Chapter 5 Periodicity and Atomic Structure
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Transcript of Chapter 5 Periodicity and Atomic Structure
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PerodicityPerodicityand Atomic Structureand Atomic Structure
Chapter 5
Chapter 5 2
Mendeleevs Periodic Table In the 1869,Dmitri Mendeleevproposed that
the properties of the chemical elements repeatat regular intervals when arranged in order ofincreasingatomic mass.
Mendeleev is the architect of the modernperiodic table.
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Chapter 5 3
Prediction of New Elements
Mendeleev noticed that there appeared to be some elementsmissing from the periodic table so he left spots for undiscovered
elements.
Based on his periodic table, he accurately predicted the properties
of some of these the unknown elements before their discovery.
Chapter 5 4
Periodic Trends The arrangement of the periodic table means that
the physical properties of the elements follow a
regular pattern.
Some trends include:
Atomic Radius (end of Chapter 5)
Ionization Energy (Chapter 6) Electron Affinity (Chapter 6)
Electronegativity (Chapter 7)
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Chapter 5 5
The Structure of Atoms
Plum Pudding Model Tootsie Pop Model
Chapter 5 6
The Dual Nature of Light:The Particle and The Wave
A great animation of the idea of light as waves and particles can be found here:
http://video.google.com/videoplay?docid=-4237751840526284618&q=quantum
From the time of the ancient
Greeks, people have thought of
light as a stream of tiny
particles - like marbles or
billiard balls
Thomas Young (in 1807)
performed the now classic
double slit experiment to testthis theory.
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Chapter 5 7
As evidenced by the double slit
experiment, light travels through space as
a wave, similar to an ocean wave.
Wavelength () is the distance light
travels in one cycle.
Frequency () is the number of wave
cycles completed each second.
Amplitude is the height measured from
the center of the wave. The square of this
value gives theIntensity.
Velocity (c) =
The Dual Nature of Light: The Wave
Light has a constant velocity (c) of3.00 108 m/s.
Chapter 5 8
Frequency - Wavelength
The red light in a laser pointer comes from a
diode laser that has a wavelength of about 632
nm. What is the frequency of the light?
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Chapter 5 9
The Dual Nature of Light: The Particle
In 1900, Max Planck proposedthat radiant energy is notcontinuous, but is emitted assmall bundles of energy
This is the quantum concept.
Planck determined that the energy of an emitted bundle is directlyproportional to the frequency of the emitted light times a constant (h)
E= h =
hc h = 6.626 10 34Js
Chapter 5 10
The Photoelectric Effect The Photoelectric Effect shows that when light is
shined on the surface of a metal, electrons are ejected
and a current can be detected.
Albert Einstein used Plancks idea ofquanta to
explain this phenomona.
If the light were waves only, then the energy of that
radiation would depend only on the intensity of the
light.
Einstein hypothesized that electrons are only ejected
if the frequency of the light exceeds a threshold
value specific to the metal, regardless of light
intensity
Even at low light intensities, electrons are ejected
immediately if the frequency exceeds the threshold
Millikan tested Einsteins hypothesis in 1914 and
proved that is was correct
Einstein wins the Nobel Prize! Yea!
Figure above from http://en.wikipedia.org/wiki/Photoelectric_Effect
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Chapter 5 11
Energy of Emission
For red light with a wavelength of about 632
nm, what is the energy of a single photon and
one mole of photons?
Chapter 5 12
Louis de Broglie suggested that if light can behave
like matter (particles) then matter can behave as light.
This concept is called wave particle duality.
For Light
=
h
mc
For a Particle
=
h
mv
Wave Particle Duality
E = mc2 E = h
h = 6.626 x 10-34 kg m2 / s h = 6.626 x 10-34 J s
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Chapter 5 13
Wave Particle Duality
How fast must an electron be moving if it has a
de Broglie wavelength of 551 nm?
How fast must an electron be moving if it has a
de Broglie wavelength of 631 nm?
me = 9.109 x 1031 kg
Chapter 5 14
The Electromagnetic Spectrum The complete electromagnetic spectrum (all possible
wavelengths and frequencies) is an un-interruptedband, orcontinuous spectrum.
The radiant energy spectrum includes most types ofradiation, most of which are invisible to the humaneye.
