1. Mathematics in Chemistry
Lab 1

2. Outline Mathematics in Chemistry MicroLAB™ Units Rounding
Digits of Precision (Addition and Subtraction)
Significant Figures (Multiplication and Division)
Order of Operations
Mixed Orders
Scientific Notation
Logarithms and Antilogarithms
Algebraic Equations
Accuracy and Precision
Statistics
Serial Dilutions
Direct Dilutions
Graphing
Calibration Curves
MicroLAB™
The Program
Reference Sheet
Pitfalls

3. Mathematics in Chemistry
Math is a very important tool, used in all of the sciences to model results and explain observations.
Chemistry in particular requires a lot of calculations before even trivial experiments can be performed. In this first exercise you will be introduced to some of the very basic calculations you will be required to perform in lab during the entire semester.
Remember, if you start memorizing rules and formulas now, you don’t have to do it the night before your exams!

4. Units Units are very important! Units give dimension to numbers.
They also allow us to use dimensional analysis in our calculations.
If a unit belongs next to a number, place it there!!!
Example: 6.23 mL
The unit “mL” indicates to us that our measurement is a metric system volume and indicates to us the order of magnitude of that volume.

5. Why is this method statistically more correct?
Rounding
When you have to round to a certain number to obey significant figure rules, remember to do the following:
For numbers 1 through 4, round down
For numbers 6 through 9, round up
For numbers with a terminal 5, round to the closest even number.
rounded to three significant digits becomes
rounded to three significant digits becomes
rounded to three significant digits becomes
rounded to three significant digits becomes ?
Why is this method statistically more correct?

6. Digits of Precision and Significant Figures
All measurements have some degree of uncertainty due to limitations of measuring devices.
Scientists have come up with a set of rules we can follow to easily specify the exact amount of significant figures, without sacrificing the accuracy of the measuring devices.

7. Digits of Precision: Addition and Subtraction
Your answer must contain no more digits after the decimal point than the number with the least number of digits after the decimal point.
= 315.6

8. Digits of Precision: Addition and Subtraction – =

9. Addition of Whole Numbers
When you add or subtract whole numbers, your answer cannot be more accurate than any of your individual terms.
– 100 = 2400
What about:
,650 = ?

10. Addition of Whole Numbers
When you add or subtract whole numbers, your answer cannot be more accurate than any of your individual terms.
– 100 = 2400
What about: ,650 = ?
The answer is 35,840

11. Significant Figures Rule #1
Numbers with an infinite number of significant digits do not limit calculations. These numbers are found in definite relationships, otherwise known as conversion factors.
100 cm = 1 m
1000 mL = 1 L

12. Significant Figures Rule #2
All non-zero digits are significant.
1.23 has 3 significant figures
98,832 has 5 significant figures
How many significant digits does have?

13. Significant Figures Rule #2
All non-zero digits are significant.
1.23 has 3 significant figures
98,832 has 5 significant figures
How many significant digits does have?
Correct! The answer is 4.

14. Significant Figures Rule #3
The number of significant figures is independent of the decimal point.
12.3, 1.23, and have 3 significant figures
and have how many significant figures?

15. Significant Figures Rule #3
The number of significant figures is independent of the decimal point.
12.3, 1.23, and have 3 significant figures
and have how many significant figures?
Correct! The answer is 4.

16. Significant Figures Rule #4
Zeros between non-zero digits are significant.
1.01, 10.1, have 3 significant figures.
How many significant digits are in 10,101?

17. Significant Figures Rule #4
Zeros between non-zero digits are significant.
1.01, 10.1, have 3 significant figures.
How many significant digits are in 10,101?
The answer is 5!

18. Significant Figures Rule #5
After the decimal point, zeros to the right of non-zero digits are significant.
has 3 significant figures has 2 significant figures.
How many significant figures are in ?

19. Significant Figures Rule #5
After the decimal point, zeros to the right of non-zero digits are significant.
has 3 significant figures has 2 significant figures.
How many significant figures are in ?
This one has 6 significant digits.

20. Significant Figures Rule #6
If there is no decimal point present, zeros to the right of non-zero digits are not significant.
3000, 50000, 20 all have only 1 significant figure
How many significant figures are in 32,000,000?

21. Significant Figures Rule #6
If there is no decimal point present, zeros to the right of non-zero digits are not significant.
3000, 50000, 20 all have only 1 significant figure
How many significant figures are in 32,000,000?
The answer is 2!

