light is a form of electromagnetic radiation

• properties of both waves and particles:

  • wavelength (λ) - the distance between adjacent wave crests, meters

    • red light (750 NM) has longest wave length

    • violet light (400 NM) has shortest wave length

    • 1 NM = 1 * 10^-9 meters

  • frequency (v) - number of cycles or crests that pass through a stationary point in one second

  • amplitude - the height of the wave from zero to crest

• wavelength and frequency are inversely/indirectly related - the shorter the wavelength, the higher the frequency

• speed of light: 2.998 * 10^8 meters/second = λv

electromagnetic radiation

• light can be viewed as a stream of particles

  • particle of light is a photon

  • photon - a single packet of light energy

    • has specific wavelength, determines what light we see

    • wavelengths of spectral lines are characteristics of the element

      • make up atomic emission spectra

    • no two elements have the same emission spectra

  • amount of energy carried in the packet depends on the wavelength of the light - the shorter the wavelength, the greater the energy

    • light waves that carry more energy in their crests are closer

    • violet light carries more energy per photon than red light

• the photoelectric effect - the emission of electrons from a metal when light shines on the metal

  • quantum of energy - the minimum quantum of energy that can be lost or gained by an atom

  • quantized: an electron has to absorb/emit a specific amount of energy to move from one energy level to another

    • ground state: the normal energy level any given electron occupies

    • excited state: the energy level an electron occupies when it has absorbed the specific quantum of energy to move up to that level

  • Planck’s Law - E=Hv

    • E - energy, joules

    • H - Planck’s constant, 6.626 * 10^-34 J*S

    • v - frequency

Bohr’s Model

• Niels Bohr changed Rutherford’s model to include newer discoveries about how the energy of an atom changes when the atom absorbs/emits energy

  • proposed electron is found only in specific circular paths/orbits around the nucleus ❌

    • incorrect - if the orbits were truly circular, the electron would spiral into the nucleus

  • each possible electron orbit has a fixed energy - energy level ✅

  • each orbit is a specific distance from the nucleus and at each specific energy

    • impossible for an electron to exist between orbits

  • amount of energy is directly related to the frequency → wavelength

de Broglie: proposed “electrons be considered as waves confined to the space around an atomic nucleus”

Heisenberg Uncertainty Principle

• Werner Heisenberg

• states that it is impossible to determine simultaneously both the position and velocity of an electron

• “we cannot know both the position and speed of a particle, such as a photon or electron, with perfect accuracy”

  • Schrödinger Wave Equation

    • Erwin Schrödinger developed an equation that treated electrons as waves

    • Quantum Theory - describes mathematically the wave properties of electrons

      • electrons exist in certain regions called orbitals

      • orbitals - 3D regions around the nucleus that indicate the probable location of an electron

        • represent probability maps showing a statistical attribution of where the electron is likely to be found

    • 4 Wave Properties

      • Energy Level: Principal Quantum Numbers - number specifying the principle shell of orbital

        • n - indicates the energy level

        • energy increases with principal quantum number

        • maximum of 7 energy levels

        • n^2 - how many orbitals in any energy level

        • 2n^2 - maxim. number of electrons possible in any energy level

      • Sub Level: Shapes of Quantum Mechanical Orbitals

        • letter indicates subshell of orbital, specifies shape

        • possible letters - s, p, d, f

        • electrons are more likely to be found closer to the nucleus than farther away

      • Orbital: Orientation

        • s - 1 orbital

        • p - 3 orbitals

        • d - 5 orbitals

        • f - 7 orbitals

      • Spin: clockwise or counterclockwise

Electron Configuration

• arrangement of electrons in an atom and the way in which the electrons are arranged in various orbitals around the nucleus

• Aufbau Principle

  • the electrons will fill the orbitals in a very specific order

  • lowest → highest energy

  • The Diagonal Rule

• Pauli Exclusion Principle

  • an individual orbital may describe at most TWO electrons

  • in order to occupy the orbital, the two electrons must have opposite spins: ⬆⬇

• EXAMPLES

  • Carbon 6e- : 1s^2, 2s^2, 2p^2

  • Aluminum 13e- : 1s^2, 2s^2, 2p^6, 3s^2, 3p^1

• Noble Gas Configuration

  • Aluminum 13e- : [Ne] 3s^2, 3p^1

Hund’s Rule

• orbitals in the same sub-level must all fill with one electron before a second electron is added to any of the orbitals: (n)p^4 - ⬆⬇ ⬆ ⬆

• the “single” electrons will all have the same spin direction

• Orbital Diagram

  

Valence and Core Electrons

• valence electrons are the electrons on the outermost energy level

• the noble gases always have full valence shells

• Selenium 34e- : 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 3d^10, 4s^2, 3d^10, 4p^4

• Silicon 14e- : 1s^2, 2s^2, 2p^6, 3s^2, 3p^2

• the chemical properties of elements are largely determined by the number of valence electrons they contain

• properties vary in a periodic fashion because the number of valence electrons is periodic

Atomic Physical Properties

• Atomic Size ⬇⬅

  • left-right decreases ⬅

    • across a period, the amount of protons in the nucleus increases which has a stronger pull on the electrons, causing them to move closer to the nucleus

  • top-bottom increases ⬇

    • size of the orbital increases with increasing principal quantum shell number

    • electrons occupying the outermost orbitals are farthest from nucleus

• Ionization Energy ⬆➡

  • amount of energy needed to remove a single electron

  • left-right increases ➡

    • electrical pull on electrons from the # protons in nucleus causes increases amount of energy needed

  • top-bottom decreases ⬆

    • electrons in outermost orbitals are less affected by the electrical pull from nucleus

• Electronegativity ⬆➡

  • ability of an atom of an element to attract electrons when the atom is in a compound

  • left-right increases ➡

  • top-bottom decreases ⬆