ROY G BIV
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Chapter 5 15
The Atomic Spectra
When a particular element isheated and the light emitted isfocused and passed through aprism, the resulting breakdownresults in anon-continuousspectrum or adiscrete linespectrum
This spectrum is called theelementsatomic spectrum
The discrete lines indicate alight is emitted or absorbedin a series of discretefrequencies rather thancontinuously
Each element has its own,unique spectrum
Chapter 5 16
In 1913, Niels Bohr suggested a new model of the atom that explained why
hydrogen had a discrete line spectrum rather than a continuous spectrum.
Bohr's basic theory: electrons in atoms can only be at certain energy levels,
and they can give off or absorb radiation only when they jump from one level
to another.
The Bohr Model of the Atom
In his model that an atom consists of an
extremely dense nucleus that is surrounded by
electrons that travel in set orbits around the
nucleus.
He hypothesized that the energy possessed by
these electrons and the radius of the orbits arequantized, meaning it is limited to specific
values and is never between those values.
These orbits were of varying energies,
dependent on their distance from the nucleusThe Gobstopper Model
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Chapter 5 17
Quantized versus Continuous
The quantum concept means that the energy of the electron andits radius of orbit around the nucleus is limited to specific values
(and cannot be anywhere in between!).
If these values were continuous, they would be free to have any
value.Energy
E1
E2
E3
hv = E3 E2
hv = E3 E1
hv = E2 E1
Chapter 5 18
Atomic Spectra:
The Balmer-Ryberg Equation In 1885, Johann Balmer determined that the pattern of the atomic
spectra of hydrogen could be predicted by a mathematical formula.
Balmer determined that the wavelength or frequency of the lines inthe spectrum could be expressed by the following equations:
In addition to the lines seen in the visible region (Balmer series),there are additional sets of lines found in the UV region (Lymanseries) and the IR region (the Paschen series).
All conform to the above equations.
(E =RH 1ni21
nf2
- )RHis the Ryberg constant (2.18 x 10
-18 J)
ni and nfare integers with nf> ni
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Chapter 5 19
The Bohr Theory: Problems
Bohrs theory only works forhydrogen atoms.
Once you have more than oneelectron, the calculations for theelectron energy and the orbit radii
breakdown.
Therefore, a new theory needed to bedeveloped for multi-electronelements.
Bohrs theory did make twoimportant contributions:
It suggested a reasonable explanationfor the discrete line spectrum of the
elements It introduced the idea of quantized
electron energy levels (orbits!)
Chapter 5 20
The Quantum Mechanical Model of the Atom
In the 1920s Erwin Schrdinger appliedthe principles of wave mechanics toatoms and developed the Quantum
Mechanical Model of the Atom
Basically, Schrdinger said to give up on theidea of literal orbits for the electrons andinstead concentrate on the electron as awave.
This theory builds on Bohrs idea ofquantized energy levels (orbits) and addsadditional requirements for electronlocation and energy.
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Chapter 5 21
Quantum Mechanics
Werner Heisenberg (19011976): supported this
idea by showing that it is impossible to know (or
measure) precisely both the position and velocity
(or the momentum) at the same time.
The simple act of seeing an electron would change its energy and
therefore its position. Think about pinpointing a fly on the wall. What happens when you try to
swat it?
)()4()(:positionselectron'inyUncertaint
4
))((:PrincipletyUncertainHeisenberg
m
hx
hmx
Chapter 5 22
Working with Heisenbergs Principle, Schrdinger
developed a compromise which calculates both the energy
of an electron and the probability of finding an electron at
any point in the molecule.
This is accomplished by solving the Schrdinger equation,
resulting in the wave function, .
These regions were termed orbitals
Quantum Mechanics
Wave
Equation
Wave Function
or Orbital ()
Probability of finding
an electron in aregion of space (2)
solve
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Chapter 5 23
Quantum Numbers
The wave function contains three variablesknown as the Quantum Numbers whichdescribe the size, energy, shape and position ofthe orbitals.
n is the principal energy level (Bohrs Orbits!)
l is the sublevel
ml is the orbital
These numbers serve as an address of the
probable location of the electron.
