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. Density worksheet answers
1=1
2=2
3=2
4=4
4-6 do not need to do
7=1.7
8b=3.9

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

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

5. 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!

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

7. Matter
Atoms are the building blocks of matter.

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

9. 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.

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

11. 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.

12. 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!!!!

13. Pure Substances and Mixtures

14. 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.

15. 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)

16. Elements

17. Metals, Nonmetals, and Metalloids

18. 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

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

20. 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

21. DIATOMIC ELEMENTS
H2O2F2Br2I2N2Cl2

23. Classification of Matter

24. Classification of Matter

25. Classification of Matter

26. Classification of Matter

27. Classification of Matter

28. Classification of Matter

29. Classification of Matter

30. Classification of Matter

31. Classification of Matter

32. Classification of Matter

33. Mixtures and Compounds

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

35. Properties and Changes of Matter

36. 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.

37. 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.

38. 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.

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

40. Chemical Reactions

41. Electrolysis of Water

42. Separation of Mixtures

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

44. Distillation

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

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

47. Units of Measurement

48. 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.

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

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

51. 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.

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

53. Temperature:
A measure of the average kinetic energy of the particles in a sample.

54. 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.

55. 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

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

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

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

59. Examples:
An object with a mass of g occupies a volume of 11.8 mL. What is its density?
A sample with a density of 3.75 g cm-3 occupies a volume of cm3. What is the mass of the sample?
A graduated cylinder is filled with 15.0 cm3 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?

60. Uncertainty in Measurement

61. 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.

62. Uncertainty in Measurement
All scientific measures are subject to error.
These errors are reflected in the number of figures reported for the measurement.
These errors are also reflected in the observation that two successive measures of the same quantity are different.
Precision and Accuracy
Measurements that are close to the “correct” value are accurate.
Measurements that are close to each other are precise.

63. 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

64. 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.

65. 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.
In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).

66. 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.

67. has 2 sf
Has 3 sf
has 5 sf

68. 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.

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