Transcript No Slide Title

e- Strucure and Periodicity
Mon 11/14 Sec 6.1 EMR
– WS #1 (1-6)
B-1 11/15 Sec 6.2 Bohr Model H atom and Quantum Mech
– WS #1 (7-9) and #3 assorted questions
B-2 11/17 Sec 6.3-4 Quantum Mech and e- config
– WS #3 assorted questions
M Lab
B-1 6.5-7 Aufbau and e- config
– WS #2 all - #3 assorted questions
• B-2 6.8 Per Tab and Rev Chap
– WS #4-8)
• B-1 Test
Electromagnetic radiation
• Emr (including visible light) can be thought
of as moving in waves
• Examples: TV and radio waves, microwaves,
infrared, visible light, ultraviolet light, xrays, and cosmic rays
• The waves have certain characteristics that
are described by math
Characteristics of
Electromagnetic Radiation (EMR)
• Radiation (such as light) can be pictured as
waves traveling through space
Characteristics of
Electromagnetic Radiation (EMR)
• Speed of EMR (waves) = c
– Physical constant in nature
• C = 3 x 108 m/s for any emr
• Amplitude
– Height of wave
• Wavelength = l
– Dist between waves
– units = nm = 1 x 10-9m
• Frequency = n
– Unit = hertz (waves/s or s-1)
– Freq and wavelength are inversely related
• n goes
l goes
Wavelengths of Visible Light
Math Relationships
• Speed of light = wavelength x frequency
– c = ln
l and n are inversely related
When 1 goes up, the other goes down
– The units of c and l must be the same
• C in m/s
l is usually in nm, convert it to m
• How many m in 350 nm?
350 nm x
1 m = 3.50 x 10-7 m
109 nm
What is the frequecy of red light with a
with a wavelength of 415 nm?
• 415 nm x 1 m
= 4.15 x 10-7 m
•
109 nm
 c = ln
• n = c/l = 3 x 108 m/s = 7.23 x 1014 s-1
•
4.15 x 10-7 m
Continuous Spectrum
• A prism bends light and splits the light into all the diff
wavelengths that make it up
• White light gets split into the diff colors of the rainbow ROYGBIV
Bright Line Spectra
A gas in a cathode ray tube glows with a color unique to
each gas (like a fingerprint) The light emitted from a
cathode ray tube, when sent thru a prism, produces a
bright line spectrum - diff for every gas
Dualism Of Light
– Basic particle = photon
– Each emr has a diff
amount of energy
• Blue has most energy, red
has least energy
• Wavelength and energy
are inversely related
•Higher Energy
• Light can be modeled as
waves OR particles
Photoelectric Effect
the emission of electrons by
substances, especially metals,
when light falls on their surfaces.
Photons
 The quantum (basic unit) of electromagnetic
energy
 generally regarded as a discrete particle
 zero mass
 no electric charge
 an indefinitely long lifetime.
Max Planck’s Equation
• Planck quantized light – particle model
– Determined the energy of it’s photons
– E = hn = hc
•
•
•
•
l
E = energy of photon of specific frequency
n = frequency
l = wavelength
H = Planck’s constant = 6.626 x 10-34 Js
– Direct relation n E (energy)
– Inverse relation l
E (energy)
• Blue light has more energy than red light
Energy of a Photon
• Calc the energy of photon with l = 630 nm
– 630 nm x
1m
= 6.30 x 10-7m
1 x 109 nm
– DE = hc = (6.626 x 10-34 Js)(3 x 108 m/s)
l
= 3.15 x 10-19 J
6.30 x 10-7m
• Calc the energy in a mol of these photons
– E = 3.15 x 10-19 J x 1 KJ x 6.02 x 1023
103J
1 mol
= 1.90 x 102 KJ/mol
Relative Energy of Photons
Types of
Spectra
– Continuous
vs Bright Line
Emission
Rydberg Equation 1888
Neils Bohr Model of Atom
• e- maintain all their energy if they remain in any 1
energy level
• e- levels are quantized – e- in any energy level have
a specific amount of energy
• e-s go to lowest energy position available and are
said to be in the ground state
• e- can absorb a photon of specific energy and jump
up to a higher energy level (quantum leap) = excited
state
• Excited e-s will emit a photon and return to a lower,
more stable, energy level as soon as possible
Neils Bohr Model of Atom
• If the emitted photon is of light in the visible
range we can see it = bright line spectrum
• Each substance has its own unique estructure and  it’s own unique bright line
