The document provides definitions and explanations of key concepts in quantum mechanics and atomic structure, including: - The formula for the energy of a photon - Definitions of wave function, probability density, electron density, orbitals, electron shells, subshells, degeneracy, Pauli exclusion principle, electron configuration, and Hund's rule - Tables listing the allowed values of the principal, angular momentum, magnetic, and spin quantum numbers - Examples of specifying orbitals using n, l, and ml values - Examples of writing electron configurations in both long and short hand notation for various elements

1. KEY
GENERAL CHEMISTRY-I (1411)
S.I. # 26
1. What is the formula for the energy of a photon? Label each variable and include
units. E = hν
E = the energy,
h = Plank’s Constant (6.626x10-34J s),
ν = the frequency of radiation
2. A photon at rest is considered to be on the _ground___ state and a photon that
has moved up in state is said to be on the __excited__ state.
3. Define the following:
a. Wave Function: (Ψ) mathematical representation that describes the
behavior of an electron within an atom
b. Probability Density: (Ψ 2) the probability of finding the location of the
electron at a particular point in space.
c. Electron Density: The probability of finding the location of the
electron at all points in space.
d. Orbitals: Allowed wave functions of they hydrogen atom
e. Electron Shell: the set of all orbitals with the same value of n, such as
3s, 3p, 3d.
f. Subshell: the set of one or more orbitals with the same n and l values.
g. Degenerate: orbitals within the same subshell, meaning they have the
same energy
h.*Pauli Exclusion Principle: states that no two electrons in an atom can
have exactly the same quantum number set
i. Electron configuration: describes how the electrons are distributed
among the orbitals of the atom.
j. Hunds Rule: states that the lowest energy is attained by maximizing
the number of electrons with the same electron spin. Same
spin must occupy different orbital of each energy level. Ex. ( 2p3 = ↑ ↑
↑ ).
K. *Valence Electrons: the outer-shell electrons that are involved in
chemical bonding. Those that are not involved are called core electrons.
4. Please Fill in the Table below:
Quantum # Name Symbol Allowed Values Functional Role
Principal Quantum # n 1, 2, 3, …∞n Determines Energy
Level
Angular Momentum l 0, 1, 2, 3, … (n-1) Defines the Shape of
Quantum # the Orbital
Magnetic Quantum # ml 0, ±1, ±2, ±3… (± l) Describes the
orientation of the
orbital in space
Intrinsic Spin ms ±½ Describes Electron
Quantum # spin of electron

2. KEY
5. Please Fill in the following Table:
Maximum Possible # of Electrons Orbital n l ml ms
2 s 1 0 0 ±½
6 p 2 1 0, ±1 ±½
10 d 3 2 0, ±1, ±2 ±½
14 f 4 3 0, ±1, ±2, ±3 ±½
18 g 5 4 0, ±1, ±2, ±3, ±4 ±½
6. How many possible values for l and ml are there when
a) n = 3 there are 3 l values (2, 1, 0) and 9 ml values
(l = 2; ml = -2,-1, 0, 1, 2)
(l = 1; ml = -1, 0, 1)
(l = 0; ml = 0)
b) n = 5 there are 5 l values (4, 3, 2, 1, 0) and 25 ml values
(l = 4; ml = -4, -3, -2, -1, 0, 1, 2, 3, 4)
(l = 3; ml = -3, -2, -1, 0, 1, 2, 3)
(l = 2; ml = -2, -1, 0, 1, 2)
(l = 1; ml = -1, 0, 1)
(l = 0; ml =0)
6. Give the values for n, l, and ml for
a) 3p n = 3, l = 1, ml = 0, ±1
b) 2s n = 2, l = 0, ml = 0
c) 4f n = 4, l = 3, ml = 0, ±1, ±2, ±3
d) 5d n = 5, l = 2, ml = 0, ±1, ±2
7. Determine the electron configuration by long hand and short hand notation:
a) Be 1s2 2s2 [He] 2s2
b) O 1s2 2s2 2p4 [He] 2p4
c) Mg 1s2 2s2 2p6 3s2 [Ne] 3s2
d) Cl 1s2 2s2 2p6 3s2 3p5 [Ne] 3s2 3p5
e) Co 1s2 2s2 2p6 3s2 3p6 4s2 3d7
[Ar] 4s2 3d7
f) Ag 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d10
[Kr]5s1 4d10
g) Rb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1
[Kr]5s1
h) Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d10 6p2
[Xe] 6s2 5d10 6p2