Electronic Structure and the Periodic Table

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Electronic Structureand the Periodic Table

• Unit 6 Honors Chemistry

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Electromagnetic Waves
Electromagnetic waves progressive, repeating
disturbances that come from the movement of
electric charges Electromagnetic Waves Light
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Wavelength and Frequency

• Wavelength (?, lambda) distance between any two
  points in a wave
• measured in any distance unit
• (mainly nm or m
• 1 nm 1x10-9 m)

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Wavelength Can be Measured in One of Two Ways
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Wavelength and Frequency

• Frequency (? pronounced nu)
• the number of cycles
• of the wave that pass through a point in a unit
  of time
• Measured in sec-1 (/sec)
• 1 sec-1 1 Hertz (Hz)

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Illustration of Frequency
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Wavelength is indirectly proportional to frequency
As Wavelength increases, frequency
_________________. As Wavelength decreases,
frequency _________________.
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Amplitude

• Note height of wave is amplitude (intensity or
  brightness of wave)
• Amplitude is INDEPENDENT of frequency or
  wavelength!

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Speed

• Speed (c) The speed of light!
• c 3.00 x 108 m/s
• (rounded to 3 sig figs)

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Equation

• One equation relates speed, frequency and
  wavelength
• c ? ?

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Example

• The wavelength of the radiation which produced
  the yellow color of sodium vapor light is 589.0
  nm. What is the frequency of this radiation?

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The electromagnetic spectrum

• complete range of wavelengths and frequencies
• mostly invisible

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What is color?

• TED Talk What is color?

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The visible/continuous spectrum

• continuous spectrum components of white light
  split into its colors, ROY G BIV
• from 390 nm (violet) to 760 nm (red)

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Line Spectra

• Pattern of lines produced by light emitted by
  excited atoms of an element
• unique for every element
• used to identify unknown elements

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How do we see color?

• TED Talk How we see color

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Max Planck

• Light is generated as a stream of particles
  called PHOTONS
• Equation
• E (Energy of a photon) h?
• (h Planks constant
• 6.626x10-34J?s)

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Relationships in Plancks Eqn.
E h ?
High frequency, low ?, high E.
Low frequency, high ?, low E.
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Photoelectric effect Nobel Prize in Physics
1921 to Einstein

• Occurs when light strikes the surface of a metal
  and electrons are ejected.
• Practical uses
• Automatic
• door openers

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Photoelectric Effect Conclusion

• Light not only has wave properties but also has
  particle properties. These massless particles,
  called photons, are packets of energy.

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Example 6.2

• Using the frequency calculated in the previous
  example, calculate the energy, in joules, of a
  photon emitted by an excited sodium atom.
  Calculate the energy, in kilojoules, of a mole of
  excited sodium atoms.

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Bohrs Hydrogen Atom A Planetary Model
Niels Bohr Proposed planetary model. Electrons
orbit the nucleus like planets around the
sun. NOT current model of atom but used to
explain some features of atom.
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Ground State vs. Excited State

• ground state all electrons in lowest possible
  energy levels
• excited state an electron that has absorbed
  energy and moved to a higher energy level
• This is a temporary state!!

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Explanation of Line Spectra Equation

• Niels Bohr
• Energy of an electron is quantized can only have
  specific values.
• Energy proportional to energy level.

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Explanation of Line Spectra
Electron will drop from excited state to ground
state and will emit energy as a photon.
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Explanation of Line Spectra

• Type of photon emitted by electron depends on
  energy difference of energy levels
• ?Elevel -RH 1 1
• (nhi)2 (nlow)2
• AND ?Elevel h? hc/?
• (h Plancks constant, 6.626 x 10-34 J
  sec/photon)

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Flaw in Bohrs Model

• Only works well for 1 electron species (H atom).
• Does not explain fine structure of line spectra.

