1900 Proposed that amounts of energy are quantized only certain E values are allowed Niels Bohr 1913 e can possess only certain amounts of energy and can therefore be only certain ID: 724891
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Slide1
Recent Atomic Models
Max Planck (1900): Proposed thatamounts of energy are quantized
only 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
–
found
here
e
–
never
found here
planetary
(Bohr) model
NSlide2
Continuous vs. Quantized Energy
Energy
A B
Zumdahl, Zumdahl, DeCoste,
World of Chemistry
2002, page 330
continuous quantizedSlide3
Bohr Atom
The Planetary Model of the AtomSlide4
Bohr’s Model
Nucleus
Electron
Orbit
Energy LevelsSlide5
ENERGY
(HEAT, LIGHT,
ELEC., ETC.)
Light
When all e
–
are in lowest possible energy state,
an atom is in the
____________
.
ground state
e.g.,
He: 2 e
-
, both in 1
st
energy level
If “right” amount of energy is absorbed by an e
–
, it can
“jump” to a higher energy level. This is an unstable,
momentary condition called the ____________.
excited state
e.g.,
He: 1 e
-
in 1
st
E level, 1 e
-
in 2
nd
E level Slide6
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 LIGHT
Any-old-value of energy to be
absorbed or released is
NOT OK. This explains
the lines of color in an
emission spectrum
.
Slide7
Frequency A
Frequency B
Frequency C
n
= 2
n
= 1
n
= 3
NSlide8
1
ST
E.L.
2
ND
E.L.
3
RD
E.L.
4
TH
E.L.
5
TH
E.L.
6
TH
E.L.
Lyman
(UV)
Paschen
(IR)
Balmer
(visible)
~
~
~
~
~
~Slide9
Continuous and Line Spectra
4000 A
o
5000
6000
7000
light
Na
H
Ca
Hg
400 450 500 550 600 650 700 750 nm
Visible
spectrum
l
(nm)Slide10
Flame Emission Spectra
Photographs of flame tests of burning wooden splints soaked in different salts.Include link to web page
http://www.unit5.org/christjs/flame%20tests.htm
methane gas
wooden splint
strontium ion
copper ion
sodium ion
calcium ionSlide11
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.Slide12
Fireworks
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.Slide13
electromagnetic radiation (i.e., light)
--
--
E
B
waves of oscillating electric (E)
and magnetic (B) fields
source is…
vibrating electric charges
Slide14
Characteristics of a Wave
frequency
: the number of cycles per
unit time (usually sec)
amplitude
A
crest
wavelength
l
trough
--
unit is Hz, or s
–1
or 1/sSlide15
radio waves
IR
visible
UV
X-rays
gamma rays
cosmic rays
electromagnetic spectrum
: contains all of the “types” of
light that vary according to frequency and wavelength
R
O
Y
G
BV
large
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 Slide16
Increasing wavelength
Increasing frequency
Increasing energySlide17
Some light humor…Slide18
Waves
Low
frequency
High
frequency
Amplitude
Amplitude
long wavelength
l
short wavelength
l
Both travel at the same speed…the speed of light!Slide19
Waves
Low
frequency
High
frequency
Amplitude
Amplitude
long wavelength
l
short wavelength
l
60 photons
162 photons
Low
energy
High energySlide20
Red and Blue Light
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 325
Photons
- particle of light that carries a quantum of energySlide21
Albert Michelson (1879)
--
first to get an accurate
value for speed of light
The speed of light in
a vacuum (and in air)is constant:
c = 3.00 x 10
8
m/s
c = f
l
-- Equation:
Albert Michelson
(1852–1931)
c = 671 x 10
6 mph Slide22
E = h f
In 1900, Max Planck assumed
that energy can be absorbed
or released only in certaindiscrete amounts, which hecalled quanta.
Later, Albert Einstein dubbeda light “particle” that carried a
quantum of energy a photon. -- Equation:
E = energy,
h = Planck’s constant
in J
= 6.63 x 10
–34
J∙s (i.e., J/Hz)Max Planck(1858–1947)
Albert Einstein
(1879–1955)Slide23
A radio station transmits
at 95.5 MHz (FM 95.5).Calculate the wavelengthof this light and the energyof one of its photons.
c = f
l
= 3.14 m
E = h f
= 6.63 x 10
–34
J/Hz (95.5 x 10
6
Hz)
= 6.33 x 10
–26
J
3.00 x 10
8
m/s
=
95.5 x 10
6
HzSlide24
quantum mechanical model
electron cloud model
charge cloud model
Schroedinger
, 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. Slide25
Electron Cloud Model
Orbital (“electron cloud”) instead of “orbits”Region in space where there is 90% probability of finding an electron…creates unique 3-D shapes
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Electron Probability vs. Distance
Electron Probability (%)
Distance from the Nucleus (pm)
100
150
200
250
50
0
0
10
20
30
40
Orbital
90% probability of
finding the electronSlide26
Models of the Atom Review
Dalton’s model
(1803)
Thomson’s plum-pudding
model (1897)
Rutherford’s model
(1909)
Bohr’s model
(1913)
Charge-cloud model
(present)
Dorin, Demmin, Gabel,
Chemistry The Study of Matter
, 3
rd
Edition, 1990, page 125
Greek model(400 B.C.)
