1.37g/mL
23.5o

21. SI Prefixes Commonly Used

22. Prefix Symbol Meaning Mega M Kilo K / k Hecta H Deca D Base Unit
106
Kilo
K / k
103
Hecta H
102
Deca D
101
Base Unit
g, m, L,sec, mol 1
deci d
10-1
centi c
10-2
milli m
10-3
micro 
10-6
nano n
10-9
pico p
10-12

23. 1 Kilometer equals how many meters?
1000 meters because the Kilometer prefix represents a thousand. So since there is 1 Kilometer, one times 1000 is 1000 meters.

24. How many meters in 1000 centimeters?
10 meters because the centi prefix represents one hundreth of a meter.
We have to divide 1000 centimeters by 100 to get our answer.
Easy trick to remember…up the chart you divide or down the chart you multiply the number of steps you move.

25. Try These 550 millimeters as meters. 3.5 seconds as milliseconds
1.6 Kilograms to grams
2500 milligrams to Kilograms
4.00 centimeters to micrometers
2800 decimoles to moles
6.1 amperes to milliamperes
3 Kilograms to milligrams
.550m
3500ms
1600g
.0025 kg
40,000 um
280 mol
6100 ma
3,000,000mg

26. Significant Figures in Measurements
Significant figures in a measurement include all of the digits that are known, plus one more digit that is estimated.
Measurements must be reported to the correct number of significant figures.

27. Which measurement is the best?
Significant Figures
Which measurement is the best?
What is the measured value?
What is the measured value?
What is the measured value?

28. Rules for Significant Figures
Zeroes are only significant when they are not a place holder.
A Place Holder
B m
not significant
C Kg All significant
D. 1,809,000 L

29. Rules for Significant Figures
Multiplication or Division
: The number with the fewest significant figures determines how many are in the answer.
3.05m x 2.10m x 0.75m
= m3
= 4.8 m3

30. Rules for Significant Figures
Addition or Subtraction
: Look for the fewest decimal places to determine the number of significant figures.
6.41g g + 11g g = g
= 31g
5.84L L L = L
= 46.0L

31. Density Which is heavier: 1 lb. of lead or 1 lb. of feathers?
Most people compare masses, but in chemistry we not only compare masses, but densities as well.
Density is a mass-to-volume ratio.
Density = mass / volume
Density depends on the composition of the substance, not how much is present.

32. Density = mass / volume

34. Density Density can be calculated 2 ways: 1. Mathmatically 2. Slope
Mathmatically D = m/v
Slope---- Line Graph of Mass vs. Volume
slope = rise / run
= mass / volume

35. Density Problems

36. Density Graphs

37. Density Lab