1. Chapter 1: Let’s Review: The Tools of Quantitative Chemistry

2. Note About Math & Chemistry
Numbers and mathematics are an inherent and unavoidable part of general chemistry.
Students must possess secondary algebra skills and the ability to recognize orders of magnitude quickly with respect to numerical information to assure success in this course.
The material presented is considered to be a prerequisite to this course.
Optional Slide

3. Units of Measure
Science predominantly uses the “SI” (System International) system of units, which is similar to the “Metric System”.

4. Units of Measure
The base units are modified by a series of prefixes which you will need to memorize.

5. Temperature Units
Temperature is measured in the Fahrenheit, Celsius, and the Kelvin temperature scale.

6. Temperature Conversion
Convert °F to °C
TC = 5/9 ( TF – )
Convert °C to K
TK = TC

7. Length, Volume, and Mass In the metric system:
The base unit of length is the meter
1 m ≡ 100 cm inch ≡ 2.54 cm
The base unit of volume is the liter
1 L ≡ 1 dm L ≡ 1000 mL
1 mL ≡ 1 cm L ≈ quart
The base unit of mass is the gram
1kg ≡ 1000 g kg ≈ lb

8. Energy Units
Energy is confined as the capacity to do work. The SI unit for energy is the joule (J).
Energy is also measured in calories (cal)
1 cal = J
A kcal (kilocalorie) is often written as Cal.
1 Cal = 103 cal

9. Making Measurements: Precision
The precision of a measurement indicates how well several determinations of the same quantity agree.

10. Making Measurements: Accuracy
Accuracy is the agreement of a measurement with the accepted, or true value of the quantity.
Accuracy is reflected by the Experimental Error, often reported as the Percent Error:
Observed – True Value
Percent Error = ———————————— x 100%
True Value

11. Making Measurements: Standard Deviation
The Standard Deviation ( s ) of a series of measurements is equal to the square root of the sum of the squares of the deviations from the mean, divided by one less than the number of measurements (n – 1).
Measurements are often reported to  the standard deviation to report the precision of a measurement.

12. Mathematics of Chemistry
Exponential or Scientific Notation:
Most often in science, numbers are expressed in a format the conveys the order of magnitude.
3285 ft =  103 ft
kg = 2.15  103 kg

13. Exponential or Scientific Notation
1.23  104
Coefficient or Mantissa
(this number is 1 and
<10 in scientific notation
Exponential part
Base
Exponent

14. Mathematics of Chemistry
Significant figures: The number of digits represented in a number conveys the precision of the number or measurement.
A mass measured to  0.1 g is far less precise than a mass measured to  g.
1.1 g vs g
(2 sig. fig. vs sig. fig)
In order to be successful in this course, you will need to master the identification and use of significant figures in measurements and calculations!

15. Counting Significant Figures
All non zero numbers are significant
All zeros between non zero numbers are significant
Leading zeros are NEVER significant. (Leading zeros are the zeros to the left of your first non zero number)
Trailing zeros are significant ONLY if a decimal point is part of the number. (Trailing zeros are the zeros to the right of your last non zero number).

16. Determining Significant Figures
Determine the number of Sig. Figs. in the following numbers
1256
4 sf
7 sf
not trapped by a decimal place.
3 sf
5 sf
Instructor Note: Each individual component to this slide may be brought in on unit at a time using the “custom animation” tools in PowerPoint.
zeros written explicitly behind the decimal are significant…
1780
3 sf
770.0
4 sf
4 sf

17. Rounding Numbers 1. Find the last digit that is to be kept.
2. Check the number immediately to the right:
If that number is less than 5, then leave the last digit alone.
If that number is 5 or greater, then increase the previous digit by one.

18. Rounding Numbers Round the following to 2 significant figures: 1056007
Instructor note: Each individual component to this slide may be brought in on unit at a time using the “custom animation” tools in PowerPoint.
1780
1800

19. Sig. Figures in Calculations
Multiplication/Division
The number of significant figures in the answer is limited by the factor with the smallest number of significant figures.
Addition/Subtraction
The number of significant figures in the answer is limited by the least precise number (the number with its last digit at the highest place value).
NOTE: Counting numbers ( integers ), such as dimes, will never limit the calculation.

20. Sig. Figures in Calculations
Determine the correct number of sig. figs in the following calculation, and express the answer in scientific notation.
from the calculator:
4 sf
4 sf
2 sf
23.50   17 =
10 sf
Your calculator knows nothing of sig. fig. !

21. Sig. Figures in Calculations
Determine the correct number of sig. figs. in the following calculation, and express the answer in scientific notation.
in sci. notation:
 103
Rounding to 2 sf:
2.0  103

22. Sig. Figures in Calculations
Determine the correct number of sig. figs in the following calculation:
391
 12.6

23. Sig. Figures in Calculations
To determine the correct decimal to round to, align the numbers at the decimal place:
One must round the calculation to the least significant decimal place.
391
 12.6
391
no digits here
12.6

24. Sig. Figures in Calculations
391
-12.6
one must round to here
(answer from calculator)
round to here (units place)
Answer: 535

25. Sig. Figures in Calculations
Combined Operations: When there are both addition & subtraction and or multiplication & division operations, the correct number of sf is determined by examination of each step.
Example: Complete the following math mathematical operation and report the value with the correct # of sig. figs.
( )  ( ) = ???

