CHAPTER ONE

1
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

• Liquid

• Solid

• Gas

6
States of Matter

• Change States
• heating
• cooling

Steam
Water
Ice
7
Substances

• Substance
• 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
• 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 ironsulfur watersand

10
Chemical Properties

• Chemical properties
• 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

• Compounds
• 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)
• www.webelements.com

14
Scientific Notation

• Use it when dealing with very large or very small
  numbers
• 42,800,000.
• 0.00000005117

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 10-1
• centi- c 10-2
• milli- m 10-3
• micro- ? 10-6
• nano- n 10-9
• pico- p 10-12

17
Metric Prefixes Examples

• 1000 m
• 0.008 s
• 30,000,000 g
• 0.07 L

18
Use of Numbers

• Exact 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
gt13.5 mm but lt13.7 mm
13.6 mm
21
Significant Figures
13.6 mm
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
• 0.000357 has three significant figures
• Zeros between nonzero digits are always
  significant
• 20.034 1509 1.0000005

23
Significant Figures - Rules
Trailing zeros

• Zeroes at the end of a number that contains a
  decimal point are always significant
• 35.7000 0.07200 40.0 41.0
• 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
• 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
x
x
9.32 yrd x
x
8.52103 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. lt 1 lighter than water
• Sp. Gr. gt 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 1.31. Calculate the mass of pure H2SO4
  in 100.0 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 100.0 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 gt T2 colder
34
Temperature Scales

• 3 common temperature scales

? Fahrenheit
? Celcius
? Kelvin
0ºF freezing (saltH2O) 30ºF freezing
H2O 90ºF human body
0ºC freezing H2O 100ºC boiling H2O
0 K absolute zero 273.15 K freezing H2O
http//home.comcast.net/igpl/Temperature.html
35
Temperature Scales Water
Melting (MP) and boiling (MP) points of water on
different temperature scales
36
Temperature Conversion
degrees Kelvin degrees Celcius
? K ?ºC 273
?ºC ? K - 273
degrees Fahrenheit degrees Celcius
?ºF (?ºC)1.8 32
?ºC (?ºF 32)/1.8

• Examples 1-16 1-17 in the book
• http//www.lenntech.com/unit-conversion-calculator
  /temperature.htm

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 4.184 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
  Jg-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 0.900
  and 4.18 Jg-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
Lets 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)
• 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)