electrons and the periodic table honors chemistry
TRANSCRIPT
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Electrons and the Periodic
TableHonors Chemistry
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The Periodic Table
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Development of the Periodic Table
1) History of the Periodic Table – By the end of the 1700’s, scientists had identified only 30 elements (ex. Cu, Ag, Au, H2, N2, O2, C).
2) By the mid 1800’s, about 60 elements had been identified.
3) Sept 1860 – chemists assembled at the First International Congress of Chemists in Germany to settle the controversial issues such as atomic mass. Standard values set for atomic mass and improved communication for research.
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Johann Dobereiner: 1817Organized the elements into
sets of three with similar properties.
He called these groups triads. The middle element is often the average of the other two.
Ex) Cl – 35.5 Br – 79.9 I – 126.9
CaAvg Sr
Ba
Cl + IAvg.
2
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Triads on the Periodic Table
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John Newlands: 1866• Arranged elements in
order of increasing atomic mass.
• Noticed repeating patterns in the elements’ properties every 8th element.
• Law of Octaves - properties of elements repeated every 8th element.
• There were 62 known elements at the time.
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Dmitri Mendeleev: 1869• Arranged elements in
order of increasing atomic mass.
• Similar properties occurred after periods (horizontal rows) of varying lengths.
• Organized the 1st periodic table according to increasing atomic mass and put elements with similar properties in the same column.
• Periodic – repeating properties or patterns
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Mendeleev’s 1st
Periodic Table
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• Mendeleeve Noticed inconsistencies in the arrangement of increasing atomic mass.
• He arranged some elements out of atomic mass order to keep them together with other elements with similar properties. (Notice Te and I)
• He also left several blanks in his table. • In 1871, he correctly predicted the
existence and properties of 3 unidentified elements – Sc, Ga and Ge
• These elements were later identified and matched his predictions.
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1st Periodic Law
Properties of the elements repeat periodically when the elements are arranged in increasing order
by atomic mass
Mendeleev is known as the Mendeleev is known as the
Father of ChemistryFather of Chemistry
Element 101 (Md) honors Element 101 (Md) honors MendeleevMendeleev
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What’s wrong with Mendeleev’s PT?
• We know that Mendeleev's periodic table was underpinned by false reasoning. – Mendeleev believed, incorrectly, that chemical
properties were determined by atomic weight.
• In 1869 the electron itself had not been discovered until 1896 - 27 years later.
• It took 44 years for the correct explanation of the regular patterns in Mendeleev's periodic table to be found...
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Henry Moseley: 19131. Fired electrons at atoms, resulting in
emission of x-rays. Each element emitted x-rays at a unique frequency. – He examined the pattern that was best
explained if the positive charge in the nucleus increased by exactly one unit from element to element.
2. Element are different from one another because their atoms have different number of positive charge (protons).
3. Analyzed data and found that the elements in the PT fit into patterns better when arranged in increasing nuclear charge, which is the Atomic Number.
The Modern Periodic
Table: When elements are arranged in
order of increasing
atomic number, their physical and chemical
properties show a periodic pattern.
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Glenn Seaborg “Seaborgium” Sg #106
• Born in 1912 in Michigan, Seaborg proposed reorganizing the Periodic Table one last time as a young chemist working on the Manhattan Atomic Bomb Project during WWII.
• He suggested pulling the “f-block” elements out to the bottom of the table.
• He was the principle or co-discoverer of 10 transuranium elements.
• He was awarded the Noble prize in 1951 anddied in 1999.
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Seaborgium is the exception…• After some argument between the USA and the rest
of the world, element 106 was named Seaborgium shortly before he died. This was a matter of some controversy because the International Union of Pure and Applied Chemistry, IUPAC, the body that deals with naming in chemistry, had previously ruled that elements should not be named after living people.
Atomic # 104 105 106 107
IUPACUnnilquadeum
UnqUnnilpentium
UnpUnnilhexium
UnhUnnilseptium
Uns
Agreed in 1995
Dubnium(Dubna, Russia)
Joliotium(Frederic Joliot)
Rutherfordium
(Earnest Rutherford)
Bohrium(Neils Bohr)
Agreed in 1996
Rutherfordium
Dubnium Seaborgium Bohrium
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Parts of the Periodic Table
A. Horizontal Rows – PERIODS– There are 7 periods in the periodic table– Elements in a period do NOT have similar
properties.