22. Significant Figures Rule #7
Zeros to the left of non-zero digits are never significant.
0.0001, 0.002, 0.3 all have only 1 significant figure
How many significant figures are in ?
How many significant figures are in ?

23. Significant Figures Rule #7
Zeros to the left of non-zero digits are never significant.
0.0001, 0.002, 0.3 all have only 1 significant figure
How many significant figures are in ?
This one has 3 significant digits.
How many significant figures are in ?
This one has 4 significant digits.

24. Significant Figures: Multiplication and Division
Your answer must contain no more digits total than the number with the least number of digits total.
5.10 x x = 172

25. Significant Figure Multiplication and Division=

26. Order of operations
1st: ( ), x2, square roots
2nd: x or /
3rd: + or –

27. Significant Figure Mixed Orders

28. Scientific Notation
The three main items required for numbers to be represented in scientific notation are:
the correct number of significant figures
one non-zero digit before the decimal point, and the rest of the significant figures after the decimal point
this number must be multiplied by 10 raised to some exponential power
123 becomes 1.23 x 102
This number has three significant digits

29. Scientific Notation
Calculators could be a significant aid in performing calculations in scientific notation.
KNOW HOW TO USE YOUR CALCULATOR
Does your calculator retain or suppress zeros in its display?
In converting between scientific and decimal notation, the number of significant digits don’t change.

30. Scientific Notation Conversions
What is the scientific notation equivalent of ? ?
What is the standard decimal notation equivalent of 3.45 x 103?
6.500 x 10-2?

31. Scientific Notation
What is the scientific notation equivalent of ? The answer is 4.32 x 10-2
? The answer is x 103
What is the standard decimal notation equivalent of 3.45 x 103? This is 3450
6.500 x 10-2? This is

32. Scientific Notation Calculations
Addition:
(4.22 x 105) + (3.97 x 106)
= (4.22 x 105) + (39.7 x 105)
= ( ) x 105
= 43.9 x 105
= 4.39 x 106
Know how to perform these types of calculations on your calculator!

33. Scientific Notation Calculations
Subtraction:
(4.22 x 105) - (3.97 x 106)
= (4.22 x 105) - (39.7 x 105)
= (4.22 – 39.7) x 105
= x 105
= x 106
Know how to perform these types of calculations on your calculator!

34. Scientific Notation Calculations
Multiplication:
(4.22 x 105) x (3.97 x 106)
= (4.22 x 3.97) x 10(5+6)
= 16.8 x 1011
= 1.68 x 1012
Know how to perform these types of calculations on your calculator!

35. Scientific Notation Calculations
Division:
(4.22 x 105) / (3.97 x 106)
= (4.22 / 3.97) x 10(5-6)
= 1.06 x 10-1
Know how to perform these types of calculations on your calculator!

36. Logarithms
Logarithms might seem strange, but they are nothing more than another way of representing exponents.
logbx = y is the same thing as x = by
Know how to use your calculator to perform these functions.

37. Logarithms
We see logarithms frequently when working with pH chemistry. If you have a solution of pH 5.2, and you need to calculate the concentration of hydrogen ions, set the problem up as follows:
pH = - log [H+]
5.2 = - log [H+]
-5.2 = log [H+]
= 10log [H+]
= [H+]
[H+] = 6.3 x 10-6

38. Did you notice anything?
Logs and Antilogs
To enter log 100 on your calculator:
Press: log  1  0  0  Enter or
Press: 1  0  0  log for reverse entry
To enter the antilog 2 on your calculator:
Press: 2nd  log  2  Enter
Press: 2  2nd  log for reverse entry
Did you notice anything?

39. Significant Figure Rules
Logarithms
log (4.21 x 1010) = 
Antilogarithms
antilog (- 7.52) = = x 10-8  3.0 x 10-8

40. Significant Figures of Equipment Electronics
Always report all the digits electronic equipment gives you.
When calibrating a probe, the digits of precision of your calibration values determine the digits of precision of the output of the data.

41. Algebraic Equations
It is important to understand how to manipulate algebraic equations to determine unknowns and to interpolate and extrapolate data. Don’t forget about significant figures.
For y = x
If x = 0.021, find y (answer = 0.024)
If y = 4.3, find x (answer = 4.0)

42. Accuracy
The accuracy of a measurement represents a comparison of the measured value (experimental value) to the “true” value.
A measure of accuracy is indicated by:
Percent Error =
Tolerances of glassware affect the accuracy of volume measurements.

