1. Unit 0: Matter, Measurement, and Chemical Equations
AP Chemistry

2. Matter
Chapter 1

3. The Classification of Matter
Matter is anything that occupies space and has mass
Classified by:
State: its physical form (solid, liquid, gas)
Composition: its components (elements, compounds, mixtures)

4. Classification of Matter by Components
Pure substance: made up of only one component
Elements (i.e. gold)
Compounds (i.e. sodium chloride)
Mixture: made of two or more components; proportions vary
Heterogeneous: composition of mixture varies; can see the parts (i.e. oil and water)
Homogeneous: all portions are uniform and have the same composition and properties (i.e. salt water)

5. Separating Mixtures
Decanting—carefully pouring off liquid from a precipitate (Remember doing this in the copper lab??)
Distillation—separates homogeneous mixtures of two liquids or a liquid & a dissolved solid

6. Separating Mixtures (cont.)
Gravity Filtration—a heterogeneous mixture of a liquid and solid are poured through filter paper and funnel
Vacuum filtration—a vacuum is set up with a sink apparatus to aid the filtration process

7. Physical vs. Chemical Changes
Physical change—alters the state or appearance, not composition
i.e. boiling, cutting magnesium
Chemical change—alters the composition of matter
i.e. rusting nails, burning magnesium

8. Physical vs. Chemical Properties
Physical property—can be observed without changing the chemical composition
i.e. odor, taste (!), color, appearance, melting & boiling point, density
Chemical property—can be observed by changing the composition (must have a chemical rxn)
Flammability, acidity, reactivity

9. Energy

10. Energy
Energy—the capacity to do work (applying a force over a given distance)
Types of energy
Kinetic—the motion of an object
Potential—the position of an object
Thermal—the temperature of an object (technically a type of kinetic energy)
Law of Conservation of Energy—cannot be created or destroyed

11. Units & measurement

12. Units of Measurement Metric System (meters, liters, etc)
English System (inches, pounds, etc.)
International System of Units (SI)
Used in science

13. SI Base Units

14. Dimensional Analysis
Used to convert between units using conversion factors
Know prefixes and how they are used to give you conversions
i.e. 1 liter contains 1000 mL
Other Useful temperature conversions:

15. Significant Figures

16. Significant Figures
Significant figures are all the certain digits and the estimated digit in a measurement.
For example, 31.7 mL has three “sig figs”.
Two certain digits (the 3 and the 1)
One estimated digit (the 7)

17. Why Significant Figures?
“Sig Figs” indicate the quality of instrument you use for your measurements.
Electronic balance may give g
Analytical balance may give g
Usually, high precision is indicated by a large number of sig figs.
Ex: A measurement of L (4 sig figs) is considered more precise than a measurement of 4 L (1 sig fig)

18. Significant Figures Rules
All non-zero digits of a measurement are significant
g  5 sig figs
Zeros written between significant digits are significant
56.06 g  4 sig figs
All final zeros past a decimal point are significant
73.00  4 sig figs
7,300  2 sig figs

19. Significant Figures Rules
Zeros written to the left of all nonzero digits are not significant.
0.09 g  1 sig fig
Numbers in scientific notation have the same number of sig figs as the portion that’s before the “x 10” part.
4.30 x 105  3 sig figs

20. Significant Figures in Calculations
Multiplication/Division
Your answer should have the same number of sig figs as the value with the fewest sig figs in the problem.
Ex: X =  round to 4.254
Because has four sig figs

21. Significant Figures in Calculations
Addition/Subtraction
Your answer should have the same number of sig figs to the right of the decimal as the value with the fewest decimal places in the problem.
Ex: =  round to 8.4
Because 3.4 has one digit to the right of the decimal

22. Rounding in Multi-Step Calculations
ROUND ONLY THE FINAL ANSWER
Underlined digits in calculations signify the least significant digit

23. Scientific Notation

24. Scientific Notation
Scientific Notation is a short way to write very large or very small numbers.
It is written as the product of a number between 1 and 10 and a power of 10.

25. Converting to Scientific Notation
Create a number between 1 and 10 by moving the decimal to the left or right.
Count the number of spaces the decimal moved to determine the exponent on 10.
The exponent is POSITIVE if you moved your decimal to the LEFT
The exponent if NEGATIVE if you moved your decimal to the RIGHT

26. Converting to Scientific Notation
Example: 3,346,000,000 = x 109
The decimal moved 9 places
Example: = 9.52 x 10-4
Exponent = number of decimal places moved to the LEFT
Number between 1 and 10
Exponent = number of decimal places moved to the RIGHT
Number between 1 and 10

27. Converting to Standard Notation
Move the decimal to the right or left the number of spaces indicated by the exponent
Move to the LEFT if exponent is NEGATIVE
Move to the RIGHT if exponent is POSITIVE
Example: x 106 = 1,312,000
Example: x 10-6 =

28. Calculations with Scientific Notation
Punch the number (the digit number) into your calculator.
Push the EE or EXP button. Do NOT use the x (times) button.
Enter the exponent number. Use the +/- button to change its sign.
Voila! Treat this number normally in all subsequent calculations.
To check yourself, multiply 6.0 x 105 times 4.0 x 103 on your calculator. Your answer should be 2.4 x 109
Instructions for scientific calculator

29. Calculations with Scientific Notation
Punch the number (the digit number) into your calculator.
Push 2nd then EE button.
Use the - button to indicate sign of exponent (if negative)
Enter the exponent number. Voila! Treat this number normally in all subsequent calculations.
To check yourself, multiply 6.0 x 105 times 4.0 x 103 on your calculator. Your answer should be 2.4 x 109
**NOTE: you may use the x (times) button and the ^ (carrot) button instead but make sure you put the number in parentheses so your calculator follows correct order of operations.
Instructions for graphing calculator

30. Precision vs. Accuracy
Accuracy—how close the measured value is to the actual value
Precision—how close a series of measurements are to each other; reproducibility
100 mL beaker gives results +/- 1 mL
100 mL graduated cylinder gives results +/- 0.1 mL

31. Practice Problems

32. Practice Problems Classification of Physical/Chemical Changes
Burning magnesium
Alcohol evaporating
Two aqueous solutions are mixed and a yellow solid forms
Boiling water
Alka Seltzer in water forms gas bubbles

33. Practice Problems Classification of Physical/Chemical Changes
Burning magnesium—chemical
Alcohol evaporating—physical
Two aqueous solutions are mixed and a yellow solid forms—chemical
Boiling water—physical
Alka Seltzer in water forms gas bubbles—chemical

34. Practice Problems
Use dimensional analysis to convert between metric units.
How many mL are in 5 kL?
How many grams are in 7.9 mg?

35. Practice Problems
Use dimensional analysis to convert between metric units.
How many mL are in 5 kL? 5,000,000 mL
How many grams are in 7.9 mg? g

36. How many Sig Figs?
5.6009
67,000,000
6.925 x 108

37. How many Sig Figs?
5.6009 5
67,000,000 2 3 8
6.925 x 108 4

38. Practice Problems Round each of these measurements to three sig figs.
5,907
67,412

39. Practice Problems Round each of these measurements to three sig figs.
5,907= 5,910
= 314 =
67,412= 67,400

40. Practice Problems Convert the following to Scientific Notation
890,000,000
706,079
Convert the following to standard notation.
8.3 x 10-3
7.902 x 107

41. Practice Problems Convert the following to Scientific Notation
890,000,000= 8.9 x 108
= 6.05 x 10-5
706,079= x 105
Convert the following to standard notation.
8.3 x 10-3 =
7.902 x 107 = 79,020,000