Electrons and their Movement Vocabulary Electromagnetic Radiation Electromagnetic Spectrum Wavelength Frequency Quantum Photon Atomic emission spectrum Ground State Excited State Principal Quantum Number Principal Energy Level Energy sublevel Electron Configuration Aufbau Principle Pauli Exclusion Principle Hund’s rule Valence electrons Electron-dot Structure

Objectives: 1.Define a quantum of energy, and explain how it is related to an energy change of matter. 2.Compare the Bohr and quantum mechanical models of the atom. 3.Use the quantum numbers to describe the electrons in an atom. 4.Apply the Pauli exclusion principle, the aufbau principle, and Hund's rule to write electron configurations using orbital diagrams and electron configuration notation. 5.Identify the relationships among a hydrogen atom's energy levels, sublevels, and atomic orbitals. 6.Compare the wave and particle natures of light. 7.Explain the impact of de Broglie's wave article duality and the Heisenberg uncertainty principle on the current view of electrons in atoms.

Review: What does a Bohr Model include? How does this differ from Rutherford’s Model? What do we know about the Quantum Mechanical Model?

Rutherford Discovered the nucleus using the Gold Foil Exp. Rutherford’s model demonstrates his findings – Mostly empty space containing e-’s – At the center is a positive sphere (nucleus) – NOTE: there aren’t e- orbitals or neutrons until further models

Bohr Model An atomic physicist, Neils Bohr developed a planetary model based on Rutherford’s findings –Included a nucleus with neutrons –Electron’s moved in definite orbits (paths) around nucleus The e - ’s have fixed energies and do not lose energy as they orbit the nucleus

Quantum Mechanical Model (our current model) The electrons no longer occur in orbits – In this model, we have an idea of where the e - ’s could be, which is our e - cloud – The electron cloud is made up of different orbitals of different shapes The orbitals are not paths but rather probability densities

Diagram: (Fig. 13.3) The energy levels in an atom are sort of like rungs (steps) of a ladder. The more energy an electron has, the farther away from the nucleus it usually will be. The energy levels are not evenly spaced. They get closer together as you travel farther away. To move from one “rung” to another requires a “quantum” of energy. Energy Levels

Quantum Numbers Describe the location of the e - ’s around the nucleus. Quantum #’s are sort of like a home address for the electron. This information about the location of the e - ’s in an atom can be used to: (1) determine chemical & physical properties for the elements. (2) show how the Periodic Table is organized. (3) show how and why elements combine to form compounds. The four quantum numbers are abbreviated: n, l, m l, m s

The Principal Quantum Number: n n, the principal quantum number a) determines the overall energy of the atomic orbital b)Tells shell number c) May have any positive integral value from 1 to infinity example: 1, 2, 3, 4…etc. d) Tells the average distance to electron is from the nucleus -as n increases, the distance between the outermost electron and nucleus increases e) The distances can be called principal energy levels 1 2 3nucleus

l : tells us which one of the orbitals in the n shell we are looking at. All of these orbital shapes are based on the probability of finding the electron in the cloud. a) l describes the shape of the electron’s momentum around the nucleus with a letter: (s, p, d, & f) These are sometimes called “ sublevels ”. s= spherical cloud p= ellipsoid d & f orbital shapes are complex criscrossed ellipsoids, and some d’s and f’s are an ellipsoid with a doughnut or two around the middle. Angular momentum quantum number: l (a.k.a orbital quantum number)

d - orbitals f - orbitals s - orbital p - orbitals

M l : tells us the orientation of the orbital in space Each orbital has a specific number of orientations it can occur in: s= 1 orientation p= 3 orientations... (x, y, and z) d= 5 orientations f= 7 orientations The orientations can be represented with a line or a box. Examples: ___This means a spherical orbital at a distance of 1s“1” (the number 1 is the n) to the nucleus. This orbital is centered about the x, y, and z axis. □ □ □ This represents an ellipsoid orbital with its 4p 3 possible orientations at a distance of “4” from the nucleus. Magnetic Quantum Number: m l

m s : spin quantum number a) there are two possible orientations (spins) for electrons b) the two electrons may occupy the same orbital if they have different spins m s describes how the electron in an orientation is spinning around the nucleus. This spin can be thought of as “up” or “down”. (Some like to imagine it spinning “clockwise” and “counterclockwise”.) The spin is represented as an arrow in the direction of the spin. Example: ↑ This represents one electron in a spherical 2s orbital with spin “up” at a distance of “2” from the nucleus. Spin Quantum Number: m s

Figure 11.31: Orbitals being filled for elements in various parts of the periodic table.

