1. The Foundations of Chemistry
CHAPTER ONE
The Foundations of Chemistry

2. Why is Chemistry Important?
Components for computers and other electronic devices
Materials for our homes
Fuel
Body functions
Cooking

3. Some definitions / Vocabulary
Chemistry
Science that describes matter – its properties, the changes it undergoes, and the energy changes that accompany those processes
Matter
Anything that has mass and occupies space.
(In other words: anything that has mass and volume)
Energy
The capacity to do work or transfer heat.
Types of energy
Kinetic and potential energy
Heat energy, light energy,
chemical energy, mechanical energy

4. Natural Laws The Law of Conservation of Mass
During a chemical or physical change the mass of the system remains constant
The Law of Conservation of Energy
Energy cannot be created or destroyed in a chemical reaction or in a physical change. It can only be converted from one form to another.
The Law of Conservation
of Matter and Energy
Read at home

5. States of Matter
Solid
Liquid
Gas

6. States of Matter
Change States
heating
cooling
Steam
Water
Ice

7. Substances Substance Examples Properties
matter all samples of which have identical composition and properties
Examples
water
sulfuric acid
Properties
physical properties – physical changes
chemical properties – chemical changes

8. Physical Properties Physical properties
changes of state
density, color, solubility
always involve only one substance
A substance cannot be broken down or purified by physical means!

9. Mixtures Mixture Homogeneous mixtures Heterogeneous mixtures
a combination of two or more substances
can be separated by physical means
Homogeneous mixtures
have uniform properties throughout
examples: salt water; air
Heterogeneous mixtures
do not exhibit uniform properties throughout
examples: iron+sulfur; water+sand

10. Chemical Properties Chemical properties Examples chemical reactions
always involve changes in composition
always involve more than one substance
Examples
burning of methane
rusting of iron
oxidation of sugar

11. Decomposition of Water
hydrogen
Element
Element
oxygen
water
Compound

12. Compounds and Elements
If a substance can be decomposed into simpler substances through chemical changes, it is called a compound
Elements
If a substance cannot be decomposed into simpler substances by chemical means, it is called an element

13. Compounds and Elements
Important to remember
both compounds and elements are substances
a compound consists of 2 or more elements
Law of Definite Proportions
different samples of any pure compound contain the same elements in the same proportion by mass
Symbols of elements
found on the periodic chart (learn Table 1-2)

14. Scientific Notation
Use it when dealing with very large or very small numbers:
42,800,000. = =

15. Measurements in Chemistry
Quantity Unit Symbol
length meter m
mass kilogram kg
time second s
current ampere A
temperature Kelvin K
amt. substance mole mol

16. Metric Prefixes Name Symbol Multiplier mega- M 106 kilo- k 103
deci d
centi c
milli m
micro 
nano n
pico p

17. Metric Prefixes: Examples
30,000,000 g =
0.07 L =

18. Use of Numbers Exact numbers Measured numbers
obtained from counting or by definition
1 dozen = 12 things for example
Measured numbers
numbers obtained from measurements are not exact
every measurement involves an estimate

19. Significant Figures Significant figures
digits believed to be correct by the person making the measurement

20. Significant Figures Side B: 13.6 mm 13.6 mm
>13.5 mm but <13.7 mm
in my judgement!
13.6 mm
certain figures
estimated figure

21. Significant Figures 13.6 mm significant figures estimated figure
certain figures +
significant figures
we always report only 1 estimated figure
the estimated figure is always the last one of the significant figures

22. Significant Figures - Rules
Exact numbers (defined quantities) have an unlimited number of significant figures. We do not apply the rules of significant figures to them.
Leading zeroes are never significant:
has three significant figures
Zeros between nonzero digits are always significant:

23. Significant Figures - Rules
Trailing zeros
Zeroes at the end of a number that contains a decimal point are always significant:
Zeroes at the end of a number that does not contain a decimal point may or may not be significant (use scientific notation to remove doubt):
173,700 may have 4, 5, or 6 significant figures

24. Significant Figures - Rules
Addition/Subtraction Rule
The position of the first doubtful digit dictates the last digit retained in the sum or difference.
Multiplication/Division Rule
In multiplication or division, an answer contains no more significant figures than the least number of significant figures used in the operation.
Study examples 1-1 & 1-2 in the book

25. The Unit Factor Method The basic idea of the method: Principles:
multiplication by unity (by 1) does not change the value of the expression
Principles:
construct unit factors from any two terms that describe identical quantity
the reciprocal of a unit factor is also a unit factor
Study examples 1-3 through 1-9 in the book

26. The Unit Factor Method 1 ft = 12 in Unit factors:
Example: Express 77.5 inches in feet
77.5 in = 77.5 in x
= 6.46 ft
See Table 1-7 for various conversion factors

27. More examples
9.32 yrd = ? mm
1. We use the following knowledge to build unit factors:
1 yrd = 3 ft 1 in = 2.54 cm
1 ft = 12 in 1 cm = 10 mm
2. Multiply 9.32 yrd by unit factors to get the value expressed in mm:
3 ft
1 yrd
12 in
1 ft
2.54 cm
1 in
10 mm
1 cm
9.32 yrd x x x x
= 8.52·103 mm

