# Modern Atomic Theory Review

• 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

• 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

ENERGY LEVEL

SUB-LEVEL

# ORBITALS (n^2)

ELECTRONS (2n^2)

n=1

1s

1

2

n=2

2s 2p

4

8

n=3

3s 3p 3d

9

18

n=4

4s 4p 4d 4f

16

32

• 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 ⬆