1. Chapter 1
The Study of Chemistry

2. Topics Introduction Scientific Method Classifications of Matter
Properties of Matter
Units of Measurement – Metric system
Temperature Conversion
Metric Conversion (Prefixes)
Accuracy vs. Precision
Significant Figures
Density

3. Scientific Method
1. Observation
2. Hypothesis – an educated guess
3. Experimentation
4. Results
5. Conclusion

4. States of Matter
Solid
Liquid
Gas
Plasma

5. ATOM
Is the simplest unit of matter.

6. Definitions
Elements – can’t be decomposed further into simpler substances elements presently Lv (element livermorium)
- Fl (element flerovium)
Compound – combination of 2 or more elements
Pure substance – has distinct properties and composition; does not vary from sample to sample (ex. Water, NaCl)

7. Definitions
Mixtures – combinations of 2 or more substances (ex. sugar in water)
2 Types of Mixtures
1. Homogenous Mixtures (solutions) = 1 phase
2. Heterogeneous Mixtures = > 2 phases

8. SOLUTIONS
Homogeneous mixtures are called SOLUTIONS.

9. Solution Solution – homogenous mixture
A solution is not necessarily a liquid. Can be gas or solid.

10. Physical vs. Chemical Properties
Physical properties – can be measured w/o changing identity and composition of substance (ex. Boiling pt.,freezing pt., color, odor, density, hardness)
Chemical properties – describe how substance reacts or changes to form other compounds (ex. Flammability, toxicity)

11. Changes of State and Properties
Physical changes – does not change composition of compound
Chemical changes – converts to a different chemical substance
Intensive Properties – independent of amt. (ex. Density, Temperature, Melting Pt)
Extensive Properties – dependent on amt. (ex. Mass, Volume)

12. Units of Measurement
Use the Metric System!

13. Units of Measurement Mass – grams; kilogram Length – centimeter; meter
Volume – milliliter or cubic centimeter (cm3)
Temperature – Celcius; Kelvin

14. Precision vs Accuracy
Accuracy – when acquired value agrees with true value
Precision – when acquired values exhibit reproducibility

15. Significant Figures
More significant figures = more certainty
Helps in determining how to round measured values and still precise

16. SIGNIFICANT FIGURES
In counting and definitions, there are an infinite number of sig figs
In measurements, the number of sig figs consists of all certain and the first uncertain digits
Unit conversions do not determine # of sig. figs.

17. Rules of Significant Figures
1. Non-zero integers always count.
Ex grams = 5 Sig. Figs.
2. Captive zeros are always significant.
Ex grams = 4 Sig. Figs.

18. Rules of Significant Figures
3. Leading zeros are NEVER significant.
Ex grams = 4 Sig. Figs.
4. Trailing zeroes are significant ONLY if there is a decimal point
Ex grams = 3 Sig. Figs
120 grams = 2 Sig. Figs

19. Rules of Significant Figures
5. Exact numbers (obtained by counting) are infinite and do not determine the number of significant figures.
Example: 4 cows = ?

20. Determine the # of Sig. Fig.
200.0
1050
3003
0.0006
10,000
0.5

21. Rules of Significant Figures
Multiplication/Division
Answer will have the same # of sig figs as the value with the least # of sig figs
Ex: 3.8 x = 2 Sig. Figs.

22. Rules of Significant Figures
Addition/Subtraction
Answer has the same # of decimal places as the number with the least # of decimal places
Ex = 11.4

23. Order of Operations
Parenthesis
Multiplication/division
Addition/subtraction

24. Rounding Look only to the right of the number you are rounding to:
- If 5 or more, round up
- If less than 5, round down

25. General Rule
Carry ALL figures through to the end of a problem. Round the final answer to the correct number of significant figures

26. Problem
Indicate the number of sig. figs. in each of the following measured quantities:
A kg
B s
C cm
D L
E x 10-3 m3

27. Problem
Round each of the following numbers to 4 sig. figs. And express the result in standard exponential or scientific notation. A
B. 656, 980 C D E

28. Problem
Carry out the following operations and express the answer with the appropriate number of sig. figs. A
B –
C. (6.21 x 103)(1.1050)
D / 0.753

29. Prefixes in Metric System
Mega - million
Kilo - 1,000
Hecto
Deka - 10
(liter, gram, meter)
Deci - 1/ or
Centi - 1/ or
Milli - 1/ or

30. Trick to Unit Conversions
Arrange the values with units such that you can cancel what you do not want.
You should only end up with the desired unit.

31. Tricks to Conversion Big to small - multiply Small to Big - divide
going down
Small to Big - divide
Going up

32. Temperature Conversions
0 oC = K
oF = oC + 32

33. Things to Remember!
1 milliliter = 1 cc
1000 milliliter = 1 liter
0 oC = 32 oF = K

34. Density Is the amount of mass in a unit volume of the substance
Is affected by Temperature.
The higher the temp., the lower the density.
D = mass of substance = grams volume of substance mL or cm3

35. Density
Density = mass volume = gram mL

36. Different ways of calculating volume
I. For solids with regular shapes:
A. For a cube: Vcube = s3
B. For a rectangular solid, V = L x W x H
C. For a cylinder: V= pr2h
D. For a sphere: V = 4/3 pr3

37. Different ways of calculating volume
II. For an Irregular Solid
Water displacement

38. Different ways of calculating volume
III. For a liquid
Use of graduated cylinder, beaker, pipet or buret.

39. Problem
A cube of osmium metal cm on a side has a mass of grams at 25 oC. What is its density in g/cm3 at this temperature?

40. Problem
The density of titanium metal is 4.51 g/cm3 at 25 oC. What mass of titanium displaces 65.8 mL of water at 25 oC?

41. Problem
The density of benzene at 15 oC is g/mL. Calculate the mass of L of benzene at this temperature.