+
-
-
-
-
-
e
e
e
+
+
+
+
+
+
+
+
e
e
e
e
e
e
e
"In science, a wrong theory can be valuable and better than no theory at all."
- Sir William L. BraggSlide27
Models of the Atom Timeline
Dalton’s model
(1803)
Thomson’s plum-pudding
model (1897)
Rutherford’s model
(1909)
Bohr’s model
(1913)
Charge-cloud model
(present)
Dorin, Demmin, Gabel,
Chemistry The Study of Matter
, 3
rd
Edition, 1990, page 125
Greek model(400 B.C.)
1800 1805 ..................... 1895 1900 1905 1910 1915 1920 1925 1930 1935 1940 1945
1803
John Dalton
pictures atoms as
tiny, indestructible
particles, with no internal structure.
1897
J.J. Thomson, a British
scientist, discovers the electron,
leading to his "plum-pudding"
model. He pictures electrons
embedded in a sphere of
positive electric charge.
1904
Hantaro Nagaoka, a
Japanese physicist, suggests
that an atom has a central
nucleus. Electrons move in
orbits like the rings around Saturn.
1911
New Zealander
Ernest Rutherford states
that an atom has a dense,
positively charged nucleus.
Electrons move randomly in
the space around the nucleus.
1913
In Niels Bohr's
model, the electrons move
in spherical orbits at fixed
distances from the nucleus.
1924
Frenchman Louis
de Broglie proposes that
moving particles like electrons
have some properties of waves.
Within a few years evidence is
collected to support his idea.
1926
Erwin Schr
ö
dinger
develops mathematical
equations 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.
+
-
-
-
-
-
e
e
e
+
+
+
+
+
+
+
+
e
e
e
e
e
e
eSlide28
Shapes of s, p, and d-Orbitals
each holds 2 electrons (s2)
e
ach of 5 orbitals holds 2 e - = 10 total d electrons (d10)
each of 3 orbitals holds 2 e - = 6 total p electrons (p6)Slide29
f-orbitals
each of 7 orbitals hold 2 e- = 14 e-How many “g-orbitals” could exist and how many e- could they hold?Slide30
theoretical g-orbitals
each of 9 orbitals hold 2 e- = 18 e-Slide31
Relative Sizes 1s and 2s
1s 2s
Zumdahl, 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- Slide32
Copyright
© 2006 Pearson Benjamin Cummings. All rights reserved.
Orbitals overlap each other as you get farther from the nucleus
Orbital Filling VideoSlide33
s, 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)Slide34
s block
p block
d block
f
block
6
7
Periodic Patterns
1
s
2
s
3
s
4
s
5
s
6
s
7
s
3
d
4
d
5
d
6
d
1
s
2
p
3
p
4
p
5
p
6
p
7
p
4
f
5
f
1
2
3
4
5
6
7
(n-1)
(n-2)
nSlide35
Sections of Periodic Table to Know
f-block
s-block
d-block
p-blockSlide36
Energy Level Diagram of a Many-Electron Atom
Arbitrary
Energy Scale
18
18
32
8
8
2
1s
2s 2p
3s 3p
4s 4p 3d
5s 5p 4d
6s 6p 5d 4f
NUCLEUS
O’Connor, Davis, MacNab, McClellan,
CHEMISTRY Experiments and Principles
1982, page 177
Each orbital can only hold 2 e-
Start from the bottom and add e-Slide37
You don’t have to memorize the order…just start at the beginning and fill in e-…Slide38
s-block
1st Period (row)
1s
1
1 e- in “1s” orbital
Periodic Patterns
Example -
Hydrogen
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chemSlide39
Writing Electron Configurations:
Where are the e–? (probably)
As
H
He
NAl
Li
Ti
Xe
1s
2
2s2
3p6
2p64s
23d10
3s
24p6…1s1
1s
2
1s
2
2s1
1s2
2s2
2p
3
1s
2
2s
2
3p
1
2p
6
3s
2
1s
2
2s
2
3p
6
2p
6
4s
2
3d
2
3s
2
1s
2
2s
2
3p
6
2p
6
4s
2
3d
10
3s
2
4p
3
1s
2
2s
2
3p
6
2p
6
4s
2
3d
10
3s
2
4p
6
5s
2
4d
10
5p
6
Filling OrderSlide40
Notable e- configuration exceptions:
Cu
Cr
1s
2
2s23p6
2p
6
4s
1
3d
53s21s2
2s2
3p62p6
4s1
3d
103s2Mo and W (theoretically Sg) will behave similarly to CrAg and Au (theoretically Rg) will behave similarly to Cu
Why? Generally, full orbitals and half-filled orbitals have lower energies and are thus more stable. So, while we might expect Cr to end with 4s
23d4, promoting an s e- to the d orbitals creates two half-filled shells.Slide41
[Ar]
4s
2
3d
10
4p2
Periodic Patterns
Example -
Germanium
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Ge
72.61
32Slide42
Electron Configuration Battleship
(your shots)Slide43
Electron Configuration Battleship
(your ships and opponent’s shots)Slide44
S
Shorthand Electron Configuration (S.E.C.)