B. Vertical Columns – GROUPS or FAMILIES– Labeled 1-18– IA-VIIIA are the Main-group or
representative elements. – Elements in a group have similar
properties. – Why?
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Family Nameswrite these on your P.T.
• Hydrogen (1)• Alkali metals (1) – most
reactive metals; reactivity increases down the group
• Alkaline earth metals (2)• Boron family (13)• Carbon family (14)• Nitrogen family (15)• Oxygen or Chalcogen
family (16)• Halogens (17)• Noble gases (18) - inert
• Transition elements or metals (3-12): d-block
• Inner transition elements or metals (f-block)– Lanthanides or
lanthanide series– Actinides or actinide
series– Transuranium elements
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1A 1
2A2
3A 4A 5A 6A 7A13 14 15 16 17
8A18
1 = Alkali Metals and Hydrogen Group
13 = Boron Group
18 = Noble Gas Group
17 = Halogen Group
16 = Oxygen or Chalcogen Group
15= Nitrogen Group
14 = Carbon Group
2 = Alkaline Earth Metals
Transition Metals
Lanthanide Series
Actinide Series
Inner Transition Metals
3 4 5 6 7 8 9 10 11 12
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Parts of the Periodic Table
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The ELECTRONHow are the
electronsarranged
around thenucleus?
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To Review• Dalton Thomson Rutherford
Bohr Quantum Mechanical (Schrödinger) Model
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• Bohr – electrons in a particular path have a fixed energy called energy levels– Rungs of a ladder
• Quantum Mechanical (Schrödinger) Model– Electrons better understood as WAVES– Does not tell where the electrons are located– Electrons have a certain amount of energy -
QUANTIZED
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Parts of a Wave
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Light as a WaveCharacteristics of a WaveA. Amplitude: Height of the wave from the
baseline. The higher the wave the greater the intensity.
B. Wavelength: (λ , “lambda”) in nanometers (1 x 10-9 m). Distance between similar points on 2 consecutive waves.
C. Frequency: (ν , “nu”) The number of waves that pass a fixed point per unit of time. Measured in cycles/second (1/s) 1 cycle/second = Hertz (Hz)
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Baseline
Crest
Wavelength Amplitude
Amplitude
Wavelength Trough
3
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D. Electromagnetic Radiation - a form of energy that exhibits wavelike behavior
as it travels through space- all forms of EM radiation move at the speed of
light
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Speed of Light (c)
E. 3.00 x 108 m/s or 186,000 miles/sec. The relationship between wavelength
and frequency can be shown with the
following equation:
This is an indirect relationship. If λ then ν .
c = λ ν
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In the US, AM stations all have longer wavelengths than the FM stations. •AM broadcast stations are licensed to operate only in the band 550-1700 kHz •FM broadcast stations are licensed to operate only in the band 88-108 MHz
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Visible Light
Radar
Microwaves
Infrared
Radio/TV Ultraviolet
X-Rays
Gamma Rays
Low
Long
High
Short
Red Orange Yellow Green
Blue Violet
Energy
Energy
Low High
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• Incandescent light bulbs give off most of their energy in the form of heat-carrying infrared light photons -- only about 10 percent of the light produced is in the visible spectrum. This wastes a lot of electricity. Cool light sources, such as fluorescent lamps and LEDs, don't waste a lot of energy generating heat -- they give off mostly visible light. For this reason, they are slowly edging out the old reliable light bulb.
http://home.howstuffworks.com/light-bulb2.htm
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Quantum Theory
A. Planck’s Hypothesis: (Max Planck 1900)
1. Studied emission of light from hot objects
2. Observed color of light varied with temperature
3. Suggested the objects do not continuously emit E, but emit E in small specific amounts
a. Light is absorbed or emitted in a little packet or bundle called a quantum (quanta –plural).
b. Quantum = minimum amount of E that can be lost or gained by an atom
c. Energies are quantized. (Think steps not a
ramp)
e-
e-
Xe-
e-
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Max Planck’s Energy Equation
4. Proposed that energy is directly proportional to frequency.
E = h Planck’s equation for each quantum
h = Plank’s constant = 6.626 x 10-34 J.s
This is a direct relationship.As energy increases, frequency
increases.
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Albert Einstein
(1879 – 1955)
German Physicist
While well-known for the equation E=mc2 , Einstein’s work on the photoelectric effect resulted in being awarded the 1921 Nobel Prize in Physics.
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Refers to the emission of electrons from a metal
when light shines on the metal
Observations:1. Electrons are ejected by light of
sufficient energy. Energy minimum is different for different metals.
