LHS Adv Chem Lecture Series - Chapter 4

Chapter 4 : 

Chapter 4 Electrons

Survey : 

Survey Frequency = Speed of Light = Electromagnetic Spectrum = Visible Spectrum = Planck’s Constant = Wave-particle duality = Ground state = Excited state= Principal energy level = E- config =

Lighten Up! : 

Lighten Up! Knowledge of e-s comes from studies of the interaction between light and matter. Newton proposed light was a particle. Huygens said it was a wave. Big fight for hundreds of years (1600-1900) Experiments now prove both = WAVE- PARTICLE DUALITY (AKA wavicle).

Slide 4: 

Newton = Particle Huygens = Wave

Lighten Up! (cont.) : 

Lighten Up! (cont.) Light as a wave: Magnetic and electric fields oscillate at 90 angles.

Lighten Up! (cont.) : 

Lighten Up! (cont.) Four features of electromagnetic waves: Amplitude = height from origin Wavelength () = crest to crest Frequency = cycles/second (or sec-1) Speed = 3.00 x 108 m/sec Wave speed = freq x wavelength Or Wave speed Nu = frequency Lambda = wavelength c = ν s = ν c = speed of light

Lighten Up : 

Lighten Up When discussing light or any type of electromagnetic wave (i.e., radio, x-ray, etc.), the term “c” is used in place of “s.” “c” comes from the latin word, celeritus, which means “speed of light.” C = ν 

Light Speed Problem : 

Light Speed Problem Using wave speed (s = ν ), we can calculate freq of radiation. Ex = What’s the freq of a photon of light that has a  of 6.0 x 10-7m? Givens: Work: c = ν =  = 3.0 x 108 m/s 6.0 x 10-7m ? s = ν  3.0 x 108 m/s = ν (6.0 x 10-7 m) (6.0 x 10-7 m) (6.0 x 10-7 m) Ans = 5.0 x 1014 ??? sec-1

Light Speed Problem 2 : 

Light Speed Problem 2 Ex = What’s the freq of a photon of light that has a  of 485 nm? Givens: Work: c = ν =  = c = ν  3.00 x 108 m/s ? 485 nm Convert!!! 485 nm 10-9 m 1 nm = 4.85 x 10-7 m Ans = 6.19 x 1014 sec-1 = 4.85 x 10-7 m 3.00 x 108 m/s = ν (4.85 x 10-7 m)

Light Speed Problem 3 : 

Light Speed Problem 3 Calculate the wavelength of a radio wave with a frequency of 9.31 x 107 sec-1. Givens: Work: c = c = ν  ν =  = Ans = 3.22 m

Electromagnetic Spectrum : 

Electromagnetic Spectrum Electromagnetic Radiation (pg. 129) Gamma, X, UV, Vis, Infra, and Radio Grama examined Uncle Victor’s Vista, informing Mike TV’s radder. Combo of mag and electrical fields that radiate into space at speed of light. Long waves vs. Short waves

Quantum Theory : 

Quantum Theory Light as just a wave didn’t cut it. If light is a wave, then why would some objects change color when heated. Ex = blacksmith “Red Hot” “White Hot”

Quantum Theory (cont.) : 

Quantum Theory (cont.) Planck answered the question He noticed there are limits to the # of energy an object can absorb. These #’s are called quanta. Planck’s Theory =  E = h Where E = energy h = 6.6261x10-34 J-sec  = freq

Quantum Theory (cont.) : 

Quantum Theory (cont.) This was a revolution vs. old physics Old Car New Car

Quantum Theory (cont.) : 

Quantum Theory (cont.) This “quantum theory” led to the photoelectric effect.

Slide 17: 

Interrupting the flow of e-s will activate the circuit.

Quantum Theory (cont.) : 

Quantum Theory (cont.) Quantum Theory = The theory that particles absorb energy in packets. Photoelectric effect = the ejection of e-s from a metal as light shines on it. Something had to cause this ejection. Einstein suggested it was a packet or “photon” of light with the right quantum of energy to free an e-

Quantum Theory (cont.) : 

Quantum Theory (cont.) Drawing = pg. 133 Red vs. Violet (less E) (more E)

Quantum Theory (cont.) : 

Quantum Theory (cont.) Photons were also proven to collide with e-s. Uncertainty Principle?

Slide 21: 

Let’s say Newton observes a donut. PHOTON

Slide 22: 

Now, let’s say Newton observes an electron. PHOTON ELECTRON

Quantum Theory (cont.) : 

Quantum Theory (cont.) Uncertainty Principle = an observer can’t know both the location and momentum of an e-. Thus, light was proven to be a wavicle. (part wave, part photon) e-s are elusive.

Spectra for elements : 

Spectra for elements Continuous Spectra White light passed through a prism will separate into a continuous spectrum. ROY G. BIV Violet has the shortest , red the longest.

Slide 26: 

Increasing Freq ()  Decreasing Wavelength (λ)  Increasing Energy (E) 

Bright Line Spectra : 

Bright Line Spectra Each element has its own fingerprint of spectra that registers particular wavelengths. Each line corresponds to the frequency at which e-s for that element “quantum leap” energy levels. Ex =

More Model Development. : 

More Model Development. Niels Bohr put together what he knew about light. Production of light: e-s must absorb a quantum of energy AND e-s then “jump” to an excited state (higher energy level) and “fall” back to the ground state.

