Chemistry

1
Chemistry

• Fall 2003
• Dr Supplee

2
Chapter 1- Definitions

• Science
• Methodical exploration of nature followed by a
  logical explanation of observations
• Scientific Method
• A systematic investigation of nature and requires
  proposing an explanation for the results of an
  experiment in the form of a general principle
  (hypothesis)

3
Chapter 1 - Definitions

• Hypothesis
• Initial explanation of observations
• Theory
• Sufficient evidence in support of the hypothesis
• Model that scientifically explains the behavior
  of nature
• Law
• Does not explain behavior
• States a measurable relationship under different
  experimental conditions

4
Chapter 1 Definition Examples

• Hypothesis
• Dalton proposed that all matter was composed of
  small individual particles (atoms)
• Theory
• 100 years later Atomic Theory which explains
  the composition of substances as well as the
  behavior of gases
• Law
• Boyles Law P1V1 P2V2 at constant temperature
• If volume decreases than pressure increases at
  constant temperature

5
Chapter 1- Definitions Summary
Scientific Theory
Natural Law
Analyze more data
Hypothesis
Analyze initial observations
Experiment
6
Chapter 1 Modern Chemistry

• Organic Chemistry
• Chemistry of carbon containing compounds
• (C, H, O, and N)
• Inorganic Chemistry
• Chemistry of all other substances
• Biochemistry
• Chemistry of substances derived from plant
  substances

7
Chapter 1 Modern Chemistry

• All three have in common
• Analytical Chemistry
• Qualitative (what) and quantitative (how much)
  analyses
• Physical Chemistry
• Theoretical and mathematical explanations of
  chemical behavior

8
Relevance to daily life
Interesting topics
Fun experiments
CHEMISTRY
Career Opportunities
Benefits to society
Applications
9
Chapter 2- Scientific Measurements

• Introduction to Laboratory
• Work alone
• Handout
• Due 9/15/03
• Measurement Uncertainty
• Plus/minus factor ( error)
• Metric versus English Units
• Conversion factors
• Significant Figures
• Rounding rules

10
Precision versus Accuracy
Precise, not accurate
True Value
Precision how close two measurements of the
same quantity are to each other
Accurate, not precise
Accuracy how close an experimental
observations to the true value
11
Chapter 2- Scientific Measurements

• Measurement
• a number with units
• Uncertainty
• the estimated unit amount
• plus/minus associated with measurement
• Mass
• Amount of matter an object possesses
• Weight
• Force exerted by gravity on an object

12
Chapter 2- Scientific Measurements

• Volume
• Amount of space occupied by a solid, gas or
  liquid

13
Significant Digits/ Figures

• Digits are significant when the do more than hold
  a decimal place
• A place holder zero is NEVER significant
• determines measurement uncertainty (error
  analysis)
• Does not apply for exact numbers, only measured
  numbers

14
Significant Digits/ Figures Rule

• Rule 1
• Count the number of nonzero digits left to right
• Do not count place holder zeros

15
Significant Figure Rounding Rules

• After all calculations are complete determine
  significant figures and then round
• 5 or greater round-up to the nearest whole number
• less than 5 truncate

16
Scientific Notation

• Exponential numbers (power of 10)
• Base 10exponent
• The number 10 is raised to the nth power
• Numbers greater than 1 the exponent is positive
• Numbers less than 1 the exponent is negative
• The decimal is placed after the first significant
  digit and sets the size of the number by using a
  power of 10.

17
Unit Equations, Factors and Conversions

• Problem Solving Technique
• Equivalent relationships
• Unit equation
• A simple statement of two equivalent quantities
• Unit Factor
• A ratio of two equivalent quantities

18
Unit Dimensional Analysis Problem Solving

• Three steps
• 1) write down the units asked for in the
  answer
• 2) write down the value given in the problem
  that is related to the required answer
• 3) Apply a unit factor to convert the units in
  the given value to the units in the answer
• Given Value x Unit units asked for
• Factor

19
Percent Concept

• amount of a single quantity compared to the
  entire sample
• one part per 100 parts
• one quantity x 100
• total sample

20
Review
21
Significant Digits/ Figures

• Digits are significant when the do more than hold
  a decimal place
• If the number is less than 1, a place holder zero
  is NEVER significant
• determines measurement uncertainty (error
  analysis)
• Does not apply for exact numbers, only measured
  numbers

22
Exact Numbers

• Infinite significant figures
• English to English conversion factors
• Metric to metric conversion factors

23
Unit Equations, Factors and Conversions

• Problem Solving Technique
• Equivalent relationships
• Unit equation
• A simple statement of two equivalent quantities
• Unit Factor
• A ratio of two equivalent quantities

24
Chapter 3 The Metric System

• Single basic unit for each quantity measured
• Decimal system that uses a system of prefixes to
  enlarge or reduce a basic unit

25
Metric System Definitions

• Meter equals one ten-millionth of the distance
  from the North Pole to the equator
• Kilogram equals the mass of one a cube of water
  one-tenth of a meter on a side
• Liter equals the volume occupied by a kilogram of
  water at 4 oC

26
The Metric System
27
Metric Prefixes
28
Metric Conversion Factors Practice

• 1 kg ? g
• k kilo 1000 basic units
• 1kg 1000g
• 2s ? ns
• nnano1 1 x 10-9
• 2s2 x 10-9 ns

29
Unit Conversion Factors

• Ratio of two equivalent quantities
• The quantity in the numerator is equal to the
  quantity in the denominator
• If 100cm 1 m, then the factor becomes
• 100 cm or 1m
• 1 m 100 cm

30
Metric- English Conversions
31
Unit Analysis

• Recall
• Problem Solving Technique

Units Given
Unit Factor
New unit
Unit Factor
Units asked for
32
Practice Problems

• Work in groups of 3-4
• One student from each group puts solution in
  board and explains to class

33
Quiz 4

• See Chemistry Current News Slides
• Presentation to be given on Oct 6, 2003.

34
Density - Review

• Lab Experiment 2
• Physical property
• Defined as mass per unit volume
• Liquids and solids expressed in g/ml (g/ cm3)
• Gases expresses in grams per liter
• Density of water is 1.00 g/ml
• Floats in water density lt1.00 g/ml
• Sinks in water density gt1.00g/ml

35
Estimating Density(page 59 and 60 )
Water, chloroform and ethyl ether are poured into
a tall glass cylinder. Three known solids are
added. Identify the liquids.
Liquid 1
Solid 1 ice
Liquid 2 water
Solid 2 rubber
Liquid 3
Solid 3 aluminum
36
TemperatureFahrenheit, Celsius and Kelvin

• Measure of the average energy of individual
  particles in a system
• Warmer temps more molecules moving thus more
  energy
• Cooler temps slow moving molecules thus less
  energy
• Fahrenheit oF
• Celsius oC
• Kelvin K

37
Temperature

• oF
• Freezing point of water 32 oF
• Boiling point of water 212 oF
• oC
• Freezing point of water 0 oC
• Boiling point of water 100 oC
• K ( SI unit)
• Absolute zero 0 K
• Equal to -273.15oC

38
Temperature Conversions

• oF to oC
• ( oF - 32 oF ) x 100 oC / 180 oF oC
• oC to oF
• ( oC x 180 oF / 100 oC ) 32 oF
• Kelvin
• oC 273

39
Heat

• Heat measures the total energy
• Temperature measures the average energy
• Heat energy units calories or kilocalories
• A calorie (cal) is defined as the amount of heat
  needed to raise 1 g of water 1 oC
• Food Calorie equals 1 kcal 1000 cal
• SI unit joule (J)
• 1 cal 4.184 J

40
Specific Heat

• Amount of heat required to bring about a given
  change in temperature
• Observed amount
• Unique for each substance
• Specific heat of water is high
• Change in temperature is minimal as water gains
  or losses heat
• Surface of earth is covered in water so water
  helps to regulate the climates

