1. Chapter One:
CHEMICAL FOUNDATIONS

2. Where would we be without Chemistry?

3. Where would we be without Chemistry?
No You!

4. A science of interest, wonder and beauty!
“If I were to present myself before you with an offer to
teach you some new game—
If I were to tell you an improved plan of throwing a ball,
of flying a kite, or
of playing leapfrog,
Oh, with what attention you would listen to me!
Well, I am going to teach you many new games.
I intend to instruct you in a science full of
interest, wonder, and beauty
a science that will afford you amusement in your youth, and
riches in your more mature years.
In short, I am going to teach you the science of
CHEMISTRY!”
—Dr. Scoffern, Devonshire, England, Chemistry No Mystery, 1848

5. CHEMISTRY: AN OVERVIEW

6. If, in some cataclysm, all of scientific knowledge were to be destroyed, and only one sentence passed on to the next generations of creatures, what statement would contain the most information in the fewest words?

7. The Atomic Hypothesis The atomic hypothesis:
“All things are made of atoms – little particles that move around in perpetual motion, attracting each other when they are a little distance apart, but repelling upon being squeezed into one another”
(Feynman, Six Easy Pieces, p. 4)

8. 1. Matter is made up of atoms
Water thought experiment (Feynman, pp. 5-10)
For the first time in history, we can actually “see” individual atoms…

9. Figure 1.1a The Surface of a Single Grain of Table Salt

10. Figure 1.1b An Oxygen Atom on a Gallium Arsenide Surface

11. Figure 1.1c Benzene molecules on a rhodium surface

12. Atoms vs. Molecules
 Everything is made from only about 100 different kinds of atoms (// “letters”)

13. Oxygen and Hydrogen Molecules
 To understand chemistry, must learn to think on the microscopic (and atomic!) level

14. 2. Atomic processes
Water thought experiment cont’d (Feynman, pp )

15. 3. Chemical reactions

16. Dancing/Levitating Raisins Experiment
Materials
mL beaker
10 g sodium hydrogen carbonate
45 mL 3% acetic acid
5 raisins

17. Dancing/Levitating Raisins Experiment
NaHCO3 (aq) + HC2H3O2 (aq)  CO2 (g) + H2O (l) + NaC2H3O2 (aq)
NaHCO3 (aq) = sodium hydrogen carbonate (aka sodium bicarbonate)
HC2H3O2 = acetic acid
NaC2H3O2 = sodium acetate
// Children’s water-wings

18. The Scientific Method

19. The Fundamental Steps of the Scientific Method

20. Law vs. Theory
A law summarizes what happens (e.g. law of conservation of mass)
A theory (model) is an attempt to explain why it happens

21. QUESTION
The difference between a scientific law and a scientific theory can, at times, be confusing. For example, we will refer to the “Atomic theory” or perhaps the “Law of Gravity.” Should the Law of Gravity be changed to the Theory of Gravity?
1. Yes, no one can see gravity, it is better described as a theory.
2. No, scientific laws are based on summaries of many observations and gravity observations are well known and
predictable.
3. Yes, gravity is better described as a theory because gravity explains why masses attract each other and theories are about explaining observations.
4. No, keep it as a law, laws offer explanations and gravity explains why masses attract each other and laws are about explaining observations.
HMClass Prep: Figure 1.5

22. ANSWER
Choice 2 follows the agreed-upon distinction between a theory and a law. Observations that consistently agree and provide the same result in a variety of systems become “Laws.” Theories are attempts to offer human interpretations and explanations about what was observed. Therefore, we should continue calling the summary about attractions the Law of Gravitation.
Section 1.2: The Scientific Method

23. Units of Measurement

24. Nature of Measurement
Measurement - quantitative observation consisting of 2 parts
Part 1 – number
Part 2 – scale (unit)
Examples:
20 grams
6.63 × Joule seconds

25. The Fundamental SI Units
Physical Quantity
Name of Unit
Abbreviation
Mass
kilogram kg
Length
meter m
Time
second s
Temperature
kelvin K
Electric current
ampere A
Amount of substance
mole
mol
Luminous intensity
candela cd

26. QUESTION
Conveniently, a U.S. nickel has a mass of approximately 5 grams. If you had one dollar’s worth of nickels what would be the mass of the nickels in milligrams?
milligrams
2. 50 milligrams
3. 1,000 milligrams
4. 100,000 milligrams
HMClass Prep: Table 1.2

27. ANSWER
Choice 4 shows the correct conversion. After determining that 20 nickels make up one dollar, then one dollar’s worth of nickels would have a mass of 100 grams. Next, the conversion between grams and milligrams can be performed by multiplying by 1,000 (because there are 1,000 milligrams per gram.)
Section 1.3: Units of Measurement

28. Figure 1.7 Common Types of Laboratory Equipment Used to Measure Liquid Volume

29. Uncertainty in Measurement

30. Uncertainty in Measurement
A digit that must be estimated is called uncertain
A measurement always has some degree of uncertainty

