Download 8/29/12 The Science of Matter Objective: Understand the

8/29/12
The Science of Matter
Objective:
 Understand the composition and their impact on the
properties of matter.
 Be able to classify chemical vs. physical changes
(How things change)
Chemistry is the study of the composition of matter and the
changes that matter undergoes.
List 5 chemicals
Exit Ticket: What did you learn today?
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8/30/12
Water, bleach, toothpaste, hairspray and deodorant are
examples of chemicals that we use everyday
Remember:
 Chemistry is the science that investigates the structure
and properties of matter.
 Matter is anything that takes up space and has mass
 Mass is the measure of the amount of matter that an
object contains
(Heat, light and radiowaves are not matter.)
Properties describe the characteristics and behavior of
matter including the change it undergoes.
Behavior of matter is determined by the elements it
contains and the arrangement of those elements
Examples: (contain different elements and properties)
Salt (Sodium and Chlorine)
Water (Hydrogen and
Oxygen)
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9/4/12
Physical change vs. Chemical change
Physical change
Chemical change
Does not change the substance
A new substance is formed and
energy is either given off or
absorbed
Can be reversed
Cannot be reversed. The
substance cannot be turned back
into its initial state.
Density:
Density is an important property of matter
It is used to identify substances
It is used to separate mixtures
Density is the mass of specific unit of volume
Density = mass/volume (D= M/V)
Volume =mass/density (V= M/D)
Mass = density x volume (M=DV)
In solids density is usually expressed grams per cubic centimeter (g/cm³) and kilograms
per cubic centimeter (kg/cm³)
In liquids density is usually expressed as grams per liter (g/L) or grams per cubic
centimeter (g/cm³)
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Example: If a piece of rock has a mass of 14g and a volume of 5cm³, what is
its density?
Steps:
 Identify and write the givens:
M=14g, V= 5cm³, D=?
 Write the formula:
D= M/V
 Substitute:
D= 14g/5cm³
 Solve: divide 14 by 5
D= 2.8 g/cm³
Homework : complete the density problems
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9/19/12
Classifying Matter
Substance is matter with constant composition

Element made of only one type of atom

Compound is 2 or more elements that are chemically combined
Mixture is matter with variable composition

Heterogeneous is a mixture made up of more than one phase

Homogeneous is also called solutions made up of only one phase
Chemical change and energy
All chemical changes involve some sort of energy
Many chemical changes (reactions) release energy

Exothermic- release energy- HOT

Endothermic - absorb energy-COLD
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9/24/12
Objective:
Students will be able to identify those scientists who were beneficial to the atomic
theory
Early model of the atom:
The atom is the smallest particle of an element
Democritus: (460BC-370BC), Greek Philosopher
He was the first to suggest the existence of atoms. . He belied that atoms are
indivisible (cannot be cut) and indestructible
His ideas was not based on experiments just philosophy
Lavoisier: (1700’s) French
“Conservation of Mass.” He changed chemistry to a quantitative science.
He measured the mass of a system before and after a reaction in a closed
system
Stoichiometry
He isolated and named hydrogen and oxygen
He discovered how respiration and combustion are related
6
His major experiment involved cinnabar: red mercury oxide
Proust: (1800’s)
“The Law of Definite Proportions”
Tested Dalton’s 3rd postulate
Copper carbonate from a variety of sources
A compound always contains the same proportion of elements in definite
proportions
Dalton’s Atomic Theory (1803), English, chemistry teacher

2000 years passed before more was known about the atom

Thought the atom looked like a marble/ball bearing (round, solid, sphere)

His findings were based on scientific experiments, not philosophy

Elements are composed of elements

Atoms of the same elements are identical and differ from the other
elements

Atoms chemically combine in whole number ratios

Atoms cannot be created or destroyed

Atoms are extremely small
A penny containing pure copper contains
2,400,000,000,000,000,000,000,000,000 atoms
These can be observed using a scanning tunneling (electron) microscope
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9/25/12
Dalton (1803)
Dalton is known as the founder of the atomic theory
He stated that the word atom is the basic unit of matter
Dalton also claimed that all atoms of a given element are identical
He discovered that the atoms of different elements have different properties and
masses
He found that combining atoms of different elements formed compounds
Dmitri Mendeleev (1869)
He arranged elements into 7 groups with similar properties. He discovered that the
properties of elements were periodic functions of their atomic weight. This became
known as the Periodic Law
8
Cathode Ray:

Vacuum tube, all gasses pumped out

Metal piece called electrodes sticking out each end

Become charged when attached to strong electrical current

Rays travel in the tube from negative electrode (cathode) to the positively
charged electrode (anode)

A magnet will deflect the cathode ray

Particles in the cathode ray are negatively charged.
JJ Thompson (1897)
While using the cathode ray tube, he discovered that the ray was deflected due to a
magnetic electrical field
From this discovery he concluded that atoms contain small negatively charged
particles called electrons.
Plum Pudding theory: Electrons are embedded within the structure of the atom just
like raisin bread
The mass of the rest of the atom (besides the electrons) was thought to be evenly
distributed and positively charged.
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9/28/12
1916 Robert Millikan
Determined the mass of the electron to be 1/1840 the the mass of the proton
or 9.11 X 1028. He used an oil drop apparatus.
1886 Eugen Goldstein
Discovered the proton, the subatomic particle with a positive charge and a
mass of 1
1932 James Chadwick
Discovered the neutron, which is a subatomic particle with no charge but a
mass equal to a proton (1).
Subatomic particles:
Particle
Charge
Mass
Location
Electron
-1
Nothing
Electron cloud
Proton
+1
1
Nucleus
Neutron
0 (no charge)
1
Nucleus
1911 Ernest Rutherford
 In 1911 he designed the Gold Foil Experiment
 He aimed a beam of alpha (α) particles at a thin piece of gold foil (only a
few atoms thick)
 Most (α) particles passed through the foil
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 A small amount of the (α) particles were deflected.
 To their surprise, some alpha particles bounced straight back
 1/8000 did not go through the foil
 He determined that the nucleus was a part of the atom
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10-1-12:
Power point due Oct. 12:
 Background of the science
 Describe the experiment
 How is it related to the atom?
 How does it effect future scientist ?
 75% slides
 25% oral presentation
Rutherford’s Atom:
Conclusions:
 There is a nucleus in the center of the atom where most of its mass is
 The nucleus is positive
 Atoms are mostly empty space
 Disproved JJ Thompson’s Plum Pudding Model
 If an atom is the size of a football field stadium, the nucleus is the size of
a marble.
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Niels Bohr 1885-1962
 Planetary model 1913
 Nucleus surrounded by orbiting electrons at different energy levels
 Electrons have definite orbits
 Utilized Planck’s Quantum energy theory
 Worked on the Manhattan Project (US Atomic bomb)
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10-2-12
Atomic Number: Element are different because they contain different number
of protons
Atomic # = the number of protons in the nucleus
Mass number: Number of protons and neutrons in the nucleus
Element
P+
N0
E-
Mass #
Oxygen
8
8
8
16
Arsenic
33
42
33
75
Phosphorous
15
16
15
31
Nuclide symbols contain the symbol of the element, the mass number and the
atomic number:
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10-4-12
Isotopes
Isotopes are atoms of the same element can have different number of
neutrons
Isotopes have different mass numbers
Isotopes change the number of neutrons and the mass number for an atom.
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10/11/12
Frederick Soddy proposed the idea of isotopes in 1912
Isotopes of the same element having different masses due to varying
numbers of neutrons
Elements occur in nature as mixtures of isotopes.
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10/22/12
Wed. Binder check
Objective: Students will be introduced to nuclear chemistry and
radioactivity
Nuclear chemistry:
Types of Radioactivity:
By the end of this section you will be able to:
 Observe nuclear changes and explain how they change an
element
 Express alpha and beta decay in nuclear equations
 Model the half-life of an isotope
 Explain how half-life is used to date materials
Vocabulary:
Radioactivity
Alpha particle
Beta particle
Alpha decay
Beta decay
Gamma decay
Half life
Radioactivity dating
Radioactive decay
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Radioactivity is the spontaneous emission of radiation by an unstable
atomic nucleus
Chemical reaction
Nuclear reaction
Occur when bonds are broken and Occurs when nuclei combine, split
formed
and emit radiation
Involve only valence electrons
Can involve protons, neutrons and
electrons
Atoms keep the same identity
Atoms of one element are often
converted into atoms of another
element
Associated with small changes in
energy
Associated with large changes in
energy
Temperature, pressure,
concentration and catalysts affect
reaction rates
Temperature, pressure,
concentration do not affect
reaction rates
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10-23-12
Nuclear Reactions

Remember that the number of protons determined the identity of an element

Changing the number of protons will change the element into another element
During nuclear reactions atoms of one element are changed into atoms of another element.
Nuclear Notation:
Different isotopes of atoms can be represented using nuclear notation:
 Nuclear reactions involve the protons and neutrons found in the
nucleus.
 During nuclear reactions a nucleus can gain or lose protons and
neutrons
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10-31-12
Radiation causes Radioactive Decay
 Radioactive decay is the release of radiation by radioactive isotopes
 Not all radioactive isotopes decay in the same way. Different types of
decay change the nucleus in different ways
 3 Types of decay
 Alpha
 Beta
 Gamma
Radioactive alpha decay

Alpha decay is the release of alpha particles (2 protons and 2 neutrons)

Alpha particles are Helium nuclei consisting of 2 protons and 2 neutrons

Alpha particles which are large in size, collide with objects around them. They
do not penetrate very deeply. They are stopped by a thin layer of material.

Alpha decay causes the decaying nucleus to lose 2 protons and 2 neutrons.
The mass number decreases by 4 (2 protons and 2 neutrons)
The atomic number decrease by 2
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11-1-12
Radioactive Beta Decay
Beta decay is the release of beta particles from a decaying nucleus

A beta particle is a high energy electron with a 1- charge

Beta particles are written as :

Beta particles pass more easily through matter than alpha particles and require
sheets of metal, blocks of wood or specialized clothing to be stopped

The electron released during beta decay is not one of the original electrons that
existed outside the nucleus.