What do we do with the quantum numbers? We said that quantum numbers are like an address for an electron but how do we write the address? Electron configurations are used to describe the location of electrons in an element. They incorporate the four quantum numbers and are specific to each neutral element. Electrons in these configurations follow four rules: 1)The Aufbau Principle 2)The Pauli exclusion principle (contains 2 rules) 3)Hund’s rule

Rule #1: Only 2 electrons can fit into each orbital Example: ___ ___ not ____ 1s 2s 1s Rule #2: Electrons in the same orientation have opposite spins. Example: ___ not ___ 1s 1s ↓ ↓↑↑↓↑↑ ↑↑↑ The Pauli Exclusion Principle

Rule #3: Electrons fill the lower energy orbitals first. -This means that energy level 1 will fill before 2. The Aufbau Principle Examples: 1s would be filled before 2s 3s would fill before 4s

↓ Hund’s Rule Rule #4 : “Bus seat rule”---> Every “ □ ” in an orbital shape gets an electron before any orientation gets a second e -. Example: □□□ not □□□ 2p 2p ↑↑↑↑↑

How to write an Electron Configuration 1)Find the element on the periodic table. (Ne) 2)List ALL of the prior orbital sublevels, l (w/energy level, n) Ne: 1s 2s 2p 3) Write the amount of electrons in each orbital -s: max of 2 e -p: max of 6 e -d: max of 10 e Ne: 1s 2 2s 2 2p 6

Electron Configurations Practice Problems: Write the electron configuration notation for each of the following atoms: H Li F Fe Br Kr

The Shorthand Shorthand Method: ex: Mg 1) Find the previous Noble Gas (last column on the periodic table) 2) Put the chemical symbol in brackets 3) Write the remaining e - configuration

Shorthand Method contd. Practice Problems Li F Fe Br Kr

Electron Configurations (Energy Level Diagram) Silicon: 1s 2, 2s 2, 2p 6, 3s 2, 3p 2 Electron Configurations can be drawn in an ENERGY LEVEL DIAGRAM -this shows energry levels rising as the n (principle Q#) increases -this diagram can aide in the order in which the energy levels and orbitals are filled AFTER Si, go to your notes and practice filling/drawing these for our previous examples.

Like all other rules, there are exceptions to 2 of the rules. They occur because of stability. THE FOLLOWING ELEMENTS DO THIS: Cr, Mo, Cu, Ag, Au (ANY OTHERS I WILL NOT HOLD YOU ACCOUNTABLE FOR) -Half-filled or completely filled d & f sublevels have lower energies and are more stable than partially filled d’s and f’s. -This means that an atom can “borrow” one of its “s” electrons from the previous orbital to become more stable. Example: ___ ___ ___ ___ ___ ___ 5s 4d becomes___ ___ ___ ___ ___ ___ 5s 4d Because the 4d sublevel is now full, the atom is at a lower energy state and therefore more stable. The Exceptions

Electron Configurations & Properties How do electron configurations relate to the chemical and physical properties of an element? -All elements w/ the same outer shell e - configurations have similar properties. -This means that elements in the same column (or group) have similar properties. Examples: (1) Li, Na, K, Rb, and Cs all have 1 lone “s” e - for their last orbital... ( 1s 1, 2s 1, 3s 1 …etc.) They all react with water to produce hydrogen gas. (2) Ne, Ar, Kr, Xe, and Rn all have the outer energy level completely filled with electrons...(2s 2 2p 6, 3s 2 3p 6, 4s 2 4p 6, etc.) This makes all of them inert. (inert means it doesn’t react)

