1. Introduction to Chemistry
Chapter 1

2. The Nature of Science and Chemistry
Chemistry is essentially the study of matter and its interactions
Includes both physical and chemical changes and the energy flow associated with them

3. Chemistry and the Scientific Method
Identify
the
problem
Make
Observations
Hypothesis
Experiment
Draw a
Conclusion

4. Parts of an Experiment
Experiment: a set of controlled observations that test a hypothesis
Consists of:
Independent variable - the variable you plan to change
Dependent variable - the variable that changes in response to a change in the independent variable
Control - a standard for comparison
Data – information/ observation recorded 1) Qualitative: uses descriptive words
Example- colors
2) Quantitative: uses numbers

5. Observation vs. Inference
is an experience perceived through one or more of the senses
What you feel, smell, see, hear, or taste
Example of observation
The temperature of a flask goes up as a reaction proceeds
Inference
an interpretation of an observation which goes one step beyond an observation.
The “story” you create about what you see, feel, hear, smell or taste which is based upon a direct observation

6. Qualitative vs. Quantitative Examples
1. Smells like apples 6. red
lbs mph
degrees C 8. round
4. Tall cold
5. Pink with purple dots

7. Scientific Method Theory & Law
a broad and extensively tested explanation of “WHY” experiments give results.
Supported by many experiments
Considered successful if it can be used to make predictions that are true
Natural Law
a concise statement that summarizes the results of many observations and experiments

8. Matter
Matter is anything that has mass and takes up space (by definition)
Weight is the force of attraction that results from an object possessing mass
They are not synonymous

9. Classification of Matter
Pure substance
Has a constant composition
It’s identified by a formula or symbol
Examples
water- H2O – always has 2 hydrogens and 1 oxygen
Helium – He
Compound
Composed of two or more different types of atoms that cannot be decomposed by physical means
Chemically combined
Sodium chloride- NaCl
Carbon dioxide- CO2

10. Classification of Matter
Element
Simple substance that cannot be decomposed into simpler substances by any chemical change

11. Mixtures of Matter
A mixture is a combination of two or more pure substances
The individual pure substances do not change their chemical properties, however
There are two types of mixtures:
Heterogeneous mixtures do not have the same consistency throughout
Homogeneous mixtures do have the same consistency

12. Examples of Mixtures Solution A on the left is
an example of a heterogeneous
mixture
Solution B on the right is an
example of a homogeneous
mixture

13. Completed Classification Flowchart
Matter
Pure Substances
Elements
Compounds
Mixtures
Heterogeneous
Homogeneous

14. Can be divided into two categories:
Properties of Matter
PHYSICAL
Properties: Can be measured or observed without changing the chemical makeup (composition) of the object being studied
Changes: Changes in physical appearance that do not result in a change in composition
CHEMICAL
Properties: Describe chemical reactivity (cannot be observed without changes in composition)
Changes: Akin to chemical properties. Occur as one substance is converted into a completely different substance
Can be divided into two categories:
Extensive Properties: Depend on the amount of sample
Intensive Properties: Depend on what the sample is; not amount

15. Measurements and Uncertainty
Accuracy involves comparison with an accepted value (Poor accuracy results in high error)
Precision involves comparison between measurements of the same quantity (Poor precision results in high uncertainty)

16. Rules for Determining Significant Figures
All nonzero digits are significant. Ex: (4 sig. figs.)
Zeros between significant digits are significant. Ex: (4 sig. figs.)
Zeros at the beginning of a number are never significant. Ex: (2 sig. figs.)
Zeroes at the end of a number are significant IF the number contains a decimal. Ex: (4 sig. figs.)

17. Number of Significant Figures?
0.709
0.4600
12000
12001
See Example 1.3 (Pg. 13) 3 4 2 5 7 6

18. Significant Figures in Calculations
“The operator must report the proper number of significant figures; the calculator will not.”
More rules regarding significant figures (why not?)
Addition/Subtraction
Answers must be reported to the fewest number of decimal places
Multiplication/Division
Answers must be reported to the fewest number of significant figures

19. Example #1
What is the volume of an object (expressed to the appropriate number of sig. figs.) with the following dimensions: length = 1.2 m, width = 1.45 m, height = m?
See Example 1.5 (Pg. 16)
Formula: L x W x H
(1.2 m)(1.45 m)(0.678 m) = m3
= 1.2 m3

20. Example #2
A gas at 25 °C fills a container whose volume is 1.05 x 103 cm3. The container plus gave have a mass of g. The container when emptied of gas, has a mass of g. What is the density of the gas at 25 °C?
See Example 1.5 (Pg. 16)
Density = mass/Volume
D = ( )g / 1.05 x 103
cm3
= (1.4 g) / 1.05 x 103 cm3
= g/cm3 OR 1.3 x 10-3 cm3

21. Quantities That Are Not Limited by Significant Figures
Counted numbers
Represent exact quantities (as opposed to measured quantities)
Ex: Number of students present in the classroom, number of eggs in a carton, etc.
Defined numbers
Typically refers to conversion factors
Ex: 1 foot = 12 inches; 1 yard = 3 feet, 1 minute = 60 seconds, etc.
The power of 10
Scientific notation
Ex: The “10” in 6.62 x 1023

22. Examples There are 67 people in this classroom
There is no range between 66.1 and 67.0 people, there are exactly 67 people
The thermometer shown here reads approximately 27 C. This is a measured quantity and must therefore contain uncertainty.

23. Measurements and Units
Every measurement is made up of two components: a quantity and a unit
The modern scientific community uses the SI (Le Système International d’Unités) system of units to express measurements
Each fundamental quantity has its own base unit
Everything else is expressed as a combination of units or derived units
Ex: cm3, g/mL, g/mol

24. SI Base Units
Some of the base units have prefixes that can be used for convenience
Base Units in italics are not required for this course and may be ignored
Time
Second (s)
Length
Meter (m)
Mass
Kilogram (kg)
Temperature
Kelvin (K)
Amount of a substance
Mole (mol)
Electrical Current
Ampere (A)
Luminous Intensity
Candela (cd)

25. SI Prefixes
Used for sake of convenience

26. Temperature Conversion Factors
There are three common temperature scales
Fahrenheit ( °F )
Celsius ( °C )
Kelvin ( K )
The conversion between the Celsius and Kelvin scales is the most important:
TK = TC

27. Celsius and Fahrenheit Conversions
Celsius to Fahrenheit Conversion:
Fahrenheit to Celsius Conversion: