1. Unit 1
Chapters 1-4

2. Introduction to Chemistry
Chapter 1
Introduction to Chemistry

3. What is Chemistry & Why should we study it?
Chemistry - The study of the structure, properties, and composition of substances, and the changes the substances undergo
Reasons to study
Materials are made using chemistry
Energy Where will it come from?
Medicine & Biotechnology
Agriculture
Environment
Astronomy & Space Exploration

4. Branches of Chemistry Organic Chemistry Inorganic Chemistry
The study of compounds that contain carbon.
Inorganic Chemistry
The study of substances that do not contain carbon.
Analytical Chemistry
The study of the composition of substances.
Physical Chemistry
The study of the behavior of chemicals.
Biochemistry
The study of the composition and changes in composition of living organisms.

5. Scientific Method

6. Scientific Method Continued
Observation
An act of seeing some fact or occurrence; often involves a measurement
Hypothesis
A descriptive model used to explain observations.
Experiment
A carefully controlled, repeatable procedure for gathering data to test a hypothesis.
Theory
A thoroughly tested model that explains why experiments give certain results.
Scientific Law
A concise statement that summarizes the result of a broad variety of observations and experiments

7. Chapter 2
Matter and Change

8. Matter
A formal definition of matter is anything that takes up space and has mass.

9. States of Matter Solid - Matter that has definite shape and volume.
Liquid - A form of matter that flows, has a fixed volume, and takes the shape of its container.
Gas - Matter that has no definite shape or volume; it adopts the shape and volume of its container.

10. Classifications of matter
Substance:
- uniform and definite compostion or pure
Element:
- simplest form of matter made of atoms
Compounds:
- two or more elements chemically combined made of molecules or formula units

11. Mixtures
A mixture is a combination of two or more substances in which the basic identity of each substance is not changed.
Homogeneous mixture – A mixture that is completely uniform in its composition; its components are not distinguishable.
Heterogeneous mixture – A mixture that is not uniform in its composition; its components are readily distinguishable.

12. Flow Chart for Classifying Substances

13. Separating Mixtures Filtration Settling Distillation Evaporation
Crystallization
Chromatography

14. Properties of Matter
The properties of matter describe the characteristics and behavior of matter, including the changes that matter undergoes.

15. Physical Properties
Physical properties: A quantity of a substance that can be observed or measured without changing the substances chemical composition.
- Color
Temperature
solubility,
melting point,
boiling point,
density,electrical conductivity
and physical state (solid, liquid, or gas).

16. Physical Change
A physical change occurs when there is an alteration of a substance that does not affect its chemical composition.
Ripping paper
Boiling water
Melting ice

17. Chemical Properties
The ability of a substance to undergo chemical reactions and form new substances.
Rust
Tarnish
Burn
Oxidize

18. Chemical Change
Formation of new chemicals or substances is the only true way.
Visual signs that could indicate chemical change.
Production of heat
Evolution of Gas
Emission of Light
Color Change
Formation of a precipitate (Solid)

19. Scientific Measurement
Chapter 3
Scientific Measurement

20. Types of Measurement
Qualitative measurement – A measurement that gives descriptive, nonnumeric results
Size
Shape
Color
Quantitative measurement – A measurement that gives definite, usually numeric results
Weight
Length
Time

21. Uncertainty in Measurements
Accuracy – The closeness of a measurement to the true value of what is being measured.
Precision – The reproducibility, under the same conditions, of a measurement
Accurate but not precise
Precise but not accurate
Accurate and Precise

22. Calculation of Percent Error
Percent error – The percent that a measured value differs from the accepted value.
Accepted Value – A quantity used by general agreement of the scientific community
Experimental Value - Measured in Lab
% Error =
|Accepted Value - Experimental Value| * 100% Accepted Value

23. International System of Units
SI - le Systéme International d’Units
Revised version of the metric system

25. Units of measurement Metric System (pg. 64)
a. 1000________ = 1 kilo________
b.  100 centi________ = 1 ________
c milli________ = 1 ________
d.  1,000,000 micro _____ = 1 _____
e. 1,000,000,000 nano____ = 1 _____

26. Scientific Notation Used to express very large or very small numbers
Written in the form N x 10n where N is greater than 1 and less than 10 and n is an integer.
If you move the decimal to the left, the power of 10 is positive and the actual number is very large.
If you move the decimal to the right, the power of 10 is negative and the actual number is very small.

27. Is there a difference between mass and weight?
YES!!!!
Mass - how much matter something has
Weight - mass with the affects of gravity
If you go to the moon do you have the same mass as on earth? How about weight?

28. Temperature
The measure of the kinetic energy of the molecules in a substance
Two main scales
- Celsius Scale - 0oC freezing point of water oC boiling point of water
Kelvin Scale - Freezing point of water is 273 K boiling point of water is 373 K
Absolute Zero - all motion and movement stops.
Converting between the two is easy!
K = oC + 273
oC = K -273

29. How to measure correctly
- Determine the smallest interval on the measuring device: scale and units.
Estimate one digit more then the smallest unit of the scale.
Example: ruler A
ruler B
ruler C

30. Significant Figures
In a measurement, they include all of the digits that are known plus one last estimated digit.

31. Rules:
) Every nonzero digit reported in a measurement is assumed to be significant.
Examples: meters and 358 grams all have 3 significant figures.
) Zeros appearing between nonzero digits are significant.
Examples: 308 cm, 5.03 liters & grams All have 3 significant figures.
) Zeros at the end of a number and to the right of a decimal point are significant.
Examples: meters & liters both have 4 sig. figs.

32. Rules Continued….
4) Zeros in front of nonzero digits are not significant
Examples: meters & grams Both have two significant figures.
Zeros at the end of a measurement and to the left of a decimal vary.
Examples: has two sig. figs
1200. Has four sig. Figs.

33. The last Rule:
) There are two situations in which measurements have an unlimited number of significant figures.
Counting - If you count then you have unlimited sig. figs as no measuring device was used.
Exactly defined quantities
60 minutes = 1 hour (both numbers have unlimited sig. figs.)

34. Calculations with Significant Figures
Rounding - less than 5, drop remaining digits greater than 5 & equal to 5, round up
Example: 3.75 grams Round to 2 sig. figs greater/= to 5 round up Answer = 3.8 grams
Addition & Subtraction
1) Add the numbers together
) Round to the least precise measurement (digit).

35. Addition Example: 90. ml + 15.5 ml __________ 105.5 ml
Answer: What digit do both measurement go to?
The ones: therefore 106 ml is correct.

36. Calculations with Significant Figures Continued...
Multiplication & Division
) Count the number of significant figures for each number.
) Do the multiplication or division
) Round answer to the least number of significant figures
Example: 3.2 cm x 5.38 cm x cm =
Calculator Answer = cm3
Correct Answer = 52 cm3

37. The ratio of mass to volume density = Mass Volume

38. Problem Solving in Chemistry
Chapter 4
Problem Solving in Chemistry

39. Dimensional Analysis
Process of changing units of a measure value

40. To solve dimensional analysis problems, use A.C. E.
A.C.E. - five step problem solving
Analyze
1. identify knowns
2. identify unknowns
3. Plan of attack: pick conversions to get from known to unknown
Calculate
4. Set up problem and show unit cancellation
Evaluate
5. Answer with label and significant figures
Does the answer make sense?

41. Conversion Factors
The measurement in the numerator is equivalent to the measurement in the denominator.
1 meter = 100 centimeters
1 meter = 1 or 100 centimeters =1 100 centimeters meter