41
Specific Heat

• Amount of heat required to raise the temperature
  of 1 g of substance 1 oC
• Units are cal/g oC

Water
Ice
Iron
Silver
1 g
1 g
1 g
1 g
1.0 oC
9.3 oC
17.7 oC
2.0 oC
42
Specific Heat

• gain or loss of heat divided by mass and
  temperature change specific heat
• How many calories are required to raise the
  temperature of a 3 inch iron nail weighing 7.05 g
  form room temperature to 100 oC?
• The specific heat of iron is 0.108 cal/g oC

43
Solution

• Specific Heat Heat/ (mass x D t)
• cal/g oC cal / g x oC
• 0.108 cal/g oC cal / 7.05 x (100-25oC)
• Solving for Heat ( energy required )
• Rearrange
• (0.108 cal/g oC) x 7.05 g x 75 oC
• 57 cal

44
Chapter 4 Matter and Energy

• Matter is any substance that has mass and
  occupies volume
• Physical State changes
• Melting solid into liquid
• Sublimation solid into gas
• Condensation gas into liquid
• Deposition gas to solid
• Freezes liquid to solid
• Vaporization liquid to gas

45
Increasing temperature
steam
ice
water
melting
vaporizing
freezing
condensing
Sublimation
Deposition
46
Chapter 4 Matter and Energy
47
Elements, Compounds and Mixtures

• Properties of matter may be consistent throughout
  or they may vary
• Melting point
• Gold (Au) 1064 oC
• Quartz 1000 1600 oC
• Gold is homogenous properties consistent
• Quartz is heterogeneous properties vary

48
Mixtures

• Heterogeneous
• Usually Solids
• Separated into pure substances by physical
  methods which take advantage of different
  physical properties
• Properties are not the same throughout the sample
• Homogeneous
• Gases or liquids
• Separated into pure substances by either chemical
  or physical methods which take advantage of
  different physical properties
• Properties a the same for any given sample, but
  can vary sample to sample

49
Mixtures

• Alloy
• Homogeneous mixture of two or more metals
• Gold ( Au)
• 10 K 14 K 18 K
• 42 75
• Substance
• Matter with definite composition and constant
  properties
• Compound or an element
• Compound
• Broken down into elements by chemical reactions
• Element
• Cannot be broken down further by chemical
  reactions

50
Matter
Mixtures
Substances
Physical Separate
Heterogeneous
Homogeneous
Compounds
Elements
51
Names and Symbols of the Elements

• 81 stable elements that occur in nature
• Only 10 account for 95 of the mass of the earths
  crust, water and atmosphere

All other elements 0.5
52
Names and Symbols

• Names are from various sources
• Hydrogen (hydro, Gr. water former)
• Carbon (carbo, Lt. coal)
• Calcium (calcis, Lt. lime)
• Chemical Symbols
• Dalton in 1803 proposed that elements are
  composed of indivisible spherical particles or
  atoms (atomos, Gr. indivisible)
• Suggested the use of circles with markings for
  symbols ( pg 83)
• Berzelius in 1813 proposed our current system of
  symbols using the first letter of the name and
  if the first letter is already in use two letters

53
Metals, Nonmetals and Semimetals

• Predict by position in Periodic Table
• Metals
• solid element
• Bright metallic luster
• Good conductor of heat and electricity
• High density
• High melting point
• Malleable ( thin sheets)
• Ductile ( fine wire)

54
Metals, Nonmetals and Semimetals

• Nonmetals
• Solid or gas element
• Dull appearance
• Low density
• Low melting point
• Poor conductor of heat and electricity
• Crush to a powder, if solid
• Semimetal
• metalloids
• midway between metal and nonmetal
• Semiconductor

55
Periodic Table of the Elements(page 86)

• Atomic number
• Number of protons
• Metals are placed on the left
• Nonmetals on the right
• Separated by semimetals starting at B
• Solids are to the left (most all elements)
• Gases to the right

56
Compounds Chemical Formulas

• 1799 Proust
• Law of Definite Composition
• Law of Constant Proportion
• Compounds always contain the same elements in a
  constant proportion by mass.

