2. LO 2.7 The student is able to explain how solutes can be separated by chromatography based on intermolecular interactions. (Sec 1.10)
LO 2.10 The student can design and/or interpret the results of a separation experiment (filtration, paper chromatography, column chromatography, or distillation) in terms of the relative strength of interactions among and between the components. (Sec 1.10)
LO 3.10 The student is able to evaluate the classification of a process as a physical change, chemical change, or ambiguous change based on both macroscopic observations and the distinction between rearrangement of covalent interactions and noncovalent interactions. (Sec 1.10)

3. A main challenge of chemistry is to understand the connection between the macroscopic world that we experience and the microscopic world of atoms and molecules.
You must learn to think on the atomic level.

4. Atoms vs. Molecules Matter is composed of tiny particles called atoms.
Atom: smallest part of an element that is still that element.
Molecule: Two or more atoms joined and acting as a unit.

5. Oxygen and Hydrogen Molecules
Use subscripts when more than one atom is in the molecule.

6. A Chemical Reaction
One substance changes to another by reorganizing the way the atoms are attached to each other.

7. Science Science is a framework for gaining and organizing knowledge.
Science is a plan of action — a procedure for processing and understanding certain types of information.
Scientists are always challenging our current beliefs about science, asking questions, and experimenting to gain new knowledge.
Scientific method is needed.

8. Fundamental Steps of the Scientific Method
Process that lies at the center of scientific inquiry.

9. Fundamental Steps of the Scientific Method
Observation:
Hypothesis:
Experiment:
Theory:
Prediction:
Law:

10. Fundamental Steps of the Scientific Method
Observation: Something that can be witnessed and recorded
Hypothesis:A possible explanation for an observation.
Experiment:
Theory: tested hypothesis
Prediction:
Law:A summary of repeatable observed (measurable) behavior.

11. Nature of Measurement Measurement
Quantitative observation consisting of two parts.
number
scale (unit)
Examples
20 grams
6.63 × joule·second

13. 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
Volume
Liter L

14. Prefixes Used in the SI System
Prefixes are used to change the size of the unit.

15. Prefixes Used in the SI System

16. Mass ≠ Weight
Mass is a measure of the resistance of an object to a change in its state of motion. Mass does not vary.
Weight is the force that gravity exerts on an object. Weight varies with the strength of the gravitational field.

17. A digit that must be estimated in a measurement is called uncertain.
A measurement always has some degree of uncertainty. It is dependent on the precision of the measuring device.
Record the certain digits and the first uncertain digit (the estimated number).

18. Measurement of Volume Using a Buret
The volume is read at the bottom of the liquid curve (meniscus).
Meniscus of the liquid occurs at about mL.
Certain digits: 20.15
Uncertain digit: 20.15

19. Precision and Accuracy
Agreement of a particular value with the true value.
Precision
Degree of agreement among several measurements of the same quantity.

20. Precision and Accuracy

21. Which set of data is the most accurate?
Accuracy & Precision
Consider 3 sets of data that have been recorded after measuring a piece of wood that is exactly m long.
Which set of data is the most accurate?
Set Y
Which set of data is the most precise?
Set X
Set Y
Set Z
5.864 m
6.002 m
5.872 m
5.878 m
6.004 m
5.868 m
Average length
5.871 m
6.003 m
5.870 m

22. Rules for Counting Significant Figures
1. Nonzero integers always count as significant figures.
3456 has ___________ sig figs (significant figures).
3456 has 4 sig figs (significant figures).
33456 has 4 sig figs

23. Rules for Counting Significant Figures
2. There are three classes of zeros.
a. Leading zeros are zeros that precede all the nonzero digits. These do not count as significant figures.
0.048 has ___________ sig figs.
2 sig figs

24. Rules for Counting Significant Figures
b. Captive zeros are zeros between nonzero digits. These always count as significant figures.
16.07 has 4 sig figs.
4 sig figs

25. Rules for Counting Significant Figures
c. Trailing zeros are zeros at the right end of the number. They are significant only if the number contains a decimal point.
9.300 has ___________ sig figs.
150 has _____________ sig figs.
9.300 has 4 sig figs.
150 has 2 sig figs.