43. Precision
Precision of a measurement reflects reproducibility of an experimental procedure.
Refer to the bull’s eye experiment on page 60.
Graduations on glassware affect the precision of the glassware in question.

44. Statistics
We use statistics in the laboratory in order to validate our results.
We evaluate the central tendency of our work by calculating the mean (related to accuracy) of our data.
We evaluate the variability in our work by calculating the standard deviation (s) (related to precision) of our data.
The relative standard deviation gives us a more meaningful number than the standard deviation.

45. Calculation of the Mean
xi = individual values
N = number of measurements
For significant figures, always keep as many digits after the decimal point as the original values. Remember units!

46. Calculation of the standard deviation of a set of numbers
xi = individual values
= the average of the individual values
N = number of measurements
For significant digits, report the same digits of precision as the xi values. The units are the same as the units for the x values.

47. Calculation of Relative Standard Deviation
RSD% =
s = standard deviation of a set of data
= average of the individual measurements
The calculation itself dictates the number of significant digits. What would the units be?

48. Dilutions
Using a solution of known concentration for the preparation of a solution with a lower concentration is commonly called dilution.

49. Solution Preparation from Solids
Determine the mass of the solid needed. You will need the following:
Molar mass of the solid
Total volume desired
Final concentration desired
Calculation:
m = M x MM x V
g = mol/L x g/mol x L
Remember the precision of your glassware!

50. Solution Preparation from Solids
Make the solution:
Weigh out the appropriate mass of solid.
Place a small volume of distilled water in the volumetric flask.
Add the solid to the volumetric flask.
Add some more distilled water to the flask, stopper, and invert several times.
Add distilled water to the calibration line (fill to volume) using a medicine dropper, stopper, and invert several times.

51. Solution Preparation from Liquids
Determine the volume of stock solution needed.
You will need the following:
Concentration of stock solution (M1)
Desired concentration of diluted solution (M2)
Desired volume of diluted solution (V2)
Calculation:
M1V1 = M2V2
Remember the precision of your glassware!

52. Solution Preparation from Liquids
Make the solution:
Obtain the appropriate volume of stock solution using a graduated cylinder. (Always add a few mL extra.)
Place a small volume of distilled water in a volumetric flask.
Use the appropriate pipet to transfer the correct volume of stock solution from the graduated cylinder to the volumetric flask.
Add some more distilled water to the flask, stopper, and invert several times.
Add distilled water to the calibration line (fill to volume) using a medicine dropper, stopper, and invert several times.

53. Serial Dilution
Serial dilution is a laboratory technique in which substance concentration is decreased stepwise in series.

54. Standard dilution
Standard dilution is a laboratory technique in which stock solution is used to prepare a diluted solution.

55. Graphing
Graphing is an important tool used to represent experimental outcomes and to set up calibration curves.
It is a modeling device.

56. Graphing: Variables Having no fixed quantitative value.
X-variable
Y-variable
Graphing in chemistry
Renamed with a chemistry label
Paired with a unit most of the time

57. Graphing: Units Give dimension to labels / variables
Give meaning to numbers
Essential!

58. Graphing: Coordinates
A coordinate set consists of an x-value and y-value, plotted as a point on a graph.
X-values: domain (independent variable)
Y-values: range (dependent variable)

59. Graphing: Axes Multiple axes on a graph
Coordinate sets determine the number of axes on a plot
Two dimensional graphs have only two axes
X-axis
Y-axis
Each axis must have a consistent scale

60. Graphing in Chemistry Graph title reflects the:
Dependent vs. Independent variables
X-axis – labeled appropriately with variable and unit
Y-axis – labeled appropriately with variable and unit
Each axis has a consistent scale

61. Graphing in Chemistry Coordinate sets are plotted
x-variable matching the x-value on the x-axis
y-variable matching the y-value on the y-axis
A single point results
A line is drawn through all the points
An equation is derived from two coordinate sets
The equation is used to find unknowns

62. Graphing: Equations Of the form y = mx + b m = slope of the graph
b = y-intercept of the graph
x = any x-value from the graph
y = corresponding y-coordinate

63. Graphing [Ni2+], M Absorbance 0.200 0.041 0.300 0.063 0.400 0.085
Label / Variable
Units
[Ni2+], M
Absorbance
0.200
0.041
0.300
0.063
0.400
0.085
0.500
0.101
Let’s look at a graphical representation of the following data:

64. Graphing
Graph Title
Best-fit Line
Graph Axis Labels

65. Graphing
Always add a title to your graph. On the previous slide, the title was
Absorbance vs. [Ni2+], M
because absorbance was graphed on the y-axis and [Ni2+], M was graphed on the x-axis.
(Always y vs. x!)
If a graph title is
Temp F, degrees vs. Temp C degrees
what should be graphed on the x-axis?

