Waves, light, and energy: Where chemistry and physics collide EMR info Waves, light, and energy: Where chemistry and physics collide http://imagers.gsfc.nasa.gov/ems/waves3.html

List as many interactions of light and matter as you can. Before we get started…. What is light? Is it matter? What forms of light exist? List as many interactions of light and matter as you can. think how light changes matter, and how matter changes light What are some uses of light?

First things first: Waves First things first: Waves a and b represent wavelength (λ)- the distance of a wave from crest to successive crest; measured in meters http://www.lbl.gov/MicroWorlds/ALSTool/EMSpec/EMSpec.html

Waves: amplitude The height of a wave from crest to midline or trough to midline; measured in meters

Terms you need to know: Wavelength (λ) Amplitude Wavelength (λ) Amplitude Frequency (ν [nu]; I know some of you have used f, move on and get with chemistry!): the number of cycles per second measured in cycles per second (s-1) or Hz (Hertz) Waves on a string

http://www.lbl.gov/MicroWorlds/ALSTool/EMSpec/EMSpec2.html

http://micro.magnet.fsu.edu/primer/lightandcolor/images/electromagneticfigure1.jpg

http://lepus.physics.ualr.edu/~tahall/EXAM2/emspec.jpg

http://www.arpansa.gov.au/images/emsline2.gif

Visible Light Violet 400---460 7.5--6.5 5.0--4.3 color wavelength(nm) f(*1014 Hz) Energy (*10-19 J) Violet 400---460 7.5--6.5 5.0--4.3 Indigo 460---475 6.5--6.3 4.3--4.2 Blue 475---490 6.3--6.1 4.2--4.1 Green 490---565 6.1--5.3 4.1--3.5 Yellow 565---575 5.3--5.2 3.5--3.45 Orange 575---600 5.2--5.0 3.45--3.3 Red 600---800 5.0--3.7 3.3--2.5

Some equations you need to know = c / ν and E = hν So…. E = hc /  And…  = h / mv* When = wavelength in m c = speed of light, 3.00E8 m/s ν (nu)= frequency in Hz (cycles/sec or s-1 or 1/s) E= energy in J h= Plank’s constant, 6.626E-34 J*s [Joule(seconds)] m= mass of particle in kg V*= velocity in m/s

What the h? Planck’s Constant What the h? Planck’s Constant When metals are heated, they glow 1800s- physicists were trying to determine the relationship between the color (wavelength) and intensity of the glow Max Planck (1900)- energy can be released or absorbed only in little chunks (packets) of energy “of some minimal size”

Max Planck and the h The chunks of energy were dubbed “quantum” (“fixed amount”), which is the smallest amount that can be emitted or absorbed as EMR. Proposed: E = hν The energy (E) of a single quantum is equal to its frequency (ν) times a constant

Planck and the Nobel (Physics) Planck determined that h= 6.626E-34 J-s Energy is always released in multiples of hv (1hv, 2hv, 3hv, etc) h is so small that we cannot see the effects of this in our daily lives Analogous to… Planck won the 1918 Nobel Prize in physics for his work

Einstein & Bohr: Perfect Together Einstein, left Bohr, above

Einstein: The Photoelectric Effect Einstein discovered that one could cause electrons to be ejected from the surface of a metal if the energy of the light wave was strong enough He treated the light needed to do this as a piece of matter- a photon, if you will This ejection of e- is the photoelectric effect

The Photoelectric Effect Only light of a certain energy could knock off an electron from the metal Intense light of a weaker wavelength would not work, but even a low intensity of the correct wavelength would work (the energy of the light is transferred to the kinetic energy of the electron) Hmmm… light acting as a particle and as a wave…..

The photoelectric effect… Online animations PhET http://www.lewport.wnyric.org/mgagnon/Photoelectric_Effect/photoelectriceffect1.htm http://www.xmission.com/~locutus/applets/Photoelectric.html

Getting to Bohr…. Light of a given wavelength is monochromatic (one color) Most common EMR sources are polychromatic, but we see only one color These can be reduced to a spectrum when the different wavelengths are separated out

Spectral Emissions Continuous spectrum: shows all colors of the rainbow

Bright line spectrum: only certain wavelengths are visible (the rest do not appear at all) Different elements have different bright line spectrum when they are heated Na is yellow Ne is orange-red

Line spectrum Ne I2

http://www.cartage.org.lb/en/themes/Sciences/Astronomy/Modenastronomy/Interactionoflight/AtomicAbsorption/AtomicAbsorption.htm

Hydrogen Spectra Emission Spectra Absorption Spectra Emission Spectra Absorption Spectra http://www.mhhe.com/physsci/astronomy/applets/Bohr/content_files/section1.html

http://www.cartage.org.lb/en/themes/Sciences/Astronomy/Modenastronomy/Interactionoflight/AtomicAbsorption/AtomicAbsorption.htm

Color and what you see: Absorption: the wavelengths that are absorbed by an object are not available for us to see, as we see the wavelengths of light that are reflected off of an object This is not the same as those wavelengths that are emitted by an object that is emitting radiant energy.

