Chapt. 11 Atomic Structure

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Chapt. 11 Atomic Structure
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From macroscopic to microscopic

• http//micro.magnet.fsu.edu/primer/java/scienceopt
  icsu/powersof10/

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History

• Greek Philosopher Democritus (460-370 B.C.)
• all matter composed of small atoms
• atomos indivisible

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• 384-322 BC
• Aristotle and Plato favored the earth, fire, air
  and water approach to the nature of matter

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400 to 0, 0 to 1600 400 1600 2000 years
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Renaissance
Medieval

• http//www.youtube.com/watch?vwrD49Jci6h8feature
  related

• http//www.youtube.com/watch?vrFI3UkPTpMsfeature
  related

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Baroque

• http//www.youtube.com/watch?v1ZhHjZLgWOs

• Boyle
• 1600s

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Avogadro who, in 1811,1 hypothesized that two
given samples of an ideal gas, at the same
temperature, pressure and volume, contain the
same number of molecules
Dalton 1803
Bernoulli in 1738
Boltzmann 1898

• Boyle published it in 1662

• about 1787,Charles had found that oxygen,
  nitrogen, hydrogen, carbon dioxide, and air
  expand to the same extent over the same 80 degree
  interval.

1687 Newton
1845, Watterson
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• We might as well attempt to introduce a new
  planet into the solar system, or to annihilate
  one already in existence, as to create or destroy
  a particle of hydrogen.
• John Dalton, A New System of Chemical Philosophy,
  1808)

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Dalton's Postulates

• 1. All matter consists of tiny
• particles.

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Dalton's Postulates

• 2. Elements are characterized by the mass of
  their atoms. All atoms of the same element have
  identical weights, Dalton asserted. Atoms of
  different elements have different weights.

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Dalton's Postulates

• 3. Atoms are indestructible
• and unchangeable.

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Daltons Postulates

• 4. When elements react, their atoms combine in
  simple, whole-number ratios.

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John Daltons Atomic Theory

• Almost right. A good start.

very small
Structure of the atom after Dalton (ca. 1810)
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our future discoveries must be looked for in the
6th decimal place, 1894, at the dedication of
Ryerson Physics Laboratory, Chicago

• Heading into the 20th century there was a feeling
  by many in the chemistry and physics communities
  that our scientific knowledge was nearly
  complete. It was universally accepted that atoms
  were the most basic constituent of matter and
  that the behavior of all matter could be
  explained through Newtonian mechanics.
• BUTseveral discoveries and observations
  contradicted these theories

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• I. Discovery of the electron

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Cathode ray tube and the electron

• http//videos.howstuffworks.com/science-channel/29
  292-100-greatest-discoveries-the-cathode-ray-tube-
  video.htm

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J.J. Thomson (1897) Cathode Rays
Atoms subjected to high voltages give off cathode
rays.
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J.J. Thomson Cathode Rays
Cathode rays can be deflected by a magnetic field.
Cathode rays are negatively charged particles
(electrons).
Electrons are in atoms.
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Thompsons Plum Pudding Model
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400 to 0, 0 to 1600 400 1600 2000 years
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Radioactivity

• Radioactivity is the spontaneous emission of
  radiation by an atom.
• First observed by Henri Becquerel
• (1852-1908).
• Marie and Pierre Curie also studied it.
• Nobel Prize in 1903 (physics).

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Radiation named by Rutherford
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2. Discovery of the nucleus
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Discovery of the nucleus

• http//www.youtube.com/watch?v5pZj0u_XMbc

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• Since the vast majority of a particles pass
  through the Au foil undeflected, the Au atoms are
  mostly ______.
• Empty space

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• A very tiny percentage of a particles hit
  something massive in the atom and backscatter
  (bounce back).
• This indicates that most of the mass of the atom
  is concentrated in a very small volume relative
  to the volume of the entire atom.
• We now call this the NUCLEUS.

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• Rutherford proposed the charge on the nucleus to
  be positive, since electrons are negatively
  charged and atoms are neutral.

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Rutherfords model
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• The primary limitation to Bohr Theory is that it
  was limited to a description of a one electron
  system, namely Hydrogen.
• The description of multiple electron systems is
  much more complex, and was only adequately
  handled by the Modern Theory of Atomic Structure.

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Bohrs Theory

• 1. The electron travels in a circular path around
  the nucleus. This path is called an orbit.

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• 2. At normal living conditions, room temperature,
  the electron resides in the orbit which is
  closest to the nucleus.
• This is the position of lowest energy content for
  the electron, and is referred to as the Ground
  State. (This statement implies that there will be
  more that one orbit available to an electron.)

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• 3. As long as the electron remains in a specific
  orbit, no energy is gained or lost by the system.
  
• 4. If energy is added to an electron, the
  electron will move to a new orbit. This orbit
  will be farther from the nucleus, and is a
  position of higher energy content. This new
  position is known as an excited state.