The beta particle (electron) is produced by the change of a neutron into a proton
and an electron
Equation for radioactive beta decay: The parent nucleus turns into a daughter with an
atomic number 1 greater. The mass number stays the same
Beta emission:
 A neutron becomes a proton which stays in the nucleus and the electron
is ejected from the atom
Add a proton and lose an
electron
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11-2-12
Radioactive Gamma Decay
Gamma decay is the release of gamma rays from a nucleus
A gamma ray is a high energy form of electromagnetic radiation
without a change in mass or charge
Gamma rays have high penetrating ability and are very dangerous to
living cells
To stop gamma rays thick blocks of lead or concrete are needed
During gamma decay only energy is released
Gamma decay does not generally occur alone, it occurs with other
modes of decay (alpha or beta)
When gamma decay is expressed in an equation it is expressed as
The following equation shows both gamma and alpha decay:
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11/5/12
Half-Life
Radioactive Decay:
Radiation can be detected with Geiger counters and scintillation counters.
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Quarter 2
11-13-12
Do Now: Complete the following problems
Sodium 24 has a half-life of 15 hours. How much Na-24 will remain in an
18.0g sample after 60 hours?
After 42 days a 2.0g sample of P-32 contains only .025g of the isotope. What
is the half life of P-32?
Po-214 has a half life of 164 seconds. How many seconds would it take for
8.0g to decay to 0.25g?
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11-14-12
The Power of the Nucleus
Vocabulary:
 Nuclear Fission
 Chain reaction
 Nuclear reactor
 Nuclear Fusion
E=mc2:
E= energy
M= mass
C= speed of light (3.0x 108m/s)
Nuclear reactions involve enormous changes in energy. During a
nuclear reaction a small amount of mass can be converted into a
large amount of energy.
Nuclear fission is the process of splitting a nucleus into 2 or more
smaller fragments. This is accompanied by release of energy.
Protons and neutrons are in the nucleus and are split. Energy is
released.
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Nuclear fission using Uranium
The sums of the mass numbers on the left and right are equal
As World War II (1939-1945) started, scientists were trying to find a
way to sustain nuclear fission in a chain reaction.
Chain Reaction is a continuing series of reactions in which each
produces a product that can react again and again.
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11/15/12
In Fission of uranium, each neutron produced has the potential to cause the
fission of another atom of Uranium 235.
In order for a chain reaction to occur there must be enough of a sample of the
material for the neutrons to collide with other atoms
Critical mass: the point where the chain reaction becomes self-sustaining.
Supercritical mass: If the amount of fissionable material is much greater
than critical mass the chain reaction escalates out of control and an explosion
results.
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11/16/12
All of the energy is released at once. This is what happens when an atomic
bomb explodes
Nuclear Fission and Nuclear energy
In order for nuclear energy to be useful the reaction must be controlled so
that the energy can be released slowly.
Nuclear power plants generate electrical energy through the controlled fission
of Uranium.
This is done in a nuclear reactor. A nuclear reactor is a device that is used to
extract energy from radioactive fuel.
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11-19-12
Nuclear reactors and Pollution:
 Nuclear reactors do not produce CO2 and other pollutants
 They do not produce radioactive waste that is difficult to safely dispose of
 New technologies allow much of the waste to be decayed, reducing the
amount of hazardous waste produced
 There is some risk of the release of this nuclear waste into the environment
Problems with nuclear reactors:
 Nuclear energy costs more to produce than energy produced through the
burning of fossil fuel
 It is more expensive than using fossil fuels
Nuclear Fusion:
Nuclear fusion is the process of combining 2 or more nuclei to form a larger
nucleus
Nuclear fusion is the process that occurs in the sun and other stars to produce
energy
Nuclear fusion….. Hydrogen to Helium
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11/20/12
The fusion of hydrogen to produce helium produces 20 X more energy than the
fission of the same amount of uranium.
 It does not produce any radioactive waste
 Fusion reactions are easier to control than fission reaction
Problems with nuclear fusion:
 Difficulty initiating and containing a fusion reaction has prevented its use as
a practical energy source.
 Nuclear fusion reactions require a large amount of energy to start the fusion
reaction
 In order to initiate a fusion reaction on earth a temperature greater than 100
million Kelvins would be required.
 No material exists on earth that could contain the reaction.
 A great goal for the future
Fission
Fusion
Splitting of an atomic nuclei caused by Combining 2 or more nuclei
a neutron
Chain reaction
Multi-step process
Produces radioactive waste
Does not produce radioactive waste
Used in nuclear plants and bombs
Not able to be initiated and contained
because it requires a great deal of
energy to start
Can generate a great deal of energy
Occurs on the sun with a tremendous
amount of energy released
