The Bohr and Quantum Models

Starts with the Study of…..Light Light is…. Made up of electromagnetic radiation. Waves of electric and magnetic fields at right angles to each other.

Parts of a wave l Wavelength l lambda Wavelength = the distance between any point on a wave and the corresponding point on the next wave (crest to crest, trough to trough); m, nm, pm

Parts of a wave Frequency (n) = number of cycles in one second; hertz, hz 1 hertz = 1 cycle/second n - nu

Frequency = n

Kinds of EM waves There are many different l and n (p. 92) Radio waves, microwaves, x rays and gamma rays are all examples. Visible light is all our eyes can detect. G a m m a R a d i o R a y s w a v e s

Electromagnetic Spectrum

Energy is Quantized - Planck Quantum – a packet of energy; smallest quantity of energy that can be emitted or absorbed

Einstein is next Light is particulate photon – a quantum of electromagnetic radiation

Which is it? Is energy a wave like light, or a particle? Yes ! Concept is called the Wave -Particle duality (Dual Nature of Light) What about the other way, is matter a wave?

Spectrum The range of frequencies present in light. White light has a continuous spectrum. All the colors are possible. A rainbow.

Electromagnetic Spectrum

Hydrogen spectrum Line spectrum – series of separated lines of different colors representing photons whose wavelengths are characteristic of one element 656 nm 434 nm 410 nm 486 nm

What this means Only certain energies are allowed for the hydrogen atom.(Recall relationship b/w energy, wavelength and frequency) Energy in the atom is quantized. Can calculate Energy based on line spectra Only the energy we can “see”

Quantum Theory: Colors emitted can be used to identify elements (absorption of energy and color emitted is a fingerprint of an element) Kind of like wearing your team colors.              Team Oxygen Team Carbon

Let’s Play!! White light - dispersion Flame Tests

Niels Bohr

The Bohr Ring Atom Quantized energy – energy levels Ground state – ion or atom at its lowest energy state Excited state – atom or molecule at any other state besides ground state Putting energy (photon) into the atom moves the electron away from the nucleus.

Bohr From ground state (lowest energy level) to excited state. When it returns to ground state (or lower energy level) it gives off light of a certain energy (photon.) Energy, frequency, wavelength – color of light

The Bohr Ring Atom n = 4 n = 3 n = 2 n = 1

Bohr Planetary Model quantum model of the hydrogen atom atom was like a solar system; protons in dense nucleus electrons orbit (specific path electron travels) ground and excited state

The Bohr Model Only works for hydrogen atoms. Electrons don’t orbit. Energy quantization correct, but not circling like planets Introduced concept of energy levels

The Quantum Model

Atomic Models: Old version = Bohr’s New version = Quantum Theory Most accepted Diagrams electrons of a atom based on probability of location at any one time Old version = Bohr’s Also known as the planetary atomic model Describes electron paths as perfect orbits with definite diameters Good for a visual

Basic Concepts of the Quantum Model Atoms and molecules can exist only in certain energy states Atoms absorb or emit radiation as they change energies Orbitals NOT orbits Energy states are described by sets of numbers called quantum numbers MATHEMATICAL model

Heisenberg’s Uncertainty Principle We can’t know how the electron is moving or how it gets from one energy level to another. Heisenberg Uncertainty Principle: It is impossible to know the exact position AND momentum of a particle simultaneously. (Marco Polo)

Electron Configuration The distribution of electrons within their atoms orbitals Distributed amongst energy levels, sublevels, and orbitals Based on a set of stated principles

AUFBAU PRINCIPLE Electrons are added one at a time to the lowest energy orbitals “building up” principle Reference sheet

Electron Configurations # of electrons in sublevel H 1s1 Energy Level sublevel

How to Determine Electron Configurations and Orbital Diagrams Determine total # of electrons Use Aufbau Principle to “fill” Make sure all electrons of the atom are accounted for!

18 electrons Argon (18 electrons) Energy

18 electrons Argon (18 electrons) Energy 1s22s22p63s23p6

Orbital Diagram – box, circle or line for each orbital in a given energy level, grouped by sublevel, with an arrow indicating electron AND its spin

HUND’S RULE Hund’s Rule (BUS RULE)- The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in the orbital. C 1s2 2s2 2p2

Details Valence electrons- the electrons in the outermost energy levels (not d). Core electrons- the inner electrons.

