1. Chapter 1: Matter and Measurements

2. What is Chemistry?
Biology vs.
Chemistry vs.
Physics

3. What is Chemistry? Biology Physics Chemistry
The study of living organisms
The study of forces & motion
The study of matter and its reactions and properties

4. Chemistry is the study of CHANGES in the “stuff” around us.
What is Chemistry?
Chemistry is the study of CHANGES in the “stuff” around us.
(We formally define “stuff” as matter!)

5. What are all the chemicals you use in your daily life?
What is Chemistry?
Think, pair, share:
What are all the chemicals you use in your daily life?

6. Review: Scientific Methods
1. Hypothesis
Suggested solution to a problem
2. Experiment
A controlled method of testing a hypothesis
3. Data
Organized observations
a. Data is always reproducible.

7. Review: Scientific Methods
4. Scientific Law
Statement which summarizes results of many observations and experiment
a. Scientific laws explain WHAT is observed.
Example of a scientific law:
5. Scientific Theory
Explanation that supports a hypothesis and which has been supported with repeated testing
b. Scientific theories explain WHY something is observed.
Example of a scientific theory:

8. Review: Scientific Methods
6. Steps of the Scientific Method—Review a. b. c. d. e. f.

9. What is Chemistry? Biology Physics Chemistry
The study of living organisms
The study of forces & motion
The study of matter and its reactions and properties

10. Chemistry is the study of CHANGES in the “stuff” around us.
What is Chemistry?
Chemistry is the study of CHANGES in the “stuff” around us.
(We formally define “stuff” as matter!)

11. Matter

12. Matter

13. Elements
Type of matter that cannot be broken down into simpler, stable substances and is made of only one type of atom

14. Compounds
A pure substance that contains two or more elements whose atoms are chemically bonded

15. Compounds Fixed compositions
A given compound contains the same elements in the same percent by mass

16. Compounds
The properties of a compound are VERY DIFFERENT from the properties of the elements they contain
Ex.) Sodium Chloride (NaCl) vs. Sodium & Chlorine
Sodium:

17. Electrolysis

18. Mixtures
A blend of two or more kinds of matter, each of which retains its own identity and properties
Homogeneous
Heterogeneous

19. Homogeneous Mixtures Composition is the same throughout the mixture
Examples: salt water, soda water, brass
A.k.a. a solution
Solute in a solvent (salt dissolved in water)

20. Heterogeneous Mixtures
Non-uniform; composition varies throughout the mixture

21. Separating Mixtures
Filtration

22. Separating Mixtures
Distillation

23. Separating Mixtures
Chromatography

24. Scientific Measurements
Chemistry is a quantitative science.
This means that experiments and calculations almost always involve measured values.
Scientific measurements are expressed in the SI (metric) system.
This is a decimal-based system in which all of the units of a particular quantity are related to each other by factors of ten.

25. Types of Observations and Measurements
We make QUALITATIVE observations of reactions — changes in color and physical state.
We also make QUANTITATIVE MEASUREMENTS, which involve numbers.
Use SI units — based on the metric system

26. SI System
Definition: modernized version of metric system; uses decimals
All units derived from base units; larger and smaller quantities use prefixes with base unit
Must memorize prefixes from nano- (10-9) to tera- (1012)

27. Prefixes (see handout & Ebook)
You will need to memorize all of the prefixes (factors, names and abbreviations from
109 (giga-) to (nano)!
One example of a memory device:

28. INSTRUMENTS & UNITS Use SI units — based on the metric system Length
Mass
Time
Temperature
Meter, m
Kilogram, kg
Seconds, s
Celsius degrees, ˚C
Kelvins, K

29. Length The standard unit of length in the metric system is the METER
which is a little larger than a YARD.
USING THE PREFIXES WITH LENGTH:
cm – often used in lab
km –
Gm –

30. Length Base unit: METER Conversions: 1 km=1000 m 1 cm = 10 -2 m
1 Gm = 106 m

31. Units of Length 1 kilometer (km) = 1000 meters (m)
10-2 meter (m) = 1 centimeter (cm)
102 meter (m) = 1 hectometer (Hm)
1 nanometer (nm) = 1.0 x 10-9 meter
O—H distance =
9.58 x m
9.58 x 10-9 cm nm

32. the liter (milliliter) and cubic centimeter (cm3)
Volume
THE COMMON UNITS OF VOLUME IN CHEMISTRY ARE:
the liter (milliliter) and cubic centimeter (cm3)
THE COMMON INSTRUMENTS FOR MEASURING VOLUME IN CHEMISTRY ARE: graduated cylinder & buret
Note that 1 cm3 = 1 mL (We will use this exact conversion factor throughout the year, so you will need to memorize it!)

