PowerPoint Presentation
Recent Atomic ModelsMax Planck (1900): Proposed thatamounts of
energy are quantizedonly certain E values are allowed Niels Bohr
(1913): e can possess only certain amounts of energy, andcan
therefore be only certain distances from nucleus.e foundheree
neverfound hereplanetary(Bohr) modelN
Continuous vs. Quantized EnergyEnergyA BZumdahl, Zumdahl,
DeCoste, World of Chemistry 2002, page 330continuous quantized2Bohr
Atom
The Planetary Model of the AtomObjectives:To describe the Bohr
model of the atom.To explain the relationship between energy levels
in an atom and lines in an emission spectrum.Bohrs
ModelNucleusElectronOrbitEnergy Levels ENERGY(HEAT, LIGHT,ELEC.,
ETC.)LightWhen all e are in lowest possible energy state, an atom
is in the ____________.ground statee.g., He: 2 e-, both in 1st
energy level If right amount of energy is absorbed by an e, it
canjump to a higher energy level. This is an unstable,momentary
condition called the ____________. excited statee.g., He: 1 e- in
1st E level, 1 e- in 2nd E level When e falls back to a
lower-energy, more stableorbital (it might be the orbital it
started out in, but itmight not), atom releases the right amount
ofenergy as light. EMITTED LIGHTAny-old-value of energy to
beabsorbed or released isNOT OK. This explainsthe lines of color in
anemission spectrum. Frequency AFrequency BFrequency Cn = 2n = 1n =
3N7n = 1 corresponds to the orbit closest to the nucleus and is the
lowest in energy.
A hydrogen atom in this orbit is called the ground state, the
most stable arrangement for a hydrogen atom.
As n increases, the radii of the orbit increases and the energy
of that orbit becomes less negative.
A hydrogen atom with an electron in an orbit with n >1 is in
an excited state energy is higher than the energy of the ground
state.
Decay is when an atom in an excited state undergoes a transition
to the ground state loses energy by emitting a photon whose energy
corresponds to the difference in energy between the two states.
The difference in energy E between any two orbits or energy
levels is E = En1 En2 , where n1 is the final orbit and n2 the
initial orbit.Substituting in Bohrs equation gives E = hc (1/n21
1/n22).
Since E = hc/, then 1/ = (1/n21 1/n22). This is the same
equation that Rydberg obtained experimentally except for the
negative sign, which indicates that energy is released as the
electron moves from orbit n2 to orbit n1 because orbit n2 is at a
higher energy than orbit n1 .
1ST E.L.2ND E.L.3RD E.L.4TH E.L.5TH E.L.6TH
E.L.Lyman(UV)Paschen(IR)Balmer(visible)~~~~~~
Continuous and Line Spectra4000 Ao500060007000lightNaHCaHg400
450 500 550 600 650 700 750 nmVisiblespectrum
l (nm)9Spectroscopy is a method of identifying unknown
substances from their spectra. Because all materials have unique
spectra, you can think of these spectra as being molecular
fingerprints.
Line spectra of selected elements showing the emission spectra
or fingerprint of that element.
The unique set of lines produced is due to the fact that
electrons are falling from different excited states in the atoms to
the ground state.
Objects at high temperature emit a continuous spectrum of
electromagnetic radiation.
A different kind of spectrum is observed when pure samples of
individual elements are heated.
When the emitted light is passed through a prism, only a few
narrow lines, called a line spectrum, are seen rather than a
continuous range of colors.
Using Plancks equation, the observation of only a few values of
(or ) in the line spectrum meant that only a few values of E were
possible only states that had certain values of energy were
possible or allowed.
Any given element has both a characteristic emission spectrum
and a characteristic absorption spectrum, which are complementary
images.
Emission spectrum: emission of light by atoms in excited states
Absorption spectrum: absorption of light by ground-state atoms to
produce an excited state
Emission and absorption spectra form the basis of spectroscopy,
which uses spectra to provide information about the structure and
composition of a substance or an object
Flame Emission SpectraPhotographs of flame tests of burning
wooden splints soaked in different salts.Include link to web
pagehttp://www.unit5.org/christjs/flame%20tests.htm
methane gaswooden splintstrontium ioncopper ionsodium ioncalcium
ion10This technique is called emission spectroscopy.
