Atoms, Ions, and the Periodic Table

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Atoms, Ions, and the Periodic Table

• What is an atom?
• It is smallest particle of an element that
  retains the elements properties.
• But how did we come to know all the information
  we have about these tiny particle?

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Democritus (460-370 BC)

• Matter is made of tiny, solid, indivisible
  particles which he called atoms (from atomos, the
  Greek word for indivisible).
• Different kinds of atoms have different sizes and
  shapes.
• Different properties of matter are due to the
  differences in size, shape, and movement of
  atoms.
• Democritus ideas, though correct, were widely
  rejected by his peers, most notably Aristotle
  (384-322 BC). Aristotle was a very influential
  Greek philosopher who had a different view of
  matter. He believed that everything was composed
  of the four elements earth, air, fire, and water.
  Because at that time in history, Democritus
  ideas about the atom could not be tested
  experimentally, the opinions of well-known
  Aristotle won out. Democritus ideas were not
  revived until John Dalton developed his atomic
  theory in the 19th century!

3
John Dalton (1766-1844)

• All matter is composed of extremely small
  particles called atoms.
• All atoms of one element are identical.
• Atoms of a given element are different from those
  of any other element.
• Atoms of one element combine with atoms of
  another element to form compounds.
• Atoms are indivisible. In addition, they cannot
  be created or destroyed, just rearranged.

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• Daltons theory was of critical importance. He
  was able to support his ideas through
  experimentation, and his work revolutionized
  scientists concept of matter and its smallest
  building block, the atom.
• Daltons theory has two flaws
• In point 2, this is not completely true.
  Isotopes of a given element are not totally
  identical they differ in the number of neutrons.
  Scientists did not at this time know about
  isotopes.
• In point 5, atoms are not indivisible. Atoms
  are made of even smaller particles (protons,
  neutrons, electrons). Atoms can be broken down,
  but only in a nuclear reaction, which Dalton was
  unfamiliar with.

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Discovery of the Electron JJ Thomson (1856-1940)

• Discovered the electron, and determined that it
  had a negative charge, by experimentation with
  cathode ray tubes. A cathode ray tube is a glass
  tube in which electrons flow due to opposing
  charges at each end. Televisions and computer
  monitors contain cathode ray tubes.
• Thomson developed a model of the atom called the
  plum pudding model. It showed evenly distributed
  negative electrons in a uniform
• positive cage.
• Diagram of plum pudding model

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Discovery of the NucleusErnest Rutherford
(1871-1937)

• Discovered the nucleus of the atom in his famous
  Gold Foil Experiment.
• Alpha particles (helium nuclei) produced from the
  radioactive decay of polonium streamed toward a
  sheet of gold foil. To Rutherfords great
  surprise, some of the alpha particles bounced off
  of the gold foil. This meant that they were
  hitting a dense, relatively large object, which
  Rutherford called the nucleus.

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Rutherford then discovered the proton, and next,
working with a colleague, James Chadwick
(1891-1974), he discovered the neutron as well.
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Models of the Atom - Niehls Bohr

• Developed the Bohr model of the atom (1913) in
  which electrons are restricted to specific
  energies and follow paths called orbits a fixed
  distance from the nucleus. This is similar to
  the way the planets orbit the sun. However,
  electrons do not have neat orbits like the
  planets.
• Diagram of Bohr model

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Quantum Mechanical Model

• This is the current model of the atom. We now
  know that electrons exist in regions of space
  around the nucleus, but their paths cannot be
  predicted. The electrons motion is random and
  we can only talk about the probability of an
  electron being in a certain region.

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Sub-Atomic ParticlesEach atom contains different
numbers of each of the three SUBatomic particles
Particle Symbol Charge Molar Mass Where found
Proton p 1 1.007 825 Nucleus
Neutron n0 0 1.008 665 Nucleus
Electron e- -1 0.000 549 Electron Cloud
A neutron walked into a bar and asked how much
for a drink. The bartender replied, For you, no
charge.
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Atomic Number

• The periodic table is organized in order of
  increasing atomic number.
• The atomic number is the whole number that is
  unique for each element on the periodic table.
  The atomic number defines the element. For
  example, if the atomic number is 6, the element
  is carbon. If the atomic number is not 6, the
  element is not carbon.
• The atomic number represents
• the number of protons in one atom of that element
• the number of electrons in one atom of that
  element (with an ion, the electrons will be
  different)
• Therefore, protons electrons in a neutral
  atom