Chapter 5 24
Principal Quantum Number (n) The Principal Quantum Number (n)
provides info about the distance ofthe electron from the nucleus As n increases, the number of allowed
orbitals also increases as does the sizeof those orbitals.
This increased size allows the electronto reside further from the nucleus
As the electron moves away from thenucleus its energy increases, therefore nalso indicates the energy of electrons
We often state that electrons andorbitals denoted by the same n valueare in the same shell
Allowed Values: n = 1, 2, 3, ... never 0
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Chapter 5 25
Angular-Momentum Quantum Number (l)
The angular momentumquantum number (l) defines the
three dimensional shape of the
orbitals found within a particular
shell
These discrete sets of orbitals
are calledsublevels
The number of sublevels within
a shell is equal to the principal
quantum number (n)
lValue: 0 1 2 3
Letter Used: s p d f
increasing energy
Allowed Values:
l= 0, 1, 2 ... n - 1
Chapter 5 26
Magnetic Quantum Number (ml) The Magnetic Quantum Number (ml) describes the orientation in 3D
space of the sublevel, thereby denoting a specific orbital.
The number of orientations (and therefore orbitals) per sublevel is
determined by the equation:
2l + 1
2l + 1
3 (f)2 (d)1 (p)0 (s)l
Allowed Values: ml= - l, , + l
x axis
y axis
z axis
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Chapter 5 27
Spin Quantum Number (ms)
ThePauli Exclusion Principle statesthat no two electrons can have the
same four quantum numbers
This results in no more than doubleoccupancy in any one orbital
only two electron per orbital and
they must have opposite spins
Chapter 5 28
Electron Occupancy in Sublevels ThePauli Exclusion Principle states that an orbital can hold
up to two electrons
The maximum number of electrons in each of the sublevels
depends on the number of orbitals within that sublevel:
Thes sublevel holds a maximum of 2 electrons (1 orbital).
Thep sublevel holds a maximum of 6 electrons (3 orbitals).
The dsublevel holds a maximum of 10 electrons (5 orbitals).
Thefsublevel holds a maximum of 14 electrons (7 orbitals)
The maximum electrons per principle quantum level (n) is
obtained by adding the maximum number of electrons in eachsublevel.
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Chapter 5 29
Quantum Number Combinations
Why cant an electron have the following quantum numbers?
(a) n = 2, l= 2, ml= 1 (b) n = 3, l= 0, ml= 3
(c) n = 5, l= 2, ml= 1
Give orbital notations for electrons with the following
quantum numbers:
(a) n = 2, l= 1, ml= 1 (b) n = 4, l= 3, ml= 2
(c) n = 3, l= 2, ml= 1
Chapter 5 30
Shapes of the Orbitals Each orbital has a specific shape determined by its
angular momentum quantum number (l).
As you increase the principle quantum number (n), theorbitals increase in size but not shape!
Remember, these orbitals represent a region in spacewhere there is a high probability of finding the electron.
These are not discrete locations!
They are also not pictures of the path that an electron followsaround a nucleus
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Chapter 5 31
l= 0 : The s Orbitals
The s orbitals are spherical, meaning the probability of finding the
electron depends only on distance from the nucleus, not direction.
The value of2 is greatest near the nucleus then drops off as you
move away.
It never reaches zero however so there is technically no definite
boundary to the atom.
Chapter 5 32
l= 1 : The p Orbitals The p orbitals are dumbbell shaped with their electron density concentrated in
identical lobes residing on opposite sides of a nodal plane.
This shape means that a p electron will never be found near the nucleus.
The two lobes of a p orbital have different phases (are opposite in sign) which
becomes important in bonding among atoms.
The three orientations (ml= -1, 0, +1) are 90 differentials along the x, y and
z axes.
The orbitals are designated px (along the x axis), py (along the y axis), and pz(along the z axis)
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Chapter 5 33
l= 2 : The d Orbitals
Chapter 5 34
Effective Nuclear Charge (Zeff) The nuclear charge (Z) of an atom is
determined by the number of protons foundin the nucleus
It is felt by the electrons as an attraction
Multiple electrons in an atom lead to ashielding effect on the outer electrons.
This electron shielding (S) leads to energy
differences among orbitals within a shell.