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• 5. When an electron moves from one orbit to
  another orbit, it does so without ever passing
  through the space between the orbits. In other
  words, the electron is only allowed to exist at
  very specific distances from the nucleus, or
  positions of very specific energy content. (This
  idea is much like climbing a ladder. The foot is
  only allowed to be placed in very specific
  locations.) This idea is known as a quantum jump,
  a transition in which the electron gains or loses
  a very specific amount of energy.

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• Part Five off the Bohr Theory is, perhaps, the
  most controversial item. It says that an electron
  is restricted to having certain specific
  quantities of energy. The electron will never be
  allowed to have energy in between the allowed
  values. This is referred to as the quantization
  of energy. The idea was first expressed by Max
  Planck. This piece of information, when given to
  Bohr, suddenly made the ideas that he expressed
  much more meaningful.

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• In essence, it now becomes clear why an atom will
  only release specific colors of light, or
  specific wavelengths of electromagnetic
  radiation. Without the Planck contribution, the
  Bohr atom would release all colors of light.

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• 6. When an electron is in an excited state, it
  will always drop down to a lower energy state,
  ultimately returning to ground state. Each
  electron transition to a lower energy state will
  be accompanied by the simultaneous release of
  energy. This energy is released as
  electromagnetic radiation. The energy of the
  released radiation will correspond to the
  difference in energy content between the two
  levels.

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EMR (electromagnetic radiation)
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3 problems with the Bohr Theory

• The theory only works for a one electron system.
  What happens when an atom has more than one
  electron?
• The theory violates the Heisenberg Uncertainty
  Principle. Bohr Theory makes the behavior of the
  electron entirely to predictable. Bohr claims it
  is possible to know exactly where an electron is
  and what it is doing. The Heisenberg Uncertainty
  Principle says that is not possible.

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• The Bohr Theory will based on trying to explain
  four visible colors in the hydrogen atomic
  spectrum. He worked with a red line, blue-green
  line, blue line, and violet line. With improved
  instrumentation, it is now known that the red
  line is actually two red lines. These lines are
  extremely close together, and are referred to as
  a doublet. The instruments that were available to
  Bohr were not sophisticated enough to distinguish
  the two red lines. To him they looked as if they
  were one wide red line.

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• The Atomic Spectrum is a series of lines of color
  produced when light from an excited atom is
  passed through a prism. It is also known as a
  line spectrum.

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• 3. (Failure of) the classical description of the
  atom

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Water wave
Sound wave
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EMR electromagnetic radiation
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Orbitals...
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Electrons are part of what makes an atom an atom
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Electrons are part of what makes an atom an atom
But where exactly are the electrons inside an
atom?
atom
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Orbitals are areas within atoms where there is a
high probablility of finding electrons.
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Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
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Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
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Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
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Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
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Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
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Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
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Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
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It has been determined where the orbitals are
inside an atom, but it is not known precisely
where the electrons are inside the orbitals
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It has been determined where the orbitals are
inside an atom, but it is not known precisely
where the electrons are inside the orbitals (as
described by Heisenburgs Uncertainty
Principle)
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Hey, where am I?
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The area where an electron can be found, the
orbital, is defined mathematically, but we can
see it as a specific shape in 3-dimensional space
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z
y
x
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z
y
The 3 axes represent 3-dimensional space
x
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z
y
For this presentation, the nucleus of the atom is
at the center of the three axes.
x
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The 1s orbital is a sphere, centered around the
nucleus (l 0)
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The 2s orbital is also a sphere.
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The 2s electrons have a higher energy than the
1s electrons. Therefore, the 2s electrons are
generally more distant from the nucleus, making
the 2s orbital larger than the 1s orbital.
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1s orbital
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2s orbital
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Dont forget an orbital is the shape of
the space where there is a high probability of
finding electrons
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Dont forget an orbital is the shape of
the space where there is a high probability of
finding electrons
The s orbitals are spheres
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There are three 2p orbitals
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The three 2p orbitals (l 1) are
oriented perpendicular to each other
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z
This is one 2p orbital (2py)
y
x
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z
another 2p orbital (2px)
y
x
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z
the third 2p orbital (2pz)
y
x
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Dont forget an orbital is the shape of
the space where there is a high probability of
finding electrons
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Dont forget an orbital is the shape of
the space where there is a high probability of
finding electrons
This is the shape of p orbitals
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z
y
x
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z
2px
y
x
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z
2px and 2pz
y
x
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z
The three 2p orbitals, 2px, 2py, 2pz
y
x
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z
The three 2p orbitals, 2px, 2py, 2pz (m -1,
0, 1)
y
x
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once the 1s orbital is filled,
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the 2s orbital begins to fill around the 1s
orbital
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once the 2s orbital is filled,
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the 2p orbitals begin to fill
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each 2p orbital intersects the 2s orbital and the
1s orbital
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each 2p orbital gets one electron before pairing
begins
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once each 2p orbital is filled with a pair of
electrons, then
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the 3s orbital gets the next two electrons
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the 3s electrons have a higher energy than 1s,
2s, or 2p electrons,
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so 3s electrons are generally found further from
the nucleus than 1s, 2s, or 2p electrons
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What does that have to do with anything??
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the billions of interactions of atoms constantly
going on around you depend on how the
electrons are arranged in each atom
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the billions of interactions of atoms constantly
going on around you depend on how the
electrons are arranged in each atom
the arrangement of an atoms electrons (its
orbitals) govern how that atom will interact with
other atoms
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the billions of interactions of atoms constantly
going on around you depend on how the
electrons are arranged in each atom
the arrangement of an atoms electrons (its
orbitals) govern how that atom will interact with
other atoms
If atoms did not interact with each other, you
would not be sitting here reading this
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An interesting place where electrons have a
specific organization within atoms, allowing for
intersting atom interactions
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Not an interesting place, where electrons have no
specific organization within atoms, where atoms
wander aimlessly about
An interesting place where electrons have a
specific organization within atoms, allowing for
intersting atom interactions
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Principal Quantum Number, n