31. Measurement of Volume Using a Buret
Record all “certain” digits plus one “uncertain” digit (aka significant figures)

32. The Difference Between Precision and Accuracy

33. QUESTION
Two chemistry students are each drinking a canned beverage. The volumes of both cans are listed as 300 milliliters. Steve remarks that the bottlers must be precise and hope to get pretty close to that in every can. Which would be a correct response from Susan?
1. If precision were the only goal, bottlers could claim any
volume in the can as long as it was always very nearly the
same volume.
2. If precision were the only goal, bottlers would have to get
exactly 300 mL in every can.
3. If bottlers wanted to be precise, all can volumes would
have to average 300 mL.
HMClass Prep: Figure 1.10

34. ANSWER
Choice 1 best fits what scientists define as precise. High precision measurements may have closeness to a set goal (such as 300 mL in a can) but precision always means closeness within a set of measurements. If the volumes of 100 cans were all from 294 to 295 mL the volumes would be precise, but not accurately 300 mL.
Section 1.4: Uncertainty in Measurement

35. Significant Figures and Calculations

36. Rules for Counting Significant Figures - Details
Nonzero integers always count as significant figures.
3456 has 4 sig figs.

37. Rules for Counting Significant Figures – Details (continued)
Leading zeros do not count as significant figures.
0.048 has 2 sig figs.

38. Rules for Counting Significant Figures – Details (continued)
Captive zeros always count as significant figures.
16.07 has 4 sig figs.

39. Rules for Counting Significant Figures – Details (continued)
Trailing zeros are significant only if the number contains a decimal point.
9.300 has 4 sig figs
150 has 2 sig figs.

40. Rules for Counting Significant Figures – Details (continued)
Exact numbers have an infinite number of significant figures.
1 inch = 2.54 cm, exactly

41. QUESTION
Which one of the following does NOT represent a result with four significant digits?
HMClass Prep: Rules for Significant Figures

42. ANSWER
Choice 4 only has three significant digits. Note that the lone zero in front of the decimal point is not based on any measurement and the next zero serves only as a place holder not as a measurement.
Section 1.5: Significant Figures and Calculations

43. Rules for Counting Significant Figures – Details (continued)
For multiplication/division, the result has the same number of significant figures as the least precise measurement used in the calculation
For addition/subtraction, the result has the same number of decimal places as the least precise measurement used in the calculation
Round at the end

44. QUESTION
If you were unloading a kg box of books from your car and a “friend” added two more 482 gram chemistry books, how much in kg and using the rules for significant digits, would you be lifting? kg kg kg kg
HMClass Prep: Table 1.2

45. ANSWER
Choice 3 provides both the conversion and proper number of significant digits. Consider that the 482 grams of mass must be doubled (to include both books) and that 482 grams is kg. When adding two measurements always report your answer to the same number of decimal places as the least precise measurement used in the calculation. In this case the answer should be reported to the hundredths place.
Section 1.5: Significant Figures and Calculations

46. QUESTION
The volume of a sample can be obtained from its density and mass. If the mass of a sample of acid from a battery were 5.00 grams and the density was 1.2 g/mL, what would you report in mL and with the proper number of significant digits, as the sample volume? mL mL mL mL

47. ANSWER
Choice 3 shows the correct volume for 5.00 grams of this sample. First, be sure to use the correct solution for solving volume from mass and density (V = m/D). Then recall that the significant digit pattern for dividing measurements is to retain the same number of significant digits in the answer as the least number in any of the related measurements.
Section 1.8: Density

48. Dimensional Analysis (aka unit factor method or factor-label method)

49. QUESTION
A student has entered a 10.0-km run. How long is the run in miles?
HMClass Prep: Figure 1.12

50. ANSWER
6.22 mi

51. Temperature

52. The Three Major Temperature Scales

53. Temperature Conversion Formulas
Tk = Tc
Tf = Tc x 9oF/5oC + 32oF

54. QUESTION
When converting temperatures between Celsius and Fahrenheit we use the equation: TF = (9°F/5°C)TC + 32°F. This indicates that:
1. Five degrees on the Celsius scale is actually the same change in temperature as nine degrees on the Fahrenheit scale.
2. A change of 1°C is smaller than a change of 1°F.
3. Zero degrees Fahrenheit would then correspond to 1.8 degrees
Celsius.
4. I do not know how to convert from Celsius to Fahrenheit (will I
need to?)
HMClass Prep: Figure 1.12

55. ANSWER
Choice 1 provides the correct relationship between Celsius and Fahrenheit temperature readings. There are 180 degrees between freezing and boiling of water on the Fahrenheit scale and only 100 for the separation on the Celsius scale. 180/100 is the same as 9/5.
Section 1.7: Temperature

56. Classification of Matter

57. The Three States of Water

58. Structure of a Solid

59. Structure of a Liquid

60. Structure of a Gas

61. Simple Laboratory Distillation Apparatus

62. The Organization of Matter

63. The Organization of Matter

64. The Organization of Matter

65. The Organization of Matter

66. Homogeneous Mixtures

67. Mixture vs. Solution

68. Mixture vs. Compound