33. Mass the gram (g) — often used in lab
THE COMMON UNIT OF MASS IN CHEMISTRY IS :
the gram (g) — often used in lab
Mass IS A MEASURE OF THE AMOUNT OF MATTER IN AN OBJECT;
Weight IS A MEASURE OF THE GRAVITATIONAL FORCE ACTING ON THE OBJECT. CHEMISTS OFTEN USE THESE TERMS INTERCHANGEABLY.
1000 g= 1 kg Mg = 10 6 g

34. TEMPERATURE IS THE FACTOR THAT DETERMINES the direction of heat flow.
Temperature Scales
Fahrenheit
Celsius
Kelvin
Anders Celsius
Lord Kelvin
(William Thomson)
TEMPERATURE IS THE FACTOR THAT DETERMINES the direction of heat flow.

35. Temperature Scales Fahrenheit Celsius Kelvin 32 ˚F 212 ˚F 180˚F 100 ˚C
Boiling point of water
32 ˚F
212 ˚F
180˚F
100 ˚C
0 ˚C
100˚C
373 K
273 K
100 K
Freezing point of water
Notice that 1 kelvin degree = 1 degree Celsius

36. Temperature Scales
100 oF
38 oC
311 K oF oC K

37. SI System
English Units (inches, feet, degrees F, etc.) are NEVER used to take measurements in the lab!

38. Calculations Using Temperature
Fahrenheit/Celsius
T (F) = 1.8 t (˚C)

39. Calculations Using Temperature
Some calculations are in kelvins (especially important for Ch 5!!)
T (K) = t (˚C) (273)
Body temp = 37 ˚C = 310 K
Liquid nitrogen = ˚C = 77 K

40. Problem
Example 1L.1 A baby has a temperature of 39.8oC. Express this temperature in oF and K.

41. SI System: Base Units
ESTABLISHMENT OF THE INTERNATIONAL SYSTEM OF UNITS (SI)—
SI UNITS AS ESTABLISHED BY THE SI:
LENGTH – meter (m)
VOLUME – cubic meter (m3)
MASS – kilogram (kg)
TEMPERATURE – Kelvin (K)

42. Time Base unit: SECOND (sec) Conversions: only non-decimal base unit
60 sec = 1 min min = 1 hr

43. Precision and Accuracy in Measurements
Precision vs. Accuracy
Definitions:
Precision—how close answers are to each other (reproducibility)
Accuracy—how close answer is to accepted (true) value (agreement to accepted value)

44. Precision and Accuracy in Measurements
Percent Error - a way to calculate accuracy in the lab
Equation:
% Error = | Accepted Value – Exp. Value | x 100
Accepted Value

45. Precision and Accuracy in Measurements
Ex1.9 A student reports the density of a pure substance to be 2.83 g/mL. The accepted value is 2.70 g/mL. What is the percent error for the student’s results?
Equation:
% Error = | Accepted Value – Exp. Value | x 100
Accepted Value

46. Scientific Notation
Exponential (Scientific) Notation—See Worksheet

47. Significant Figures: Why are they Important?
Numbers in math: no units, abstract, no context, can read calculator output exactly for answer.
vs.
Numbers in chemistry: measurements – include units.
SIG FIGS WILL BE IMPORTANT THROUGHOUT THIS COURSE!

48. Graduated Cylinder Example

49. What are significant figures? (aka sig figs)
Significant figures are all the digits in a measurement that are known with certainty plus a last digit that must be estimated.
With experimental values your answer can have too few or too many sig figs, depending on how you round.

50. How Rounding Influences Sig Figs
1.024 x 1.2 = Too many numerals (sig figs)
Too precise
1.024 x 1.2 = 1 Too few numerals (sig figs)
Not precise enough

51. Why This Concept is Important
We will be adding, subtracting, multiplying and dividing numbers throughout this course.
You MUST learn how many sig figs to report each answer in or the answer is meaningless.
You must report answers on lab reports & tests/quizzes with the correct number of sig figs (+/- 1) or else you will lose points!!