More Practice Problems (1) Which element has its last electron as a 4p 5 ? ___________ (2) Which elements are similar in properties as Bromine? __________ (4) Which electron is added after 6s 2 ? ________ (5) Which element would “borrow” a 5s electron to get a half-filled “d” sublevel? ___________ (6) What is the shape of the last orbital filled in Calcium, (Ca)? _____ (7) How many electrons are in the last “p-orbital” of Sulfur, (S)? ____ Bromine F, Cl, I, At 4f 1 Mo sphere 4

What is Electromagnetic Radiation? Any wave of energy traveling at a speed of light is called electromagnetic radiation. - Waves are made of photons - equivalent to a quantum of light - Electromagnetic radiation can be broken down by frequency and wavelength into seven types of radiation - frequency is the number of repeating periods over a specified time. A period is the duration of one wavelength - Wavelength is the distance between the crests of a wave

Electromagnetic Radiation Below is the electromagnetic spectrum. It shows the types of radiation from the longest to the shortest wavelength. -This order also demonstrates the lowest to highest energies.

Electromagnetic Radiation (1) Radio Waves – longest wavelength, lowest energy: used in communications (2) Microwaves-- broadcasts TV signals and used to cook food. (3) Infrared (IR) -- we feel this as heat; Snakes & owls can “see” this. infrared image of heating pipes under a floor

Infrared Vision

Electromagnetic Radiation (4) Visible Light -- the only radiation we can detect with our eyes. It can be separated into the colors of the spectrum with a prism ROYGBIV (5) Ultraviolet (UV) -- gives you a sunburn; bees can “see” this; some of this radiation from the sun gets blocked by the ozone layer flower photo under normal lightflower photo under UV light

(6) X-Rays: used for medical imaging Electromagnetic Radiation

(7) Gamma Rays– In the universe, gamma rays are produced by neutron stars, pulsars, supernova explosions and is found near black holes. -ON EARTH- gamma rays are generated by nuclear explosions, lightning and radioactive decay UNFORTUNATELY… YOU WILL NOT BECOME HULK FROM EXCESSIVE EXPOSURE! Credit: NASA/DOE/Fermi LAT Collaboration, CXC/SAO/JPL- Caltech/Steward/O. Krause et al., and NRAO/AUI

COSMIC RAYS The last type of radiation is sometimes grouped with gamma rays… (8) Cosmic Rays– not true radiation; they are comprised of charged particles. Almost all of this radiation from the sun is blocked by the ozone layer and our magnetic field. Most is made up of matter, a little is made of anti-matter.

How is Light Produced? Light is produced by the movement of electrons in between orbitals within an atom - e - movement is caused by atoms getting hit with energy (heat/electricity) -e - absorb the energy jump energy levels - the electron falling back to its original state cause a release of energy as light or other forms of EMR

How Light is Produced -Each photon emitted has a specific frequency -The color of the light that is given off depends on how far the electron falls -The farther the fall, the greater energy

All the Photons Produced by Hydrogen

How hydrogen produces the four visible photons

How Light is Produced Electrons are located in certain energy levels (and orbitals) around the nucleus, only certain specific colors of light are emitted. Scientists use a spectroscope to separate these colors into bands of light. These bands of color look like a bar code of color which is characteristic of that element. No two elements produce the same spectrum of colors. This can be used to distinguish one element from another contained in a sample. (See Fig )

Chemical Spectra Chemists use machines called spectrometers to analyze substances. -All matter absorbs and reflects electromagnetic radiation differently. Our eyes see the visible light that was reflected by objects. -Chemists can compare information, called spectra, to previously recorded information to identify substances and predict molecular structure..

Emission Spectrum Hydrogen Spectrum Neon Spectrum

Chemical Spectra (IR spectra) caffeine

NMR Spectra NMR is similar to an MRI but it is used to see molecular structure - the spectra below is for benzoin