Net nuclear charge felt by an electron is
called the effective nuclear charge (Zeff).
Zeff= Z + S
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Chapter 5 35
Effective Nuclear Charge (Zeff)
Note that the 3d
sublevel is actually
higher in energy than
the 4s sublevel.
WHY??
Chapter 5 36
The Modern Periodic Table
H.G.J. Moseley discovered that the nuclear charge increased by 1 for each element in theMendeleevs table.
He concluded that the changingatomic number rather than the changing mass explained therepeating trends of the elements
Theperiodic law states that the properties of elements recur in a repeating pattern when arrangedaccording to increasing atomic number.
With the introduction of the concept of electron energy levels byNiels Bohr, the periodic table tookits current arrangement.
http://www.webelements.com/
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Chapter 5 37
Electron Configurations
Many of an elements chemical
properties depend on its electron
configuration
The electron configuration of an atom is a shorthand method ofwriting the location of electrons by sublevel.
The principal quantum level (n) is written first, followed by the
letter designation of the sublevel (l) then a superscript with the
number of electrons in the sublevel.
There are rules for the order and manner that each sublevel isfilled called the Aufbau Principle.
Chapter 5 38
Electron Configurations: Aufbau Principle
Pauli Exclusion Principle:No two electrons in an atom can
have the same quantum numbers (n, l, ml, ms).
Hunds Rule: When filling orbitals in the same sublevel,
maximize the number of parallel spins (so fill then pair!).
Rules of Aufbau Principle:
1. Lower n orbitals fill first.
2. Each orbital holds two electrons; each with different ms.
3. Half-fill degenerate orbitals before pairing electrons.
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Chapter 5 39
Using the Periodic Table for Electron Configurations
The periodic table took on its current shape once the quantum model of the
atom was developed.
You can use it to fill up your sublevels and orbitals to build your electron
configurations.
Chapter 5 40
Writing Electron Configurations Step 1: Locate the element on the periodic table.
Step 2: Determine the number of electrons the element
has: Iron has 26 electrons
Step 3: Starting at the beginning of the Periodic Table,
move left to right across the periods, filling each sublevel
with electrons until you reach the location of your
element:
Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
Step 4: Check that the sum of the superscripts equals the
atomic number of iron (26).
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Chapter 5 41
An Alternative Method
1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d7s 7p
Increasing
Energy
[He]
[Ne]
[Ar]
[Kr]
[Xe]
[Rn]
Core
Chapter 5 42
Li 1s2 2s1
1s 2s
Be 1s2 2s2
1s 2s
B 1s2 2s2 2p1
1s 2s 2px 2py 2pz
C 1s2 2s2 2p2
1s 2s 2px
2py
2pz
Writing Electron Configurations
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Chapter 5 43
Writing Electron Configurations
N 1s2 2s2 2p3
1s 2s 2px 2py 2pz
O 1s2 2s2 2p4
1s 2s 2px
2py
2pz
Ne 1s2 2s2 2p5
1s 2s 2px
2py
2pz
S [Ne] [Ne] 3s2 3p4
3s 3px
3py
3pz
Chapter 5 44
Give the ground-state electron configurations for:
Ne (Z= 10) Mn (Z= 25) Zn (Z= 30)
Eu (Z= 63) W (Z= 74) Lw (Z= 103)
Identify elements with ground-state configurations:
1s2
2s2
2p4
1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d6
1s2 2s2 2p6 [Ar] 4s2 3d1 [Xe] 6s2 4f14 5d10 6p5
Writing Electron Configurations
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Chapter 5 45
Exceptions to the Filling Order
When filling the d sublevel, exceptions occur for thechromium (Cr) and copper (Cu) families:
4s 3d 4p
4p4s 3d
4s 3d 4p
4s 3d 4p
Cr
Cu
Chapter 5 46
Periodic Trends The arrangement of the periodic table means that
the physical properties of the elements follow a
regular pattern.
Some trends include:
Atomic Radius (end of Chapter 5)
Ionization Energy (Chapter 6) Electron Affinity (Chapter 6)
Electronegativity (Chapter 7)
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Chapter 5 47
Atomic Radius
An atomsatomic radius is the distance from the nucleus to theoutermost electrons.
Why do you think the radiusincreases in this way?