1. Chapter 1 Introduction: Matter and Measurement
Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Chapter 1 Introduction: Matter and Measurement
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall

2. Matter:
Anything that has mass and takes up space. 2

3. Properties and Changes of Matter3

4. Properties of Matter Physical Properties: Chemical Properties:
Can be observed without changing a substance into another substance.
Boiling point, density, mass, volume, etc.
Chemical Properties:
Can only be observed when a substance is changed into another substance.
Flammability, corrosiveness, reactivity with acid, etc. 4

5. Properties of Matter Intensive Properties: Extensive Properties:
Independent of the amount of the substance that is present.
Density, boiling point, color, etc.
Extensive Properties:
Dependent upon the amount of the substance present.
Mass, volume, energy, etc. 5

6. Changes of Matter Physical Changes: Chemical Changes:
Changes in matter that do not change the composition of a substance.
Changes of state, temperature, volume, etc.
Chemical Changes:
Changes that result in new substances.
Combustion, oxidation, decomposition, etc. 6

7. Chemical Reactions
In the course of a chemical reaction, the reacting substances are converted to new substances. 7

8. Units of Measurement 8

9. There are two types of units:
SI Units
There are two types of units:
fundamental (or base) units;
derived units.
There are 7 base units in the SI system.

10. SI Units Système International d’Unités
Uses a different base unit for each quantity 10

11. Metric System
Prefixes convert the base units into units that are appropriate for the item being measured. 11

12. The units for volume are given by (units of length)3.
SI unit for volume is 1 m3.
We usually use 1 mL = 1 cm3.
Other volume units:
1 L = 1 dm3 = 1000 cm3 = 1000 mL.

13. Uncertainty in Measurements
Different measuring devices have different uses and different degrees of accuracy. 13

14. Physical property of a substance
Density:
Physical property of a substance d= m V 14

15. Do now: A graduated cylinder is filled with 15. 0 mL of water
Do now: A graduated cylinder is filled with mL of water. An object with a mass of g causes the total volume to increase to 23.4 mL. What is the density of the sample?

16. Vocabulary review Mass : amount of matter in an object.
It is measured with a balance.
Unit in the SI system: grams g
Weight: a measure of the pull that the gravity exerts over an object.
If we stay in the same planet (same gravity) is used like the mass.

17. Volume The space that an object occupies.
Unit of volume in SI Liter = L

18. What is density?
Density is a comparison of how much matter there is in a certain amount of space.
IT IS AN INTENSIVE PHYSICAL PROPERTY. It is used to identify a substance.

19. DENSITY OF WATER
1g/ml
That means that a gram of water has a volume of 1 milliliter or 1 cubic centimeter.

20. Which one is more dense?
Now which one is more dense?

21. What is density? Density = mass volume or mass ÷ volume.
Units for density: g
cm3
Why are these the units for density?
ALWAYS REMEMBER UNITS!

22. Let’s try a density problem together
Find the density of a yellow rock has a mass of 8 g and a volume of 4 cm3.
Use your table S to determine what element it could be!

23. IMMISCIBLE LIQUIDS
If you pour together liquids that don’t mix and have different densities, they will form liquid layers.
Liquids that don’t mix are said to be IMMISCIBLE
The liquid with the highest density will be on the bottom.
The liquid with the lowest density will be on the top.

24. Liquid Layers . Which layer has the highest density?
Which layer has the lowest density?
Imagine that the liquids have the following densities:
10g/cm3. 3g/cm3.
6g/cm3. 5g/cm3.
Which number would go with which layer?
Is any of the liquids water?

25. To measure the volume of an object
If is a regular object measure the dimensions needed and use the formula
cube= LxWxH
Cylinder = p h r2
For an irregular object use the water displacement method.

26. Liquid Layers – Try with your neighbor
Which liquid has the highest density?
Which liquid has the lowest density?
Which liquid has the middle density?

27. Liquid Layers – Try on your own!
Imagine that the liquids on the right have the following densities:
15g/cm g/cm3
3g/cm g/cm3
7g/cm g/cm3
Match the colors to the correct densities.
3g/cm3
7g/cm3
9g/cm3
10g/cm3
12g/cm3
15g/cm3

28. Review What is the formula for density?
What happens if you pour together liquids that have different densities?
Will the liquid on the top have the highest or lowest density?
Will the liquid on the bottom have the highest or lowest density?

29. Density worksheet answers
1=1
2=2
3=2 D of Na g/mL
4=4 D of Mg g/cm3
5=2
g/24.4L = 0.82 g/L

30. HOMOGENEOUS : SAME PROPERTIES THROUGHOUT THE SAMPLE
HETEROGENEOUS :
DIFFERENT PROPERTIES IN DIFFERENT PARTS OF THE SAMPLE

31. PURE SUBSTANCES have a constant composition
ELEMENTS – Made up of same kind of atoms. Could not be decomposed.
COMPOUNDS – Made up of different kind of atoms CHEMICALLY COMBINED. Can be decomposed.
Recognizable by formulas!