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12/3/12
Periodic Table Geography
What information does the periodic table tell me?
The horizontal rows of the periodic table are called periods
The vertical columns of the periodic table are called groups or families. The
elements in any group of the periodic table have similar physical and chemical
properties
There are 18 groups and 7 periods
Periodic Law:
When elements are arranged in order of increasing atomic number, there is a
periodic pattern in their physical and chemical properties
Elements are arranged based on similar properties
Elements are arranged based on increasing atomic number
Atomic number is the number of protons
Atom mass is the number of protons + number of neutrons
Most elements are metals
Metals:
 Good conductors of heat and electricity
 Malleable and able to be bent
 Ductile
 Loose electrons becoming positive ions
Metalloids:
 Elements that touch the staircase
 Metalloids have properties of both metals and nonmetals
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 More brittle than metals but less brittle than nonmetals
 Semiconductors of electricity (anything that’s computerized or uses radio
waves depends on semi conduction. Today, most semiconductor chips and
transistors are created with silicon (14) is the heart of any electron device).
 High tensile strength
 Loose and gain electrons to get to eight
Non-metals
Poor conductors of heat and electricity
Tend to be brittle
Many are gaseous at room temperature
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12-6-12
 Alkali metals: Group 1
 Alkaline Earth metals: Group 2
 Transition Metals: groups 3-12
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 Inner Transition Metals/ Lanthanide and Actinide (Fit in Groups 3-12)
 Halogens: Group 17
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12/10/12
 Noble Gases: Group 18
Practice:
What elements are in the halogen family?
How many elements are in the alkali group?
What is the first element in the 3rd period?
Name 3 noble gases.
Why are elements arranged in this way on the periodic table?
35
Mendeleev:
In 1896 Dimitri Mendeleev created the first accepted version of the
periodic table. He grouped elements according to their atomic mass
and as he did he found that the families had similar chemical
properties. Blank spaces were left open to add the new elements
predicted would occur.
Elements:
Science has come a long way since Aristotle’s theory of air, water, fire
and earth.
Scientist have identified 90 naturally occurring elements and created
about 28 others.
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1/2/13
Students will be able to identify parts of the periodic table
There are 7 periods and 18 groups in the periodic table
Valence Electrons:
 The number of valence electrons an atom has may also appear in a
square
 Valence electrons are the electrons in the outer energy level of an
atom
 These are the electrons that are transferred or shared when atoms
bond together.
 The elements that are further away from 8 the more dangerous the
element
Hydrogen:
 The hydrogen square sits on top of the family A1, but it is not a member
of the family. Hydrogen is in a class of its own
 It’s a gas at room temperature
 It has one proton and one electron in its one and only energy level
 Hydrogen only needs 2 electrons to fill up its valence shell
37
Alkali Metals:
Found in the first column of the periodic table
Atoms of the alkali metals have a single electron in their outer most level, in
other words, 1 valence electron
They are shiny, and have consistency of class, and are easily cut with a knife
38
Halogens: (Group 17)
 The elements in this family are Fl, Cl, Br, I and Astatine
 The halogens have 7 valence electrons which explain why they are the
most active non-metals. They are never found free in nature
 Halogen atoms only need to gain 1 electron to fill their outermost
energy level
 They react with alkali metals to form salts
Noble Gases : (Group 18)
 Noble gases are colorless gases that are extremely unreactive
 One important property is their inactivity
 Noble gases are called inert because they do not combine with other
elements
 Noble gases are: He, Ne, Ar, Kr, Xe, Ra
 All noble gases are found in small amounts in the earth’s atmosphere
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1/7/13
Metallic Trends:
Fr (most metallic), He (lest metallic)
Atomic Radii Trend: Radius is the distance from the center of the nucleus to
the edge of the electron cloud. Atomic radii are usually measured in
picometers (pm) or angstrums (A)
A= 1x 10-10
40
Ionic Radii Trend: For the metals the nuclear charge is greater than the
number of electrons pulling them in tighter making the radius smaller at the
nonmetals the radius gets larger because the ion gained electrons.
Ionization Energy Trend: The larger the atom is the easier its electrons are to
remove. Ionization energy and atomic radius are inversely proportional.
Ionization energy is always endothermic, that is energy is added to the atom
to remove the electron.
Electron Affinity Trend: A measure of the energy change when an electron is
added to a neutral atom to form a negative ion.
41
Electronegativity Trend: A measure of the tendency of an atom to attract a
bonding pair of electrons
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1/14/13
Electron Configuration: Filling of atomic orbital and energy levels
 Electron configuration is a form of notation which shows how the electrons
are distributed among various atomic orbital and energy levels
