1. Chapter 1 Introduction: Matter and Measurement
Lecture Presentation
Chapter 1 Introduction: Matter and Measurement
Dr. Subhash C. Goel
South GA State College
Douglas, GA

2. Paper Chromatography Demonstration

3. Chemistry
Chemistry is the study of the properties and behavior of matter.
It is central to our fundamental understanding of many science-related fields.

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Chemistry
The atom is the basic unit of chemistry and a fundamental component of matter. Chemistry is concerned with atoms and their interactions with other atoms, with particular focus on the properties of the chemical bonds formed between species.
Chemistry is also concerned with the interactions between atoms or molecules and various forms of energy (e.g. photochemical reactions, oxidation-reduction reactions, changes in phases of matter, separation of mixtures, properties of polymers, etc.)
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Matter
is the physical material of the universe.
has mass.
occupies space.
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Matter
Atoms are the building blocks of matter.
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ELEMENTS
are made up of unique atoms, the building blocks of matter.
Names of the elements are derived from a wide variety of sources (e.g., Latin or Greek, mythological characters, names of people or places).
For example:
Hydrogen (H) Iron (Fe)
Carbon (C) Zinc (Zn)
Oxygen (O)
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MOLECULES
are combinations of atoms held together in specific shapes.
Macroscopic(observable) properties of matter relate to submicroscopic realms of atoms.
Properties relate to composition (types of atoms) and structure (arrangement of atoms) present.
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10. Compound
A substance composed of two or more elements chemically combined.
For example:
Water (H2O) Glucose
Carbon dioxide (CO2)
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11. Compounds and Composition
Compounds have a definite composition. That means that the relative number of atoms of each element that makes up the compound is the same in any sample.
This is The Law of Constant Composition (or The Law of Definite Proportions).

12. Mixture
A material that can be separated by physical means into two or more substances
For example:
Italian salad dressing
Saltwater
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13. Heterogeneous Mixture
A mixture that consists of physically distinct parts, each with different properties.
For example:
Salt and iron filings
Oil and vinegar
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14. Homogenous Mixture
A mixture that is uniform in its properties; also called a solution.
For example:
Saltwater Sugar Solution
Air
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15. Methods of Classification
State of Matter
Composition of Matter

16. Classification of Matter Based on Composition
If you follow this scheme, you can determine how to classify any type of matter.
Homogeneous mixture
Heterogeneous mixture
Element
Compound
Matter
And
Measurement

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States of Matter
Solids: fixed shape and volume
Liquids: fixed volume but no shape
Gases: No fixed shape and volume
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Matter
A Property is any characteristic that allows us to recognize a particular type of matter and distinguish it from others.
There are >100 substances called Elements that combines to give all kind of matter in the universe.
We will see that properties of matter relates to both the kind of atoms in the matter contains (composition) & their arrangement in matter (structure).
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19. Physical Change
A change in the form of matter but not in its chemical identity.
For example:
Melting
Dissolving
Density
Mass
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20. Chemical Change = Chemical Reaction
A change in which one or more kinds of matter are transformed into a new kind of matter or several new kinds of matter.
For example:
Rusting
Burning
Reactivity with acids
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Types of Properties
Intensive Properties…
Are independent of the amount of the substance that is present.
Density, boiling point, color, etc.
Extensive Properties…
Depend upon the amount of the substance present.
Mass, volume, energy, etc.
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22. Physical Property
A characteristic that can be observed for a material without changing its chemical identity.
For example:
Physical state
Boiling point
Color
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23. Chemical Property
A characteristic of a material involving its chemical change.
For example:
Ability to react with oxygen
Ability to react with fluorine
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24. Changes in State of Matter
Converting between the three states of matter is a physical change.
When ice melts or water evaporates, there are still 2 H atoms and 1 O atom in each molecule.

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Chemical Reactions
In the course of a chemical reaction, the reacting substances are converted to new substances.
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26. Potassium is a soft, silvery-colored metal that melts at 64°C
Potassium is a soft, silvery-colored metal that melts at 64°C. It reacts vigorously with water, with oxygen, and with chlorine. Identify all of the physical properties and chemical properties given in this description.
Physical Property
Chemical Property
Soft
Reacts with water
Silvery-colored
Reacts with oxygen
Melting point (64°C)
Reacts with chlorine
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27. Separation of Mixtures
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Filtration
In filtration, solid substances are separated from liquids and solutions.
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Distillation
Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components.
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30. Chromatography
This technique separates substances on the basis of differences in the ability of substances to adhere to the solid surface, in this case, dyes to paper.

31. Numbers and Chemistry
Numbers play a major role in chemistry. Many topics are quantitative (have a numerical value).
Concepts of numbers in science
Units of measurement
Quantities that are measured and calculated
Uncertainty in measurement
Significant figures
Dimensional analysis

32. Units of Measurement
1960: An international agreement was reached to use metric units (called SI units) in scientific measurements.

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SI Units
Système International d’Unités
A different base unit is used for each quantity.
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Metric System
Prefixes convert the base units into units that are appropriate for the item being measured.
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Using Prefixes
103 grams = 1_____grams
2x10-3 meters = 2______meters
2.6 x 10-1 L = 2.6______L
6 x 10-2 g = 6______g
1 gigajoules = 1 x _______joules
9 µm (micrometer) = 9 x _______m
8 cm = 8 x __________m
1 ng = 1 x ________g
Convert 0.30 gigameters/second to a value without a prefix in scientific notation.
1. Kilo 2. milli 3. d 4. c x 108 meters/second
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37. Volume
Note that volume is not a base unit for SI; it is derived from length (m × m × m = m3).
The most commonly used metric units for volume are the liter (L) and the milliliter (mL).
A liter is a cube 1 decimeter (dm) long on each side.
A milliliter is a cube 1 centimeter (cm) long on each side, also called 1 cubic centimeter (cm × cm × cm = cm3).