32. Mixtures and Compounds

33. Compounds
Compounds can be broken down into more elemental particles.

34. Matter
Atoms are the building blocks of matter.

35. Matter Atoms are the building blocks of matter.
Each element is made of the same kind of atom.

36. Matter Atoms are the building blocks of matter.
Each element is made of the same kind of atom.
A compound is made of two or more different kinds of elements.

37. SEPTEMBER 21 ELEMENT , COMPOUNDS AND MIXTURES REVIEW FOR TEST
PHYSICAL AND CHEMICAL PROPERTIES
PHYSICAL AND CHEMICAL CHANGES
DENSITY

38. MIXTURES
Combination of two or more pure substances. Can be separated by physical means. They do not have a fixed composition.
Can be homogeneous or heterogeneous.
SOLUTIONS ARE HOMOGENEOUS MIXTURES.

39. AQUEOUS SOLUTIONS The solution is prepared using water as the solvent
(aq) means DISSOLVED IN WATER!!!
Na Cl (s) is a compound
Na Cl (aq) is a mixture!!!!

40. Pure Substances and Mixtures

41. Pure Substances and Mixtures
If matter is not uniform throughout, then it is a heterogeneous mixture.
If matter is uniform throughout, it is homogeneous.
If homogeneous matter can be separated by physical means, then the matter is a mixture.
If homogeneous matter cannot be separated by physical means, then the matter is a pure substance.
If a pure substance can be decomposed into something else, then the substance is a compound.

42. Elements
If a pure substance cannot be decomposed into something else, then the substance is an element.
There are 114 elements known.
Each element is given a unique chemical symbol (one or two letters).
Elements are building blocks of matter.
The earth’s crust consists of 5 main elements. (O, Si, Al, Fe, Ca)
The human body consists mostly of 3 main elements. (O, C, H)

43. Elements

44. Metals, Nonmetals, and Metalloids

45. Symbols First letter of element in CAPITAL letter
Second or third letter in lower case.
Some elements have symbols different from the english name
SODIUM Na
POTASSIUM K
CUPPER Cu
LEAD Pb

46. IRON Fe
MERCURY Hg
GOLD Au
SILVER Ag
TIN Sn

47. Symbols from Latin Names
Element Symbol Latin name
Copper Cu cuprum
Gold Au aurum
Lead Pb plumbum
Mercury Hg hydrargyrum
Potassium K kalium
Silver Ag argentum
Sodium Na natrium
Tin Sn stannum

48. DIATOMIC ELEMENTS
H2O2F2Br2I2N2Cl2

50. Classification of Matter

51. Classification of Matter

52. Classification of Matter

53. Classification of Matter

54. Classification of Matter

55. Classification of Matter

56. Classification of Matter

57. Classification of Matter

58. Classification of Matter

59. Classification of Matter

60. MC ANSWERS
1 . A
2 . A
3. A
4. D
5. A
6. B
7. C
8. D
9. D

61. Chemical Reactions

62. Electrolysis of Water

63. Separation of Mixtures

64. Distillation:
Separates homogeneous mixture on the basis of differences in boiling point.

65. Distillation

66. Filtration:
Separates solid substances from liquids and solutions.

67. Chromatography:
Separates substances on the basis of differences in solubility in a solvent.

68. Uncertainty in Measurement

69. Significant Figures
The term significant figures refers to digits that were measured.
When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

70. Uncertainty in Measurement
All scientific measures are subject to error.
These errors are reflected in the number of figures reported for the measurement.
Precision and Accuracy
Measurements that are close to the “correct” value are accurate.
Measurements that are close to each other are precise.

71. PRECISSION DEALS WITH REPRODUCIBILITY OF A MEASUREMENT.
ACCURACY DEALS WITH THE EXACTESNESS OF THE MEASUREMENT, HOW CLOSE IT IS TO THE , TRUE, ACCEPTED OR STANDARD VALUE
PRECISSION DEALS WITH REPRODUCIBILITY OF A MEASUREMENT.
IF SEVERAL MEASUREMENTS GIVE A SIMILAR RESULT IT IS SAID THAT THE MEASUREMENT IS PRECISE

72. Accuracy versus Precision
Accuracy refers to the proximity of a measurement to the true value of a quantity.
Precision refers to the proximity of several measurements to each other.

73. October 5 UNCERTAINTY IN MEASUREMENT
SIGNIFICANT FIGURES – Rules and examples
DO NOW
Calculate the density of an object that has a mass of 10.0 g and a volume of 3.0 mL.

74. Measuring Volume by water displacement

75. Significant Figures
The number of digits reported in a measurement reflect the accuracy of the measurement and the precision of the measuring device.
All the figures known with certainty plus one extra figure (estimated digit) are called significant figures.