 The format consists of a series of numbers, letters and superscripts as shown
below
Helium: He: 1S2
Information provided by electron configuration:
 The large number 1 stands for the energy level (refers to the principle
quantum number). It tells us that the electrons of Helium occupy the first
energy level of an atom
 The letter “S” tells us that the electrons of the Helium element occupy an S
orbital or spherical orbital
 The exponent 2 in the example refers to the total number of electrons in that
orbital or sub shell
 There are 2 electrons in the spherical orbital at the first energy level
Information needed to understand electron configuration:
The number of sublevels that an energy level can contain is equal to the principle
quantum number of that level (n)
For example, the second energy level has two sublevels
The third energy level has 3 sublevels
43
Names of sublevels:
1st= S
2nd= p
3rd= d
4th= f
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1/15/13
An orbital is a space that can be occupied by up to 2 electrons
Each type of sublevel holds a different number of orbitals and therefore a different
number of electrons
 “s” sublevels have 1 orbital and can hold up to 2 electrons
 “p” sublevels have 3 orbitals each of which can hold 2 electrons for a total of
6 electrons
 “d” sublevels have 5 orbitals for a possible total of 10 electrons
 “f “ sublevel with 7 orbitals can hold up to 14 electrons
Sublevel
S
P
D
F
Electron capacity of Sublevels
# of orbital x 2 electrons Maximum # of electrons
1 x 2
2
3 x 2
6
5 x 2
10
7 x 2
14
Total number of orbital and electrons per energy level
An easy way to calculate the number of orbitals found in an energy level is to use
the formula n2, (n= energy level)
For example, the third energy level (n=3) has a total of 32 or 9 orbitals.
This makes sense because we know that the third energy level would have 3
sublevels.
An “s” sublevel with 1 orbital, a “p” sublevel with 3 orbitals and a “d” sublevel
with 5 orbitals. 1 + 3+ 5 = 9.so the formula n2 works
45
1/18/13
Different Shapes of Orbitals:
Order of filling the sublevels with electrons:
 The energy sublevels are filled in a specific order
 You must follow the rules to understand when electrons will fill each level
 Once you know this you can write them out in electron configuration
46
3rd Quarter
1/28/13
Electrons In Atoms
2.2 Configurations, Lewis Dot diagram and the electromagnetic
spectrum
Electrons in motion
 Electrons are negative and the nucleus is positive
Why aren’t electrons pulled into the nucleus and held there?
The planetary model
Neils Bohr (1885-1962) worked with Rutherford and they suggested that
electrons have enough energy to keep them in constant motion
Bohr’s explanation:
 Electrons have energy of motion that enables them to overcome the
nucleus’ attractive forces. This is called the planetary model
 Experiments show that electrons occupy orbits with defined amounts
of energy, not random amounts, like satellites in orbit
47
The electromagnetic spectrum
 High voltage electricity can increase the energy of an electron
 Electromagnetic radiation or radiant energy can also supply energy to
electrons
 Electromagnetic radiation travels in waves with both electric and
magnetic properties. (Ex. Radiant energy from the sun)
48
2/1/13
Frequency: number of waves vibrations per second/ Unit of measure is hertz (Hz)
Wavelength: The distance between corresponding points on two consecutive
waves
Low frequency- Longer wavelength
High frequency- shorter wavelength
Waves and energy
Waves transfer energy from one place to another
Electromagnetic waves have similar characteristics
However they can travel through empty space:
 Radio waves
 Microwave radiation
 Visible light
All of these forms of radiant energy are part of the electromagnetic spectrum
Higher frequency most difficult:
 UV light
 X-rays
 Gamma rays
Visible light- red, orange, yellow, green, blue, indigo, violet
49
Radio waves has longest wavelength
10-2 microwaves
Gamma rays have highest energy
Emissions Spectrum:
 The spectrum of light released from excited atoms of an element
 It is like a fingerprint
 Evidence of energy levels
 Electrons absorb only enough energy needed to move to a specific higher
energy level
 When falling back they give off only the amount of energy to drop them
back to that lower energy: giving off a specific color of light.
50
2/4/13
Electrons and Energy Levels
 Like your feet on a ladder, electrons cannot be between energy
levels
 They must absorb the amount of energy required to land on an
energy level
 Energy levels are not evenly spaced like rungs on a ladder
 As n increases the energy levels get closer together
Electron Cloud
Continuing research has shown that energy levels are not neat. They are
regions of space around the nucleus in which electrons are mostly to be
found. The space around the nucleus where the electrons are most likely
to be found is called the electron cloud.
51
Valence Electrons: the electrons in the outer most or highest energy
level
 Many physical and chemical properties of an element are directly
related to the number and arrangement of valence electrons.
 Valence electrons are the ones that interact during a chemical
reaction
Lewis Dot Diagrams
 Because valence electrons are very important to the behavior of
atoms we represent them in symbols
 Lewis dot diagrams represent valence electrons as dots around
the chemical symbol of the element
 Each dot represents 1 valence electron
Writing Lewis Dot diagram
He: 2 valence electrons
52