Chapter 1 Chemical Foundations

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Chapter 1Chemical Foundations
2
Why Chemistry?
Chapter 1 Section 1

• Everything in this universe is made out of
  approximately 100 different kinds of atoms.
• Sand (Silicon, Oxygen)
• Table Salt (Sodium, Chloride)
• Water (Oxygen, Hydrogen)
• They are as letters in an alphabet.

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Why Chemistry?
Chapter 1 Section 1
4
Scientific Method
Chapter 1 Section 1

• One of lifes most important activities is
  solving problems that affects the people and the
  world they live in.
• The scientific method is the way used by
  scientists to understand the universe and its
  changes.
• The more creative youre at solving problems, the
  more effective you will be in your career and
  your personal life.
• Chemistry helps to develop solving problems
  capabilities.

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Scientific Method
Chapter 1 Section 2

• To understand the universe and its changes.

• Steps of scientific methods
• Observation.
• Can be qualitative or quantitative.
• Prediction (Hypothesis).
• Trying to explain the observation.
• Experiment.
• Performed to test the validity of the hypothesis.
• Experiments always produce new information.

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Scientific Model
Chapter 1 Section 2

• A set of valid hypotheses is assembled into a
  theory (model).
• This theory is modified as more observations are
  recorded. Thus, it is a continuous process.
• Observation is something that is witnessed from
  experiments.
• Theory is an explanation why nature behaves in a
  particular way.

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Theory vs. Law
Chapter 1 Section 2

• Theory changes over time as more observations
  from experiments are recorded.
• Some observations are found to apply to many
  different systems.
• Law
• A summary of observed behaviors applied to
  different systems.
• Law of conservation of mass.
• Law of conservation of energy.
• Theory
• An attempt to explain these observations.

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Measurements
Chapter 1 Section 3

• Making observations can be done quantitatively or
  qualitatively.
• A quantitative observation is called a
  measurement. It must include two important pieces
  of information

Number
Unit
(1) Metric Units International System (SI Units)
There are two major systems of measurements
(2) English Units
Used in Science
9
Units of Measurements
Chapter 1 Section 3
There are two major systems of measurements
Number
Unit
(1) Metric Units International System (SI Units)
10
kilometers
1
gram
5
kelvin
Used in Science
20
miles
(2) English Units
1
pound
60
Fahrenheit
10
The Fundamental SI Units
Chapter 1 Section 3
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Using Prefixes in the SI System
Chapter 1 Section 3

• The distance between Dammam and Jubail is 90,000
  meters.
• 90103 meters.
• 90 kilo-meters (km)
• The capacity of this computer is 80,000,000,000
  bites.
• 80109 bites
• 80 giga-bites (GB)

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Using Prefixes in the SI System
Chapter 1 Section 3

• 0.0001 kg
• 110-4 kg
• 0.1 g
• 1 with 35 zeros kg
• 11035 kg

Scientific Notation
How much does that pin weigh?
How much does the earth weigh?
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Table of Prefixes in the SI System
Chapter 1 Section 3
Must be memorized!
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Measurement of Volume
Chapter 1 Section 3

• Volume is not an SI unit, but it is extremely
  important in chemical measurements.
• Volume 1m 1m 1m 1m3
• 1m 10dm
• (1m)3 (10dm)3
• 1m3 1000dm3
• 1dm3 1L
• 1L 1000mL 1000cm3
• 1mL 1cm3
• (milli)Liter milli 10-3
• (centi)Meter centi 10-2

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Measurement of Volume
Chapter 1 Section 3
Common types of laboratory equipment used to
measure liquid volume.
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Measurement of Mass (Weight)
Chapter 1 Section 3
An electronic analytical balance.
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Uncertainty in Measurements
Chapter 1 Section 4

• A measurement always has some degree of
  uncertainty.

• Uncertainty is 0.01 ml.
• Certain and uncertain digits are known as
  significant figures.

20.16 ml 20.17 ml 20.15 ml 20.18 ml 20.16 ml
certain digit
uncertain digit (must be estimated)
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Uncertainty in Measurements
Chapter 1 Section 4

• Now you should be able to tell how many digits
  you need to include in your reading
  (measurement).
• 20 ml
• 20.1 ml
• 20.16 ml
• 20.160 ml
• 20.1600 ml

Which one??
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Uncertainty in Measurements
Chapter 1 Section 4

• An apple is measured using a bathroom scale and
  an electronic balance
• Bathroom scale Electronic balance
• 0.3 kg 0.332 kg
• Does the apple have two different masses?
• Different equipments have different uncertainties
  in their measurements.