Copyright 2007 Pearson Benjamin Cummings. All rights
reserved.
11Fireworks
Copyright 2007 Pearson Benjamin Cummings. All rights
reserved.12The chemistry of fireworks Colors of fireworks due to
atomic emission spectra A typical shell used in a fireworks display
contains gunpowder to propel the shell into the air and a fuse to
initiate a variety of redox reactions that produce heat and small
explosions Thermal energy excites the atoms to higher energy
states, and as they decay to lower energy states, the atoms emit
light that gives the familiar colors
electromagnetic radiation (i.e., light) -- -- EBwaves of
oscillating electric (E)and magnetic (B) fields source is vibrating
electric charges Characteristics of a Wavefrequency: the number of
cycles perunit time (usually sec)amplitude A crest wavelength l
trough-- unit is Hz, or s1 or 1/s
radio waves IR visible UV X-rays gamma rays cosmic rays
electromagnetic spectrum: contains all of the types oflight that
vary according to frequency and wavelengthROYGBVlarge l low f low
energy small l high f high energy -- visible spectrum ranges from
only ~400 to 750 nm (a very narrow band of spectrum)microwaves 400
nm 750 nm
Increasing wavelengthIncreasing frequencyIncreasing energySome
light humor
WavesLow frequencyHigh frequencyAmplitudeAmplitudelong
wavelength lshort wavelength lBoth travel at the same speedthe
speed of light!18WavesLow frequencyHigh
frequencyAmplitudeAmplitudelong wavelength lshort wavelength l60
photons162 photonsLow energyHigh energy19Einsteins photons of light
were individual packets of energy that had many characteristics of
particles.
Einsteins hypothesis that energy is concentrated in localized
bundles was in sharp contrast to the classical notion that energy
is spread out uniformly in a wave.Red and Blue Light
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page
325Photons- particle of light that carries a quantum of
energy20
Albert Michelson (1879)-- first to get an accuratevalue for
speed of light The speed of light ina vacuum (and in air)is
constant:c = 3.00 x 108 m/s c = f l-- Equation: Albert
Michelson(18521931)
c = 671 x 106 mph E = h fIn 1900, Max Planck assumedthat energy
can be absorbedor released only in certaindiscrete amounts, which
hecalled quanta. Later, Albert Einstein dubbeda light particle that
carried aquantum of energy a photon. -- Equation:E = energy,h =
Plancks constantin J = 6.63 x 1034 Js (i.e., J/Hz)Max
Planck(18581947)Albert Einstein(18791955)
A radio station transmitsat 95.5 MHz (FM 95.5).Calculate the
wavelengthof this light and the energyof one of its photons.c = f
l
= 3.14 mE = h f= 6.63 x 1034 J/Hz (95.5 x 106 Hz)= 6.33 x 1026
J
3.00 x 108 m/s= 95.5 x 106 Hzquantum mechanical modelelectron
cloud modelcharge cloud modelSchroedinger, Pauli, Heisenberg, Dirac
(up to 1940):According to the QMM, we never know for certain where
the e are in an atom, but the equations of the QMM tell us the
probability that we will find an electron within a certain distance
from the nucleus.
Electron Cloud ModelOrbital (electron cloud) instead of
orbitsRegion in space where there is 90% probability of finding an
electroncreates unique 3-D shapesCourtesy Christy Johannesson
www.nisd.net/communicationsarts/pages/chem
Electron Probability vs. DistanceElectron Probability
(%)Distance from the Nucleus (pm)100150200250500010203040
Orbital90% probability offinding the electronModels of the Atom
ReviewDaltons model (1803)Thomsons plum-pudding model
(1897)Rutherfords model (1909)Bohrs model (1913)Charge-cloud model
(present)Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd
Edition, 1990, page 125Greek model(400 B.C.)
+-----eee++++++++eeeeeee"In science, a wrong theory can be
valuable and better than no theory at all."- Sir William L.
