Chapter 3

Features of Waves (3.1)

• Electromagnetic radiation
  • a wave that travels at the speed of light
  • the electric field is the sine wave and is measured in voltage
  • magnetic field is the cosine wave
• quantized waves
  • fixed ends of stationary wave restrict number and types of waveforms
  • only discrete values of a continuous set are allowed
  • node - region where amplitude is zero
    • no motion

Energy, wavelength, frequency of a photon (3.1)

• wavelength goes in the opposite direction of frequency and energy
• frequency is represented by v
• wavelength is represented by lambda
• the visible light spectrum is 400-800nm
  • blue is 400nm and red is 800nm
• https://www.cyberphysics.co.uk/graphics/diagrams/waves/EMSpectrumcolor.png
• one photon is equal to one light unit
• Energy is measured in joules

Photoelectric effect (3.1)

• minimum threshold to eject electrons is related to the frequency of light
  • classical physics suggested the amplitude of a wave influences the energy but experimentally, it was discovered that it is the frequency
• there is a certain amount of energy required to hold the atom together
• in order to eject an electron, you must overcome the threshold energy

Orbit energy of e-, delta E, Bohr Model (3.1, 3.2)

• line spectra
  • elements and molecules have emission spectra that correspond to a series of line
    • each line is a single wavelength
    • Rydberg equation (included in sheet)
      • n refers to which orbit an electron might be in 
      • 1 over lambda
    • https://i.stack.imgur.com/oDIES.jpg
• wave particle duality
  • light behaves both wavelike and particle-like
    • wavelike property - interference patterns obtained when light passed through a double slit
    • particle like property - photoelectric effect
• Bohr's atomic theory
  • Bohr used the results of Planck and Einstein's experiments to develop a better atomic theory, making the following assumptions about an atom's electrons
    • electrons travel in circular orbit around the nucleus
    • energy proportional to distance
    • orbits are quantized
      • equation is included on the sheet 
      • E sub n
      • Delta E
      • Z is the nuclear charge
    • energy is absorbed/emitted when orbits change
  • Orbit transitions
    • an electron moves to a higher level (n) when a specific quantity of energy is absorbed
    • an electron moves to a lower level (n) when a specific quantity of energy is emitted
  • the Bohr model fails with more than one electron, but was right about the following
    • e- attracted to nucleus by the opposing charges of the particles
    • e- energies are quantized 
    • e- energies increase as the distance from the nucleus increases
    • discrete energies in spectra of elements result from quantized energy transitions of e- between levels

DeBroglie Wavelength (3.3)

• wavelength is equal to h over linear momentum
  • linear momentum is equal to mass multiplied by velocity
• quantization of angular momentum, L, is equal to n multiplied by h
• wavelike properties are only important for small particles (electrons)
  • Electron location and momentum
    • Heisenberg Uncertainty Principle
      • We can only know either the position or momentum of a particle at any given instant
      • we can not know both in one instant
      • consequence of wave-particle duality
        • there are limitations governing the motions of quantum particles
• schrodinger
  • expanded on de Broglie's work and developed a wave equation that described and electron as a wave function, represented by psi
    • psi are not physical waves but complex probability amplitudes
    • psi squared is equal to the probability distribution (e- distribution)

Quantum Theory, Quantum Numbers (3.3)

• a series of four quantum numbers describe the energies, locations and other properties of an electron in a given atom 
  • principle quantum number (n)
    • discrete energy level
    • energies increase as the distance from the nucleus increase
    • atomic orbital
    • which orbital
    • corresponds to the row on the periodic table
  • angular momentum quantum number (l)
    • shape
    • integers from 0 to n-1
    • s, p, d, and f orbitals are 0, 1, 2, and 3, respectively
    • when l is not equal to 0, the orbitals have directionality about the z axis described by the magnetic quantum number
  • magnetic quantum number (m sub l)
    • orientation
    • integers from -l to l
    • number of degenerate orbitals in subshell
  • spin quantum number (m sub s)
    • electronic spin
    • overall one magnitude, two quantized states 
    • +1/2, -1/2
    • spin states degenerate unless in external magnetic field

Orbitals and Wave functions (3.3)

• pauli exclusion principle
  • no two electrons can have the same two sets of electrons
  • they have different quantum numbers in some way or another
• radial nodes
  • there are certain distances from the nucleus at which the probability density of finding an electron located at a particular orbital is zero
  • the number of nodes is equal to the energy level minus l minus 1

Electron Configurations (3.4)

• the total number of electrons in an atom populate these orbitals following a set of rules
• the end result is an electron configuration or orbital diagram
• hund's rule
  • electrons spread out within a sub shell and have the same spin
  • filled and half filled subshells are especially stable 
• aufbau principle
  • build up from lowest energy to highest
• valence electrons
  • the electrons in the outermost principle quantum level of an atom (largest value of n)
  • the 3de- are not counted as part of the valence shell
    • full d and f shells are not considered part of the valence shell because they are not involved in reactions if they are full
    • therefore, determination of the number of valence e- for any main group atom is simply the group number
    • HOWEVER, atoms in the transition series do count the de- toward the valence shell

Periodic Trends (3.5, 3.6)

• periodic law - the properties of the elements are periodic functions of their atomic numbers
• nuclear charge
  • many of the periodic trends involving e- can be qualitatively explained using the effective nuclear charge (Z sub eff) of the atom
    • this is the pull exerted on a specific electron by the nucleus, taking into account any electron-electron repulsions
    • Z sub off is equal to Z minus the shielding 
      • shielding is the e- repulsions
• radii
  • covalent radius - one half the distance between the nuclei of two identical atoms when they are joined by a covalent bond
  • ionic radius - size of an ion
  • increases as you go down and to the left
• ionization energy - Energy required to remove the most loosely bound e- in formation of cation
  • as size increases, IE decreases
    • it becomes easier to remove electrons
  • increases as you go up and to the right 
  • watch for filled and half filled subshells
• electron affinity - energy change for adding an electron to gaseous atom in formation of an anion
  • increases going up and to the right
• unit prefixes and conversions
  • pico 10^-12
  • nano 10^-9
  • micro 10^-6
  • milli 10^-3
  • centi 10^-2
  • kilo 10^3

Ionic and Covalent Bonds, including polyatomic (3.7)

• monatomic ions
  • the majority of chemical reactions do not change the nucleus of an atom 
  • electrons can be transferred or shared
  • groups
    • alkali metals
    • alkaline earth metals
    • chalcogens
    • halogens
    • noble gases
• polyatomic ions
  • usually have a negative charge
• binary bonding
  • ionic
    • metal and nonmetal
    • groups of atoms held together by electrostatic interactions 
    • transfer of electrons
  • molecular/covalent
    • nonmetal and nonmetal
    • groups of atoms held together by sharing electrons