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Atomic Mass

• mass of an element measured in amu (atomic mass
  units)
• all compared to C-12 (the mass of carbon 12,
  which has a mass of exactly 12 amu
• listed on the periodic table
• Mass number of protons of neutrons

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Isotopes

• Isotopes are atoms of an element with the same
  number of protons but different numbers of
  neutrons.
• Most elements on the periodic table have more
  than one naturally occurring isotope.
• There are a couple of ways to represent the
  different isotopes. One way is to put the mass
  after the name or symbol Carbon-12 or C-12
• Another way is to write the symbol with both the
  mass number and atomic number represented in
  front of the symbol

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Determining Average Atomic Mass

• The atomic mass on the periodic table is
  determined using a weighted average of all the
  isotopes of that atom.
• In order to determine the average atomic mass,
  you convert the percent abundance to a decimal
  and multiply it by the mass of that isotope. The
  values for all the isotopes are added to together
  to get the average atomic mass.

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Example of Average atomic mass calculation

• Given
• 12C 98.89 at 12 amu
• 13C 1.11 at 13.0034 amu
• Calculation
• (98.89)(12 amu) (1.11)(13.0034 amu)

(0.9889)(12 amu) (0.011)(13.0034 amu)
12.01 amu
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Now you try one

• Neon has 3 isotopes  Neon-20 has a mass of
  19.992 amu and an abundance of 90.51.  Neon-21
  has a mass of 20.994 amu and an abundance of
  0.27.  Neon-22 has a mass of 21.991 amu and an
  abundance of 9.22.  What is the average atomic
  mass of neon?
• The answer is
• (0.9051)(19.992 amu) (0.0027)(20.994 amu)
  (0.0922)(21.991 amu)
• 20.179 amu
• Now compare this mass for Neon to the mass on the
  periodic table!

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Electromagnetic Radiation

• Electromagnetic radiation is a form of energy
  that travels through space in a wave-like
  pattern. eg. Visible light
• It travels in photons, which are tiny particles
  of energy that travel in a wave like pattern.
  Although we call them particles, they have no
  mass. Each photon carries one quantum of energy.
• These photons of energy travel at the speed of
  light (c) 3.00 x 108 m/s in a vacuum

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What is a wave and how do we measure it?

• Frequency (?) number of waves that passes a
  given point per second (measured in Hz)
• Wavelength (?) shortest distance between two
  equivalent points on a wave (measured in m, cm,
  nm)

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Electromagnetic spectrum (EM)

• The electromagnetic spectrum shows all
  wavelengths of electromagnetic radiation the
  differences in wavelength, energy and frequency
  differentiates the different types of radiation.
  
• Note that as the wavelength increases, the energy
  and the frequency decrease.

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Ground state vs. Excited state

• Electrons generally exist in the lowest energy
  state they can. We call this the ground state.
• However, if energy is applied to the electrons,
  they can be excited to a higher energy and we
  call this an excited state.
• The excited state electron doesnt
• stay excited. It will fall back to
• the ground state quickly. When
• the electron returns to the ground
• state, energy is released in the
• form of light. One example of this
• is lasers.

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Electrons in Atoms

• We are most concerned with electrons because
  electrons are the part of the atom involved in
  chemical reactions.
• Electrons are found outside the nucleus, in a
  region of space called the electron cloud.
• Electrons are organized in energy levels of
  positive integer value (n 1, 2, 3,...).
• Within each energy level are energy sublevels,
  designated by a letter s, p, d, or f.
• Each sublevel corresponds to a certain electron
  cloud shape, called an atomic orbital.

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The electron cloud is like an apartment building.
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• The energy levels are like floors in the
  apartment building.

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The sublevels are like apartments on a floor of
the building. Just like there are different
sizes of sublevels, there are different sizes of
apartments 1 bedroom, 2 bedroom, etc.
The orbitals are like rooms within an apartment.
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• The electrons are like people living in the
  rooms.

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What do these orbitals look like?

• The s, p, d and f orbitals look different and
  increase in complexity (f-orbitals not shown
  they are very complex)

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Number of electrons in each sublevel depends on
number of orbitals!

