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
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
l lambda
Parts of a wave
Frequency ( )n = number of cycles in one second; hertz, hz1 hertz = 1 cycle/second - n nu
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.
GammaRays
Radiowaves
Energy is Quantized - PlanckQuantum – a packet of energy;
smallest quantity of energy that can be emitted or absorbed
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? Yes !
Spectrum
The range of frequencies present in light.
White light has a continuous spectrum.–All the colors are possible.–A rainbow.
Hydrogen spectrum
Line spectrum – series of separated lines of different colors representing photons whose wavelengths are characteristic of one element
410 nm
434 nm
486 nm
656 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
The Bohr Ring AtomQuantized energy – energy levelsGround state – ion or atom at its
lowest energy stateExcited 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
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
Atomic Models: Old version = Bohr’s Also known as the planetary
atomic model Describes electron paths as
perfect orbits with definite diameters
Good for a visual
New version = Quantum Theory
Most accepted Diagrams electrons of a
atom based on probability of location at any one time
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” principleReference sheet
How to Determine Electron Configurations and Orbital Diagrams
1. Determine total # of electrons
2. Use Aufbau Principle to “fill”
3. Make sure all electrons of the atom are accounted for!
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.
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 ConfigurationNa = 1s22s22p63s1
Al = 1s22s22p63s23p1
Ne = 1s22s22p6
Shorthand Electron ConfigurationNa = [Ne] 3s1
Al = [Ne] 3s23p1
EX: NaStep 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
Noble Gas Configuration
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
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!!!
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)
m l
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
4. Spin Quantum Number
m s
Electrons are negatively charged Behave as if spinning on axis – repel so… either +1/2 or -1/2