To write S.E.C. for an element:
1. Put symbol of noble gas that precedes
element in brackets.
2. Continue writing e– config. from that pointCo
In
Cl
Rb
[ Ne ]
3s
2 3p4 [ Ar ]
4s2 3d7
[ Kr ]
5s2 4d10 5p1
[ Ne ] 3s2 3p5 [ Kr ]
5s
1
Ge
72.61
32Slide45
Shorthand Configuration Review
[Ar] 4
s
2
Electron configuration
Element symbol
[Ar] 4
s
2
3
d
3
[Rn] 7
s
2
5
f
14
6
d4
[He] 2
s
2
2p5
[Kr]
5
s
1
4
d
10
[Kr] 5
s
2
4
d
10
5
p
5
[Kr] 5
s
2
4
d
10
5
p
6
or [Xe]
[He] 2
s
2
2
p
6
3
s
2
3
p
6
4
s
2
3
d
6
Ca
V
Sg
F
Ag
I
Xe
Fe
[Ar] 4
s
2
3
d
6Slide46
Periodic Patterns ReviewPeriod # (1-7) primary energy level (n) (subtract for d & f)
Group # (1-8…excluding d block) total # of valence e-Column within Sublevel block # of e- in sublevel
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chemSlide47
Three Principles about Electrons
aufbau 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
Hund’s
Rule
:
two e–
in same orbitalhave different spins
1s22s23p
62p6
4s2
3d10…
3s
2Friedrich Hund
Wolfgang PauliSlide48
General RulesAufbau Principle Electrons fill the lowest energy
orbitals first. “Lazy Tenant Rule”Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
2
s
3
s
4
s
5
s
6
s
7
s
1
s
2
p
3
p
4
p
5
p
6
p
3
d
4
d
5
d
6
d
4
f
5
f
1
s
2
s
2
p
3
s
3
p
4
s
4
p
3
d
4
d
5
s
5
p
6
s
7
s
6
p
6
d
4
f
5
f
5
d
Energy
NO ELEVATOR!Slide49
General Rules
Hund’s
Rule
Within a sublevel, place one electron per orbital before pairing them.
“Empty Bus Seat Rule”
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
2p
2p
ORSlide50
General RulesPauli Exclusion Principle Each orbital can only hold TWO electrons and the must have opposite spins.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Wolfgang Pauli
1sSlide51
O
Orbital Diagrams
…show spins of e
– and orbital location
1s
2s
2p
3s
3p
1s
2s
2p
3s
3p
P
1s
2
2s
2
2p
4
1s
2
2s
2
2p
6
3s
2
3p
3Slide52
HAVE MORE
ENERGY
ARE FURTHER
FROM NUCLEUS
AND
The Importance of Electrons
orbitals
:
In a generic e
–
config (e.g., 1s
2 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
–
…Slide53
Shorthand Configuration
S 16e
-
Valence Electrons
(Highest energy level)
Kernel (Core) Electrons
S 16e
-
[Ne]
3s
2 3p4
1s2
2s2
2p6
3s
23p4Electron ConfigurationLonghand Configuration
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
S
32.066
16Slide54
INVOLVED IN
CHEMICAL
BONDING
kernel (core) electrons
:
valence electrons:
in inner energy level(s);
close to nucleus
in outer energy level
He:
Ne:
Ar:
Kr:
1s
2
[ He ] 2s2 2p6 [ Ne ] 3s2
3p6
[ Ar ] 4s2 3d10 4p
6
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
–
) Slide55
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…
doesn’t 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
F
–
chlorine atom, Cl
17 p+
, 17 e– 17 p+, 18 e– Cl is more stable asa Cl– ion
Cl
–
How to be like
a noble gas…?
[He] 2s
2
2p
5
[Ne] 3s
2
3p
5
gain 1 e
–
or lose 7 e
-
? Slide56
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] 2s
1
[Ne] 3s
1Slide57
1+
Know charges on these columns of Table:
Group 1:
Group 2:
Group 3:
Group 5:
Group 6:
Group 7:
Group 8:
2+
3+
3–
2–
1–
0
1+
2+
3+
3–
2–
1–
0
…also called oxidation numbers. Slide58
s
p
d
f
6
7
Periodic Patterns and Charge Trends
1
s
1
2
3
4
5
6
7
+1
+2
(n-2)
n
+3
- 3
- 2
- 1
Variable ChargeSlide59
Naming Ions
e.g., Ca2+
Cs
+
Al3+
Cations use element name and then say “ion”
e.g., S
2–
P
3–
N
3–O2–Cl–
Anions
change ending of element name to “ide”and then say “ion”
calcium ion
cesium ion
aluminum ionsulfide ion
phosphide ion
nitride ion
oxide ion
chloride ion