CHEM 1040 Exam Review
Atomic and Molecular Structure
● S orbitals (left 2 groups), P orbitals (right 6 groups), D orbitals (middle 10 groups), F orbitals
(bottom groups)
● Spectroscopy Line Spectrum - every line (energy) represents an electron moving from one
orbital to another
● Quantum Numbers:
○ First three quantum numbers identify the position of an electron in an orbit
○ The last quantum number accounts for the spin of an electron around its axis
○ 2 electrons in each orbital
○ Not all electrons follow the same path
○ Principal Quantum Number(n)
■ Determines the energy of an electron
● The smaller n is, the lower the energy
■ Potential energy
○ Angular Momentum Quantum Number (l):
■ Defines the number of shapes possible within a major energy level
■ Kinetic energy
■ L = n-1
■ Shape of orbital can affect the type of energy
○ Magnetic Quantum Number (ml)
■ Indicates how many orbitals that have the same n and l value but different
orientations
■ Ex. if l = 2, ml = -2, -1, 0, +1, +2; there is 5 orientations in space
○ Spin Quantum Number (ms)
■ Describes possible directions of the spin of an electron around its axis
■ Can be +/- ½
■ Explains magnetic properties of certain metals
● Electron Distribution Rules:
○ Order is determined by defining potential and kinetic energies, and by considering
spectroscopic data, ionization energies, magnetic properties
○ Electrons always seek the lowest total energy level
○ Maximum number of electrons is 2n2
● Positive ion always removes the electron from the highest n value (further shell away from the
nucleus)
● Outer shell electron / nucleus (protons) = e- / p
● Atomic and Ionic Radii:
○ Distance of the outer electron from the nucleus
○ As you go down a column and across a period e- / p increases
○ As you go across a period the atomic and ionic radii decreases
● Ionization Energy:
○ Energy required to remove an electron from a gaseous atom
○ As you go down a column e- / p decreases
○ As you go across a period e- / p increases
○ It becomes harder to remove an electron moving from left to right
● Electron Affinity:
○ As you go down a column e- / p, electron affinity decreases
, ○ As you go across a period e- / p, and electron affinity increases
● Electronegativity:
○ As you go down a column, electronegativity decreases
○ As you go across a period, electronegativity increases
● Lewis Structures:
○ Account for all the electrons and charges on the structure
○ Pick a central atom and draw a skeletal diagram
○ Draw single lines (covalent bonds) between each pair of atoms
○ Complete the octets of all the atoms by adding free electrons
○ If your electron has more electrons than it should according to the 1st step, go back
and make multiple bonds
○ If your structure does not have enough electrons, add enough electrons to the central
atom
○ Account for formal charges
● Molecular Shapes/Polarity:
○ 3 kinds of electron pairs:
■ Bonded, Free Electrons, Pi Bonded
○ Molecular Shapes:
■ Valence Shell Electron Pair Repulsion (VSEPR)
● Electron pairs try to get away from each other
● Only free electron pairs and sigma bonded pairs count (no pi electrons)
■ Framework Geometry
● Accounts for all electron pairs (sigma, pi, and free electron pairs)
■ Free electron pairs require more space than bonded electron pairs
■ Each free electron pair causes the other angles to be less than expected
● Hybridization:
○
Electron Pairs Hybridization Framework Geometry
2 sp Linear
3 sp2 Trigonal Planar
4 sp3 Tetrahedral
5 sp3d Trigonal Bipyramid
6 sp3d2 Octahedral
● The Wave Nature of Light:
○ Visible light extends from 400 nm to 800 nm
○ For any two waves travelling with a given speed, wavelength and frequency are
inversely related
● Quantum Effects and Photons:
○ Electrons are ejected however, only when the frequency of light exceeds a certain
threshold value characteristic of the particular metal
● The Bohr Theory of the Hydrogen Atom:
○ When an electron undergoes a transition, the electron goes from a higher energy level
to a lower energy level, the electron loses energy which is emitted as a photon