2. The current (# of electrons emitted/s) increases with brightness of the light.
+-
Albert Einstein and the Photoelectric
Effect
my.hrw.com
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Conclusions: 1. Proposed that light consists of quanta of
energy that behaves like particles. 2. Quantum of light = photon = massless
particle that carries a quantum of energy.
3. Proposed the Dual Nature of Light: its wave and particle nature.
a) Light travels through space as wavesb) Light acts as a stream of particles when it
interacts with matter.
Albert Einstein and the Photoelectric Effect
my.hrw.com
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Light (Electromagnetic Radiation)
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SpectroscopyDefinition: a method of studying substances
that are exposed to some sort of continuous exciting energy.
A. Emission Line Spectra: contains only certain colors or wavelengths ( ) of light.
1. Every element has its own line spectrum (fingerprint).
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Continuous Spectrum – White Light
Line Spectrum – Excited Elements
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White light gives off a Continuous Spectrum
a blending of every possible wavelength
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Gas Discharge Tubes
• Electricity is added to the gas which causes the electrons to jump to a higher or excited state. They immediately fall back to the ground state and give off particular wavelengths of light. We see a blending of wavelengths without the spectroscopes.
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Flame Tests
• used to test qualitatively for the presence of certain metals in chemical compounds.
• the heat of the Bunsen flame excites electrons that emit visible light.
Copper(II) sulfate Lithium chloride
Potassium chloride Barium nitrate
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Spectroscope• Uses a diffraction grating to diffract the
light into particular wavelengths of light.
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A Line Spectra result from excited elements - as electrons of an element gain energy and
rise to an excited state they then fall back to their ground state in the same pattern
producing the same energy drop each time which we see as individual wavelengths of
light.
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Atomic Spectra and the Bohr Model of Hydrogen (1913)
Neils Bohr - Danish ScientistExplained the bright-line spectrum
of hydrogenStudy: • Added E as electricity to H gas
at low pressure in a tube.• Emitted E as visible light, was
observed through a prism
Result: Hydrogen emitted 4 distinct bright lines of color, aka bright line spectrum
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Electrons release energy as they fall back to a lower
energy level
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Electrons absorb energy to rise to a higher or excited state and emit energy in the form of a photon of light as they fall back to their
ground states.
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Path of an excited electron as it “falls” back to the Ground State
• When electrons gain energy, they jump to a higher energy level (excited state).
• Electrons are not stable at the excited state and will immediately fall back to a lower level or ground state.
• As they fall, they emit electromagnetic radiation.
• Depending on how far they fall determines the type of radiation (light) released.
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Bohr Model of HydrogenConclusion: • *Unique line spectrum is due to quantized electron energies.• *Electrons are in specific orbits related to certain amounts of
energy known as stationary states. • *Orbits are related to energy levels.
• *Energy levels are identified as E1, E2, E3, … (n = 1, 2, 3, …)
• *Lowest energy level = ground state• *Electrons absorb certain amounts of energy to move to a higher
energy level farther away from the nucleus = excited state• *Electrons return to the more stable ground state and release a
photon that has energy equal to the difference in energy between the energy levels.
– from E2 to E1: Ephoton = E2 – E1 (difference in energy)
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The Bohr Atom for Hydrogen -a Model
1. Successful in calculating the wavelength, frequency, & energy of hydrogen’s line spectrum.
2. Successful in calculating the energy needed to remove hydrogen’s electron
H(g) + energy H+1(g) + 1e-
Calculated ionization E = observed ionization E = 1312.1 kJ/mol
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Lyman, Balmer and Paschen series of the Hydrogen Atom
• Lyman series: electrons fall to n = 1 and give off UV light.
• Balmer series: electrons fall to n = 2 and give off visible light.
• Paschen series: electrons fall to n = 3 and give off infrared light.
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When electrons absorb energy they jump to a higher (excited) state.
n=2 n=3 n=4 n=5 n=6 n=7
Electrons are not stable. Radiation (light) is emitted when an electron falls back from a higher level to a lower level.
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Ultraviolet Light
Visible Light
Infrared Light
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Atomic Spectra
Hydrogen
Lithium
Mercury
Helium
Although Bohr’s atomic model explained the line spectra of hydrogen, it failed for heavier elements.