spectrum
Line Spectra of Hydrogen
Lyman series => ultraviolet
n > 1 to n = 1
Balmer series => visible light
n > 2 to n = 2
Paschen series => infrared
n > 3 to n = 3
Line Spectra of Hydrogen
•Electron Transitions
According to the energy diagram below for
the Bohr model of the hydrogen atom, if an
electron jumps from E1 to E2, energy is
absorbed
emitted
not involved
Energy of e- Transitions
• Energy of e- in an energy level
– En = -Rh Rh = Rydberg const = 2.18 x 10-18 J
n2 n = principle energy level (1,2,3,…)
– En = -2.180 x 10-18J
n2
• To calc energy or l of a transition
– Ehi – Elo = DE = hn = hc/l
Energy of Emitted Photons
 Energy of an emitted photon =

En = -Rh/n2
Rh = 2.180 x 10-18 j/s
 DE = hn = Ehi - Elo
 DE = hn = -Rh
-Rh
nhi2
nlo2 OR
DE = Rh 1
nlo2
1 = hn
nhi2
Rh = 2.18 x 10-18 J
h = 6.626 x 10-34 Js
Calc l and n for transition of
e-from n=6 to n=3
DE = Rh
1
1
nlo2 nhi2
DE = 2.18 x 10-18 J
Rh = 2.18 x 10-18 J
1
32
1
62
= 1.82 x 10-19J
DE = hn
1.82 x 10-19J = 6.626 x 10-34Js n
n = 2.74 x 1014 s-1
DE = hc/l
l = hc/ DE
= (6.626 x 10-34Js)(3 x 108m/s)/1.82 x 10-19J
= 1.09 x 10-6 m x 109 nm/m = 1092 nm
Quantum Mechanics
• De Broglie
– If wave can have particle characteristics particles
can have wave character
large mass = very small wave character or function
• Heisenberg
– Uncertainty principle
• Cannot determine both the position and movement of an
electron
• Probability cloud - cd 6.9
– Regions in space where e-s are most probably found cd 6.8
Wave Function = Y2
• Erwin Schroedinger
– Electrons occupy regions in space determined by
electromagnetic forces
– Wave function
• Set of equations that determine regions in space
• Quantum numbers
– 3 numbers
– When placed into equations the quantum numbers describe
areas in space occupied by electrons
Quantum Numbers
• Describe e- orbitals - regions in about the nucleus where eare most likely located
• 4 quantum numbers
– n = principle quantum # - values n =1,2,3,4,…∞
• Energy of electrons and distance from nucleus
– l = sublevel (azimuthal Q#) = shape of orbital (area) where e-s are
most likely found
• There are as many sublevels in a level as the n value for the level
• Values for l
–
–
–
–
l = 0 or s orbital is a sphere
l = 1 or p orbital is 2 lobed
l = 2 of d orbital is 4 lobed
l = 3 or f orbital is 8 lobed
Quantum Numbers
– ml = orbital (magnetic Q#)
•
•
•
•
Orientation of each sublevel in space
Each orientation is called an orbital
Each orientation (orbital) can hold 2 e-s
# e-s that fill the orbitals
–
–
–
–
s = 1 orbital holds 2 e-s
p = 3 " holds 6 "
d = 5 " holds 10 "
f = 7 " holds 14 "
– ms = spin quant #
– Pauli exclusion principle - if 2 e-s are in the same orbital they must
have opposite spins
– Ms values = + ½ or – ½
Levels, Sublevels and Orbitals
Levels, Sublevels, Orbitals, and # of e-s
Quantum Number Values
Orbitals
• region of probability of finding an electron
around the nucleus
• 4 types sublevels => s p d f
Atomic Orbitals
shape
s
p
d
f
•
•
# of orbitals(orientations)
per sublevel
spherical
1
dumbbell
3
4 lobes
5
8 lobes
7
# of Orbitals within a sublevel
– s=1
p=3 d=5 f=7
maximum of 2 electrons per orbital therefore
– Number of e-s per sublevel
– s=2
p = 6 d = 10 f = 14
Quantum Numbers based on
the periodic table
Probability Cloud s Orbitals
Probability Cloud
p Orbitals
s and p Orbitals
Probability Cloud d Orbitals
s, p, d, and f orbitals
s and p Orbitals
Energy of Orbitals
For prin. quant # n
1 < 2 < 3 < 4…
For sublevels
s < p < d < f
Overlaps occur between d sublevel and
next levels s sublevel
Double overlap between f sublevel and s
sublevels
Full and half full sublevels are extra stable
(lower energy)
Energy Ladder
Filling order for sublevels
•Single
overlap
between s
and d
sublevels,
double
overlap
between s
and f
sublevels
Energy Ladder
Filling order for sublevels
Blocks of Periodic Table
•1
•l =1
•2
•l =2
•3
•4
•5
•6
•n=
•7
•l =0
•l =3
Electron Structure and Periodic Table
What you need to know!