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Wave-Particle Duality

• Light has properties of both WAVES and PARTICLES.
• most matter has undetectable wavelengths (1000
  kg car at 100 km/hr has ? 2.39 x 10-38 m)
• This work led to the development of the electron
  microscope

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Quantum Mechanics

• Quantum mechanics
• atomic structure based on wave-like properties
  of the electron
• Schrödinger wave equation that describes
  hydrogen atom

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Heisenberg Uncertainty Principle

• The exact location of an electron cannot be
  determined (if we try to observe it, we interfere
  with the particle)
• You can know either the location or the velocity
  but not both
• Electrons exist in electron clouds
• and not on specific rings or orbits

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Quantum Numbers

• Four quantum numbers are a mathematical way to
  represent the most probable location of an
  electron in an atom
• analogy...
• state energy level, n
• city sublevel, l
• address orbital, ml
• house number spin, ms

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Principal Quantum Number n

• Always a positive integer (1,2, 37)
• Indicates size of orbital, or how far electron is
  from nucleus
• Similar to Bohrs energy levels or shells
• Larger n value larger orbital or distance from
  nucleus

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The Periodic Table and Shells
n row number on periodic table for a given
element
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Angular Momentum Quantum Number l

• positive integer from zero to n-1
• Sublevel within an energy level indicates shape
  of orbital
• 0 s
• 1 p
• 2 d
• 3 f

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Types of Sublevels
s
p
d
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Magnetic Quantum Numbers ml

• integer from -l to l
• Indicates orientation of orbital in space
• Orbital electron containing area

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Spin Quantum Number ms

• Two values only ½ or -½
• 2 electrons max. allowed in each orbital
• (Pauli Exclusion Principle)
• Indicates spin of electron spins of each
  electron must be opposite

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REVIEW QUANTUM NUMBERS
Every Electron has four!

• n ---gt level 1, 2, 3, 4, ...
• l ---gt sublevel 0, 1, 2, ... n - 1
• ml ---gt orbital -l ... 0 ... l
• ms ---gt electron spin ½ and -½

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Orbitals

• No more than 2 e- assigned to an orbital
• Orbitals grouped in s, p, d (and f) subshells

s orbitals
p orbitals
d orbitals
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Capacities of levels, sublevels, and orbitalssee
packet
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Example

• Example 6.6 Give the n and l values for the
  following orbitals
• a. 3p
• b. 4s

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Example

• Example 6.8 What are the possible ml values for
  the following orbitals
• a. 3p
• b. 4f

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Shapes of Atomic Orbitals
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Shapes of Atomic Orbitals
s spherical p peanut d dumbbell (clover) f
flower
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Multielectron Atoms
In the hydrogen atom the subshells (sublevels) of
a principal energy level or shell are at the
same energy level. Previous Equation En RH
/n2
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Multielectron Atoms
In a multielectron atom, only the orbitals are at
the same energy level the sublevels are at
different energy levels!
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The increasing energy order of sublevels is
generally

• s lt p lt d lt f

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Overlapping subshells
At higher energy levels, sublevels overlap.
Note 4s vs. 3d!
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Introduction to Electron Configuration
Definition describes the distribution of
electrons among the various orbitals in the atom
Represents the most probable location of the
electron!
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Electron Configurations

• The system of numbers and letters that designates
  the location of the electrons
• 3 major methods
• Full electron configurations
• Abbreviated/Noble Gas configurations
• Orbital diagram configurations

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Full or Complete Electron Configuration (uses
spdf)
Uses numbers to designate a principal energy
level and the letters to identify a sublevel a
superscript number indicates the number of
electrons in a designated sublevel.
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Rules for Electron Configurations
The Aufbau principle Electrons fill from the
lowest energy level to the highest (they dont
skip around) 1s22s22p63s23p64s23d10etc.
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Pauli Exclusion Principle

• No two electrons in the same atom can have the
  same set of 4 quantum numbers.
• That is, each electron has a unique address

In other words, the maximum of electrons an
orbital can hold is 2 e- (one with ms 1/2 and
one with ms -1/2)
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HUNDS RULE

• Orbitals of equal energy in a sublevel must all
  have 1 electron before the electrons start
  pairing up
• a.k.a creepy person on the bus rule
• also electrons in half-filled orbitals have
  same spin

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Why are these incorrect?
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Why are these incorrect?
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Why are these incorrect?
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Full Electron Configuration