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Limitations of the Bohr Model
a. Model could not calculate the wavelengths of observed spectra of multi-electron atoms.
b. Model could not explain the chemical behavior of atoms.
c. Bohr used classical mechanics to understand the behaviors of small particles.
d. The Bohr model is also known as the planetary, solar system, or satellite model.
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Quantum Mechanical Model of the Atom
A. Louie De Broglie (1924-5) 1. Took Einstein’s idea that light can
exhibit both wave and particle properties2. Very small particles (like electrons)
display properties of waves.3. Behavior of electrons in Bohr’s quantized
orbits was similar to behavior of waves
Known: any wave confined to a space can only have specific frequenciesDe Broglie suggested electrons are waves confined to the space around the atomic nucleus. Electrons could exist only at specific frequencies which correspond to specific energies (E = h quantized E of Bohr)
French scientist
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4. Experimentally proven in 1927 by diffraction of electrons by Davisson & Germer (showed diffraction of electrons by a crystal of Ni)
B. Wave-Particle Duality of Naturea. Light and electrons (very small particles like
electrons, atoms, molecules) have properties of waves and particles QUANTUM MECHANICS (based on WAVE properties)
**Large objects obey the laws of classical mechanics**
Quantum Mechanical Model
of the Atom
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C. Werner Heisenberg: (1927)
1. Heisenberg’s Uncertainty Principle: states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.
2. You cannot predict future locations of particles.
3. He found a problem with the Bohr Atom - no way to observe or measure the orbit of an electron.
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D. Erwin Schrödinger Wave Equation (1926)
1. Wave nature of an electron is described by a mathematical equation.
2. Four quantum numbers in the equation are used to describe an electron’s behavior – location and energy.
3. Electron is treated as a wave with quantized energy.
4. Describes the probability of the electrons found in certain locations around the nucleus.
(1887 – 1961) Austrian Physicist
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Electron DensityAn orbital is a region inwhich an electron with aparticular energy is likelyto be found. Where the density of anelectron cloud is high
there is a high probability that
iswhere the electron islocated. If the electrondensity is low then there
isa low probability.
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E. Atomic Orbitals - region around the nucleus where an electron with a particular energy is likely to be found (not the same as Bohr’s orbits!)1. Orbitals have characteristic shapes, sizes,
& energies.2. Orbitals do not describe how the electron
moves.3. The drawing of an orbital represents the
3-dimentional surface within which the electron is found 90% of the time.
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4. Sublevels can have 4 different shapes
s – orbital spherical
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1s, 2s & 3s orbitals Superimposed on one another
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Electron-Cloud Models
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p-orbital – dumbbell shaped
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p-orbital - dumbbell shaped
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d-orbital - double dumbbell or fan blades
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s,p and d orbitals
For a more complete representation and presentation of atomic orbitals go to http://winter.group.shef.ac.uk/orbitron/
x
y
z
x
y
z
x
y
z
x
y
z
x
y
z
x
y
z
x
y
z
x
y
z
x
y
z
s orbital p orbitals
d orbitals
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Models of d-orbitals
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f-orbital – more
complex!
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f orbitals
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f – orbitals (3D)
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Quantum Numbers
• Each quantum number provides more specific information on the probable location of an electron.
• Each electron within an atom can be described by a unique set of 4 quantum numbers.
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Quantum Numbers - Finding an address for each
electron:1. “state” Principle Quantum Number (n)
or the energy level; a. Describes the relative size of the electron cloud.b. Positive integer values (n = 1 to n = 7)
2. “city” Sublevel (l) a. Describes the shape of the electron cloud.b. The maximum number of sublevels within a
level = nc. Shapes are s, p, d,or f.d. Lowest energy = s Highest energy = f
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Quantum numbers cont.
3. “street” Orbital (ml) odd # of orbitals1. Describes the orientation or direction in
spacea) s – 1 orbitalb) p – 3 orbitals (x, y, z)c) d – 5 orbitals (xy, yz, xz, x2 – y2, z2)d) f – 7 orbitals (y3 – 3yx2, 5yz2-yr2, x3-3xy2,
zx2-zy2, xyz, 5xz2-3xr2, 5z3-3zr2)
2. Orbitals within the same sublevel have the same energy are called degenerate orbitals
3. An orbital can hold a maximum of 2 electrons
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4. “house” Spin (ms) 1. Describes the direction of electron spin in
an orbital. 2. The clockwise or counterclockwise motion
of electrons.3. Only electrons with opposite spins can
occupy the same orbital. 4. The opposite spin is written as+1/2 or -1/2 or or
Quantum numbers cont.
