1. CHEM IS TRY: AN INTRODUCTION
Chapter 1
CHEM IS TRY:
AN INTRODUCTION

2. The science that deals with the materials of the universe and the changes these materials undergo.
The central science.
Understanding most other fields of science requires an understanding of chemistry.
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3. Memorize important information. Learn and practice processes.
Learn the vocabulary.
Memorize important information.
Learn and practice processes.
Keep working and learning from your mistakes.
Ask questions!
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4. Steps in the Scientific Method
Process that lies at the center of scientific inquiry.
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5. Scientific Models Summarizes what happens. Law Hypothesis
A possible explanation for an observation.
Theory (Model)
An attempt to explain why it happens.
Set of tested hypotheses that gives an overall explanation of some natural phenomenon.
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6. Measurements and calculations
Chapter 2
Measurements and calculations

7. SI System Internationale
Decimal system -- multiply or divide by 10’s.
SI conversions

8. SI BASE UNITS Length meter m Mass kilogram kg Time second s
Electric current ampere A
Temperature kelvin K
Amount mole mol
Light intensity candela cd

9. SI TO ENGLISH UNITS 1.0 m = 39.37 inches 1.0 inch = 2.54 cm
1.0 kg = lb
1.0 lb = g
1.0 L = qt

11. ACCURATE = CORRECT PRECISE = CONSISTENT
Accuracy vs. Precision
Accuracy - how close a measurement is to the accepted value
Precision - how close a series of measurements are to each other
ACCURATE = CORRECT
PRECISE = CONSISTENT

12. A digit that must be estimated is called uncertain.
A measurement always has some degree of uncertainty.
Record the certain digits and the first uncertain digit (the estimated number).
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13. Rules for Counting Significant Figures
1. Nonzero integers always count as significant figures.
3456 has 4 sig figs (significant figures).
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14. Rules for Counting Significant Figures
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.
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15. 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.
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16. 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.
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17. Rules for Counting Significant Figures
3. Exact numbers have an infinite number of significant figures.
1 inch = 2.54 cm, exactly.
9 pencils (obtained by counting).
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18. Exponential Notation Example Two Advantages 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.
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19. Rules for Rounding Off
1. If the digit to be removed is less than 5, the preceding digit stays the same.
5.64 rounds to 5.6 (if final result to 2 sig figs)
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20. Rules for Rounding Off
1. If the digit to be removed is equal to or greater than 5, the preceding digit is increased by 1.
5.68 rounds to 5.7 (if final result to 2 sig figs)
3.861 rounds to 3.9 (if final result to 2 sig figs)
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21. Rules for Rounding Off
2. In a series of calculations, carry the extra digits through to the final result and then round off. This means that you should carry all of the digits that show on your calculator until you arrive at the final number (the answer) and then round off, using the procedures in Rule 1.
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22. Significant Figures in Mathematical Operations
1. For multiplication or division, the number of significant figures in the result is the same as that in the measurement with the smallest number of significant figures.
1.342 × 5.5 =  7.4
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23. Significant Figures in Mathematical Operations
2. For addition or subtraction, the limiting term is the one with the smallest number of decimal places.
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24. 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.
Choose the appropriate conversion factor by looking at the direction of the required change (make sure the unwanted units cancel).
Multiply the quantity to be converted by the conversion factor to give the quantity with the desired units.
Check that you have the correct number of sig figs.
Does my answer make sense?
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26. Three Systems for Measuring Temperature
Fahrenheit
Celsius
Kelvin
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27. HEAT vs TEMPERATURE Heat is a form of energy.
System – whatever you want it to be
Open system – heat can cross boundaries.
Closed system – no heat crosses boundaries.

28. The Three Major Temperature Scales
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29. Converting Between Scales
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30. Density Mass of substance per unit volume of the substance.
Common units are g/cm3 or g/mL.
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31. Chapter 3
Matter

32. Matter Anything occupying space and having mass.
Mass -- amount of matter
Weight -- measure of the force of attraction of gravity.
Inertia -- resistance to movement.
Matter exists in three states.
Solid
Liquid
Gas
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33. The Three States of Water
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34. Solid Rigid Has a fixed volume and shape.
Very low kinetic energy- vibrate but cannot move around
Examples:
Ice cube, diamond, iron bar
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35. Liquid Has a definite volume but no specific shape.
Assumes shape of container.
Low kinetic energy: particles move around but still are close together.
Examples:
Gasoline, water, alcohol, blood
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36. Gas Has no fixed volume or shape.
Takes the shape and volume of its container.
High kinetic energy: particles can separate and move about the container.
Examples:
Air, helium, oxygen
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37. Physical Properties
The characteristics of matter that can be changed without changing its composition.
Characteristics that are directly observable.
Examples:
Odor, color, volume, state (s, l, or g), density, melting point, and boiling point
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38. Chemical Properties A substance’s ability to form new substances.
The characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy.
Characteristics that describe the behavior of matter.
Examples:
Flammability, rusting of steel, and the digestion of food
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39. Physical Change
Change in the form of a substance, not in its chemical composition.
Example:
Boiling or freezing water
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40. 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)
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41. Element
A substance that cannot be broken down into other substances by chemical methods.
Examples:
Iron (Fe), aluminum (Al), oxygen (O2), and hydrogen (H2)
All of the matter in the world around us contains elements.
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42. Compound
A substance composed of a given combination of elements that can be broken down into those elements by chemical methods.
Examples:
Water (H2O), carbon dioxide (CO2), table sugar (C12H22O11)
A compound always contains atoms of different elements.
A compound always has the same composition (same combination of atoms).
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43. Pure Substances Always have the same composition.
Either elements or compounds.
Examples:
Pure water (H2O), carbon dioxide (CO2), hydrogen (H2), gold (Au)
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44. Mixtures Have variable composition. Examples Wood, wine, coffee
Can be separated into pure substances: elements and/or compounds.
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45. Homogeneous Mixture Same throughout.
Having visibly indistinguishable parts.
A solution.
Does not vary in composition from one region to another.
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46. Heterogeneous Mixture
Having visibly distinguishable parts.
Contains regions that have different properties from those of other regions.
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47. Different Physical Property
Mixtures can be separated based on different physical properties of the components.
Evaporation
Volatility
Chromatography
Adherence to a surface
Filtration
State of matter
(solid/liquid/gas)
Distillation
Boiling point
Technique
Different Physical Property
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48. The Organization of Matter
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