57
Chemical Formulas

• Most elements occur in nature as collection of
  individual atoms
• Diatomic molecules
• Oxygen (O2), Hydrogen (H2), Nitrogen (N2)
• Halogens Chlorine (Cl2),Bromine Br2, Iodine (I2)
• A chemical formula expresses the number and type
  of each atom in a compound
• The number of the each atom is indicated with a
  subscript. The number 1 in the subscript is
  implied and therefore is omitted.
• Parentheses are used to help clarify the
  structure of the compound and

58
Examples

• Water H2O
• 2 hydrogen atoms and 1 oxygen atom
• Calcium Chloride CaCl2
• 1 calcium atom and 2 chlorine atoms
• Propylene Glycol C3H8O2
• 3 carbon atoms, 8 hydrogen atoms, 2 oxygen atoms
• Lead acetate PbC2H3O2
• 1 lead atom, 2 carbon atoms, 2 oxygen atoms, 3
  hydrogen atoms
• 4-amino-2-hydroxytoluene C7H9NO
• 7 carbon atoms, 9 hydrogen, 1 nitrogen, 1 oxygen

59
Physical and Chemical Properties

• Substances ( Compounds or Elements)
• Physical and chemical properties are the
  consistent throughout
• No two substances have all the same physical and
  chemical properties
• 2 Na Cl2 2 NaCl
• metal yellow gas white solid

60
Physical and Chemical Properties

• Physical properties are measured and observed
  without changing the chemical composition of the
  substance
• Examples
• Appearance, color, melting point, density,
  solubility, boiling point, freezing point
• Chemical properties describes how a substances
  reacts with other substances (reaction chemistry)
  and always involve a chemical change

61
Periodic Table and Reaction Chemistry

• Elements with similar reaction chemistry are
  Grouped into families
• Columns in the Periodic table
• Group IA Alkali Metals
• Group IIA Alkaline Earth Metals
• Group VIIA Halogens
• Group VIIIA Noble Gases
• Group IIlB to IIB Transition Metals
• Lanthanides
• Actinides

62
Physical and Chemical Changes

• Physical Change chemical composition does not
  change, but the physical state does
• Example ?
• Chemical Change chemical composition changes
  and the physical state may or may not change (
  formation of a new substance)
• Example ?

63
Physical and Chemical Changes

• Chemical changes are often detected by
• Gas formation bubbles or odor
• Color change permanent
• Release of energy light or heat
• Precipitate formation solids

64
Conservation of Mass and Energy

• Matter is nether created or destroyed during a
  chemical reaction
• Energy can not be created or destroyed. It can
  however, be converted from one form to another
• Total mass and energy of the universe is constant

65
Potential vs. Kinetic Energy

• Key concepts for understanding chemical reactions
• Potential energy is stored energy of matter at
  rest
• Kinetic energy is energy that is a result of
  motion
• Potential energy can be (and is) converted into
  kinetic energy

66
Kinetic Energy as a function of Physical State
67
Chemical Reactions

• Potential Chemical Energy
• Kinetic Heat Energy
• Exothermic reactions
• Reactions that give off or produce heat
• Reactants have more potential energy than
  products
• Endothermic reactions
• Reactions that take in or absorb heat
• Reactants have less potential energy than
  products

68
Example

• Exothermic
• Reactants Products heat
• ( high P.E.) (low P.E. ) ( K.E.)
• Endothermic
• Reactants heat Products
• ( low P.E.) (K.E.) (high P.E. )

69
Forms of Energy

• Six Forms
• Light
• Heat
• Chemical
• Electrical
• Mechanical
• Nuclear

70
Energy Examples

• Radioactive Ur vaporizing water
• nuclear heat
• Steam driving a turbine
• heat mechanical
• Lead acid battery
• chemical electrical

