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Atoms, Ions, and the Periodic Table

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Title: Atoms, Ions, and the Periodic Table


1
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?

2
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.

4
  • 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.

5
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

6
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.

7
Rutherford then discovered the proton, and next,
working with a colleague, James Chadwick
(1891-1974), he discovered the neutron as well.
8
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

9
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.

10
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.
11
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

12
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

13
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

14
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.

15
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
16
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!

17
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

18
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)

19
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.

20
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.

21
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.

22
The electron cloud is like an apartment building.
23
  • The energy levels are like floors in the
    apartment building.

24
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.
25
  • The electrons are like people living in the
    rooms.

26
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)

27
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28
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

29
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

30
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!
31
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.

32
  • 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!

33
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
34
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
35
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 .

36
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
37
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

38
  • 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.

39
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





40
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.

44
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)

45
Radiation
  • Radiation-its the transfer of energy
  • Radioactivity-The spontaneous emission of
    radiation by an unstable nucleus.

46
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

47
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
48
B. Nuclear Decay
  • Alpha Emission

Numbers must balance!!
49
B. Nuclear Decay
  • Beta Emission

50
B. Nuclear Decay
  • Gamma Emission
  • Usually follows other types of decay.
  • Transmutation
  • One element becomes another.

51
B. Nuclear Decay
  • Why nuclides decay
  • need stable ratio of neutrons to protons

DECAY SERIES TRANSPARENCY
52
C. Half-life
  • Half-life (t½)
  • Time required for half the atoms of a radioactive
    nuclide to decay.
  • Shorter half-life less stable.

53
F ission
  • splitting a nucleus into two or more smaller
    nuclei
  • 1 g of 235U 3 tons of coal

54
F ission
  • chain reaction - self-propagating reaction
  • critical mass - the minimum
  • amount of
  • fissionable
  • material needed
  • to sustain a chain reaction

55
Fission
  • Uranium-235 is the only naturally occurring
    element that undergoes fission.

Uranium - 235
56
Fission
  • Why does fission produce so much energy?
  • Small quantities of mass are converted into
    appreciable quantities of energy.

E mc2
57
Fission
Energy
1 gram matter
700,000 Gallons of high octane gasoline
58
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

59
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

60
Nuclear Power
  • Fission Reactors

61
Nuclear Power
  • Fission Reactors

62
Nuclear Power
  • Fusion Reactors (not yet sustainable)

63
Nuclear Power
  • Fusion Reactors (not yet sustainable)

National Spherical Torus Experiment
Tokamak Fusion Test Reactor Princeton University
64
Synthetic Elements
  • Transuranium Elements
  • elements with atomic s above 92
  • synthetically produced in nuclear reactors and
    accelerators
  • most decay very rapidly
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