Bohr’s Model Rutherford’s model didn’t explain the arrangement of electrons around the nucleus.

Nature of Light Before 1900, scientists thought that light behaved only as a wave

Electromagnetic Radiation form of energy that exhibits wave-like characteristics as it travels through space λ = wavelength (m) ν = frequency (Hertz) = (s -1 ) c = speed of light (3.0 x 10 8 m/s) c = λν

Electromagnetic Spectrum

German Physicist Max Planck Studied emission of light by hot objects Proposed that – Object doesn’t emit continuous energy (waves) – It emits energy in small specific “packets” called quanta – A quantum of energy is the minimum energy that can be lost or gained by an atom

E = hν E = energy in Joules (J) h = Planck’s constant = x J·s ν = frequency in Hertz (1/s OR s -1 ) PROPORTIONALITY? Direct or Indirect? – Higher energy means higher frequency

The Photoelectric Effect = the emission of e- from metal when high frequency (high energy) light shines on it Photon = a “particle” of light having no mass as a quantum of energy

Einstein Dual nature of light – WAVE – Particle (PHOTON) E photon = hν In order for an electron to be ejected from a metal surface, the e- must be struck by a single photon possessing at least the minimum amt of energy required to knock it loose (will differ for each element). REMEMBER: energy and frequency are directly proportional, SO minimum energy corresponds to minimum frequency.

Ground State and Excited State

Ground → ExcitedExcited → Ground When an e - returns to ground state it gives off energy When an e - becomes excited it gains energy

Line Emission Spectrum -unique to each element and unique for each type of EMR -visible light spectrum created when light is passed through a gas Hydrogen Visible light Line emission spectrum

Line Emission Spectra for H Balmer Series – visible light Lyman Series - UV Paschen Series - infrared These series show that H atoms are only “excited” by certain frequencies of light AND therefore give off only certain frequencies of light, NOT a continuous spectrum of light like was previously supposed Like the EM spectrum

Bohr’s Atom Because H emitted only certain frequencies, the e - of H atoms exist in only very specific energy states Scientists needed to provide a model for the H atom that explained this

Niels Bohr Modeled the H atom Electrons exist at definite energy levels at ground state Electrons must become excited to move to higher energy levels bohr gizmo

Quantum Model of Atom – De Broglie Electrons also have a dual particle-wave nature beams of e - can be refracted like beams of light beams of e - can cause interference like beams of light If electrons are both particles AND waves, then where are they in the atom?

Locating Electrons in the Atom e- are only detected by their interactions with light (photons) When you hit an electron with photons, it knocks it off course (interference) SO…… Heisenberg Uncertainty Principle = it is impossible to determine simultaneously both the position and velocity of an e - or any other particle

Schrödinger Wave Equation and Quantum Theory = mathematically describes the wave properties of electrons and other very small particles Quantum numbers = based on Schrödinger Wave Equation and specify the properties of atomic orbitals and the e - in them

Electrons in an Atom Don’t exist in orbits like the planets Electrons are in orbitals, which exist in energy levels (energy shells). There can be more than 1 orbital in an energy shell You can think of the energy shell like a planet’s orbit, but there can be more than 1 orbital in a shell and more than 1 electron in an orbital

Quantum Numbers 1.Principal Quantum Number (n) = indicates the main energy level occupied by the e - – Positive integers – As n ↑, energy and distance of e- from nucleus ↑ – Ex. An e - with n=1 is in the 1 st energy level (shell) away from the nucleus – Multiple e - can have the same value for n

2. Angular Momentum Quantum Number (l) = indicates the shape of the orbital – Orbitals of different shapes known as sublevels exist for each value of n – The # of possible shapes = n – Values for l can be 0-3 and are given corresponding letters

s-orbitals are Spherical

p-orbitals are Dumbbells

d-orbitals are more complex

F-orbitals Exist on 3 planes simultaneously VERY COMPLEX

3.Magnetic Quantum Number (m) = indicates the orientation of the orbital around the nucleus – Whole numbers from -1 to 1 4.Spin Quantum Numbers (+1/2, -1/2) – A single orbital holds only 2 e - with opposite spin numbers

n = energy level sublevel# of orbitals per sublevel # of e- per sublevel Number of e- per main energy level 1s122 2s12 p368 3s12 p36 d s12 p36 d510 f71432 Quantum Number Relationships

Rules for Orbital Filling Aufbau Principle = an e- occupies the lowest- energy orbital that can receive it Pauli Exclusion Principle = no 2 e- in the same atom can have the same 4 quantum numbers Hund’s Rule = orbitals of equal energy are each occupied by 1 e- before any orbital is occupied by a 2 nd e-, and all e- in singly occupied orbitals must have the same spin state

Order of Orbital Filling 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 Draw diagonal lines from top right to bottom left to determine order of filling. Memorize this!!! Expect a QUIZ! Quantum # notation: nl x n=energy level l = orbital shape x = maximum # of e- held in that type of orbital

Elements of the 2 nd period Li1s 2 2s 1 Be1s 2 2s 2 B1s 2 2s 2 2p 1 C1s 2 2s 2 2p 2 N1s 2 2s 2 2p 3 O1s 2 2s 2 2p 4 F1s 2 2s 2 2p 5 Ne1s 2 2s 2 2p 6 ↑↓ ↑_ ↑↓ ↑↓ ↑↓ ↑_ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑_ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑_ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ e - configuration notationOrbital Notation 2s 2s 2p 2p 2p

Noble Gas Notation 1.Assume the e- configuration for the last noble gas (by atomic number) before the element 2.Put the symbol for that NG in brackets 3.Write only the e- configuration that comes after the NG Ex. Ca = [Ar] 4s 2 Al = [Ne] 3s 2 3p 1 Si = [Ne] 3s 2 3p 2

Electron Dot Notation Lewis dot notation = symbol of element surrounded by number of valence e - Valence e - = e - in outer energy level Ex. Nitrogen Oxygen Fluorine