2. 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.

3. 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.

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

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

6. 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.

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

8. Scientific Models
Law
A summary of repeatable observed (measurable) behavior.
Hypothesis
A possible explanation for an observation.
Theory (Model)
Set of tested hypotheses that gives an overall explanation of some natural phenomenon.

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

10. 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

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

12. Prefixes Used in the SI System

13. 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.

14. 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).

15. 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

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

17. Precision and Accuracy

18. Next, we study Significant Figures
The term significant figures refers to all certain digits and one uncertain digit in a measurement .
When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

19. All nonzero digits are significant.
Based on these, please indicate number of significant figures for each:
1) 244.6
2) 5006 3) 4) 5) 6)
7) 36,000
All nonzero digits are significant.
Zeroes between two significant figures are themselves significant.
Zeroes at the beginning of a number are never significant.
Why some zeroes are not significant? Because they are not related to your reading. They are related to unit conversion. Take example of 2.5 km (draw a ruler with 2 and 3 km, put the reading in between), one would read 2.5km, with 5 estimated. Convert it to m, one gets 2500m, but, still, 5 is estimated, not the two 0 after.
Zeroes at the end of a number are significant only if a decimal point is written in the number, or when the zero carries an overbar.

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

21. 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 2 sig figs.

22. 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.

23. 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 4 sig figs.
150 has 2 sig figs.

24. 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.

25. Determining the number of significant figures in each of the following .
Four (SF)
6,600,000,000,000,000,000,000, Two (SF)
Two (SF)
4850
Three (SF)
41.0

26. Determine the number of significant in each od the following
205,000
Three (SF)
0ne (SF)
Four (SF)
Six (SF)

27. Multiplication and Division of SF
In multiplication or division, the final answer is written so it has the same number of significant as the measurement with the least significant figure.
Example: Multiply the following
24.66×0.35 = = 8.6 (2 SF)
Multiply 4SF by 2SF gives us 2SF

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. Multiplication and Division of SF
Multiply and Divide the following
1800.0÷ = 300 = (5SF)
Add one zero to give 2 SF

30. 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.

31. Addition and Subtraction
In addition or subtraction , the final answer is written so it has the same number of decimal place as the measurement with the fewest decimal places.
Example: =
66.1

32. Subtract the following. 65.09 – 3.0 = 62.09 = 62.1
Subtraction
Subtract the following.
65.09 – 3.0 = = 62.1

33. Rounding off numbers
Rule 1: If the first digit to be dropped is less than 5, that digit and all digits that follow it are simply dropped
Example rounded off to three significant figures is 62.3

34. Rounding off Numbers
Rule 2: If the first digit to be dropped is a digit greater then 5, or a 5 followed by digits other than zero then all zeros, the excess digits are dropped and the last retained digit is increased in value by one unit.
Example : and rounded off to three significant figures become respectively
62.8 and 62.6.

35. Rounding off Numbers
Rule3: If the first digit to be dropped is a 5 not followed by any other digits or a 5 followed only by zeros, an odd-even rule applies.
Drop the 5 and any zeros that follow it and then
A increase the last retained digit by one unit if it is odd or
B leave the last retained digit the same if it is even.

36. Example: 62.650 and 62.350 rounded to three (SF) becomes 62.6 ( even).
Rounding off Numbers
Example: and rounded to three (SF) becomes ( even).
becomes 62.4 (odd Rule)
Express 15,500 to two significant Figures
15,500 becomes 16,000 ( odd Rule)

37. 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?
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.

38. Questions to ask when approaching a problem
What is my goal?
What do I know?
How do I get there?

39. 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.

40. Problem solving skills
Step 1: state the given and needed quantities Step 2:write a plan to convert the given unit to the needed unit. Step 3: sate the equalities and conversion factors needed to convert units. Step 4: set up problem to cancel units and calculate answer.

41. Example 1

42. Example 1

43. Example 2

44. Example 2

45. Example 3

46. Example 3

47. Example #4
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:

48. Example #4
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).

49. Example #4
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.

50. Example #5
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)

51. Three Systems for Measuring Temperature
Fahrenheit
Celsius
Kelvin

52. The Three Major Temperature Scales

53. Converting Between Scales

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

55. 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?

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

57. Example 3 Density
Suppose a sample of aluminum is placed in a 25ml graduated cylinder containing 10.5 ml of water. The level of the water rises to 13.5ml. What is the mass of the aluminum sample.
Step 1 :State the given and needed quantities.
Given : Density of aluminum = 2.7g/cm3
Volume= 13.5ml ml = 3.0 ml
Needed Mass =g

58. Example 3 Density

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

60. The Three States of Water

61. Solid
Rigid
Has fixed volume and shape.

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

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

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

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69. 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

70. 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

71. 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)

72. 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)

73. Identify each of the following as physical (P) or chemical (c) change
1. Nacl (Table salt ) dissolve in water
Answer physical
2. An apple is cut.
3.Heat changes water to steam
Answer Physical
4. Baking soda reacts to vinegar.
Answer chemical
5. Fe(iron) rusts.
6. Ice melts
7.Wood rots
8. Food digested
9. Grass grows
10. Pancakes cook

74. The Organization of Matter