BraggModels of the AtomDescription: This slide shows he evolution
of the concept of the atom from John Dalton to the present.Basic
Concepts The model of the atom changed over time as more and more
evidence about its structure became available. A scientific model
differs from a replica (physical model) because it represents a
phenomenon that cannot be observed directly.Teaching Suggestions
Use this slide as a review of the experiments that led up to the
present-day view of the atom. Ask students to describe the
characteristics of each atomic model and the discoveries that led
to its modification. Make sure that students understand that the
present-day model shows the most probable location of an electron
at a single instant. Point out that most scientific models and
theories go through an evolution similar to that of the atomic
model. Modifications often must be made to account for new
observations. Discuss why scientific models, such as the atomic
models shown here, are useful in helping scientists interpret heir
observations.QuestionsDescribe the discovery that led scientists to
question John Daltons model of the atom ad to favor J.J. Thomsons
model.What experimental findings are the basis for the 1909 model
of the atom?What shortcomings in the atomic model of Ernest
Rutherford led to the development of Niels Bohrs model?A friend
tells you that an electron travels around an atoms nucleus in much
the same way that a planet revolves around the sun. Is this a good
model for the present-day view of the atom? Why or why not?Another
friend tells you that the present-day view of an electrons location
in the atom can be likened to a well-used archery target. The
target has many holes close to the bulls-eye and fewer holes
farther from the center. The probability that the next arrow will
land at a certain distance from the center corresponds to the
number of holes at that distance. Is this a good model for the
present-day view of the atom? Why or why not?Suppose that, it the
future, an apparatus were developed that could track and record the
path of an electron in an atom without disturbing its movement. How
might this affect the present-day model of the atom? Explain your
answer.How does developing a model of an atom differ from making a
model of an airplane? How are these two kinds of models the
same?Drawing on what you know in various fields of science, write a
general statement about the usefulness of scientific models.
Bragg and his father, Sir W.H. Bragg, shared the 1915 Nobel
prize in physics for studies of crystals with X-rays.Models of the
Atom TimelineDaltons model (1803)Thomsons plum-pudding model
(1897)Rutherfords model (1909)Bohrs model (1913)Charge-cloud model
(present)Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd
Edition, 1990, page 125Greek model(400 B.C.)1800 1805
..................... 1895 1900 1905 1910 1915 1920 1925 1930 1935
1940 19451803 John Dalton pictures atoms astiny, indestructible
particles, with no internal structure.1897 J.J. Thomson, a
Britishscientist, discovers the electron,leading to his
"plum-pudding" model. He pictures electronsembedded in a sphere
ofpositive electric charge.1904 Hantaro Nagaoka, aJapanese
physicist, suggests that an atom has a centralnucleus. Electrons
move in orbits like the rings around Saturn.1911 New Zealander
Ernest Rutherford statesthat an atom has a dense,positively charged
nucleus. Electrons move randomly in the space around the
nucleus.1913 In Niels Bohr'smodel, the electrons move in spherical
orbits at fixed distances from the nucleus.1924 Frenchman Louis de
Broglie proposes thatmoving particles like electronshave some
properties of waves. Within a few years evidence is collected to
support his idea.1926 Erwin Schrdinger develops
mathematicalequations to describe the motion of electrons in atoms.
His work leads to the electron cloud model.1932 James Chadwick, a
British physicist, confirms the existence of neutrons, which have
no charge. Atomic nuclei contain neutrons and positively charged
protons.+-----eee++++++++eeeeeee
Models of the AtomDescription: This slide shows he evolution of
the concept of the atom from John Dalton to the present.Basic
Concepts The model of the atom changed over time as more and more
evidence about its structure became available. A scientific model
differs from a replica (physical model) because it represents a
phenomenon that cannot be observed directly.Teaching Suggestions
Use this slide as a review of the experiments that led up to the
present-day view of the atom. Ask students to describe the
characteristics of each atomic model and the discoveries that led
to its modification. Make sure that students understand that the
present-day model shows the most probable location of an electron
at a single instant. Point out that most scientific models and
theories go through an evolution similar to that of the atomic
model. Modifications often must be made to account for new
observations. Discuss why scientific models, such as the atomic
models shown here, are useful in helping scientists interpret heir
observations.QuestionsDescribe the discovery that led scientists to
question John Daltons model of the atom ad to favor J.J. Thomsons
model.What experimental findings are the basis for the 1909 model
of the atom?What shortcomings in the atomic model of Ernest
Rutherford led to the development of Niels Bohrs model?A friend
tells you that an electron travels around an atoms nucleus in much
the same way that a planet revolves around the sun. Is this a good
model for the present-day view of the atom? Why or why not?Another
friend tells you that the present-day view of an electrons location
in the atom can be likened to a well-used archery target. The
target has many holes close to the bulls-eye and fewer holes
farther from the center. The probability that the next arrow will
land at a certain distance from the center corresponds to the
number of holes at that distance. Is this a good model for the
present-day view of the atom? Why or why not?Suppose that, it the
future, an apparatus were developed that could track and record the
path of an electron in an atom without disturbing its movement. How
might this affect the present-day model of the atom? Explain your
answer.How does developing a model of an atom differ from making a
model of an airplane? How are these two kinds of models the
same?Drawing on what you know in various fields of science, write a
general statement about the usefulness of scientific models.