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Precision and Accuracy
Chapter 1 Section 4
Random error
Systematic error
Precise but not accurate reproducible
Neither precise nor accurate
Accurate
Poor technique
Good technique but needs calibration
Good technique
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Precision and Accuracy
Chapter 1 Section 4

• Accuracy Agreement of a particular value
  (measurement) with the true value.
• Precision Agreement among several values
  (measurements), not necessarily agreeing with the
  true value.

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Sample Exercise 1.2
Chapter 1 Section 4

• A student used a buret to check the accuracy of a
  graduated cylinder at the 25-ml mark. He got the
  following results
• Trial Volume shown by Volume shown
  by graduated cylinder the buret 1 25
  mL 26.54 mL 2 25 mL 26.51
  mL 3 25 mL 26.60 mL 4 25
  mL 26.49 mL 5 25 mL
  26.57 mL
• Average 25 mL 26.54 mL
• It is precise but not accurate (systematic
  errors)

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Significant Figures
Chapter 1 Section 5

• In many cases, important physical quantities are
  obtained from measured values.
• Volume l w h
• Density mass / volume
• Calculations need to be done on the basis of
  Significant Figure (S.F.) Rules
• Rules for counting S.F.
• Rules of mathematical operations on S.F.
• Implication of the word Significant. It is to
  have the correct degree of uncertainty in the
  resultant physical quantities.

Mathematical operations
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Rules for Counting Significant Figures
Chapter 1 Section 5

• 1- Nonzero integers are always counted as S.F.
• Example Give the number of S.F. for the
  following
• 34
• 236
• 17296.1
• 12.1102 Exponential (Scientific) notation

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Rules for Counting Significant Figures
Chapter 1 Section 5

• 2- Zeros (leading zeros, captive zeros, and
  trailing zeros)
• a) Leading zeros are not counted as S.F.
• Example Give the number of S.F. for the
  following
• 00121.1
• 0.0025

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Rules for Counting Significant Figures
Chapter 1 Section 5

• b) Captive zeros are always counted as S.F.
• Example Give the number of S.F. for the
  following
• 1.008
• 701.1 10-4
• 3.000000008
• 0.0901

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Rules for Counting Significant Figures
Chapter 1 Section 5

• c) Trailing zeros are counted as S.F. only if the
  number contains a decimal point.
• Example Give the number of S.F. for the
  following
• 1.000
• 320.00 10-1
• 100
• 100.
• 100.0

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Rules for Counting Significant Figures
Chapter 1 Section 5

• 3- Exact numbers are assumed to have an infinite
  number of S.F.
• Examples
• 3 Apples is 3.00000000 (zeros are all the way to
  8)
• 2 in 2pr (the circumference of a cycle).
• 1 km 1000 m
• 1 in 2.54 cm

Mathematical relationships
Definitions
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Scientific Notations
Chapter 1 Section 5

• 123.1 1.231 102 1.231 100
• 0.00013 1.3 10-4 1.3 / 10000
• 0.13 10-3
• Avogadros Number
• 602,214,000,000,000,000,000,000
• The number of atoms contained in 12 g of carbon
  and is equal to 1 mole.

6.022 1023 Scientific notation is a very
convenient way to express the number of atoms in
chemistry problems.
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Mathematical Operations
Chapter 1 Section 5

• Multiplication or division
• 4.56 1.4 6.38 6.4

From calculator before correction
Number of S.F.
3
2
2 S.F. (After correction)

• Addition and Subtraction
• 12.11
• 18.0
• 1.013 31.1
• 31.123

3 S.F. (After correction)
before correction
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Rule for Rounding
Chapter 1 Section 5

• 6.38
• The digit to be removed
• If 5, then round up, i.e. the 3 becomes 4.
• 6.4
• If lt 5, then the digit stays unchanged.
• 6.34 becomes 6.3

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Exercises
Chapter 1 Section 5

• Perform the following mathematical operation and
  express the result to the correct number of
  significant figures

Rounding off should be carried out for the final
answer and NOT to the intermediate answers.
However, you must keep track of the significant
figures in the intermediate steps.
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Dimensional Analysis
Chapter 1 Section 6

• Used to convert from one unit to another.