76. Sig fig in calculations
In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).

77. Significant Figures All nonzero digits are significant.
Zeroes between two significant figures are themselves significant.
Zeroes at the beginning of a number are never significant.
Zeroes at the end of a number are significant if a decimal point is written in the number or if they are to the right of a decimal point.

78. has 2 sf
Has 3 sf
has 5 sf

79. EXAMPLE FOR ADDITION
Copy and perform the following operation indicating the right number of sig fig 12
1.2
0.2

80. MULTIPLICATION AND DIVISION
5 x 100 = x 745 = 3469/ 5799=

81. Examples: How many significant figures are in each of the following?kg s
507 people
230,050 cm A

82. Tell the number of significant digits in each of the following measurements.
1. 48 cm __________
g __________
m __________
°C __________
mm __________
cm3__________
g __________
mm __________
kg __________
× 1015 sec __________
× 10-4 m __________
g __________

83. Tell the number of significant digits in each of the following measurements.
1. 48 cm __________
g __________
m __________
°C __________
mm __________
cm3__________
g __________
mm __________
kg __________
× 1015 sec __________
× 10-4 m __________
g __________

84. How do scientist express the accuracy of a measurement?
DO NOW : Observe the two instruments in my desk to measure volume, determine which would determine the volume of an Al cylinder with greater accuracy and explain in your notebook why.

85. Percent Error
To determine the accuracy of a measurement. It tells us how far our measured stands from an accepted or known value.
% error = I measured value – accepted value I X 100
___________________________________________
accepted value

86. Example
Calculate the percent error of the measurement for a student that determined that the density for Al is 2.5 g/mL. Hint use table T to determine the accepted value!

87. 1) 3.482 cm + 8.51 cm + 16.324 cm ____________________
2) g g g ____________________
3) 80.4 cm cm ____________________
4) 106.5mL mL ____________________
5) 48.2 cm × 1.6 cm × 2.12 cm ____________________
6) 8.3 m × 4.0 m × m

88. 7) 64.34 cm3 ÷ 8.149 cm ____________________
8) 4.93 mm2 ÷ mm ____________________
9) 0 57 mL x 760 mm/740 mm x 273K/250 K
10) g x amu/ a m u

89. Answers
28.32 cm g
63.9 cm
76 mL
160 cm3
33 m3
7.895 cm2
0.263 mm

90. MULTIPLE CHOICE 3 2 1 4

91. 1. According to an accepted chemistry reference
1. According to an accepted chemistry reference. the heat of vaporization of water is 540. calories per gram. A student determined in the laboratory that the heat of vaporization of water was 620. calories per gram. The student's results had a percent error of (1) 12.9, (2) 80.0, (3) 14.8, (4) 87.1
2. Which measurement contains a total of three significant figures?
(1) 0.01 g (2) g (3) g (4) g

92. 3. In an experiment the gram atomic mass of magnesium was determined to be Compared to the accepted value 24.3, the percent error for this determination was (1) , (2) 24.7, (3) 1.65, (4) 98.4
4. A student determined the melting point of a substance to be 55.2°C. If the accepted value is 50. 1°C the percent error in her determination is (1) 5.10, (2) 10.2, (3) 9.24, (4) 12.0
5. Using the rules for significant figures, the sum of gram and gram should be expressed as (1) gram, (2) 0.03 gram, (3) gram, (4) gram

93. 6. Which milligram quantity contains a total of four significant
figures? (1) mg (2) 3100 mg (2) 3010 mg (4) mg

94. Scientific Notation
Numbers written in scientific notation include a numeral with one digit before the decimal point, multiplied by some power of ten (6.022 x 1023)
In scientific notation, all digits are significant.
You should be able to convert from non-scientific notation to scientific and vice-versa.

95. Temperature
Definition
Instrument
Scales

96. TEMPERATURE
Is associated with heat but it is NOT HEAT. IT IS NOT A FORM OF ENERGY!!!! ( Heat is)
Review: What is KINETIC ENERGY?

97. KINETIC ENERGY (KE) Is associated with movement.
If an object is moving fast has high KE
If an object is moving slowly it has low KE

98. Temperature
In scientific measurements, the Celsius and Kelvin scales are most often used.
The Celsius scale is based on the properties of water.
0C is the freezing point of water.
100C is the boiling point of water.

99. Temperature:
A measure of the average kinetic energy of the particles in a sample. If an object is at HIGH temperature its particles are moving FAST At LOW temperature particles move SLOWLY

100. Instrument to measure temperature THERMOMETER

101. FIXED POINTS OF A THERMOMETER
BOILING POINT OF WATER
FREEZING POINT OF WATER

102. Temperature The Kelvin is the SI unit of temperature.
It is based on the properties of gases.
There are no negative Kelvin temperatures.
K = C

103. Temperature
The Fahrenheit scale is not used in scientific measurements.
F = 9/5(C) + 32
C = 5/9(F − 32)

104. Examples:
What is 35ºC in Kelvin? In ºF?
What is 183 K in ºC? In ºF?