1. 7.1: Electromagnetic Radiation
Types of Electromagnetic Radiation: Gamma, X-ray, UV, visible, IR, micro, radio
Properties of Electromagnetic Waves:
Wavelength ( λ ) :
Distance Between two consecutive peaks or troughs in a wave
Measured in meters
Frequency (ν)
Number of waves that pass a given point per second
Measured in Hertz
Speed (c)
Speed of light
Measured in meters/ second
Relationship Between Properties
Shortest wavelength = highest frequency
Longest wavelength = lowest frequency
INVERSE RELATIONSHIP

2. 7.2 The Nature of Matter Einstein’s Photoelectric Effect
Max Plank & Quantum Theory: Energy is gained/lost in whole numbers multiples of the quantity hv ( frequency=v, Planck’s constant =h) Planck’s Constant: h = × m2 kg / s ( J/s) Planck discovered that energy is transferred to matter in packets of energy called quantum, rather than energy of matter being continuous. ΔE = hν : The quantum of energy can be calculated from this equation
Einstein’s Photoelectric Effect
Phenomenon in which electrons are emitted from the surface of a metal when light strikes it His observations are explained by assuming electromagnetic radiation is quantized (photons) and the threshold frequency is the minimum energy required to remove the electron.

3. 7.2: The Dual Nature of Matter
Dual Nature of Light:
Light travels through space as a wave
Light transmits energy as a particle
Particles have wavelength, exhibited by diffraction patterns
De Broglie’s Equation: Allows calculation of wavelength for a particle
λ = h/mv
Diffraction: results when light is scattered from a regular array of points or lines
Diffraction Patterns: The interference pattern that results when a wave or a series of waves undergoes diffraction, as when passed through a diffraction grating or the lattices of a crystal. The pattern provides information about the frequency of the wave and the structure of the material causing the diffraction.

4. 7.3: The Atomic Structure of Hydrogen
Continuous Spectrum: results when white light is passed through a prism. Contains all wavelengths of visible light Line Spectrum: only see a few lines, each of which corresponds to discrete wavelength when passed thorough a prism. (Hydrogen emission spectrum)

5. 7.4: The Bohr Model
Quantum Model : electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits
Bright line spectra confirms that only certain energies exist in the atom, and atom emits photons with definite wavelengths when the electron returns to a lower energy state.
Energy levels available to the electron in the hydrogen atom:
n= an integer
z= nuclear charge
J= energy in Joules

6. 7.4: The Bohr Model Calculating the energy of the emitted photon
Calculate electron energy in outer level
Calculate electron energy in inner level
Calculate the change in the energy
ΔΕ= energy of final state- energy of initial state
hc/ ΔΕ : to calculate the wavelength of emitted photon
Energy Change in Hydrogen atoms
Calculate the energy change between any two energy levels:
Limitations of the Bohr Model
Bohr’s model does not work for atoms other than hydrogen
Electrons do not move in circular orbits

7. 7.5: Quantum Mechanical Model
Electron bound to nucleus similar to standing waves
The exact path of the electron is not known
Heisenberg Uncertainty Principle- a limitation to the position and momentum of a particle at a given time
Physical Meaning of ψ
Square of the function is the probability of finding an electron near a particular point
Represented as a probability distribution
aka electron density map, electron density, electron probability
Radial Probability Distribution Since the orbital size cannot be calculated, the size of the orbital is the radius of the sphere that an electron is in for 90% of the time

8. 7.6: Quantum Numbers Principal quantum number (n) Main energy level
1, 2, 3, …
Size and energy of orbital
When n increases: orbital becomes larger, electron is further from the nucleus, higher energy b/c electron is less tightly bound to the nucleus so the energy is less negative
Angular momentum quantum number/Azimuthal QN (l)
Sublevels, subshell
0...n-1 for each value of n
Shape of atomic orbitals
Magnetic quantum number ( 𝒎 𝒍 )
Integral values from l to –l
Orientation of the orbital in space

9. 7.7: Orbital Shapes and Energies
s Orbitals
Spherical shape
Nodes for s orbitals of n=2 or greater
p Orbitals
Two lobes each
Occur in levels n=2 and greater
Each orbital lies along an axis
d Orbitals
Occur in levels n=3 or greater
Four orbitals with four lobes each centered in the plane indicated in the orbital label
Fifth orbital has two lobes along z axis and a belt centered in the xy plane
f Orbitals
Occur in levels n=4 and greater
Complex shapes
Usually not involved in bonding in compounds
Size of orbital:
Defined as the surface that contains 9-% of the total electron probability.
As n increases orbitals of the same shape grow larger.
Orbital Energies
All orbitals with the same value of n have the same energy for hydrogen atoms (Degenerate)
The lowest energy state = ground state
When the atom absorbs energy the electrons can move to higher energy orbitals
“excited state”

10. 7.8: Electron Spin & the Pauli Principle
Electron Spin Quantum Number
An orbital can only hold two electrons, must have opposite spins.
Spin can have +1/2 or -1/2
Pauli Exclusion Principal
In a given atom no two electrons can have the same set of four quantum numbers

11. 7.9: Polyatomic Atoms
Polyatomic Atoms: Atoms with more than one electron
3 energy contributions must be considered in description of the atom:
Kinetic energy of electrons as they move around the nucleus
Potential energy of attraction between nucleus and electrons
The potential energy of repulsion between the two electrons
Electron correlation problem: Electron pathways are not known, so electron repulsive forces cannot be calculated exactly
Average repulsions are approximated by...
Treat each electron as it were moving in a field of charge that is the net result of the nuclear attraction and average repulsions of all other electrons
Screening or Shielding
Electrons are attracted to the nucleus
Electrons are repulsed by other electrons
Electrons would be bound more tightly if other electrons weren’t present
Closer proximity to the nucleus = lower energy

12. 7.10 History of the periodic table
Originally constructed to represent patterns observed in chemical properties of elements
Mendeleev and Meyer both independently conceived present periodic table
Mendeleev also corrected several atomic masses

13. 7.11: Aufbau Principle & the Periodic Table
Aufbau Principle: “As protons are added one by one to the nucleus to build up elements, electrons are similarly added to these hydrogen like orbitals”
Hunds Rule: “The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by Pauli principle in a particular set of degenerate orbitals
Periodic Table Vocab:
Valence electrons: electrons in outermost principal quantum level of an atom
Transition metals: “d” Block
Lanthanide and Actinide Series : “f” block
Representative Elements: Group 1A through 8A
Metalloids: Border between metals and nonmetals, exhibit properties of both

14. 7.12 Periodic Trends in Atomic Properties
Ionization energy: energy required to remove an electron from an atomic (increase across period, decreases with increasing atomic number within a group) Electron affinity: energy change associated with the addition of an electron (decrease down period, increase across period) Atomic Radius: Determination of radius (increases down group, decreases across period)
7.13: The Properties of a Group: Alkali Metals
Easily lose valence electrons
Reducing agents
React with water
Large hydration energy
Positive ionic charge