Atomic and Molecular Structure
● S orbitals (left 2 groups), P orbitals (right 6 groups), D orbitals (middle 10 groups), F orbitals
(bottom groups)
● Spectroscopy Line Spectrum - every line (energy) represents an electron moving from one
orbital to another
● Quantum Numbers:
○ First three quantum numbers identify the position of an electron in an orbit
○ The last quantum number accounts for the spin of an electron around its axis
○ 2 electrons in each orbital
○ Not all electrons follow the same path
○ Principal Quantum Number(n)
■ Determines the energy of an electron
● The smaller n is, the lower the energy
■ Potential energy
○ Angular Momentum Quantum Number (l):
■ Defines the number of shapes possible within a major energy level
■ Kinetic energy
■ L = n-1
■ Shape of orbital can affect the type of energy
○ Magnetic Quantum Number (ml)
■ Indicates how many orbitals that have the same n and l value but different
orientations
■ Ex. if l = 2, ml = -2, -1, 0, +1, +2; there is 5 orientations in space
○ Spin Quantum Number (ms)
■ Describes possible directions of the spin of an electron around its axis
■ Can be +/- ½
■ Explains magnetic properties of certain metals
● Electron Distribution Rules:
○ Order is determined by defining potential and kinetic energies, and by considering
spectroscopic data, ionization energies, magnetic properties
○ Electrons always seek the lowest total energy level
○ Maximum number of electrons is 2n2
● Positive ion always removes the electron from the highest n value (further shell away from the
nucleus)
● Outer shell electron / nucleus (protons) = e- / p
● Atomic and Ionic Radii:
○ Distance of the outer electron from the nucleus
○ As you go down a column and across a period e- / p increases
○ As you go across a period the atomic and ionic radii decreases
● Ionization Energy:
○ Energy required to remove an electron from a gaseous atom
○ As you go down a column e- / p decreases
○ As you go across a period e- / p increases
○ It becomes harder to remove an electron moving from left to right
● Electron Affinity:
○ As you go down a column e- / p, electron affinity decreases
, ○ As you go across a period e- / p, and electron affinity increases
● Electronegativity:
○ As you go down a column, electronegativity decreases
○ As you go across a period, electronegativity increases
● Lewis Structures:
○ Account for all the electrons and charges on the structure
○ Pick a central atom and draw a skeletal diagram
○ Draw single lines (covalent bonds) between each pair of atoms
○ Complete the octets of all the atoms by adding free electrons
○ If your electron has more electrons than it should according to the 1st step, go back
and make multiple bonds
○ If your structure does not have enough electrons, add enough electrons to the central
atom
○ Account for formal charges
● Molecular Shapes/Polarity:
○ 3 kinds of electron pairs:
■ Bonded, Free Electrons, Pi Bonded
○ Molecular Shapes:
■ Valence Shell Electron Pair Repulsion (VSEPR)
● Electron pairs try to get away from each other
● Only free electron pairs and sigma bonded pairs count (no pi electrons)
■ Framework Geometry
● Accounts for all electron pairs (sigma, pi, and free electron pairs)
■ Free electron pairs require more space than bonded electron pairs
■ Each free electron pair causes the other angles to be less than expected
● Hybridization:
○
Electron Pairs Hybridization Framework Geometry
2 sp Linear
3 sp2 Trigonal Planar
4 sp3 Tetrahedral
5 sp3d Trigonal Bipyramid
6 sp3d2 Octahedral
● The Wave Nature of Light:
○ Visible light extends from 400 nm to 800 nm
○ For any two waves travelling with a given speed, wavelength and frequency are
inversely related
● Quantum Effects and Photons:
○ Electrons are ejected however, only when the frequency of light exceeds a certain
threshold value characteristic of the particular metal
● The Bohr Theory of the Hydrogen Atom:
○ When an electron undergoes a transition, the electron goes from a higher energy level
to a lower energy level, the electron loses energy which is emitted as a photon