• Each orbital can hold a maximum of 2 electrons.
• An s sublevel contains 1 s orbital. How many
  total electrons can fit in an s sublevel?
• 2
• A p sublevel contains 3 p orbitals. How many
  total electrons can fit in a p sublevel?
• 6
• A d sublevel contains 5 d orbitals. How many
  total electrons can fit in a d sublevel?
• 10
• An f sublevel contains 7 f orbitals. How many
  total electrons can fit in an f sublevel?
• 14

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The Aufbau Principle

• Three rules govern the filling of atomic
  orbitals. The first is
• The Aufbau Principle Electrons enter orbitals
  of lowest energy first. The Aufbau order lists
  the orbitals from lowest to highest energy
  (Aufbau is from the German verb aufbauen to
  build up)
• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
• 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10

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The Pauli Exclusion Principle

• An atomic orbital may hold at most 2 electrons,
  and they must have opposite spins (called paired
  spins).
• When we draw electrons to show these opposite
  spin pairs, we represent them with arrows drawn
  in opposite directions.

Write this down in your notes if you havent!
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Hunds Rule
Write this down in your notes if you havent!

• When electrons occupy orbitals of equal energy
  (such as three p orbitals), one electron enters
  each orbital until all the orbitals contain one
  electron with spins parallel (arrows pointing in
  the same direction). Second electrons then add
  to each orbital so that their spins are paired
  (opposite) with the first electron in the
  orbital.

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• An electron configuration uses the Aufbau order
  to show how electrons are distributed within the
  atomic orbitals.
• How to read a segment of an electron
  configuration
• Example 3p6
• 3 energy level
• p sublevel
• 6 of electrons
• Now, lets look at how to put these together for
  a specific element!

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Electron Configurations

• This is one way to represent the electrons of an
  atom. We will try a few together

Element Total of electrons Electron Configuration
carbon
fluorine
magnesium
argon
6
1s2 2s2 2p2
9
1s2 2s2 2p5
1s2 2s2 2p6 3s2
12
1s2 2s2 2p6 3s2 3p6
18
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Orbital Diagrams

• Orbital diagrams show with arrow notation how the
  electrons are arranged in atomic orbitals for a
  given element.

Element Total of electrons Orbital Diagram
carbon
fluorine
magnesium
argon
?? ?? ? ? . 1s 2s
2p
6
?? ?? ?? ?? ? . 1s 2s
2p
9
?? ?? ?? ?? ?? ??. 1s
2s 2p 3s
12
?? ?? ?? ?? ?? ?? ?? ??
?? 1s 2s 2p 3s
3p
18
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Valence electrons

• Electrons in the outer energy level of an atom.
  They are like the front lines of an army, because
  they are the ones involved in chemical reactions
  (valence electrons get shared or transferred
  during reactions).
• The number of valence electrons that an atom has
  is directly responsible for the atoms chemical
  behavior and reactivity.
• We can represent the number of valence electrons
  pictorially by drawing the electrons around the
  symbol in a dot diagram. The electrons are
  drawn in on each side of the symbol and are not
  paired up until they need to be.
• Eg. . Be .

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Element Electron Configuration Valence Electrons Electron Dot Structure
Li
Be
B
C
N
O
F
Ne
Li.
1
1s2 2s1
. Be .
1s2 2s2
2
. B . ?
1s2 2s2 2p1
3
. . C . ?
1s2 2s2 2p2
4
. . N ?
5
1s2 2s2 2p3
. O ?
1s2 2s2 2p4
6
.. F ?
7
1s2 2s2 2p5
.. Ne ?? ??
1s2 2s2 2p2
8
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The Periodic Table

• The rows on the periodic table are called periods
• The columns on the periodic table are called
  groups or families
• Elements within a group or a family have similar
  reactivity. What do you know about all elements
  in a period that could explain this?
• They have the same number of valence electrons

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• Since many of the families on the periodic table
  have such similar properties, they some have
  specific names that you need to know. Get out
  your periodic table and label each section as we
  look at them together.