52. How Do We Find the Correct Number of Sig Figs In an Answer?
First, we will learn to count number of sig figs in a number. You must learn 4 rules and how to apply them.
Second, we will learn the process for rounding when we add/subtract or multiply/divide. We will then apply this process in calculations.

53. Rules for Counting Sig Figs
Rule #1: Read the number from left to right and count all digits, starting with the first digit that is not zero. Do NOT count final zero’s unless there is a decimal point in the number!
3 sig figs
4 sig figs
5 sig figs
23.4
234
0.234
2340
203
345.6
3.456
34560
3405
678.90
6789.0
67008
60708

54. Rules for Counting Sig Figs
Rule #2: A final zero or zero’s will be designated as significant if a decimal point is added after the final zero.
3 sig figs
4 sig figs
5 sig figs
2340
23400
234000
2340.
2000.
20000.
23400.

55. Rules for Counting Sig Figs
Rule #3: If a number is expressed in standard scientific notation, assume all the digits in the scientific notation are significant.
2 sig figs
3 sig figs
4 sig figs
2.3 x 102
2.0 x 103
2.30 x 102
2.00 x 103
2.300 x 102
2.000 x 103

56. Rules for Counting Sig Figs
Rule #4: Any number which represents a numerical count or is an exact definition has an infinite number of sig figs and is NOT counted in the calculations.
Examples:
12 inches = 1 foot (exact definition)
1000 mm = 1 m (exact definition)
25 students = 1 class (count)

57. Practice Counting Sig Figs
How many sig figs in each of the following? mm cm
1.200 x 105 mL m
0.02 cm
8 ounces = 1 cup
30 cars in the parking lot

58. Answers to Practice How many sig figs in each of the following?
mm (5)
cm (6)
1.200 x 105 mL (4)
m (3)
0.02 cm (1)
8 ounces = 1 cup (infinite, exact def.)
30 cars in the parking lot (infinite,
count)

59. General Rounding Rule
When a number is rounded off, the last digit to be retained is increased by one only if the following digit is 5 or greater.
EXAMPLE: rounds to
5 (ones place)
5.35 (hundredths place)
5.355 (thousandths place)
5.4 (tenths place)
You will lose points for rounding incorrectly!

60. Process for Addition/Subtraction
Step #1: Determine the number of decimal places in each number to be added/subtracted.
Step #2: Calculate the answer, and then round the final number to the least number of decimal places from Step #1.

61. Addition/Subtraction Examples
Round to tenths place.
Example #2:
Round to hundredths place.
Example #3:
Round to ones place.
23.456
24.706
Rounds to:
24.7
3.56
2.0699
2.07 14
26.735 27

62. Process for Multiplication/Division
Step #1: Determine the number of sig figs in each number to be multiplied/divided.
Step #2: Calculate the answer, and then round the final number to the least number of sig figs from Step #1.

63. Multiplication/Division Examples
Round to
1 sig fig.
Example #2:
2 sig figs.
Example #3:
3 sig figs.
23.456 x x
Rounds to: 1
3.56 x x
0.67
14.0/
11.73
1.19

64. Practice Write the answers to the following computations using the correct number of sig figs
a g g g =
b m m =
c × 4.20 × =
d. 17 ÷ =

65. Important Rounding Rule
When you are doing several calculations, carry out all the calculations to at LEAST one more sig fig than you need (I carry all digits in my calculator memory) and
only round off the FINAL result.

66. Use of Conversion Factors
Also known as dimensional analysis or factor-label method (or unit conversions)
Dimensional analysis/ Use of conversion factors
Definition: technique to change one unit to another using a conversion factor
Ex.) # in original unit x new unit = New # in new unit
original unit

67. Using Dimensional Analysis
Express the quantity 1.00 ft in different dimensions (inches, meters).
Conversion factors:

68. Using Dimensional Analysis
Example 1L.5 Calculate the following single step conversions:
a. How many Joules are equivalent to 25.5 calories if 1 cal = joules?
b. How many liters gasoline can be contained in a 22.0 gallon gas tank if L = 1 gal?

69. Using Dimensional Analysis
Example 1L.6 The following multiple step conversions can be solved, knowing that 1 in = 2.54 cm. Convert the length of ft to millimeters.