26. Sig. Figures in Calculations
Example: Complete the following mathematical operation and report the value with the correct # of sig. fig.
( )  ( ) = ???
2nd determine the correct # of sf in the denominator (bottom)
1st determine the correct # of sf in the numerator (top)
The result will be limited by the least # of sf (division rule)

27. Sig. Figures in Calculations
3 sf
26.05
+ 32.1 =
58.150
7.7032
0.0032
+ 7.7
2 sf
The result may only have 2 sf

28. Sig. Figures in Calculations
Carry all of the digits through the calculation and round at the end.
3 sig fig
58.150
7.7032
2 sig fig =
= 7.5
Round to here
2 sf

29. Problem Solving and Chemical Arithmetic
Dimensional Analysis:
Dimensional analysis converts one unit to another by using a conversion factor (CF).
The resulting quantity is equivalent to the original quantity, it differs only by the units.
= unit (2)
unit (1)
 conversion factor

30. Problem Solving and Chemical Arithmetic
Dimensional Analysis:
Dimensional analysis converts one unit to another by using a conversion factor (CF).
Conversion factors come from equalities:
1 m = 100 cm
1 m
100 cm or

31. Problem Solving and Chemical Arithmetic
Dimensional Analysis:
Dimensional analysis converts one unit to another by using a conversion factor (CF).
Wanted Units
CF = —————————
Given Units

32. Examples of Conversion Factors
Exact Conversion Factors: CF in the same system of units, and CF determined by definition
1 m ≡ 100 cm
1 inch ≡ 2.54 cm
1 mile ≡ 5280 feet
Use of exact CF’s will not affect the significant figures in a calculation, because exact CF’s contain an infinite number of SF’s

33. Examples of Conversion Factors
Inexact Conversion Factors: CF that relate quantities in different systems of units
1.000 kg = lb
SI units
British Std.
(4 sig. figs.)
Use of inexact CF’s will affect significant figures.

34. Problem Solving and Chemical Arithmetic
Problem solving in chemistry requires “critical thinking skills”.
Most questions go beyond basic knowledge and comprehension.
You must first have a plan to solve a problem before you plug in numbers.
You must evaluate the result to see if it makes sense. (units, order of magnitude)
You must also practice to become proficient because...
Chem – is – try

35. Problem Solving and Chemical Arithmetic
Before starting a problem, devise a “Strategy Map”.
Use this to collect the information given to work your way through the problem.
Solve the problem using Dimensional Analysis.
Check to see that you have the correct units along the way.

36. Problem Solving and Chemical Arithmetic
Most importantly, before you start...
PUT YOUR CALCULATOR DOWN!
Your calculator wont help you until you are ready to solve the problem based on your strategy map.

37. Problem Solving and Chemical Arithmetic
Example: How many meters are there in 125 miles?
First: Outline of the conversion:
miles 
ft 
in 
cm  m
Each arrow indicates the use of a conversion factor.

38. Problem Solving and Chemical Arithmetic
Example: How many meters are there in 125 miles?
Second: Setup the problem using Dimensional Analysis: m
Miles
ft 
in 
cm  =

39. Problem Solving and Chemical Arithmetic
Example: How many meters are there in 125 miles?
Third: Check your sf and cancel out units. m =
miles 
ft 
in 
cm 
3 sf
exact / / / / / / / /

40. Problem Solving and Chemical Arithmetic
Example: How many meters are there in 125 miles?
Fourth: Now use your calculator: m / =
miles 
ft 
in 
cm 
3 sf
exact
Carry though all digits, round at end
x 105 m Round to: x 105 (3 sf)

41. Problem Solving and Chemical Arithmetic
Example: How many square feet are in 25.4 cm2 ?
Map out your conversion:
ft2 /
 10-2 ft2 =
cm2 
in2 
Round to:  ft2 (3 sf)
exact
3 sf

42. Problem Solving and Chemical Arithmetic
Example: How many cubic feet are in 25.4 cm3 ?
Map out your conversion:
cm3 
in3 
ft3 / / =
 10-4 ft3 / /
3 sf
exact
exact
Round to:  ft3 (3 sf)

43. Problem Solving and Chemical Arithmetic
Example: What volume in cubic feet would g of air occupy if it’s density is 1.29 g / L ?
Map out your conversion:
ft3
L 
in3 
cm3 
g  /
3 sf
exact

44. Problem Solving and Chemical Arithmetic
Example: How many picometers are in 25.4 nm? How many yards are in 25.4 m?

45. Problem Solving and Chemical Arithmetic
Unit conversions: How many cm3 are there in 25.4 L ?
1 L = 103 mL mL = 1 cm3

46. Problem Solving and Chemical Arithmetic
Unit conversions: How many grams are there in 5.67 pounds ?
1 kg ≈ lb kg ≡ 1000 g
lb → kg → g
1.000 kg g
5.67 lb x ———— x ——— = x 103 g
2.205 lb kg