Chemistry 1, H/G
1.
2.
3.
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5.
Atomic Structure
Atomic Models
Dalton’s Theory- follows Democritus’ view; atoms are indivisible
J.J. Thomson- ”plum-pudding” model; e- in positively charged mass
E. Rutherford- “nuclear” model; e- found in mostly empty space surrounding a small, positive nucleus
N. Bohr- “planetary” model; e- found in definite energy orbitals
Schrödinger- Quantum Mechanical Model; (mathematical wave equations) 90% probability e- cloud
Periodic Table and Element Key
1{Group # = no. of Valence e-}
1
1
1
{period # = no. of energy levels}
1.
2.
3.
4.
H
hydrogen
1.00794
G
Atomic Orbitals
Electrons enter orbitals of lowest energy first (Aufbau Principle).
Each atomic orbital can be occupied by a maximum of 2e- (Pauli-Exclusion Principle).
Each e- must have an opposite spin [+1/2 spin, -1/2 spin].
Electrons do not pair up until all orbitals contain one e- each (Hund’s Rule).
e- cloud
energy levels
sublevels
s
p
d
f
atomic orbitals
1
2
1.
2.
3.
4.
5.
6.
7.
3
5
Electrons
6
10
7
14
Behavior of Atoms
Atoms are most stable when energy levels are full or half-full, especially the outermost energy level, also called
the VALENCE SHELL.
Noble gases (group zero, 8A (VIIIA), or 18) have completely filled valence shells, therefore are already stable
elements that do not engage in chemical reactions; their former name was the Inert Gases.
Atoms will tend to follow the electron configuration (e- arrangement within the orbitals) of the Noble gases.
Hydrogen and Helium have only one principal energy level, therefore their only orbital will fill with 2e-; in He
this orbital is already full; this is called a duet of electrons.
All other elements’ valence shell will fill with 8e-; this is called an octet of electrons.
Atoms gain stability by attaining Noble gas e- configuration (8 valence shell electrons or a full orbital) or halffull valence shells (pseudo-Noble Gas configuration).
Atoms of elements will attain chemical stability by engaging in chemical reactions; this will provide the extra
electrons needed or will remove extra electrons not needed, or will provide a sharing of electrons that will add
to the 8 electrons needed for STABILITY.
©T. M. Chipi, chemistry instructor
1
Chemistry 1, H/G
Atomic Structure
6.
Energy levels, Sublevels, and Orbitals.
An energy level has as many sublevels as its number (ie. Energy level 1 has one sublevel, energy level 2 has two sublevels,
etc.). The principal quantum number (n = ?) refers to the principal energy levels 1-7.
The energy sublevels or subshells are designated by the letters s, p, d, and f. These are subenergy levels that
take on shapes according to the paths with greatest electron density and highest electron probability (90%).
a. s-shell (sharp)- spherical shape
the lowest energy subshell.
b. p-shell (principal)- dumbbell shape
the next higher energy subshell.
c. d-shell (diffuse)- four of the orbitals have 4-leaf clover shape, one has a dumbbell-doughnut combination
shape. The third highest energy subshell.
d. f-shell (fundamental)- various complex shapes (8-lobed), the highest energy subshell.
The s-sublevel contains 1 spherical orbital, with a maximum capacity of 2e-.
The p-sublevel contains 3 dumbbell (double-lobed) orbitals, with a maximum capacity of 6e-.
The d-sublevel contains 5 total orbitals (four 4-leaf clover, 1 dumbbell-doughnut), with a maximum capacity of
10e-.
The f-sublevel contains 7 complex-shaped orbitals, with a maximum capacity of 14e-.
1.
2.
Orbital Filling Guidelines
Aufbau Principle: Electrons enter orbitals of lowest energy first (follow periodic table or diagonal rule).
Pauli Exclusion Principle: An atomic orbital can be occupied at most by 2 e- and these must have opposite
1.
2.
3.
4.
5.
3.
1.
2.
3.
4.
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spins (indicated by
or
, for half-filled orbitals and
for filled orbitals).
