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Chapter 5: Models of the Atom
How are atoms like the
solar system?
5.1 Atomic Models
• In 1911, Rutherford described the atom as a small, dense positively charged nucleus with electron that orbit it.
• However, Rutherford’s model could only explain simple atomic properties.
The Bohr Model (1913﴿
• Bohr’s model of the atom described how electrons could only be found in specific paths, or orbits, around the nucleus.
• These orbits are called energy levels.
• There is only space for two electrons in each orbital. • A quantum of energy is how much energy is required to move an electron between energy levels.
The Bohr Model con’t
• Electrons in energy levels that are close to the nucleus have less energy than those farther away. • Electrons may only exist in energy levels, they are never found between energy levels.
• How can this be???
The Quantum Mechanical Model!
• This model uses mathematical equations to calculate the probability of finding an electron in a certain location around the nucleus.
• Described by Erwin Schrodinger in 1926.
• Solving Schrodinger’s equation describes the types of atomic orbitals where we have a high probability of finding an electron.
• The energy level of these atomic orbitals is described by the principal quantum (energy) number, n.
• At each energy level, there are a specific amount of sublevels:
(orbitals) Atomic Orbitals
• At each principal energy level, there can be several orbitals of different shapes and at different energy levels. What are orbitals shaped like?
Types of Atomic Orbitals
• Type 1: The s Orbital
Photos page 131­132
• The s orbital has a spherical shape centered around the origin of the three axes in space.
• Type 2: The p orbital
• There are three dumbbell‐shaped p orbitals in each energy level where n=2 and up, each assigned to its own axis (x, y and z) in space.
• They are named px, py and pz.
• So, all three p orbitals will exist at the same energy level, they make up this sublevel.
Type 3: The d orbital There are five d orbitals:
All five d orbitals will exist at the same energy level.
Type 4: the f orbitals
There are 7 types of f orbitals all at the same level
• http://www.d.umn.edu/~pkiprof/ChemWebV2/AOs/ao1.html
àHigher Energy Level • Questions:
• Page 149, # 22‐27, 29.
5.2 Electronic Arrangements
• Electrons are arranged into energy levels three specific ways:
• The Aufbau Principle
• The Pauli Exclusion Principle
• Hund’s Rule
• Each element has its own unique electronic configuration that describe how the electron are arranged in energy levels around the nucleus.
The Aufbau Principle
Electrons in an atom will occupy the lowest­energy orbitals first.
The “Lazy Tenant” Rule
Diagonal Rule
Steps:
• Write the energy levels top to bottom.
• Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level.
• Draw diagonal lines from the top right to the bottom left.
• To get the correct order, • follow the arrows!
By this point, we are past the current periodic table so we can stop.
Pauli Exclusion Principle
• An orbital can only hold two electrons.
• Two electron in one orbital must have opposite spins. These electrons are called
spin pairs.
• Vertical arrows are used to designate spins.
Hund’s Rule
• When an atom has two or more unpaired electrons in the same energy level, they will each occupy their own orbital and will have the same spin.
Example:
Oxygen 2p
2p
2s
2s
1s
1s
Writing Electron Configurations
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Determine the number of electrons an atom has (equal to the number of protons).
Fill the atom’s orbitals in the order of increasing energy. (Aufbau)
Remember at each principal energy level there can be:
One s orbital
Three p orbitals (starting at n = 2)
Complete questions 10­13 page 136
Five d orbitals (starting at n = 3)
Seven f orbitals (starting at n = 4)
Write the full electron configurations for:
Carbon
Argon
Nickel
Boron
Silicon
Write the e­ configurations for:
Na:
Se:
Y:
Example: Sulfur
• Full Electron Configuration:
S 16e­
Core Electrons
Valence Electrons
Ne:
Germanium: – Example
Shorthand:
1. Find the noble gas in the row above and put it
in [square brackets] for the core electrons.
2. Write the remaining electron config as before
for the valence electrons.
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Questions:
Page 149 # 32‐37, 39.
Page 150 # 57.
Write the shorthand electron configuration for:
• Nitrogen
• Magnesium
• Phosphorus
• Lead
Exceptional Electron Configurations
• Writing electron configurations using the aufbau principle is correct up to the element vanadium.
• Consider copper:
9
6 2
6 2
2 2
• 1s 2s 2p 3s 3p 4s 3d
• This configuration is not correct since it is more stable for copper to exist with a full d orbital subshell:
10 1
6
6 2
2 2
• 1s 2s 2p 3s 3p 3d 4s
• Similarly for chromium:
• 1s22s22p63s23p64s23d4 is not correct.
• The electrons will reorganize themselves so that d orbitals each have one unpaired electron:
• 1s22s22p63s23p64s13d5 is correct.
• Half‐filled subshells are not as stable as filled subshells but will be of lower energy than other arrangements .