1. Chemistry
an introduction

2. Why is Chemistry important?
In our daily lives:
New materials
New pharmaceuticals
New energy sources
Food supplies
Can you think of anything else?

3. Chemistry
Is the science that deals with the materials of the universe and the changes that those materials undergo

4. Chemical Changes What are some examples of chemical changes?
Iron rusting
Wood burning
Food cooking
Grape juice fermenting
Plants growing
How do we know that these are chemical changes?

5. Steps in the Scientific Method
Observations
Quantitative vs Qualitative
Quantitative – measurement involves a number and a unit
Formulating Hypotheses
Possible explanation for the observation
Performing Experiments
Gathering new information to decide whether the hypothesis is valid

6. Quantitative & Qualitative Observations
Qualitative Quantitative
red book 4 quarters
round tire 6 wheels
wooden desk 24 students
metal chair 5 atoms
aluminum foil 65°C
glass square 2” x 4” x 8”
rough board 2 graduated cylinders

7. Outcomes over the Long Term
Theory (Model)
A set of tested hypotheses that give an overall explanation of some natural phenomenon
Natural Law
The same observation applies to many different systems
Ex. Law of Conservation of Mass

8. Law vs Theory
A law summarizes what happens; a theory (model) is an attempt to explain why it happens

9. The various parts of the Scientific Method

10. Problems with the scientific method
Scientists must be objective when using the scientific method. The scientific method is affected by:
Profit motives Religious Beliefs
Wars Misinterpretation of Data
Budgets Emotions
Fads Prejudices
Politics Peer Pressure

11. Scientific Terminology
What is the difference between a hypothesis and a theory?
What is the difference between an observation and a theory?
What is the difference between a natural law and a theory?

12. The Fundamental SI Units
Physical Quantity Name Abbreviation
mass kilogram kg
length meter m
time second s
temperature Kelvin K
Electric Current Ampere A
Amount of Substance mole mol
Luminous Intensity candela cd

13. SI Prefixes Common to Chemistry
Unit Abbr.
Exponent
Mega M
106
Kilo k
103
Deci d
10-1
Centi c
10-2
Milli m
10-3
Micro
10-6
Nano n
10-9
Pico p
10-12

14. Uncertainty in Measurement
A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.
Measurements are performed with instruments
No instrument can read to an infinite number of decimal places.

15. Precision and Accuracy
Accurate and precise
Precise, but not accurate
Neither accurate not precise
Accuracy refers to the agreement between the measure quantity and the accepted value
Precision refers to the degree of agreement of several repeated measurements (made in the same manner) to each other.

16. Types of Error Random Error (Indeterminate Error) –
Measurement has an equal probability of being high or low
Systematic Error (determinate Error) –
Occurs in the same direction each time (high or low), often resulting from poor technique or incorrect calibration.
This can result in measurements that are precise, but not accurate

17. Rules for Counting Significant Figures
Non-zero integers always count as sig. fig.
3456
4 sig figs

18. Rules for Counting Significant Figures
Zeros
Leading Zeros do not count as sig figs
0.0486
3 sig figs

19. Rules for Counting Significant Figures
Zeros
Captive Zeros always count as sig figs
16.07
4 sig figs

20. Rules for Counting Significant Figures
Zeros
Trailing Zeros are significant only if the number contains a decimal point.
9.300
4 sig figs

21. Rules for Counting Significant Figures
Exact Numbers have an infinite number of significant figures.
1 inch = 2.54 cm

22. Practice Counting Significant Figuresm
17.10 kg
100,890 L
3.29 x103 s cm
3, 200, 000
5 sig figs
4 sig figs
3 sig figs
2 sig figs

23. Rules for Significant Figures in Mathematical Operations
Multiplication and Division
number of sig figs in the results equals the number of sig figs in the least precise measurement used n the calculation (the one with the lowest number of sig figs).
6.38 x 2.0 = 12.76
13 (2 sig figs)

24. Practice for Significant Figures in Mathematical Operations
Answer
23 m2
4.22 g/cm3
0.05 cm2
240 m/s
5870 lb·ft
2.96 g/mL
Calculation
Calculator Says
3.24 m x 7.0 m
22.68 m2
100.0 g ÷ 23.7 cm3
g/cm3
0.02 cm x cm
cm2
710 m ÷ 3.0 s
m/s
lb x 3.23 ft
lb·ft
1.030 g ÷ 2.87 mL
g/mL

25. Rules for Significant Figures in Mathematical Operations
Addition and Subtraction
The number of decimal places in the result equals the number of decimal places in the least precise measurement =
18.7 (3 sig figs)

26. Practice for Significant Figures in Mathematical Operations
Calculation
Calculator Says
3.24 m m
10.24 m2
100.0 g cm3
76.27 g/cm3
0.02 cm cm
2.391 cm2
713.1 m – s
m/s
lb lb
lb·ft
2.030 mL mL
0.16 g/mL
Answer
10.2 m
76.3 g
2.39 cm
709.2 L lb
0.160 mL