• What info do all Q #’s give us?
– What are their symbols?
– What do they represent?
– What are the 4 Q #’s for any electron?
• Capacities
– A level contains how many
• Sublevels? Orbitals? Electrons?
– A sublevel contains how many
• Orbitals? Electrons?
– An orbital contains how many electrons?
Aufbau Process
• Building process for atoms noting the position
of e-s by level #, sublevel letter, and the
number of e- in each as a superscript
• Start with H and work up the periodic table
• Hund's Rule - each orbital within a sublevel is
filled with 1 e- before any of the sublevels get
a second, paired, e• Pauli Exclusion Principle - if 2 e-s occupy the
same orbital they must have opposite spins
e- Configurations
Electronic Configurations
• The shorthand representation of the
occupancy of the energy levels (shells and
subshells) of an atom by electrons
• Show level number, sublevel letter, and the #
of e-s in the sublevel as a superscript after the
sublevel letter
Electron Filling Order Diagram
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
3d
4d 4f
5d 5f
6d
Electronic Configuration
H atom
1 electron
1s
e- configuration notation
H = 1s1
Electronic Configuration
He atom
2 electrons
1s
e- configuration notation
He = 1s2
Electronic Configuration
Li atom
3 electrons
1s
2s
e- configuration notation
Li = 1s2 2s1
Electronic Configuration
N atom
7 electrons
1s
2s
2p
e- configuration notation
N = 1s2 2s2 2p3
Electronic Configuration
Cl atom
17 electrons
1s
2s
2p
3s
3p
e- configuration notation
Cl = 1s2 2s2 2p6 3s2 3p5 or = [Ne] 3s2 3p5
•Oddities in e- configuration
4p
•
3d
•
4s
•
3p
•
3s
•
2p
•
•
• 1s
2s
Electronic Configuration
As has
33 electons
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p3
Or
Abbreviated e- config
Use the symbol of the noble gas with the same config
as the atom’s kernel
2
10
3
As = [Ar] 4s , 3d , 4p
Section Review
Mn: [Ar]4s2 3d?
How many d electrons does Mn have?
Electronic Configuration of Ions
Negative Ions
add 1 electron for each negative charge
continue up the energy ladder
Positive Ions
Remove 1 electron for each positive charge
largest n value e- are first to go
Electronic Configuration
S atom = 16 e-s
S = 1s2, 2s2, 2p6, 3s2, 3p4
S-2 ion (16 + 2) electrons
S-2 = 1s2, 2s2, 2p6, 3s2, 3p6 or = [Ar]
Ca atom = 20 e-s
Ca = 1s2, 2s2, 2p6, 3s2, 3p6, 4s2
Ca2+ ion (20-2) electrons
Ca2+ = 1s2, 2s2, 2p6, 3s2, 3p6 or = [Ar]
Electronic Configuration
+2
Mg ion
(12-2)electrons
1s2, 2s2, 2p6, 3s2
Section Review
How many valence electrons are in
Cl = [Ne]3s2 3p5?
2, 5, 7
Section Review
For Cl to achieve a noble gas configuration, it
is more likely that
electrons would be added
electrons would be removed
SURVEY OF ELEMENTS
• Atoms of most elem are hard to find.
Electron configuration
1s1
1s2
(stable)
1s22s1
1s22s2
1s22s22p 1
1s22s22p 2
...
1s22s22p 6
(stable)
1s22s22p 63s1
1s22s22p 63s2
1s22s22p 63s23p 1
...
1s22s22p 63s23p 6
(stable)
...
1s22s22p 63s23p 63d 10 4s246
(stable)
• Why? Valence (outer) shell usually not filled completely.
They react to form chem bonds to gain an octet!!!