26. Rules for Counting Significant Figures
3. Exact numbers have an infinite number of significant figures.
1 inch = 2.54 cm, exactly.
9 pencils (obtained by counting).

27. Exponential Notation Example 300. written as 3.00 × 102
Contains three significant figures.
Two Advantages
Number of significant figures can be easily indicated.
Fewer zeros are needed to write a very large or very small number.

28. Significant Figures in Mathematical Operations
1. For multiplication or division, the number of significant figures in the result is the same as the number in the least precise measurement used in the calculation × 5.5 =  7.4

29. Significant Figures in Mathematical Operations
2. For addition or subtraction, the result has the same number of decimal places as the least precise measurement used in the calculation.

30. CONCEPT CHECK!
You have water in each graduated cylinder shown. You then add both samples to a beaker (assume that all of the liquid is transferred).
How would you write the number describing the total volume?
Refer to the previous slide for adding
3.1 mL
What limits the precision of the total volume?
The total volume is 3.1 mL. The first graduated cylinder limits the precision of the total volume with a volume of 2.8 mL. The second graduated cylinder has a volume of 0.28 mL. Therefore, the final volume must be 3.1 mL since the first volume is limited to the tenths place.

31. Rules for Rounding / not as important as Sig Figs
1. Round at the end of a series of calculations, NOT after each step
2. Use only the first number to the right of the last sig fig to decide whether or not to round
a. Less than 5, the last significant digit is unchanged
b. 5 or more, the last significant digit is increased by 1

32. The answer is always "Yes!"
1. Significant figures rules will be observed in all calculations throughout the year in this course. You need never ask, "Do we have to watch our
sig figs?"
The answer is always "Yes!"

33. SI Prefixes Common to Chemistry
Unit Abbr.
Exponent
Mega M
106
Kilo k
103
Deci d
10-1
Centi c
10-2
Milli m
10-3
Micro 
10-6
Nano n
10-9
Pico p
10-12

34. 1.0070 m  5 sig figs 17.10 kg  4 sig figs 100,890 L  5 sig figs
Sig Fig Practice #1
How many significant figures in each of the following?
m 
5 sig figs
17.10 kg 
4 sig figs
100,890 L 
5 sig figs
3.29 x 103 s 
3 sig figs
cm 
2 sig figs
3,200,000 
2 sig figs

35. Rules for Significant Figures in Mathematical Operations
Multiplication and Division:
# sig figs in the result equals the number in the least precise measurement used in the calculation.
6.38 x 2.0 =
12.76  13 (2 sig figs)

36. Sig Fig Practice #2 Calculation Calculator says: Answer 3.24 m x 7.0 m
100.0 g ÷ 23.7 cm3
g/cm3
4.22 g/cm3
0.02 cm x cm
cm2
0.05 cm2
710 m ÷ 3.0 s
m/s
240 m/s
lb x 3.23 ft
lb·ft
5870 lb·ft
1.030 g ÷ 2.87 mL
g/mL
2.96 g/mL

37. Rules for Significant Figures in Mathematical Operations
Addition and Subtraction: The number of decimal places in the result equals the number of decimal places in the least precise measurement. =
 18.7 (3 sig figs)

38. Sig Fig Practice #3 Calculation Calculator says: Answer 3.24 m + 7.0 m
100.0 g g
76.27 g
76.3 g
0.02 cm cm
2.391 cm
2.39 cm
713.1 L L L
709.2 L
lb lb lb lb
2.030 mL mL
0.16 mL
0.160 mL

39. Uncertainty in Measurement
A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.
Measurements are performed with
instruments
No instrument can read to an infinite number of decimal places

40. Questions to ask when approaching a problem
1. What is my goal? Or you might phrase it as: Where am I going?
2. Where am I starting? Or you might phrase it as: What do I know?
3. How do I proceed from where I start to where I want to go? Or you might say: How do I get there?

44. Use when converting a given result from one system of units to another.
To convert from one unit to another, use the equivalence statement that relates the two units.
Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units).
Multiply the quantity to be converted by the unit factor to give the quantity with the desired units.

45. Example #1
A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent?
To convert from one unit to another, use the equivalence statement that relates the two units.
1 ft = 12 in
The two unit factors are:

46. Example #1
A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent?
Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units).