• Example Notation
• 1s2 2s1 (Pronounced one-s-two, two-s-one)
• A. What does the coefficient mean?
• Principle energy level
• B. What does the letter mean?
• Type of orbital (sublevel)
• C. What does the exponent mean?
• of electrons in that orbital

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Steps to Writing Full Electron Configurations

• 1. Determine the total number of electrons the
  atom has (for neutral atoms it is equal to the
  atomic number for the element).
• Example F
• atomic of p of e-
• 2. Fill orbitals in order of increasing energy
  (see Aufbau Chart).
• 3. Make sure the total number of electrons in the
  electron configuration equals the atomic number.

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Aufbau Chart (Order of Energy Levels)

• When writing electron configurations
• d sublevels are n 1 from the row they appear in
• f sublevels are n 2 from the row they appear in

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Writing Electron Configurations

• Nitrogen
• Helium
• Phosphorous
• Rhodium
• Bromine
• Cerium

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Abbreviated/Noble Gas Configuration

• i. Where are the noble gases on the periodic
  table?
• ii. Why are the noble gases special?
• iii. How can we use noble gases to shorten
  regular electron configurations?

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Abbreviated/Noble Gas Configuration

• Example Barium
• 1. Look at the periodic table and find the noble
  gas in the row above where the element is.
• 2. Start the configuration with the symbol for
  that noble gas in brackets, followed by the rest
  of the electron configuration.

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Abbreviated/Noble Gas Configuration

• Practice! Write Noble Gas Configurations for the
  following elements
• Rubidium
• Bismuth
• Arsenic
• Zirconium

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Writing Electron Configurations

• Another way of writing configurations is called
  an orbital diagram.
• (also called orbital notation or orbital diagrams)

One electron has n 1, l 0, ml 0, ms
½ Other electron has n 1, l 0, ml 0, ms
- ½
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Orbital Diagrams

• Orbital diagrams use boxes (sometimes circles) to
  represent energy levels and orbitals. Arrows are
  used to represent the electrons.

orbital
sublevels
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Orbital Diagrams

• Dont forget - orbitals have a capacity of two
  electrons!! Two electrons in the same orbital
  must have opposite spin so draw the arrows
  pointing in opposite directions.
• Example oxygen 1s22s22p4

2p
Increasing Energy ?
2s
1s
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Drawing Orbital Diagrams

• First, determine the electron configuration for
  the element.
• Next draw boxes for each of the orbitals present
  in the electron configuration.
• Boxes should be drawn in order of increasing
  energy (see the Aufbau chart).
• Arrows are drawn in the boxes starting from the
  lowest energy sublevel and working up. This is
  known as the Aufbau principle.
• Add electrons one at a time to each orbital in a
  sublevel before pairing them up (Hunds rule)
• The first arrow in an orbital should point up
  the second arrow should point down (Pauli
  exclusion principle)
• Double check your work to make sure the number
  of arrows in your diagram is equal to the total
  number of electrons in the atom.
• of electrons atomic number for an atom

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Electron Configurations for Nitrogen
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Electron Configurations for Nickel
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Lithium

• Group 1A
• Atomic number 3
• 1s22s1 ---gt 3 total electrons

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Beryllium

• Group 2A
• Atomic number 4
• 1s22s2 ---gt 4 total electrons

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Boron

• Group 3A
• Atomic number 5
• 1s2 2s2 2p1 ---gt
• 5 total electrons

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Carbon

• Group 4A
• Atomic number 6
• 1s2 2s2 2p2 ---gt
• 6 total electrons

Here we see for the first time HUNDS RULE.
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Nitrogen

• Group 5A
• Atomic number 7
• 1s2 2s2 2p3 ---gt
• 7 total electrons

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Oxygen

• Group 6A
• Atomic number 8
• 1s2 2s2 2p4 ---gt
• 8 total electrons

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Fluorine

• Group 7A
• Atomic number 9
• 1s2 2s2 2p5 ---gt
• 9 total electrons

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Neon

• Group 8A
• Atomic number 10
• 1s2 2s2 2p6 ---gt
• 10 total electrons

Note that we have reached the end of the 2nd
period, and the 2nd shell is full!
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Exceptions to the Filling Order Rule (Cr,
Cu)these will not be on test!
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Valence electrons