• Indicates main energy levels
• n 1, 2, 3, 4
• Each main energy level has sub-levels

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Energy Sublevels

• s p d f g

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• The principle quantum number, n, determines the
  number of sublevels within the principle energy
  level.

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Orbital Quantum Number, l(Angular Momentum
Quantum Number)

• Indicates shape of orbital sublevels
• l n-1
• l sublevel
• 0 s
• 1 p
• 2 d
• 3 f
• 4 g

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Orbital

• The space where there is a high probability that
  it is occupied by a pair of electrons.
• Orbitals are solutions of Schrodingers equations.

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Orbitals in Sublevels

• Sublevel Orbitals electrons
• s 1 2
• p 3 6
• d 5 10
• f 7 14
• g 9 18

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Three rules are used to build the electron
configuration

• Aufbau principle
• Pauli Exclusion Principle
• Hunds Rule

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Aufbau Principle

• Electrons occupy orbitals of lower energy first.

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Aufbau Diagram
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-Pauli Exclusion Principle(Wolfgang Pauli,
Austria, 1900-1958)-Electron Spin Quantum Number

• An orbital can hold only two electrons and they
  must have opposite spin.
• Electron Spin Quantum Number (ms)
• 1/2, -1/2

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Hunds Rule

• In a set of orbitals, the electrons will fill the
  orbitals in a way that would give the maximum
  number of parallel spins (maximum number of
  unpaired electrons).
• Analogy Students could fill each seat of a
  school bus, one person at a time, before doubling
  up.

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Aufbau Diagram for Hydrogen
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Aufbau Diagram for Helium
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Aufbau Diagram for Lithium
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Aufbau Diagram for Beryllium
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Aufbau Diagram for Boron
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Aufbau Diagram for Carbon
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Aufbau Diagram for Nitrogen
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Aufbau Diagram
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Notations of Electron Configurations

• Standard
• Shorthand

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Aufbau Diagram for Fluorine
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Standard Notation of Fluorine
Number of electrons in the sub level 2,2,5
1s2 2s2 2p5
Main Energy Level Numbers 1, 2, 2
Sublevels
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Shorthand Notation

• Use the last noble gas that is located in the
  periodic table right before the element.
• Write the symbol of the noble gas in brackets.
• Write the remaining configuration after the
  brackets.
• Ex Fluorine He 2s2 2p5

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Blocks in the Periodic Table
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General Periodic Trends

• Atomic and ionic size
• Ionization energy
• Electron affinity

Higher effective nuclear charge.
Electrons held more tightly
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Atomic Size

• Size goes UP on going down a group.
• Because electrons are added farther from the
  nucleus, there is less attraction.
• Size goes DOWN on going across a period.

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Atomic Radii
Figure 8.9
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Trends in Atomic SizeSee Figures 8.9 8.10
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Ion Sizes
Does the size go up or down when losing an
electron to form a cation?

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Ion Sizes
Forming a cation.
Li,152 pm
3e and 3p

• CATIONS are SMALLER than the atoms from which
  they come.
• The electron/proton attraction has gone UP and so
  size DECREASES.

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Ion Sizes

• Does the size go up or down when gaining an
  electron to form an anion?

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Ion Sizes
Forming an anion.

• ANIONS are LARGER than the atoms from which they
  come.
• The electron/proton attraction has gone DOWN and
  so size INCREASES.
• Trends in ion sizes are the same as atom sizes.

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Trends in Ion Sizes
Figure 8.13
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Ionization EnergySee Screen 8.12

• IE energy required to remove an electron from
  an atom in the gas phase.

Mg (g) 738 kJ ---gt Mg (g) e-
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Ionization EnergySee Screen 8.12

• Mg (g) 735 kJ ---gt Mg (g) e-
• Mg (g) 1451 kJ ---gt Mg2 (g) e-

Mg2 (g) 7733 kJ ---gt Mg3 (g) e-
Energy cost is very high to dip into a shell of
lower n. This is why ox. no. Group no.
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