5

into groups with similar properties
 Know properties of elem based on e- config
Periodic Table
 Elements above and below each other have
similar e- config and therefore similar properties
Several properties show a periodic change across
or down the table = trends
e- Config and the Periodic Table
Discoveries
• Newlands Octaves early 1800's
– Placed elements in order by mass
• Many elements were unknown
– Every 8th element has similar properties
• F and Cl
• Li, Na, K
• Mendeleev
– Ordered elements in table by mass, placing elem with similar
properties above one another. He left gaps assuming some
elements were unknown
– Predicted properties of missing elements
• Moseley
– Determined # of + charges in the nucleus
– Order of elements on periodic table is now by increasing
atomic number
Structure of Per Table
• Periods = rows (1 -7)
• Groups or families = columns
– 1 - 18
– Know
•
•
•
•
Alkali metals = grp 1 (very active metals)
Alkaline earth metals = grp 2 (active metals)
Halogens = grp 17 (very active nonmetals)
Noble or inert gases = grp 18 (inactive gases)
– Relate character of each group to it's econfiguration
Regions of Per Table
• Blocks
– s, p, d, f
• Main table elem = s and p blocks
• Transition elem = d block
• Rare earth elem = f block
– Lanthanides and actinides
• Metals, nonmetals, metalloids
– Know general characteristics of each
• e- sea model of metals
Periodic Trends
• Atomic radius - Size of the atom
– Effects of shielding of valence e-s by kernel e-s
• Ionization energy - meas of ability to lose an e– energy change when a valence e- is emitted from an atom
– High for nonmetals low for metals
• e- affinity - meas of ability to gain an e– Energy change associated with the acceptance of an e- into the
valence shell
• Usually (+) for metals and (-) for nonmetals electronegativity
• Electronegativity - meas of ability of atom to attract e-s in a
covalent bond
– High for nonmetals, low for metals
Periodic Trends
• Atomic radius
• Ionization energy
• Electronegativity
Section Review
Mn: [Ar]4s2 3d?
How many 3d electrons does Mn have?
4, 5, 6
Periodic Trends
Trends in the Periodic Table
•
•
•
•
•
Atomic radius
Ionic radius
Ionization energy
Electron affinity
Metallic character
Atomic Radius
• Hard to measure
• decrease left to right across a period
Zeff = Z - S
where
Zeff => effective nuclear charge
Z => nuclear charge, atomic number
S => shielding effect
Shielding Effect
•Coulombic (electrostatic) forces are affected by
• distance betwee the charges and size of the charges
Na
11+
2e-
8e-
1e-
Al
13+
2e-
8e-
3e-
Cl
17+
2e-
8e-
7e-
The inner (kernel) e-s shield
the valence shell e-s from
the protons in the nucleus
Sodium’s valence e-s see
only 1 proton ([11+] – [10-]).
Aluminum’s valence e-s see
3+ charges from the nucleus
([13+] – [10-]). How many
positive charges from the
nucleus do the valence e-s in
Cl see?
Atomic Radius
• increase top to bottom down a group
– each additional electron “shell” increases the
diameter of the atom
• increases from upper right corner to the
lower left corner
– across a period as shielding is constant while
charges in nucleus and valence shell increase
Atomic Radii
Ionic Radius
• same trends as for atomic radius
• positive ions smaller than atom
• negative ions larger than atom
Ions are isoelectric with a noble gas atom
Al
Na
11+
2e-
8e-
O
1e-
13+
2e-
8e-
Ne
8+
2e-
6e-
10+
2e-
8e-
3e-
Ions are isoelectric with a noble gas atom
Al3+
Na1+
11+
2e-
8e-
O2-
8+
13+
2e-
8e-
Ne
2e-
8e-
10+
2e-
8e-
•Atomic radius vs Ion radius
•Ionic Radii
Ionic Radius
Isoelectronic Series
• series of negative ions, noble gas atom, and
positive ions with the same electronic
configuration as a noble gas
– Cations move backwards to previous noble gas
– Anions work forward to the next noble gas
• size decreases as “positive charge” of the
nucleus increases
e- Configuration of Elements
Octet Rule
• Atoms of elements gain, lose, or share
electrons in order to obtain an octet (full
outer e- level)
– Eight e-s for all atoms except H and He
– Extra stability – lower energy
– Metals want to lose, and nonmetals gain e-s
M + energy  M+ + e- (endo)
NM + e-  NM- + energy (exo)
Ionization Energy
• energy necessary to remove an electron to form
a positive ion
• low value for metals, e-s easily removed
• high value for non-metals, e-s difficult to
remove
• increases from lower left corner of periodic
table to the upper right corner
1st Ionization Energy
Ionization Energies
» Abnormalities caused by extra stability of full and half full sublevels
Ionization Energies
first ionization energy
• energy to remove first electron from an atom
second ionization energy
• energy to remove second electron from a +1
ion etc.
Successive Ionization Energies
Electron Affinity
• energy released when an electron is added to
an atom
• same trends as ionization energy, increases
from lower left corner to the upper right
corner
• metals have low “EA”
• nonmetals have high “EA”
M + e-  M- + energy
Electronegativity
• All atoms want an octet
– Nonmetals want to add e-s to complete their
outer level
• Nonmetals attract e-s strongly
• They have high electronegativities
– Metals want to lose their outer level e-s to obtain
an octet
• Metals do NOT attract e-s strongly
• They have low electronegativities
Electronegativity
• The atom’s attraction for the electrons in a
covalent bond
Electronegativity