• Importance and definition
• Definition Electrons in the outermost energy
  levels they determine the chemical properties of
  an element.
• Write the noble gas configuration...the valence
  electrons are the ones beyond the core.
• Example Sulfur

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Valence Electrons and Core Configuration
(Shorthand)
What is the shorthand notation for S?
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Configurations of Ions
Cations Formed when metals lose e in highest
principal energy level. Example
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Configurations of Ions
Anions Formed when non-metals gain e to
complete the p sublevel. Example
Z 18 Cl-
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Transition Metals
Transition metals (and p block metals) lose e
from the highest principal energy level (n)
FIRST, then lose their d electrons!
Zr Kr 5s24d2 Zr2 Kr 4d2
EOS
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Isoelectronic Species

• Definition Ions or atoms that have the same
  number of electrons
• Example Neon, O2-, F-, Na, Mg2, Al3
• all have the same configuration (1s22s22p6) and
  are isoelectronic

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Electron Spin and Magnetism

• Diamagnetic NOT attracted to a magnetic field
• Paramagnetic substance is attracted to a
  magnetic field.
• Substances with unpaired electrons are
  paramagnetic.

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Examples

• Mg
• Cl
• Write orbital notation if it has an unpaired e-
  it is paramagnetic.

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Periodic Properties Trends

• Electronegativity
• Ability of an atom to pull e- towards itself
• Increases going up and to the right
• Across a period ? more protons in nucleus more
  positive charge to pull electrons closer
• Down a group ? more electrons to hold onto
  element cant pull e- as closely

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Periodic Properties Trends

• Electronegativity
• Ability of an atom to pull e- towards itself
• Across a period ? more protons in nucleus more
  positive charge to pull electrons closer
• Down a group ? more electrons to hold onto
  protons in nucleus cant pull e- as closely

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Atomic Radius

• Definition
• ½ experimental distance between centers of two
  bonded atoms

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Atomic Radius

• Trend in a family
• Size increases
• down a group.
• (More principal
• energy levels)

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Atomic Radius

• Trend in a period
• Size decreases across a period, e- more strongly
  attracted to nucleus.

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Atomic Radius

• Transition metals
• Size stays relatively constant across a period
  e- added to inner energy level.

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Memory Device

• LLLL Lower Left, Larger Atoms

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Sizes of Ions
Li,152 pm
3e and 3p

• CATIONS are SMALLER than the atoms from which
  they are formed.
• Size decreases due to increasing he
  electron/proton attraction.

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Sizes of Ions

• ANIONS are LARGER than the atoms from which they
  are formed.
• Size increases due to more electrons in shell.

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Trends in Ion Sizes
Trends in ion sizes are the same as atom sizes.
Active Figure 8.15
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First Ionization Energy

• Definition 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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First Ionization Energies
Trend in a group Decreases going down a group
(e- further away easier to remove) Trend in a
period Increases going across a period (e- held
more tightly).
EOS
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Memory Device

• LLLL Lower Left,
• Larger Atoms
• Looser electrons

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Second Ionization Energy
Definition energy required to remove 2nd
electron from an atom in the gas phase. Takes
more energy because e- is removed from
increasingly positive ion.

• Mg (g) 738 kJ ---gt Mg (g) e-

Mg (g) 1451 kJ ---gt Mg2 (g) e-
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Electron Affinity

• Some elements GAIN electrons to form anions.
• Electron affinity is the energy involved when an
  atom gains an electron to form an anion.
• A(g) e- ---gt A-(g) E.A. ?E

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Trends in Electron Affinity
Trend in a group Affinity for e- decreases going
down a group Trend in a series or
period Affinity for e- increases going across a
period
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Electron Affinity
Note that the trend for E.A. is the SAME as for
I.E.!
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Trends in Metallic Properties
Most metallic means easiest loss of
electrons! Metals are on left, nonmetals on right
of p.t.
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A Summary of Periodic Trends
Remember LLLL!!