1. Scientific Measurement
Part 2:
Scientific Measurement

2. Importance of Measurement
Not all measurements give the same amount of information.
Qualitative Measurement
Quantitative Measurement
Descriptive
Cannot be measured (color, feel)
Nonnumerical
Numerical
Definite value
(numbers & units)

3. Scientific Notation
In chemistry, you will often encounter very small and very large numbers.
Scientific Notation
Makes numbers easier to work with and to write
Makes long numbers shorter

4. Scientific Notation What happened?
36,000 becomes 3.6 X 104
What happened?
The decimal point was moved either right or left to get only 1 number in front of decimal.
(Hint: If no decimal showing, it is after the last number!)

5. Scientific Notation If move RIGHT…. 0.00789 = 7.89 X 10-3
Number of places decimal point moves is the power of 10 and it is NEGATIVE
= X 10-3

6. Scientific Notation If move LEFT…. 366,000,000 = 3.66 X 108
Number of places decimal point moves is the power of 10 and it is POSITIVE
366,000,000 = X 108

7. Try it! Can you express the mass of a single atom of gold, which is
gram,
in scientific notation?

8. 3.27 X 10-22 gram It should be Try it! Answer

9. Try It! Can you express the distance to Mars, which is roughly
34,800,000 miles,
in scientific notation?
Hint: If no decimal point is showing, it is assumed to be behind the last number.

10. Try it! Answer
34,800,000 miles
It should be
3.48 X 107 miles!

11. Your Turn! Complete the Scientific Notation Worksheet.
Due next class period.
Be prepared for Scientific Notation Quiz next class period.

12. Uncertainty in Measurements
Accuracy
Precision
How “close” a measurement comes to the actual or true value
How close repeated measurements are to one another

13. Precision Increased Precision = More Decimals Most Precise 35111.001
When you report a measurement, you should only include digits if you’re really confident about their values.
The more significant figures in a measurement, the more precise that measurement must be!
For Example:
Increased Precision = More Decimals
Most Precise
35111
Least Precise

14. Significant Figures (aka) Sig Figs
Significant figures tell us:
something about the data
are important for measurements
are all the possible digits plus one guess on a measurement
We know 1 mm and 0.2 mm and then we guess at last digit 
So, we record
1.25 mm

15. Sig Figs Tell people how good our measuring tool was
PURPOSE:
Tell people how good our measuring tool was
Also shows how precise the measuring tool used was

16. Rules for Sig Figs
Rule #1: Any non-zero digits are significant seconds contains 3 sig figs

17. Rules for Sig Figs Rule #2:
All zeros between significant figures are significant.
307 mm = 3 sig figs

18. Rules for Sig Figs Rule #3: Leading zeros are not significant.
mL and mg = 3 sig figs each

19. Rules for Sig Figs Rule #4:
Trailing zeros are significant ONLY if after a decimal and after a significant figure.
L = 5 sig figs

20. Try it! 4 sig figs 2 sig figs 3 sig figs 2 sig figs 3 sig figs
Determine the number of significant figures in each of the following:
6.751 g
0.070 kg
28.0 mL
2500 m
400. L
30.07 g
4 sig figs
2 sig figs
3 sig figs
2 sig figs
3 sig figs
4 sig figs

22. + 67.68 mL 90.78 mL Should record 90.8 mL Try It!
For the results of an experiment you need to add 23.1 mL of HCl acid and mL of HCl acid. What results should you record? mL mL
90.78 mL
Should record 90.8 mL

23. Try It! 14.5 cm3 5.091034483 You should record 5.09 g/cm3
For the results of an experiment, you need to divide the mass of g by 14.5 cm3 to find the density of a material. What is the density you should record for the results?
73.82 g =
14.5 cm3
You should record 5.09 g/cm3

24. Your Turn! Complete the Sig Fig Worksheet with partner.
Be prepared for Sig Fig Quiz next class period.
When done with Sig Fig WS, turn in and work on Sig Fig Lab/Hwk with partner
Due next class period at start of class.