Timeline: Wysession, Frank, Yancopoulos Physical Science
Concepts in Action, Prentice Hall/Pearson, 2004 pg 114Shapes of s,
p, and d-Orbitals
each holds 2 electrons (s2)each of 5 orbitals holds 2 e - = 10
total d electrons (d10)each of 3 orbitals holds 2 e - = 6 total p
electrons (p6)f-orbitals
each of 7 orbitals hold 2 e- = 14 e-How many g-orbitals could
exist and how many e- could they hold?theoretical g-orbitals
each of 9 orbitals hold 2 e- = 18 e-Relative Sizes 1s and 2s1s
2sZumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page
334Remember: s, p, d, and f refer to the orbital shapeAs you add
more e-, progressively larger orbitals are needed to accommodate
all the of e-
Copyright 2006 Pearson Benjamin Cummings. All rights
reserved.Orbitals overlap each other as you get farther from the
nucleusOrbital Filling Videos, p, and d-orbitals overlap
s orbitals:Each holds 2 electrons(outer orbitals ofGroups 1 and
2)p orbitals:Each of 3 sets holds 2 electrons = 6 electrons(outer
orbitals of Groups 3 to 8)d orbitals:Each of 5 sets holds 2
electrons= 10 electrons(found in elements in third periodand
higher)
s blockp blockd blockf block67Periodic
Patterns1s2s3s4s5s6s7s3d4d5d6d1s2p3p4p5p6p7p4f5f1234567(n-1)(n-2)n34Sections
of Periodic Table to Know f-blocks-blockd-blockp-blockEnergy Level
Diagram of a Many-Electron Atom ArbitraryEnergy Scale1818328821s2s
2p3s 3p4s 4p 3d 5s 5p 4d6s 6p 5d 4fNUCLEUSOConnor, Davis, MacNab,
McClellan, CHEMISTRY Experiments and Principles 1982, page 177Each
orbital can only hold 2 e- Start from the bottom and add e-You dont
have to memorize the orderjust start at the beginning and fill in
e-
s-block1st Period (row)1s11 e- in 1s orbital
Periodic PatternsExample - HydrogenCourtesy Christy Johannesson
www.nisd.net/communicationsarts/pages/chem
38Writing Electron Configurations: Where are the e?
(probably)AsHHeNAlLiTiXe1s22s23p62p64s23d103s24p61s11s21s22s11s22s22p31s22s23p12p63s21s22s23p62p64s23d23s21s22s23p62p64s23d103s24p31s22s23p62p64s23d103s24p65s24d105p6Filling
OrderNotable e- configuration exceptions:
CuCr1s22s23p62p64s13d53s21s22s23p62p64s13d103s2Mo and W
(theoretically Sg) will behave similarly to CrAg and Au
(theoretically Rg) will behave similarly to CuWhy? Generally, full
orbitals and half-filled orbitals have lower energies and are thus
more stable. So, while we might expect Cr to end with 4s23d4,
promoting an s e- to the d orbitals creates two half-filled
shells.[Ar]
4s23d104p2Periodic PatternsExample - GermaniumCourtesy Christy
Johannesson www.nisd.net/communicationsarts/pages/chem
Ge72.613241Electron Configuration Battleship (your
shots)Electron Configuration Battleship (your ships and opponents
shots)SShorthand Electron Configuration (S.E.C.) To write S.E.C.
