1. Chemistry Chapter 3 Scientific Measurement
WARNING:
Learn it now...it will be
used all year in Chemistry!!!

2. He is tall Electrons are tiny 3.1 Qualitative measurements Examples:
measurements that give results in a descriptive, non-numerical form.
Examples:
He is tall
Electrons are tiny

3. Quantitative measurements
measurement that gives results in a definite form, usually as numbers and units.
Examples:
He is 2.2 m tall
Electrons are 1/1840 times the mass of a proton

4. Scientific Notation
a number is written as the product of two numbers: a coefficient and 10 raised to a power.
Examples:
= 5.67 X 105
= 2.31 X 10-3

5. Examples: Convert to or from Scientific Notation: 2.41 x 102 241 =
662
.0034
241 =
6015 = =
.512 =
6.62 x 102 =
3.4 x 10-3 =

6. (This is important to master!!!)
CALCULATOR PRACTICE
(This is important to master!!!)
6.25 x x 102 =
(2.15 x 103)(6.1 x 105)(5.0 x 10-6) =
3.25 x =
3.6 x 107
FYI: “EE” button on calc= typing “X10^”
6.05 x 103
6.6 x 103
9.03

7. 3.2
Accuracy
the measure of how close a measurement comes to the actual or true value of whatever is measured.
how close a measured value is to the accepted value.
Precision
the measure of how close a series of measurements are to one another.

8. think about a
TARGET
Accurate and Precise
Precise
Neither

9. Percent Error Formula:
accepted value- experimental value x 100
accepted value
*always a positive number- indicated by the absolute value sign*
You will use this formula when checking the accuracy of your experiment.

10. Significant Figures – includes all of the digits that are known plus a last digit that is estimated.!
FYI: These rules are IMPORTANT and they will save you many points in the future if you learn them NOW!

11. Rules for determining Significant Figures 1
Rules for determining Significant Figures 1. All non-zero digits are significant.
1, 2, 3, 4, 5, 6, 7, 8, 9

12. 2. Zeros between non-zero digits are significant. (AKA zero sandwich)
102
7002

13. 3. Leading zeros (zeros at the beginning of a measurement) are NEVER significant.
00542
0.0152

14. 4. Trailing zeros (zeros after last integer) are significant only if the number contains a decimal point.
210.0
0.860
210

15. 5. All digits are significant in scientific notation.
2.1 x 10-5
6.02 x 1023
Time to practice!!

16. Examples: How many significant digits do each of the following numbers contain:
a) d) b) e) c) f) x 1023 2 2 2 5 3 4

17. Exact numbers have unlimited Significant Figures
Do not use these when you are figuring out sig figs…
Examples:
1 dozen = exactly 12
29 people in this room

18. < 5 round down (don’t change)
Rounding Rules:
 5 round up
< 5 round down (don’t change)
Examples:
Round to 1 significant digit =
Round to 3 sig. digs. =
Round to 2 =
Round 65,002 to 2 sig. digs. = 40
61.6
0.016
65,000

19. BellWork Thursday How many sig figs? 45,000 7.80 x 10-3 20040 0.00243
700
940.0
Round to 2 sig figs
Round to 3 sig figs

20. Addition and Subtraction
The measurement with the fewest significant figures to the right of the decimal point determines the number of significant figures in the answer.

21. Examples: Solve using correct significant figures 45.756 m + 62.1 m =
107.9

22. Multiplying and Dividing Measurements
The measurement with the fewest significant figures determines the number of significant figures in the answer.

23. Examples: Solve using correct significant figures:
3.43 m X m =
m X m =
45.01 m2 / m =
22.0 m2
55 m2
20. m
Notice the decimal!

24. Uncertainty
In lab, you record all numbers you know for sure plus the first uncertain digit. The last digit is estimated and is said to be uncertain but still considered significant.
Graduated cylinders have markings to the nearest mL (milliliter) and you will determine volume to the nearest 0.1 mL… because that is ONE DIGIT OF UNCERTAINTY.

25. International System of Units
revised version of the metric system
abbreviated SI
All units, their meanings and values can be found on pgs. 63,64,65.
Meter (m) –
Liter (L) –
Gram (g) –
SI unit for length
SI unit for volume
SI unit for mass

26. Mass – (g) amount of matter in an object
Density
Mass – (g) amount of matter in an object
Volume – (mL) amount of space occupied by an object
Density – (g/mL) a ratio of mass to volume

27. m D = Formula: Rewrite this formula to solve for m & v!
What is the unit for Density????
Remember: A material has the same density no matter how big or small it is!

28. Example:
A piece of metal has a volume of 4.70 mL and a mass of 57.3 g. What is the density?
M = 57.3 g
V = 4.70 mL
D = M / V
D = 57.3 g / 4.70 mL
D = g/mL
D = 12.2 g/mL

29. Graphs of mass and volume can be used to find the density of a substance. The slope of the line formed when mass and volume are plotted is the density. Remember “rise over run”.

30. Temperature
– measurement of the average kinetic energy of a system.

31. Temperature Scales Celsius
Sets the freezing point of water at 0C and the boiling point at 100C
Kelvin
Absolute zero is set as the zero on the
Kelvin scale. It is the temperature at
which all motion theoretically ceases.

32. To convert: K = ºC + 273 (Kelvin does not use “degrees”.)
-273 º C = 0 K = absolute zero

33. Examples: Convert 25 º C to Kelvin. K = ºC + 273 K = 25ºC + 273

34. Bell Work Friday Give Answer in proper sig figs 1. 1.36 - 0.253 ==
m m m =
3. ( 2.3 m)( 2.8 m) =
s2 / 1.25 s =