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E. Electron Configurations:
1. Shorthand notation for indicating the number of electrons in each level, sublevel, and orbital.
1s2
2. Shows the distribution of electrons among the orbitals. Describes where the electrons are found & what energy they possess.
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Electron Configuration Rules
1. The Aufbau Principle: electrons are added one at a time to the lowest energy orbital available.
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Pauli Exclusion Principle:1. Each orbital can only hold 2 electrons.
2. The electrons must have opposite spins.
s-sublevel = max 2 electronsp-sublevel = max 6 electronsd-sublevel = max 10 electronsf-sublevel = max 14 electrons
incorrect: ↑↑↑ incorrect: ↑↑ correct: ↑↓
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Hund’s Rule:• Electrons will
remain unpaired in degenerate orbitals before they pair up.
incorrect ↑↓ ↑ __
correct ↑ ↑ ↑
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Electron Blocks on the
Periodic Table
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Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Pauli Exclusion Principle: No more than 2 e- are put in each orbital and they must have opposite spin.Hund’s Rule: electrons spread out among equal energy orbitals in a sublevel (like charges repel)Aufbau Principle: Electrons fill lowest energy levels first (n=1)
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Electron Configuration Examples:
Ex) electron configuration for Na:
1s2 2s2 2p6 3s1
Ex) orbital filling box diagram for Na:
yx z
1s 2s 2p 3s
_
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Electron Dot Diagrams:
Write the symbol for the element.
Place dots around the symbol to represent the
valence s & ps & p electrons only.
Do NOT include d & f orbitals in diagram.
p orbital electrons s orbital electrons
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Electron Configuration Orbital Box Diagram Electron-dot Diagram
yx z
1s 2s 2p
y yx z x
1s 2s 2p 3s 3p
z
168O
3517 Cl 1s22s22p63s23p5
12752Te 1s22s22p63s23p64s23d104p65s24
d105p4
1s22s22p4
y y y yx z x z x z x z
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p
What does the Tellurium electron-dot resemble???
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Mark your Periodic Tables
1 2 13 14 15 16 17 18
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Unpaired vs. Paired Electrons
Filled and Half-filled orbitals• Atoms with unpaired electrons are said to
be paramagnetic. These are weakly attracted to a magnetic field.
• Atoms with all paired electrons are said to be diamagnetic. These are weakly repelled from a magnetic field.
• ½ filled and filled orbitals have special stability
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Noble Gas or Shorthand Electron Configurations
• Rb
• Se
• At
1[Kr]5s
2 10 4[Ar]4s 3d 4p
2 14 10 5[Xe]6s 4f 5d 6p
Draw the Dot Diagrams for these elements
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Exceptions to the Rules
•Max stability - ½ filled and filled orbitals–Cr–Mo–Cu–Ag–Au
2 4[Ar]4s 3d 1 5[Ar]4s 3d
1 10[Ar]4s 3d2 9[Ar]4s 3d
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Exceptions to the Rules
•Max stability - ½ filled and filled orbitals–Cr–Mo–Cu–Ag–Au
2 4[Ar]4s 3d 1 5[Ar]4s 3d
1 5[Kr]5s 4d
1 10[Ar]4s 3d
1 10[Kr]5s 4d
1 14 10[Xe]6s 4f 5d
2 9[Ar]4s 3d
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Electron Configuration for Ions
• K
• P
• Al
• Se
1[Ar]4s • K+1
• P-3
• Al+3
• Se-2
2 2 6 2 6 11s 2s 2p 3s 3p 4s2 2 6 2 61s 2s 2p 3s 3p
[Ar]
2 2 6 2 31s 2s 2p 3s 3p
2 3[Ne]3s 3p2 2 6 2 61s 2s 2p 3s 3p
2 6[Ne]3s 3p
2 2 6 2 11s 2s 2p 3s 3p
2 1[Ne]3s 3p2 2 61s 2s 2p
[Ne]
2 2 6 2 6 2 10 41s 2s 2p 3s 3p 4s 3d 4p
2 10 4[Ar]4s 3d 4p2 2 6 2 6 2 10 61s 2s 2p 3s 3p 4s 3d 4p
2 10 6[Ar]4s 3d 4p
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Excited vs. Ground State
•If an electron absorbs energy, it is in an EXCITED state
Ne: 1s22s22p53s1
•How is this different from the ground state configuration?
Ne: 1s22s22p6