Color and what you see…

Chlorophyll absorption spectra

Perception of color

Line spectra formation- go to….. http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

Bohr Model and Spectral Emissions Bohr proposed that the emission of light energy from an (electrically or thermally) excited atom corresponds to the orbit of the electron around the nucleus of the atom That energy can only be achieved by being a specific distance from the nucleus

What you’ve seen so far…. Model of a Nitrogen (z=7) atom

Bohr Model and moving electrons http://www.colorado.edu/physics/2000/quantumzone/bohr.html

Energy levels- Bohr Model Electrons travel within set energy levels that have a particular energy associated with each level After all, the e-s are moving around the nucleus think KE here Each shell has a number Closest to the nucleus is n=1 For each successive level add 1 to n n=2, n=3, ect….

Energy increases as the distance from the nucleus increases

Bohr Model and moving electrons http://www.colorado.edu/physics/2000/quantumzone/bohr.html

Electron config in energy level

SO… We know that the e-’s are free to move around the nucleus We know that the e-’s are free to move around the nucleus They also can move from one energy level to the next (and fall) back when energy is added Move from ground state (“home” level) to a higher level (the “excited” state) Returning back to the ground state releases energy

This emission is how we see colors: This emission is how we see colors: the wavelengths of EMR released from an atom when it has been excited by Heat energy Electrical energy Chemical energy Think glowing red hot metal, or fireworks

Determining Energy for n To determine the energy for a given energy level, use the equation: En=(-RH)(Z/n2) Where n=1, 2, 3, 4…. And RH = 2.18E-18J, So En=(-2.18E-18J)(Z/n2)

To determine E emitted or absorbed: To determine the change in energy for a given energy transition: ΔE=Ef-Ei *Remember E=hν, so ΔE=hν if ΔE=[(-2.18E-18J)(Z/n2)]f- [(-2.18E-18J)(Z/n2)]i

E changes continued *Remember E=hν, so ΔE=hν to get the frequency of the light emitted or absorbed If ΔE is positive since Ef >Ei E is absorbed The e- was going from ground state to a excited state If ΔE is negative since Ef < Ei E is released The e- was going from excited state to a ground state

Also…life after Einstein and Bohr We know that electrons have characteristics of both light (waves) and matter, so we say that they have a dual nature

De Broglie De Broglie proposed that an electron moving about the nucleus had a wave-like behavior, so it has a particular wavelength associated with it. This wavelength depends upon the mass and velocity of the electron.  = h / mv mv = the momentum of the particle Mass* velocity = p momentum = p so p = mv therefore  = h / p

This matter-wave idea applies to all matter, not just to electrons However, the mass is so large, and the wavelength so small, that we cannot see it in macroscale objects This matter-wave theory led to applications like the electron microscope

De Broglie wavelength

Heisenberg: The Uncertainty Principle We can’t determine information about small scale objects the same way we can for large scale objects Case in point: a ball rolling down a ramp- we can get position, direction, and speed at the same time We can’t for electrons Hence, the uncertainty principle

Heisenberg, cont’d It is inherently impossible for us to simultaneously know both the exact momentum and exact location of an electron This is because anything we do to determine the location or momentum of the electron moves it from its original path and location; this can’t be reduced past a certain minimal level We can know only momentum or location- not both We can talk probability of the location/ momentum of an electron

Schrodinger’s Wave Equation .. Continued wave-particle theory Treated Hydrogen’s electron as a wave Also worked for other elements’ electrons (not just hydrogen like Bohr’s model). Known as the quantum mechanical model Limits electrons to certain energy levels (values) Does not try to describe an electron’s path around the nucleus

Schrodinger’s Wave Equation .. Very complex mathematical equation Solutions to equations are know as wave functions Is the probability of finding an electron in a certain space around the nucleus Predicts a 3 dimensional region known as an atomic orbital Looks like a fuzzy cloud Density of cloud is determined by probability of finding the e- there high probability = dense cloud (dark)

Sublevels and Orbitals The electrons are spread out in orbitals that have varying Shapes Energy (distance from nucleus) The orbitals are described in regards to their quantum numbers Descriptions that are descriptive and hierarchical There are 4 numbers that describe an orbital Written as follows: (#, #, #, ±#)

Principal quantum number (n) The first number (1, #, #,±#) Describe the distance from the nucleus of the orbital The energy of the orbital Values for n are integers The smallest possible value is 1 As the distance from the nucleus (and therefore energy) increases, the number increases

Quantum numbers

There periodic table and n The 7 periods on the periodic table correspond to n values Each period has a unique n value For the 1st period, n=1 For the 2nd period, n=2 And so on….