38. Temperature
In general usage, temperature is considered the “hotness and coldness” of an object that determines the direction of heat flow.
Heat flows spontaneously from an object with a higher temperature to an object with a lower temperature.

39. 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.
The kelvin is the SI unit of temperature.
It is based on the properties of gases.
There are no negative Kelvin temperatures.
The lowest possible temperature is called absolute zero (0 K).
K = C

40. Temperature
The Fahrenheit scale is not used in scientific measurements, but you hear about it in weather reports!
The equations below allow for conversion between the Fahrenheit and Celsius scales:
F = 9/5(C) + 32
C = 5/9(F − 32)

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Exercise
Convert 40o F to K.
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42. In winter, the average low temperature in interior Alaska is -30
In winter, the average low temperature in interior Alaska is -30.°F (two significant figures). What is this temperature in degrees Celsius and in kelvins?
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43. Density Density is a physical property of a substance.
It has units that are derived from the units for mass and volume.
The most common units are g/mL or g/cm3.
D = m/V

44. Oil of wintergreen is a colorless liquid used as a flavoring. A 28
Oil of wintergreen is a colorless liquid used as a flavoring. A 28.1-g sample of oil of wintergreen has a volume of 23.7 mL. What is the density of oil of wintergreen?
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45. A sample of gasoline has a density of 0. 718 g/mL
A sample of gasoline has a density of g/mL. What is the volume of 454 g of gasoline?
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46. Uncertainty in Measurement
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47. Numbers Encountered in Science
Exact numbers are counted or given by definition. For example, there are 12 eggs in 1 dozen.
Inexact (or measured) numbers depend on how they were determined. Scientific instruments have limitations. Some balances measure to ±0.01 g; others measure to ±0.0001g.

48. Uncertainty in Measurements
Different measuring devices have different uses and different degrees of accuracy.
All measured numbers have some degree of inaccuracy.

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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.
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51. S. Goel, SGSC

52. S. Goel, SGSC

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54. 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.
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Dimensional Analysis
We use dimensional analysis to convert one quantity to another.
Most commonly, dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm)
1 in.
2.54 cm or
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Dimensional Analysis
Use the form of the conversion factor that puts the sought-for unit in the numerator:
Given unit 
 desired unit
desired unit
given unit
Conversion factor
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Dimensional Analysis
Example:
Convert 8.00 m to inches.
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Exercise
How many mL are in 1.63 L?
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Exercise
A liquid helium tank has a volume of 275 L. What is the volume in m3?
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Execrcise
The speed of sound in air is about 343 m/s. What is this speed in miles per hour?
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62. Notation Terminology
Fixed Decimal or Fixed Notation: numbers such as , , or
Exponential Notation: very large or very small numbers are written more clearly in exponential notation. Example: instead of writing that one drop of water has 1,500,000,000,000,000,000,000 molecules, we write 1.5x1021 molecules or atomic radius of Ne atom is centimeter, we write 7.0x10-9 centimeter.

63. Power of 10
106 = = = = = = =

64. Multiplying/Dividing by 10, 100, 1000
When multiplying by 10, 100, 1000, etc move the decimal to the right by the number of zeros in 10, 100, or 1000
72x100 = x1000 = -62.4
When dividing by 10, 100, 1000, etc move the decimal to the left by the number of zeros in 10, 100, or 1000
72/100 = x1000 =

65. Converting Exponential Notation to Fixed Decimal Notation
Rule 1. The sign in front never changes.
Rule 2. In Exponential notation, if the power of 10 is positive then move decimal point to the right by the same number of places as the value of exponent.
Rule 3. If the power of 10 is negative then move decimal point to the left.
Examples: 2 x 102 = 200; x 103 = -3.3
2 x 10-2 = 0.02; x 10-3 =

66. Changing Exponential to Scientific Notation
To convert an exponential to scientific notation, move the decimal so that significand is one or greater but less then ten.
Examples:
0.045 x 105 = 4.5 x 103
x 10-2 = -5.7 x 10-5
-8544 x 10-7 = x 10-4

67. Converting Numbers to Scientific Notation
Any number to the zero power is equal 1.
Examples: 20 = 1; 420 = 1; 100 = 1
Conversions to scientific notation:
42 = 42 x 100 = 4.2 x 101
943 = 943 x 100 = 9.43 x 102
= x 100 = -3.6 x 10-4
= x 100 = 9.3 x 10-3

68. MULTIPLYING AND DIVIDING POWERS OF 10
When multiplying exponential, add exponent
103x102 = 105; 10-5 x10-2 = 10-7; 105 x10-2 = 103
When dividing exponential, subtract exponent
104/102 = 102; 10-5/102 = 10-7; 10-4/10-2 = 10-2
When you take reciprocal of an exponential, change the sign of exponent
1/103 = 10-3; /105x10-2 = 10-6

69. Multiplying and Dividing in Exponential Notation
Do number math using number rules and exponential math using exponential rules. Then combine the two parts.
Example: (2x103)(4x1023) = 8x1026
2x4 = 8
103x1023 = 1026
Also memorize: ½ = 0.5; 1/3 = 0.33; ¼ = 0.25; 1/5 = 0.2; 2/3 = 0.67, ¾ = 0.75; 1/8 = 0.125