25. The SI System
Chemists measure physical properties all the time, such as mass, volume, length, or temperature.
They need a way to communicate their measurements to other chemists that is exact and agreed upon to use for a comparison.
SI System is based on metric system (base 10-units)

26. The SI System Units of Measurement
Quantity
SI Base Unit
Symbol
Non-SI Unit
Length
meter m
Volume
cubic meter m3
Liter L
Mass
kilogram kg
Density
grams/cubic centimeter
g/cm3
grams/milliliter
g/mL
Temperature
kelvin K
celsius 0C
Time
second s
Pressure
pascal Pa
atmosphere
atm
Energy
joule J
calorie
cal
Amt of substance
mole
mol

27. The SI System

28. How do I convert units?
Sometimes you start with one set of units and need to convert to another unit. How do we do this?
Use Conversion Factors (Dimensional Analysis)
Watch this helpful YouTube video!

29. The SI System Common Conversions Others 1 kg = 1000 g 1 m = 10 dm
1 m = 100 cm
1 m = 1000 mm
Others
1 mile = 5280 ft
1 inch = 2.54 cm
1 lb = 454 g

30. Now you Try It! x = 1000m 1 km 8.5 km 8500 km
How many meters are in 8.5 kilometers?
x =
1000m
1 km
8.5 km
8500 km

31. Try It!
Express a mass of grams in kilograms.
x =
1 kg
1000 g
5.712 g kg

32. Try One More! x x x = 24 m/s 86km hr 1000m 1 km 1 hour 60 min 1 min
Convert 86km/hr to m/s.
x x x =
86km hr
1000m
1 km
1 hour
60 min
1 min
60 sec
24 m/s

33. Your Turn! Be prepared for SI System Quiz next class period.
Now complete the SI System Worksheet #1 with a partner.
When finished, have Mrs. Breeding check it.
Then complete the SI System Worksheet #2 as homework. (Get it from Mrs. Breeding when you are done with Wksh #1)
Be prepared for SI System Quiz next class period.

34. AND apply it to the Effervescent Launcher Lab!
Your Turn!
Now you need to take all the information you learned about:
Scientific Method
Scientific Measurement (SI System)
Scientific Notation
Sig Figs
AND apply it to the Effervescent Launcher Lab!

35. Your Turn! Complete SI System Homework Sheet
Due at start of next class

36. Derived SI Units
Some units in the SI System are derived from multiplying or dividing standard units.
One such unit is DENSITY

37. Density Density = mass = m volume V
Density is the ratio of the mass of an object to its volume
i.e. how much “stuff” there is in some amount of space
Density = mass = m
volume V

38. Density
A 10.0 cm3 piece of lead has a mass of 114 g. What is the density of lead?
D = m = = V
114 g
10.0 cm3
11.4 g/cm3

39. Densities of Some Common Materials
Solids & Liquids
Density at 200C
(g/cm3)
Gases
(g/L)
Gold
19.3
Chlorine
2.95
Mercury
13.6
Carbon dioxide
1.83
Lead
11.4
Oxygen
1.33
Aluminum
2.70
Air
1.20
Table sugar
1.59
Nitrogen
1.17
Water (40C)
1.000
Neon
0.84
Corn oil
0.922
Ammonia
0.718
Ice (00C)
0.917
Methane
0.665
Ethanol
0.789
Helium
0.166
Gasoline
Hydrogen
0.084
Cork
0.24

40. You Try It! D = m = = 612 g 245 cm3 2.50 g/cm3
A student finds a shiny piece of metal that she thinks is aluminum. In the lab, she determines that the metal has a volume of 245 cm3 and a mass of 612 g. Calculate the density. Is the metal aluminum?
D = m = = V
612 g
245 cm3
2.50 g/cm3
So, her shiny metal is NOT aluminum.

41. You Try It!
The density of silver at 200C is 10.5 g/cm3. What is the volume of a 68 g bar of silver?
D =m so m = D * V then V = m
V D
V = 68 g = 6.5 cm3
10.5 g/cm3

42. Your Turn!
Complete the Density Worksheet which is due next class period.
Complete the “How Sweet It Is!” Lab.