More Model Dev (Cont.) : 

More Model Dev (Cont.) This fall will result in a release of energy. This energy could be vis light, UV, infra, etc. Bohr concluded that each orbit in an atom corresponded to different energies.

Bohr Model of the atom : 

Bohr Model of the atom Each electron orbit got a quantum number (n) Ground state (n = 1 ), excited state (n = 2 and up) Larger jumps in energy corresponded to higher frequency wavelengths

Matter as a wave? : 

Matter as a wave? When scientists found out that light was a wave and particle, they wanted to find if matter was also a wave. This study led to the idea that all objects have wave properties that can be observable. An electron “wave” is bigger than its mass!

Slide 33: 

Wind induced resonance at Tacoma Narrows Bridge in November 1940

Slide 34: 

= 3 x 10-34 m, mass = 50 g Can’t observe wave. Golf Ball

Slide 35: 

 = 2 x 10-5 m, mass = 9.1 x 10-28 g An e-

e- Wave Cloud Model : 

e- Wave Cloud Model This is the model accepted by scientists e-s were found to not travel in straight lines (Uncertainty Principle). e-s now travel in probability areas (clouds) where the e- is most likely to be. Within these clouds are energy sublevels (kind of like Bohr) and orbitals where one or two e-s occupy with similar energy.

e- Wave Cloud Model (cont.) : 

e- Wave Cloud Model (cont.) The probability areas are quantified by their average energy level. We call this = the principal quantum number (n) The higher the quantum number, the farther away from the nucleus the e- is likely to be found.

More on energy levels : 

More on energy levels Any energy level is the “average” level of any electron. However, there are sublevels for each level and the # of them is equal to “n” This was an improvement on the Bohr model. Ex = n = 2, then 2 sublevels

Energy levels (cont.) : 

Energy levels (cont.) Sublevels occur in this order: S P D F  Increasing energy , if n = 2, then 2 sublevels will be present (S and P) So the F sublevel does not appear until n = 4

Slide 41: 

n = 1 n = 2 n = 3

Orbitals : 

Orbitals e-s tend to occupy a certain region within sublevels that are called orbitals. Orbital Rules No more than 2e- in an orbital Max number of e-s in each sublevel S = 2 P = 6 D = 10 F = 14 3. A jump in energy level corresponds to the rows of the PT.

Orbital Shapes (remember PROBABILITY!!) : 

Orbital Shapes (remember PROBABILITY!!) S = “sharp,” spherical P = “principal,” dumbbell D = “diffuse,” double dumbbell F = “fundamental,” messed up Also, corresponding PT areas that are filling their valence shells with e-s: S = Fam IA and IIA (w/ He), P = IIIA  VIIIA, D = Trans, and F = Rare Earth

S orbital : 

S orbital

3 orbitals of the P sublevel : 

3 orbitals of the P sublevel

5 orbitals of the D sublevel : 

5 orbitals of the D sublevel

Electron Spin : 

Electron Spin All e-s exhibit spin properties. Pauli Exclusion Principle says: 2 e-s in the same orbital must exhibit opposite spins. These e-s are said to be paired. Ex = Spinning Tops (these were popular before X Box or PS2)

Slide 48: 

Chemists represent spin by:

e- configurations : 

e- configurations The e-s config of an element tells: Energy the atom possesses. How they react. Tendency of cation or anion formed. Shorthand for the way in which the e-s fill their orbitals.

e- configurations (cont.) : 

e- configurations (cont.) Ex of a ground state Helium atom = 1s2 Coefficient of 1 = energy level, s = type of sublevel, 2 = # of e-‘s , expression refers to 2 e-‘s in the “s” sublevel of energy level # 1 Notation for a neutral ground state oxygen atom = 1s22s22p4

e- configurations (cont.) : 

e- configurations (cont.) Notation for a neutral ground state sodium atom = 1s22s22p63s1

Electron configs (cont.) : 

Electron configs (cont.) A hair more difficult e- configuration Element e- config. Valence e-s Calcium 1s22s22p63s23p64s2 2 How many e-s in the highest energy level?

E- configs (cont.) : 

E- configs (cont.) Element E- config. Valence e-s Iron Rubidium Technetium Xenon Neodymium [Ar]4s23d6 2 [Kr]5s1 1 [Kr]5s24d5 2 [Xe] 8 [Xe]6s24f4 2

Orbital diagrams (more on) pg. 348 : 

Orbital diagrams (more on) pg. 348 To make an orbital diagram, follow these rules: Aufbau Rule = Fill in e-s one at a time into the lowest energy orbitals until you account for all e-s. Pauli Exclusion = see above Hund’s Rule = a sublevel with more than one orbital will fill up unpaired e-s first.

Orbital diagrams (cont.) : 

Orbital diagrams (cont.) Recall that some energy levels of electrons overlap, especially in heavier elements Ex = Argon (#18)  Potassium (#19) Argon = 1s22s22p63s23p6 Potassium will fill which sublevel next? 3d? (NO) Diagram 4-31 Potassium = 1s22s22p63s23p64s23d1 There are several exceptions to these rules (Cr, Cu, Nb, etc.)