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Alkali Metals are group 1 and are the most
reactive metals. They form 1 ions by losing
their highest energy s1 electron. 1 valence
electron.
Alkaline Earth Metals are in group 2. the form
2 ions by losing both of the electrons in the
highest energy s orbital. 2 valence electrons.
Halogens are in group 17 and they are the most
reactive nonmetals. The form -1 ions by gaining
1 electron to fill the highest energy p orbital.
They have 7 valence electrons.
The transition metals include groups 3 through 12
and these metals all lose electrons to form
compounds
Noble Gases are in group 18. They do not form
ions because they have a full outer shell of
electrons and do not need any more electrons.
They do not form compounds.8 valence electrons





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Electromagnetic Radiation

• Electromagnetic radiation is a form of energy
  that travels through space in a wave-like
  pattern. eg. Visible light
• It travels in photons, which are tiny particles
  of energy that travel in a wave like pattern.
  Although we call them particles, they have no
  mass. Each photon carries one quantum of energy.
• These photons of energy travel at the speed of
  light (c) 3.00 x 108 m/s in a vacuum

41
What is a wave and how do we measure it?

• Frequency (?) number of waves that passes a
  given point per second (measured in Hz)
• Wavelength (?) shortest distance between two
  equivalent points on a wave (measured in m, cm,
  nm)

42
Electromagnetic spectrum (EM)

• The electromagnetic spectrum shows all
  wavelengths of electromagnetic radiation the
  differences in wavelength, energy and frequency
  differentiates the different types of radiation.
  
• Note that as the wavelength increases, the energy
  and the frequency decrease.

43
Ground state vs. Excited state

• Electrons generally exist in the lowest energy
  state they can. We call this the ground state.
• However, if energy is applied to the electrons,
  they can be excited to a higher energy and we
  call this an excited state.
• The excited state electron doesnt
• stay excited. It will fall back to
• the ground state quickly. When
• the electron returns to the ground
• state, energy is released in the
• form of light. One example of this
• is lasers.

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Nuclear Forces

• The force that holds the protons together within
  the nucleus even though there are repulsive
  forces that would otherwise push the positive
  protons away from each other. (also known as
  strong force)

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Radiation

• Radiation-its the transfer of energy
• Radioactivity-The spontaneous emission of
  radiation by an unstable nucleus.

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Good vs. Bad

• Ionizing
• Has enough energy to kick off an ion.
• Very high energy
• Non ionizing
• Does not have enough energy to kick off an ion
• Low energy

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A. Types of Radiation

• Alpha particle (?)
• helium nucleus

paper
2

• Beta particle (?-)
• electron

1-
cardboard

• Positron (?)
• ly charged e-

1
concrete

• Gamma (?)
• high-energy photon

thick lead
0
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B. Nuclear Decay

• Alpha Emission

Numbers must balance!!
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B. Nuclear Decay

• Beta Emission

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B. Nuclear Decay

• Gamma Emission
• Usually follows other types of decay.
• Transmutation
• One element becomes another.

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B. Nuclear Decay

• Why nuclides decay
• need stable ratio of neutrons to protons

DECAY SERIES TRANSPARENCY
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C. Half-life

• Half-life (t½)
• Time required for half the atoms of a radioactive
  nuclide to decay.
• Shorter half-life less stable.

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F ission

• splitting a nucleus into two or more smaller
  nuclei
• 1 g of 235U 3 tons of coal

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F ission

• chain reaction - self-propagating reaction
• critical mass - the minimum
• amount of
• fissionable
• material needed
• to sustain a chain reaction

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Fission

• Uranium-235 is the only naturally occurring
  element that undergoes fission.

Uranium - 235
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Fission

• Why does fission produce so much energy?
• Small quantities of mass are converted into
  appreciable quantities of energy.

E mc2
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Fission
Energy
1 gram matter
700,000 Gallons of high octane gasoline
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Fusion

• combining of two nuclei to form one nucleus of
  larger mass
• thermonuclear reaction requires temp of
  40,000,000 K to sustain
• 1 g of fusion fuel 20 tons of coal
• occurs naturally in stars

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Fission vs. Fusion
FISSION
FUSION

• 235U is limited
• danger of meltdown
• toxic waste
• thermal pollution

• fuel is abundant
• no danger of meltdown
• no toxic waste
• not yet sustainable

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Nuclear Power

• Fission Reactors

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Nuclear Power

• Fission Reactors

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Nuclear Power

• Fusion Reactors (not yet sustainable)

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Nuclear Power

• Fusion Reactors (not yet sustainable)

National Spherical Torus Experiment
Tokamak Fusion Test Reactor Princeton University
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Synthetic Elements

• Transuranium Elements
• elements with atomic s above 92
• synthetically produced in nuclear reactors and
  accelerators
• most decay very rapidly