for an element: 1. Put symbol of noble gas that precedeselement in
brackets.2. Continue writing e config. from that pointCoInClRb[ Ne
] 3s2 3p4 [ Ar ] 4s2 3d7 [ Kr ] 5s2 4d10 5p1 [ Ne ] 3s2 3p5 [ Kr ]
5s1 Ge72.6132Shorthand Configuration Review[Ar] 4s2Electron
configurationElement symbol[Ar] 4s2 3d3[Rn] 7s2 5f14 6d4[He] 2s2
2p5[Kr] 5s1 4d10[Kr] 5s2 4d10 5p5[Kr] 5s2 4d10 5p6 or [Xe][He]
2s22p63s23p64s23d6CaVSgFAgIXeFe[Ar] 4s23d6Periodic Patterns
ReviewPeriod # (1-7)primary energy level (n) (subtract for d &
f)Group # (1-8excluding d block)total # of valence e-Column within
Sublevel block# of e- in sublevel
Courtesy Christy Johannesson
www.nisd.net/communicationsarts/pages/chem
46Three Principles about Electronsaufbau Principle:for
equal-energy orbitals (p, d)each must have one e beforeany take a
second Pauli Exclusion Principle:e will fill the
lowest-energyorbital available Hunds Rule:two e in same orbitalhave
different spins 1s22s23p62p64s23d103s2Friedrich Hund
Wolfgang Pauli
OOrbital Diagrams show spins of e and orbital
location1s2s2p3s3p1s2s2p3s3pP1s22s22p41s22s22p63s23p3HAVE
MOREENERGYARE FURTHERFROM NUCLEUSANDThe Importance of
Electronsorbitals: In a generic e config (e.g., 1s2 2s2 2p6 3s2
3p6): regions of space where an e may be found # of energy level #
of e in those orbitals coefficient superscript In general, as
energy level # increases, e
Shorthand ConfigurationS 16e-Valence Electrons(Highest energy
level)Kernel (Core) ElectronsS16e-[Ne] 3s2
3p41s22s22p63s23p4Electron ConfigurationLonghand Configuration
Courtesy Christy Johannesson
www.nisd.net/communicationsarts/pages/chem
S32.0661653INVOLVED INCHEMICALBONDINGkernel (core) electrons:
valence electrons: in inner energy level(s);close to nucleusin
outer energy levelHe: Ne:Ar: Kr:1s2 [ He ] 2s2 2p6 [ Ne ] 3s2 3p6 [
Ar ] 4s2 3d10 4p6 Noble gas atoms have FULL valence orbitals. They
are stable, low-energy, and unreactive.(2 valence e) (8 valence e)
(8 valence e) (8 valence e)
octet rule: the tendency for atoms to fill valence orbitals
completely with 8 e (outer E level) Other atoms want to be like
noble gas atoms doesnt apply to He, Li, Be, B (which require 2) or
to H (which requires either 0 or 2)duet rule fluorine atom, F 9 p+,
9 e ** So, they lose or gain e...gain 1 e or lose 7 e-? 9 p+, 10 e
F is more stable asan F ion Fchlorine atom, Cl 17 p+, 17 e 17 p+,
18 e Cl is more stable asa Cl ion ClHow to be likea noble gas?
[He] 2s22p5[Ne] 3s23p5gain 1 e or lose 7 e-? lithium atom, Li 3
p+, 3 e lose 1 e or gain 7 e-? 3 p+, 2 e Li is more stableas the
Li+ ion. Li+sodium atom, Na 11 p+, 11 e lose 1 e or gain 7 e-? 11
p+, 10 e Na is more stableas Na+ ion Na+How to be likea noble
gas?
[He] 2s1[Ne] 3s11+Know charges on these columns of Table: Group
1:Group 2:Group 3:Group 5:Group 6:Group 7:Group 8:
2+3+32101+2+3+3210also called oxidation numbers. spdf67Periodic
Patterns and Charge Trends1s1234567+1+2(n-2)n+3- 3- 2- 1Variable
Charge58Naming Ions e.g., Ca2+ Cs+Al3+Cations use element name and
then say ion e.g., S2P3N3O2ClAnions change ending of element name
to ideand then say ioncalcium ion cesium ionaluminum ionsulfide
ionphosphide ionnitride ionoxide ionchloride ion1
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