47. Example #1
A golfer putted a golf ball 6.8 ft across a green. How many inches does this represent?
Multiply the quantity to be converted by the unit factor to give the quantity with the desired units.

48. Example #2
An iron sample has a mass of 4.50 lb. What is the mass of this sample in grams?
(1 kg = lbs; 1 kg = 1000 g)

49. CONCEPT CHECK!
What data would you need to estimate the money you would spend on gasoline to drive your car from New York to Los Angeles? Provide estimates of values and a sample calculation.
This problem requires that the students think about how they will solve the problem before they can plug numbers into an equation. A sample answer is:
Distance between New York and Los Angeles: miles
Average gas mileage: 25 miles per gallon
Average cost of gasoline: $2.75 per gallon
(3200 mi) × (1 gal/25 mi) × ($2.75/1 gal) = $352
Total cost = $350

50. Three Systems for Measuring Temperature
Fahrenheit
Celsius
Kelvin

51. The Three Major Temperature Scales

52. Converting Between Scales

53. At what temperature does °C = °F?
EXERCISE!
At what temperature does °C = °F?
The answer is -40. Since °C equals °F, they both should be the same value (designated as variable x). Use one of the conversion equations such as °C = (°F-32)(5/9), and substitute in the value of x for both °C and °F. Solve for x.

54. EXERCISE!
Since °C equals °F, they both should be the same value (designated as variable x).
Use one of the conversion equations such as:
Substitute in the value of x for both TC and TF. Solve for x.

55. EXERCISE!

56. Mass of substance per unit volume of the substance.
Common units are g/cm3 or g/mL.

57. Example #1
A certain mineral has a mass of 17.8 g and a volume of 2.35 cm3. What is the density of this mineral?

58. Example #2
What is the mass of a 49.6-mL sample of a liquid, which has a density of 0.85 g/mL?

59. AP Learning Objectives, Margin Notes and References
LO 2.7 The student is able to explain how solutes can be separated by chromatography based on intermolecular interactions.
LO 2.10 The student can design and/or interpret the results of a separation experiment (filtration, paper chromatography, column chromatography, or distillation) in terms of the relative strength of interactions among and between the components.
LO 3.10 The student is able to evaluate the classification of a process as a physical change, chemical change, or ambiguous change based on both macroscopic observations and the distinction between rearrangement of covalent interactions and noncovalent interactions.
AP Margin Notes
Solutions do not contain components large enough to scatter visible light. See Tyndall effect in Chapter 11.
The components of a solution can be separated by distillation.
The components of a solution are so small that they cannot be separated by filtration.
Additional AP References
LO 2.10 (see APEC Lab 5, "Thin-Layer Chromatography”)
LO 3.10 (see APEC Lab 9, "Actions, Reactions, and Interactions”)

60. Matter Anything occupying space and having mass.
Matter exists in three states.
Solid
Liquid
Gas

61. The Three States of Water

62. Solid
Rigid
Has fixed volume and shape.

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64. Liquid Has definite volume but no specific shape.
Assumes shape of container.

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66. Gas Has no fixed volume or shape.
Takes on the shape and volume of its container.

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68. Mixtures Have variable composition. Homogeneous Mixture
Having visibly indistinguishable parts; solution.
Heterogeneous Mixture
Having visibly distinguishable parts.

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70. Which of the following is a homogeneous mixture?
CONCEPT CHECK!
Which of the following is a homogeneous mixture?
Pure water
Gasoline
Jar of jelly beans
Soil
Copper metal
gasoline

71. Physical Change
Change in the form of a substance, not in its chemical composition.
Example: boiling or freezing water
Can be used to separate a mixture into pure compounds, but it will not break compounds into elements.
Distillation
Filtration
Chromatography

72. Chemical Change
A given substance becomes a new substance or substances with different properties and different composition.
Example: Bunsen burner (methane reacts with oxygen to form carbon dioxide and water)

73. Which of the following are examples of a chemical change?
CONCEPT CHECK!
Which of the following are examples of a chemical change?
Pulverizing (crushing) rock salt
Burning of wood
Dissolving of sugar in water
Melting a popsicle on a warm summer day
1 (burning of wood)

74. The Organization of Matter