70. Using Dimensional Analysis
Example 1L.7 The average velocity of hydrogen molecules at 0oC is 1.69 x 105 cm/s. Convert this to miles per hour.

71. Using Dimensional Analysis
Example 1L.8 A piece of iron with a volume of 2.56 gal weighs lbs. Convert this density to scruples per drachm with the following conversion factors:
1.00 L = gal, kg = lb, scruple = g, mL = drachm.

72. Using Dimensional Analysis: Area Conversions
Example 1L.14 Express the area of a 27.0 sq yd carpet in square meters.
Conversion factors needed:

73. Using Dimensional Analysis: Volume Conversions
Example 1L.15 Convert 17.5 quarts to cubic meters. (1 L = qt, 1 ft3 = L)

74. Properties of Substances
1. Every pure substance has its own unique set of properties that serve to distinguish it from all other substances.
2. Properties used to identify a substance must be intensive; that is, they must be independent of amount.
Extensive properties depend on the amount.
Classify the following as either intensive (I) or extensive (E):
a. density
b. mass
c. melting point
d. volume

75. Properties of Substances
Density is an INTENSIVE property of matter, which does NOT depend on quantity of matter.
Contrast with EXTENSIVE properties of matter, which do depend on quantity of matter.
Examples of extensive properties: mass and volume.
Styrofoam
Brick

76. Chemical and Physical Properties
Chemical property – Observed when the substance changes to a new one.
Example of a chemical property:
Copper reacts with air to form copper (II) oxide.
Physical property – Observed without changing the substance to a new one.
Example of a physical property:
Water boils at 100oC.

77. Physical Changes Physical changes do not result in a new substance:
boiling of a liquid
melting of a solid
dissolving a solid in a liquid to give a SOLUTION.

78. Physical vs. Chemical Change
Another name for a Chemical change is a chemical reaction — change that results in a new substance.

79. Example:
Classify the following as either physical (P) or chemical (C) changes:
a. ice melting
b. gasoline burning
c. food spoiling
d. log of wood sawed in half

80. Density
Density is an INTENSIVE property of matter, which does NOT depend on quantity of matter.
Contrast with EXTENSIVE properties of matter, which do depend on quantity of matter.
Examples of extensive properties: mass and volume.
Styrofoam
Brick

81. DENSITY : Review Definition: ratio of mass to volume for an object=
mass ( g )
volume cm 3
Aluminum
Platinum
Mercury
2.7 g/cm3
13.6 g/cm3
21.5 g/cm3

82. Sample Problem A piece of copper has a mass of 57. 54 g. It is 9
Sample Problem A piece of copper has a mass of g. It is 9.36 cm long, 7.23 cm wide, and 0.95 mm thick. Calculate density (g/cm3).

83. Density as a Conversion Factor
Density is a “bridge” between mass and volume, or vice versa
Volume (cm3) x density g = mass (g)
cm3
Mass (g)  density cm3 = Volume (cm3) g

84. Solve the problem using DENSITY AS A CONVERSION FACTOR.
SAMPLE PROBLEM: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds?
Solve the problem using DENSITY AS A CONVERSION FACTOR.

85. Ex1L. 9 What is the density of Hg if 164. 56 g occupy a volume of 12
Ex1L.9 What is the density of Hg if g occupy a volume of 12.1cm3?

86. Ex1L.10 What is the mass of 2.15 cm3 of Hg?

87. Ex1l.11 What is the volume of 94.2 g of Hg?

88. Example 1L. 12: Given the following densities: chloroform 1
Example 1L.12: Given the following densities: chloroform 1.48 g/cm3 and mercury 13.6 g/cm3 and copper 8.94 g/cm3. Calculate if a 50.0 mL container will be large enough to hold a mixture of 50.0 g of mercury, 50.0 g of chloroform and a 10.0 g chunk of copper.

89. Example 1L. 13 How many kilograms of methanol (d = 0
Example 1L.13 How many kilograms of methanol (d = g/mL) does it take to fill the 15.5-gal fuel tank of an automobile modified to run on methanol?

90. Density of Water Density of water changes with temperature
(As water temperature changes, volume changes)
Maximum density of water is at
4oC = g/cm3
(often rounded to 1.00 g/cm3)

91. Derived Units Definition: derived from base units
Example: m/sec (unit of speed)
Divide meters by seconds
Volume examples
m3 (m x m x m) or cm3 (cm x cm x cm)