Example 1 How many centimeters in 25.5 inches
(in)?
64.8 cm
Example 2 How many inches in 25.5 centimeters?

10.0 in
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Dimensional Analysis
Chapter 1 Section 6

• 1L 1000 ml
• 1 ml 0.001 L
• 1 10-3 L

Example 3 How many ml are in 1.63 L?

Which direction you choose? 1.63 L
1.63 L
L2
1 L
0.00163
ml
1000 ml
?
1000 ml
1.63103 ml
1 L
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Dimensional Analysis
Chapter 1 Section 6

• A complete list of conversion factors, that you
  will need in solving some homework and quizs
  problems are available in Appendix 6 in your book
  (page A26). Please study it very carefully.
• Also carefully study sample exercises 1.5 thru
  1.9. You must understand them before moving to
  Chapter 2.

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Dimensional Analysis
Chapter 1 Section 6

• How many centimeters in 0.25 megameters?
• 0.25 megameters
  0.25106 m
• 0.25106 m 0.25108
  cm
• 
  2.5107 cm
• 
  25.106 cm

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Temperature
Chapter 1 Section 7

• Three systems are used to measure temperatures
• Celsius scale (C)
• Kelvin scale (K)
• Fahrenheit (F)
• You have to be able to convert from one scale to
  another.

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The Three Major Temperature Scales
Chapter 1 Section 7
Celsius (oC)
Kelvin (K)
Fahrenheit (F)
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Celsius Scale vs. Kelvin Scale
Chapter 1 Section 7

• Temperature scales for C and K are identical,
  but their zeros are different.
• TK TC 273.15
• TC TK 273.15

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Fahrenheit Scale vs. Celsius Scale
Chapter 1 Section 7

• Both unit temperature size and zero locations are
  different.
• Since
• 180F 100C gt 9F 5C
• and
• 32F 0C
• Then to convert from F to C
• Tf (F) 32 (F) Tc (C)

5C
9F
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Fahrenheit Scale vs. Celsius Scale
Chapter 1 Section 7

• Tc (a) Tc (b) C (5C/9F) Tf (a) Tf
  (b) F
• When Tc (b) - 40C, then Tf (b) -40F
• giving that
• Tc (a) (-40) C (5C/9F) Tf (a)
  (-40) F
• Thus
• or

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Density
Chapter 1 Section 8

• It is the mass of substance per unit volume.

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Exercise 65 from Chapter 1

• Methodology to solve such problems
• Start with the quantity given in the question.
• Use possible conversion factors to convert the
  unit of the given quantity in the question to the
  desired/needed unit.

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Classification of Matter
Chapter 1 Section 9

• Matter is anything occupying the space and having
  a mass.
• Classes of matter
• Solid rigid, has fixed volume and shape.
• Liquid has definite volume but not fixed shape.
• Gas has no fixed volume or shape. and is
  compressible.

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Mixtures
Chapter 1 Section 9
Mixtures have variable composition.

• Gasoline
• Juices
• Air (N2,O2,H2,CO2, etc.)

• Wood
• Sand
• Rocks

Physical methods
Are those with constant composition.

• H2O
• NaCl
• H2SO4

• Carbon (C)
• Hg (mercury)

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Separation (Physical) Methods
Chapter 1 Section 9

• Based on the physical properties of the
  substances (boiling point, adsorption,
  solubility, etc.)
• Separation methods discussed in the text
• Distillation.
• Filtration.
• Chromatography.

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Distillation
Chapter 1 Section 9
Separation of mixture components by distillation
depends on the difference in boiling point
temperatures of the components contained in that
mixture.
Mixture
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Distillation
Chapter 1 Section 9
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Filtration
Chapter 1 Section 9

• For mixtures of solids and liquids.

Gravity Filtration
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Chromatography
Chapter 1 Section 9
Stationary phase A piece of film or paper the
road.
The mixture whose components are to be separated
the riders.
Mobile phase the driver.
Paper Chromatography Experiment
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Chromatography
Chapter 1 Section 9
3
2
1
(a)
(b)
(c)
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Chemical Methods
Chapter 1 Section 9
Substances that cant be decomposed into simpler
substances by physical or chemical means
Substances with constant composition that can be
broken down into elements by chemical processes

• Chemical changes (decomposition of compounds or
  recombining elements), such as electrolysis,
  heating, and photolysis.

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Components of Chemical Elements
Chapter 1 Section 9