TRY! Calcium and Manganese

Calcium (20 electrons) Energy 1s22s22p63s23p64s2

Try Oxygen, Iron, Strontium, Molybdenum, Tungsten! Orbital Diagrams AND Electronic Configuratiions

Using the Periodic Table! Noble Gas Configuration!

Blocks and Sublevels P E R I  O D S 1-7 1 2 3 4 5 6 7 d (n-1) 4 5

Label your blank periodic table. Read it “like a book”

WRITE the Electron Configuration using the periodic table: Kr Ca Fe Hg

Electron Configuration – Noble Gas Configuration Electron Configuration demonstrates a periodic trend, so you can write shorthand electron configuration using the electron configuration of the noble gases in Group 18 of the periodic table. Noble gases have stable configurations.

Noble Gas Configuration When writing shorthand e- config for an element, refer to the noble gas in the energy level (period) just above the element. Write the symbol of the noble gas in brackets. Write out the remaining e-config based on the energy filling diagram. Electron Configuration Na = 1s22s22p63s1 Al = 1s22s22p63s23p1 Ne = 1s22s22p6 Shorthand Electron Configuration Na = [Ne] 3s1 Al = [Ne] 3s23p1

Noble Gas Configuration EX: Na Step 1: Na is in period 3 so refer to the noble gas in period 2 which is Neon. Step 2: Write Ne in brackets. [Ne] Step 3: Now write remaining electrons in standard form. 3s1. Step 4: Combine. [Ne]3s1

Nobel Gas Configuration Now try: I Kr Na Cu

Electron Configuration with Ions When we write the electron configuration of a positive ion, we remove one electron for each positive charge: Na → Na+ 1s2 2s2 2p6 3s1 → 1s2 2s2 2p6 When we write the electron configuration of a negative ion, we add one electron for each negative charge: O → O2- 1s2 2s2 2p4 → 1s2 2s2 2p6

Electron Configuration with Ions Now try: Ca+2 Fe-3

Details Elements in the same column have the same electron configuration. Put in columns because of similar properties. Similar properties because of electron configuration. Noble gases have filled energy levels.

Exceptions Ti = [Ar] 4s2 3d2 V = [Ar] 4s2 3d3 Cr = [Ar] 4s1 3d5 Mn = [Ar] 4s2 3d5 Half filled orbitals. Scientists aren’t sure of why it happens same for Cu [Ar] 4s1 3d10

More exceptions Lanthanum La: [Xe] 6s2 5d1 Cerium Ce: [Xe] 6s2 4f1 5d1 Promethium Pr: [Xe] 6s2 4f3 5d0 Gadolinium Gd: [Xe] 6s2 4f7 5d1 Lutetium Pr: [Xe] 6s2 4f14 5d1 We’ll just pretend that all except Cu and Cr follow the rules.

Quantum Numbers Quantum Numbers represent 4 solutions to Schroedinger’s Equation Describe the distribution of electrons in space BASED ON PROBABILITY!!!

1. Principal Quantum Number (n) Describes main energy level an electron occupies n=1,2,3,4,5,6,7

2. Angular Momentum Quantum (l) sublevel designates the shape of the region in space an electron occupies (orbital). integer values from 0 to n-1 l = 0 is called s Sublevels include s (l=0, p=1, d=2, f=3) Not all energy levels have all sublevels

3. Magnetic quantum number (Orbitals ml) designates the spatial orientation of orbital integer values between - l and + l s: ml = 0 p: ml = -1, 0, +1 (i.e. px, py, pz) s-1, p-3, d-5, f-7

S orbitals

P orbitals

P Orbitals

D orbitals

F orbitals

F orbitals

4. Spin Quantum Number m s Electrons are negatively charged Behave as if spinning on axis – repel so… either +1/2 or -1/2

Pauli Exclusion Principle NO TWO ELECTRONS MAY HAVE THE SAME FOUR QUANTUM NUMBERS (Spin must be different!)