Angular Momentum (l) (this is a script l, as in llama) Angular Momentum (l) (this is a script l, as in llama) Is the shape of the sublevel It is the second number in the description (#,1, #, ±#) Range from _____________(although we never deal with anything above l=3) s =0 p =1 d =2 f = 3

The s sublevel l = 0 http://www.sfu.ca/~nbranda/28xweb/images/s_orbital.gif

p sublevel l = 1

d sublevel l = 2

Another look at d sublevel

f sublevel l = 3

General tutorials for electron configuration stuff some slides in this PowerPoint are from this site already http://www.wwnorton.com/chemistry/overview/ch3.htm See key equations and concepts (select from menu on the left), as well as the looking through the overview where to the tutorials are listed (links for just those are on the left, too)

Magnetic number (ml) Denote the orbital sublevel that is filled Denote the orbital sublevel that is filled It is the third number in the description (#,#,1, ±#) s sublevesl have one orbital; a sphere has one orientation in space p sublevels have three orbitals; 3 orientations in space d sublevels have five orbitals; 5 orientations in space f sublevels have seven orbitals; 7 orientations in space

“Flavors” of orbitals (ml) s sublevels have one orbital; a sphere has one orientation in space

“Flavors” of orbitals (ml) p sublevels have three orbitals; 3 orientations in space

“Flavors” of orbitals (ml) d sublevels have five orbitals; 5 orientations in space

“Flavors” of orbitals (ml) f sublevels have seven orbitals; 7 orientations in space

Spin Spin is +½ or -½ Up or down (could say clock wise and counter-clock wise)

Summary, excluding spin

How we use this…. There is a specific order to how the e- fill the orbitals; it is not random Although there are exceptions to the rules (last thing we do)

“The aufbau diagram shows the energy of each sublevel relative to the energy of other sublevels. Each box on the diagram represents an atomic orbital.” Excerpt From: Thandi Buthelezi, Laurel Dingrando, Nicholas Hainen & Cheryl Wistrom. “Chemistry.” McGraw-Hill Education, 2013. iBooks. https://itun.es/us/I0WrD.n

The principles of e- configuration The Aufbau (next) Principle: That e- fill the lowest energy sublevel before going to the next sublevel The Pauli Exclusion Principle: That e-s are paired according to opposite spins Hund’s Rule: e-s spread out in equal energy orbitals before pairing electrons

The first level to fill is the 1s level It is the lowest energy sublevel It holds two electrons They are oppositely paired (up and down- ↑↓) Each sublevel (each __) holds 2 electrons

Next… The second sublevel is the 2s sublevel It also holds 2 electrons (because s holds 2, not because of the number), also oppositely paired ↑↓

1s2, 2s2,then comes 2p6 So, as it states above 1s fills, 2s fills ,then comes 2p It holds up to six electrons Because p orbitals hold 6 electrons

Next… From 2p, Notice, you follow the arrows From 2p, 3s fills with 2e-, then onto 3p, with 6e- then 4s with 2e- followed by 3d with 10e- (because d holds 10e-) Then 4p with 6e- Notice, you follow the arrows Remember, the number of electrons comes from the letter (the orbital’s momentum, m)

The sublevels of the orbitals are first filled, then you continue onto the next level (Aufbau) Also be sure to place one electron in each sublevel prior to filling the level (↑ ↑ ↑ and not ↑↓ ↑ _) (Hund) e-s must be paired with e-s of opposite spin (↑↓, not ↑↑ or ↓↓) (Pauli)

Putting it all together… Carbon (neutral, so 6 electrons) What this would look like: ↑↓ ↑↓ ↑ ↑ _ 1s 2s 2p (notice there are 6 arrows for 6 electrons) This can also be written as 1s2 2s2 2p2 Notice the superscripts add up to 6

Nobel gases Last column of P.T. All end the same way with s2 p6 Last column of P.T. All end the same way with s2 p6 Neon (8 electrons) would look like: ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 1s 2s 2p Argon (18 electrons) would look like: ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 1s 2s 2p 3s 3p Can use this trend as a shortcut for writing out long e- configurations

Nobel gas shortcut Non-Nobel gas elements have inner e- configurations identical to certain Nobel gases Ne is 1s2 2s2 2p6 Na is 1s2 2s2 2p6 3s1 Write Nobel gas in brackets [Ne] instead of 1s2 2s2 2p6 After the brackets, write the electrons left over [Ne] 3s1 * When using the shortcut, use the Nobel gas from the period (row) before the element in question*

There are some exceptions… This is because some energy levels are very close together electrons are able to move between close orbitals in order to minimize repulsion Example: the 4s and 3d orbitals are very close in energy So exceptions for some period 4 d block elements occur Cr is not 1s2 2s2 2p6 3s2 3p6 4s2 3d4 Cr is 1s2 2s2 2p6 3s2 3p6 4s1 3d5 Because it takes less energy to split the electrons between the 5 sublevels than it does to put them together in the 4s and 3d

Valence electrons The electrons in the outermost orbital (usually highest energy level) Be is 1s2 2s1 has 1 valence electrons (2s1) N is 1s2 2s2 2p3 has 5 valence electrons (2s2 + 2p3 ) Determine the chemical properties of elements Used in forming chemical bonds

Valence electrons Can represent valence electrons visually using electron dot structure Dots stand for valence electrons Placed one at a time on the four sides of the symbol (they may be placed in any sequence) and then paired up until all are shown