Review for Chapter 1: Chemistry: The Study of Change
1. Matter is anything that has mass and occupies space.
2. Matter can be classified into pure substances or mixtures.
Pure substances include elements and compounds. Elements are substances that
cannot be made any simpler by chemical means, such as hydrogen, carbon, and oxygen.
Compounds are substances made of atoms of two or more elements united in fixed
proportions, such as water or carbon dioxide.
Mixtures are combinations of two or more substances in which the substances retain
their identities. Mixtures can be either homogeneous (the same throughout) or
heterogeneous (non-uniform composition).
3. Substances can exist in three states of matter: gas, liquid, and solid. The states of
matter can be interconverted without changing the composition of the substance.
Changing the temperature is a common way of changing the state.
4. Substances are characterized by their physical and chemical properties.
A physical property can be measured without changing the composition of a substance.
A chemical property is measured during a chemical change of the substance, such as
burning.
5. SI units are used in chemistry. The base units are:
Length:
meter (m)
Mass:
kilogram (kg)
Time:
second (s)
Temperature: Kelvin (K)
6. Know the prefixes used with SI units (Table 1.3): tera-, giga-, mega-, kilo-, deci-,
centi-, milli-, micro-, nano-, and pico-.
7. Volume can be measured in cubic meters (m3) or in liters (L). Volumes used in the
lab are usually expressed in cubic centimeters (cm3) or in milliliters (mL). Remember
that:
1 mL = 1 cm3
1 L = 1000 mL
8. Density = mass
volume
For example, water at 25°C has a density of 1.00 g
cm3
9. Three temperature scales are commonly used: °F (degrees Fahrenheit), °C (degrees
Celsius), and K (Kelvin). Know how to convert from one temperature scale to another:
?°C = (___°F – 32°F) x 5°C
9°F
?°F = (___°C x 9°F) + 32°F
5°C
? K = (___°C + 273.15°C) x 1 K
1°C
10. Scientific notation makes it easier to handle numbers that are very large or small.
1. Move the decimal point so that you have a number that is between 1 and 10.
2. Count the number of places the decimal point is moved to get the power of 10.
3. If the original number is smaller than 1, the power of 10 should be negative.
If the original number is larger than 1, the power of 10 should be positive.
Examples: 0.0168 = 1.68 x 10-2
5278 = 5.278 x 103
11. Know how to determine the number of significant figures:
a. A digit that is non-zero is significant.
b. A zero between nonzero digits is significant: 404 has 3 significant digits.
c. Any zeros to the left of the first nonzero digit are not significant: 0.001 has only 1
significant digit.
d. For numbers greater than 1, zeros to the right of a decimal point are significant:
3.0 has two significant digits.
e. For numbers that do not contain decimal points, trailing zeros may or may not be
significant, depending on the measurement. For example, 5000 may have 1, 2, 3, or 4
significant digits. We can avoid this confusion by writing numbers in scientific notation
to express the correct number of significant digits.
Example: 5 x 103 or 5.0 x 103 or 5.00 x 103 or 5.000 x 103
12. When you add or subtract numbers, the answer cannot have more digits to the
right of the decimal point than either of the original numbers.
Example: 28.332 + 1.1 = 29.432  round to 29.4
13. When you multiply or divide numbers, the number of the significant figures in the
answer is determined by the original number that has the smallest number of significant
figures.
Example: 1.62 x 3.2853 = 5.322186  round to 5.32
14. Accuracy refers to how close a measurement is to the true value of the quantity
being measured while precision refers to how closely two or more measurements of the
same quantity agree with each other. Scientists seek both accuracy and precision.
15. Always write the units with any number! During calculations, show how certain
units cancel out and verify that the answer has the proper units.
Example: ? g = 100 lb
? g = 100 lb x 453.6 g = 454 g
1 lb