Chapter 5 Chemistry
Electrons in Atoms
I. Quantum Theory and the Atom
A. Bohr Model
1. Proposed that energy states were involved with atoms
a. Ground state is lowest allowable energy for an atom
b. Excited state occurs as energy is gained
2. Also proposed electrons travel in orbits around the nucleus
a. The smaller the orbital, the lower the energy state (energy
level)
b. Bohr assigned each level a quantum number. The quantum
number corresponds to the energy level
1. 1 was closest to the nucleus
2. 2 was second closest
c. Electrons can move from energy level to energy level
d. Electrons dropping from a higher energy level to a lower
energy level release energy in the form of light (spectrum)
B. Quantum Mechanical Model
1. de Broglie concluded if waves can have particle like
motion, particles can have wave like motion
2. This meant all particles (cars) have wavelike motion
3. Also called the wave mechanical model
4. Does not describe the electron path around the
nucleus
5. Heisenberg Uncertainty Principle
a. States it is impossible to know velocity and
location of electrons at the same time
C. Atomic orbitals
1. Three dimensional area around the nucleus
2. Indicate where electrons are located 90% of the time
3. Principal quantum numbers
a. Indicate size and energy of atomic orbitals
b. Are numbered 1 to 7
c. Same as the principal energy levels
4. Energy sublevels are found in
principal energy levels
5. Sublevels are labeled s, p, d or f as determined by their
shape (illustrations of p. 133)
a. Each sublevel has a number of orbitals. Orbitals can
hold up to two electrons each
1. s has one orbital
2. p has three orbitals
3. d has five orbitals
4. f has seven orbitals
b. Energy level 1 has one sublevel (s)
c. Energy level 2 has two sublevels (s and p)
d. Energy level 3 has three sublevels (s, p, and d)
e. Energy level 4 has four sublevels (s, p, d, and f)
II. Electron configurations
A. Arrangement of electrons in an atom
B. Three rules define an electron’s location
1. Aufbau principle
a. States electrons occupy the lowest energy
orbital available
b. All orbitals of a sublevel have equal energy and
different sublevels have different energy
c. Orbitals from one sublevel can overlap orbitals
from other energy levels
2. Pauli exclusion principle
a. each electron has a spin
b. electrons in orbitals have opposite spins
3. Hund’s rule
a. negative charges play a role in
distribution of electrons
b. electrons with the same spin will occupy
equal energy orbitals before another
electron will fill an orbital (each orbital of p
will have one electron before one orbital has
two electrons)
C. Electron configurations and orbital diagrams
1. Orbital diagrams
a. Have a box for each orbital with the
orbital written under it
b. Spins are shown through arrows pointing
up or down
2. Electron configuration notation
a. Uses the orbitals and a superscript to
represent electrons
b. 1s2 2s2 2p2 is carbon
D. Electrons filling
1. Electrons fill into the lowest energy level and sublevel
available
2. Diagram on page 138 shows order
3. Exceptions
a. Metals can have d orbitals fill before s is
completed (chromium, copper)
b. Look to see if the d orbitals will have an
electron in each orbital
E. Valence electrons
1. Are the electrons in the outermost energy level of an
atom
2. Determine the properties of an atom
3. Look at the outermost energy level
4. Noble gases have full outer energy levels (very stable)
5. Determine electron dot structures
a. Demonstrates the number of valence electrons for
an atom
b. Is the chemical symbol with a dot for each valence
electron
c. Pattern involved is by filling one orbital at a time