28. Density mass volume density =
tells us how heavy a unit volume of matter is
usually expressed as “g/ml” for liquids and solids and as “g/L” for gases
Table 1-8 lists densities of some common substances

29. Density: Example
Example: Calculate the density of a substance if 742 grams of it occupies 97.3 cm3.
Learn examples 1-11 through 1-13 in the book

30. Specific Gravity d (substance) d (water) Sp. Gr. =
tells us how much heavier or lighter a substance is compared to water:
Sp. Gr. < 1 – lighter than water
Sp. Gr. > 1 – heavier than water
specific gravity has no units – it is a dimensionless quantity
See example 1-14 in the book

31. Specific Gravity: Example
Example 1-15: Battery acid is 40% sulfuric acid, H2SO4, and 60% water by mass. Its specific gravity is Calculate the mass of pure H2SO4 in mL of battery acid.
What do we know?
1. The mass percentage of H2SO4 and H2O in the sample of battery acid.
2. Specific gravity of battery acid.
3. Density of water (1.00 g/mL).
To find the mass of H2SO4, we need to know the mass of mL of battery acid.

32. Specific Gravity: Example
Therefore,

33. Heat and Temperature Heat and Temperature are not the same thing:
Heat is a form of energy
T is a measure of the intensity of heat in a body
Heat always flows spontaneously from a hotter body to a colder body – never in the reverse direction
Body 1 T1
Body 2 T2
Heat
hotter T1 > T2 colder

34. Temperature Scales 3 common temperature scales
 Fahrenheit
 Celcius
 Kelvin
0ºF – freezing (salt+H2O)
30ºF – freezing H2O
90ºF – human body
0ºC – freezing H2O
100ºC – boiling H2O
0 K – absolute zero
K – freezing H2O

35. Temperature Scales & Water
Melting (MP) and boiling (MP) points of water on different temperature scales
MP BP
Fahrenheit 32 oF 212 oF
Celsius oC 100 cC
Kelvin K 373 K

36. Temperature Conversion
degrees Kelvin degrees Celcius
? K = ?ºC + 273
?ºC = ? K - 273
degrees Fahrenheit degrees Celcius
?ºF = (?ºC)·
?ºC = (?ºF – 32)/1.8
Examples 1-16 & 1-17 in the book

37. Heat Chemical and Physical changes:
evolution of heat (exothermic processes)
absorption of heat (endothermic processes)
Units of measurement:
joule (J) – SI units
calorie (cal) – conventional units
1 cal = J
A “large calorie” (1 large cal = 1000 cal = 1 kcal) is used to express the energy content of foods

38. Specific Heat The specific heat (Cp) of a substance:
the amount of heat (Q) required to raise the temperature of 1 g of the substance 1ºC (or 1 K)
Units of measurement:

39. Specific Heat: Example 1
Knowing specific heat, we can determine how much energy we need in order to raise the temperature of a substance by T = T2 – T1:
Calculate the amount of heat necessary to raise the temperature of 250 mL of water from 25 to 95ºC given the specific heat of water is 4.18 J·g-1 ·ºC-1.
What do we know?
the temperature change
the specific heat of water
the volume of water
the density of water

40. Specific Heat: Example 1
Examples 1-18 through 1-20 in the book

41. Specific Heat: Example 2
Given specific heats of two different substances, we can also calculate the heat transfer between them:
0.350 L of water at 74.0ºC is poured into an aluminum pot at room temperature (25.0ºC). The mass of the pot is 200 g. What will be the equilibrium temperature of water after it transfers part of its heat energy to the pot? The specific heats of aluminum and water are and 4.18 J·g-1 ·ºC-1, respectively.
You might encounter this kind of problem at your first exam

42. Specific Heat: Example 2
What do we know?
the pot and water come to equilibrium, that is eventually they have the same temperature
the specific heat of aluminum and water
the mass of aluminum
the volume of water
the density of water
finally, the Law of conservation of energy which tells us that the amount of heat lost by water is the same as the amount of heat gained by the aluminum pot

43. Specific Heat: Example 2
Let’s denote the final temperature as Tf. Then the changes in temperature for water and aluminum are:
Note that we used the unit factor method to convert L to mL

44. Specific Heat: Example 2
Solving this equation with respect to Tf, we obtain
Tf = 68.6ºC
Try to solve the equation yourself and analyze why the answer is given with 3 significant figures

45. Reading Assignment Read Chapter 1 Learn Key Terms (pp. 40-41)
Go through Chapter 2 notes available on the class web site
If you have time, read Chapter 2

46. Homework Assignment Textbook problems (optional, Chp. 1): OWL:
11, 13, 15, 18, 27, 29, 30, 32, 36, 41, 43, 47, 49, 57, 62, 68, 80
OWL:
Chapter 1 Exercises and Tutors – Optional
Introductory math problems and Chapter 1 Homework problems – Required (homework #1; due by 9/13)