71
Chapter 5 Models of the Atom

• Atom indivisible
• Dalton
• Proposed that all matter was composed as tiny
  particles
• Based on the Laws of Conservation of Mass and
  Definite Proportion
• Compounds are simply the combination of two or
  more atoms of different elements

72
Daltons Atomic Theory

• An element is composed of tiny, indivisible
  particles called atoms
• All atoms of an element are identical and have
  the same properties
• Atoms of different elements combine to form
  compounds
• Compounds contain atoms in small whole number
  ratios
• Atoms can combine in more than one ratio to form
  different compounds

73
Thomson Model (Plum Pudding Model)

• Cathode Ray experiment
• Glass tubes containing a low pressure amount of
  gas emitted light when electricity was applied to
  one end of the tube (Fluorescence light energy)
• Ray emanates form the negative cathode in the
  tube, the radiation is referred to as a cathode
  ray
• Placed tube in magnetic field, light or ray
  curved towards the positive
• Concluded that cathode rays were composed of tiny
  negatively charged particles ( electrons, e- )
• Further experiments showed that certain rays
  contained small particles that had an equal but
  opposite in sign charge to electrons protons
  (p)

74
Relative Charges and Masses
75
Thomson Model of the Atom

Homogeneous sphere plum pudding

-

-
-
-

-

-
-
76
Rutherford Model of the Atom

• Alpha ray
• particles identical to He w/o electrons ( He 2 )
  
• most passed through thin Au foil
• suggested that most of the atoms were empty
  space, with electrons moving about a center (
  nucleus)
• Nucleus contains protons and is tiny and dense
• Beta ray -
• Gamma ray not affected by magnetic fields

77
Rutherford Model of the Atom
Neutrons heaviness of the atom

- -

n 0
1 X 10 -8 cm
1 x 10 -13 cm
- - - -
78
Subatomic Particles
79
Atomic Notation
Mass Number (protons and neutrons)
Sy
A
Atomic Number ( protons)
symbol
Z
80
Isotopes

• Same element with different amount of neutrons
• Number of Protons and Electrons are the same
• Mass of the element is different, so the mass
  number is different
• For example, Hydrogen and deuterium

1 1
H
H
2 1
81
Isotopes

• Naming
• Some are common (hydrogen, deuterium, tritium)
• Name of the element followed by the mass number
• For example,
• carbon -14
• oxygen -18
• cobalt-60
• Hint should be able to determine number of
  protons electrons, neutrons atomic notation,
  isotope names

82
Atomic Mass

• Atoms are to small to weigh on balance
• Masses are relative to each other
• Specifically, mass is relative to carbon-12
• Carbon-12 has 12 amu (atomic mass unit)
• So, an amu 1/12 the mass of carbon
• Weighted average

83
Atomic Mass

• Weighted average of all the naturally occurring
  isotopes
• Given the natural abundance of the atom, the amu
  can be calculated for a given atom
• Since its a calculated value no atom will
  actually weigh this number

84
Periodic Table

• Does not tell about the number of naturally
  occurring isotopes
• Atomic Number Protons
• Atomic Mass weighted average protons and
  neutrons (whole numbers)
• Mass number is in parentheses, then the element
  is unstable. Mass is given for the best known
  isotope.
• Hint Should be able to tell form periodic table
  which elements are stable and which are
  not.

85
Wave Nature of Light

• Wavelength the distance the light wave travels
  to complete one cycle
• Frequency - the number of wave cycles completed
  in 1 s
• Velocity of light is constant 3.00 x 103 m/s
• If wavelength decreases, frequency increases
• Low frequency low light energy long
  wavelengths
• High frequency-high energy light short
  wavelengths ( page 124)

86
Light- A Continuous Spectrum

• White light passes through a prism it separates
  into all the colors
• ROY G BIV
• Light radiant energy that is visible
• Visible spectrum 400-700 nm
• Radiant energy spectrum a continuous spectrum
  of visible and invisible light that ranges form
  short to long wave lengths

87
Radiant Energy Spectrum
Wave Length Increases
Cosmic Rays