Hund’s Rule: When electrons occupy orbitals of equal energy, one e- enters each orbital until all the orbitals
contain one e- each (spins parallel), then the second e- enters and pairs up with the other e- (at this time spins
are opposite and are said to be paired).
Quantum Theory
Light energy is electromagnetic energy that travels as waves in quantized packets of energy called photons that
behave as particles.
Light, as well as matter, has a dual behavior: it behaves as waves and as particles.
When light of a suitable amount of energy strikes the surface of a metal, electrons are emitted. This
phenomenon, known as the photoelectric effect, could not happen unless the light striking the surface of the
metal were made up of particles-like species called photons.
The concept of light being made up of photons (discrete energy packets or quanta) is the basis of the Quantum
Theory, attributed to Max Planck.
The energy in a photon of light is directly proportional to the frequency of the radiation emitted.
(low energy)
1.
2.
3.
4.
Red Orange Yellow Green Blue Indigo Violet
(high energy)
Electrons in energy levels near the nucleus have lower energies than those in energy levels at a greater distance
from the nucleus.
Electrons in one energy level have definite, discrete amounts of energy different from the energy values of
electrons in a different energy level.
When an electron moves from one energy level to another it absorbs or releases a discrete amount of energy
that is quantized called a photon of energy (a packet).
In higher energy levels, energy differences decrease and energy sublevel complexity increases.
Quantum Numbers
Schrödinger developed a mathematical equation, the solution of which gives the probability (90%) that an electron is
present in a region in the atom around the nucleus. This equation supplies 4 factors that are called quantum
numbers. These are designated as a set {n,l,ml ,ms} consisting of the following:
1. n = principal quantum number- refers to the energy level (shell) where the electron is found. These energy
levels go from n = 1 to n = 7(∞), and may be also designated letters K-Q, respectively.
2.
l = orbital or azimuthal (spatial orientation) quantum number- refers to the spatial region (shape) formed as the
electron moves (angular momentum) in an elliptical orbit within the atom. These regions are called sublevels
(subshells) and are designated each value from l = 0 to l = n-1. For n=1, l = 0 and only one sublevel is allowed
©T. M. Chipi, chemistry instructor
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Chemistry 1, H/G
Atomic Structure
(s), for n=2, l = 0,1 (s,p); for n=3, l = 0,1,2 (s,p,d); for n=4, l = 0,1,2,3 (s,p,d,f). So, for l = 0, s is the sublevel;
for l = 1,p; for l = 2, d; for l = 3,f.
3.
ml = magnetic quantum number- refers to the magnetic properties (energies) or the orientation in space of the
orbitals belonging to the sublevels within the energy levels of the atom. The values for ml are from –l to +l and
the number of orbitals per sublevel is given by (2l + 1).
4.
ms = spin quantum number- refers to the spin of an electron and has a value of +1/2 or –1/2. Where there are
two e- in an orbital, one has a clockwise spin and the other a counterclockwise spin.
Principal
Energy
Level
n=
1
2
3
4
5
6
7
Number Of
Sublevels In
Principal
Energy Level
(s, p, d, f…)
1
2
3
4
5
6
7
Type(s) Of
Sublevel(s) In
Principal Energy
Level
l = 0,1,2,3…
S
s,p
s,p,d
s,p,d,f
s,p,d,f,g
s,p,d,f,g,h
s,p,d,f,g,h,i
Number Of
Orbitals In
Sublevel (s,
p, d, f)
ml = - l t o + l
s=1
p=3
d=5
f=7
g=9
h=11
i=13
Maximum
Number Of eIn Sublevel (s,
p, d, f)
s=2
p=6
d=10
f=14
g=18
h=22
i=26
Total Number
Of Orbitals In
Principal
Energy Level
n2
1
4
9
16
25
36
49
Maximum
Number Of eIn Energy
Level
2n 2
2
8
18
32
50
72
98
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©T. M. Chipi, chemistry instructor
3
Chemistry 1, H/G
Atomic Structure
The Eight Diatomic Elements
7
N
O
F
Cl
Br
1
H
=8
octet
I
At
©T. M. Chipi, chemistry instructor
4