66. If an axis is labeled with Temp F, degrees
Graphing
Always label your axis appropriately. The label for the y-axis is Absorbance. Absorbance has no units, so none are listed. The label for the x-axis is [Ni2+] and the units M. What does “M” stand for?
If an axis is labeled with Temp F, degrees
which one is the unit?

67. Graphing
If your data points look like they fall on a line, be sure to add a “linear” calibration curve to them. If they don’t appear linear, DO NOT add a linear line.
When you add a calibration curve, an equation results. This equation describes the line and can help you solve unknown values. The equation on our graph was: y = 0.20 x

68. Calibration Curves
A calibration curve gives you a graphical representation of an instrument’s response to a particular analyte.
If we were to declare your 1992 Ford Escort an “instrument” and the gas it uses an “analyte,” we could construct a calibration curve for:
Distance Driven, miles vs. Gas Consumed, gallons

69. Calibration Curves The measurements that are made are all made:
with the same vehicle
using the same set of tires
driving under similar road and environmental conditions
using the same type of gas

70. Data Table of Standards
Distance Driven, miles
Gas Consumed, gallons
25.3
1.0
49.2
2.0
73.9
3.0
98.2
4.0
122.8
5.0

71. Calibration Curve
Notice that the average gas consumption of this vehicle is 24 mpg!

72. Unknowns
When we talk about unknown “analytes,” we are referring to an unknown measurement, not an unknown identity.
If we were to analyze three unknowns related to our previous example…we are still talking about gas, the unknown measurement refers to either the gallons of gas consumed, or the distance driven.
No matter which unknown we are trying to determine, our analysis must be made under the same conditions as previously, in other words, unknown measurements should be made:
with the same vehicle
using the same set of tires
driving under similar road and environmental conditions
using the same type of gas

73. Data Table of Unknowns Distance Driven, miles Gas Consumed, gallons
41.5 ?
82.1
103.6

74. Distance = 24 (Gas Consumed) + 0.7
Calibration Equation
We can use the previously determined calibration equation to determine how many gallons of gas it would take to drive the number of miles indicated on the previous slide.
Since Distance Driven was plotted on the y-axis and Gas Consumed was plotted on the x-axis, the new equation becomes:
Distance = 24 (Gas Consumed) + 0.7

75. Calibration Equation Distance = 24 (Gas Consumed) + 0.7
Let’s solve this equation for Gas Consumed:
Gas Consumed = (Distance – 0.7) / 24
Let’s solve for our unknowns:

76. Solving Unknowns Gas Consumed = (Distance – 0.7) / 24
= 1.7 gallons
Gas Consumed = (82.1 – 0.7) / 24
= 3.4 gallons
Gas Consumed = (103.6 – 0.7) / 24
= 4.3 gallons

77. Laboratory Computer Etiquette
Do not surf the web
Do not check unrelated to this course
Do not print materials unrelated to this course
Do not connect a USB mass storage drive
No social networking!!!
Do not open any attachments, unless directly from your lab Blackboard shell or lab instructor.
You may access Blackboard from your lab computer once given permission.

78. MicroLAB™
MicroLAB™ is a computerized system that allows us to collect experimental data in real-time.
MicroLAB™ also has a spreadsheet program with several data analysis features.
Your lab manual has the instructions for this part of Lab 1.
Double check your work onscreen before printing out anything.
MicroLAB™ does not give you the correct number of significant figures for statistics. Use your data set to determine these.
Don’t save any of your files!

79. Important…
Always type a “0” before a decimal point for numbers smaller than 1, e.g
Label columns correctly the first time. If you don’t, you will have to delete and redo them, which will result in a loss of data.
If you are told to label a column [Fe], M the “[Fe]” refers to the label and the “M” the unit. If there are no units present, then the particular variable in question does not have any units.

80. Also Important… Do not give two columns the same label.
Do not label a column for which you will need to create a formula.
The digits of precision you set your column properties to should reflect the digits of precision of your tabulated data.

81. More Important…
Select “Accept Data” often! Always select “Accept Data” right before looking up column statistics!
Column statistics can be accessed by right clicking on a column and selecting “Column Statistics.”
When asked to look at a graph to determine certain values, always reference the spreadsheet for the exact values instead of visually estimating from the graph.