Chapter 5 CH1g,1h,2g,2i Electrons in Atoms

Actually, the Chemical History powerpoint talked about a lot of the stuff from Chapter 5 too, specifically part of section 1 and most of section 3.

This powerpoint will talk about what is still left over in Chapter 5.

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Quantum Mechanics…where we left off from Chem History

http://www.meta-synthesis.com/webbook/30_timeline/310px-Bohr-atom-PAR.svg.png

S Better than any previous model, quantum mechanics does explain how the atom behaves. S Quantum mechanics treats electrons not as particles (which they are), but as waves (like light) which can gain or lose energy.

S But they can’t gain or lose just any amount of energy. They can only gain or lose a “quantum” of energy.

A quantum

this idea.

is just an amount of energy that the electron needs to gain (or lose) to move to the next energy level. Max Planck, another German Nobel Prize winning scientist first came up with

What the heck is a Quantum?

http://www.blogcdn.com/www.slashfood.com/media/2008/08/splenda425.jpg

http://upload.wikimedia.org/wikipedia/commons/e/e9/Sucralose2.png

S Think of a quantum as a “packet” of energy, much like a sugar packet at a restaurant. A sugar packet contains a teaspoonful of sugar.

S If the electron absorbs energy, it moves to a higher energy level. If it emits (loses) energy, it moves to a lower energy level.

S But like Bohr suggested in his model, the electron has to gain or lose exactly the right amount. That amount is a quantum of energy.

Neils Bohr: The Planetary Model & Energy Levels

(http://www.usd.edu/phys/courses/phys300/gallery/clark/bohr.html) S You can’t just step anywhere.

S You have to step on the rungs of a ladder.

S An electron has to jump from one level to another.

S The steps on a ladder are all the same distance apart. S But in Bohr’s model, the energy levels get closer and closer the further away you get from the nucleus.

Quantum Mechanics “borrowed” the concept of energy level.

S The electron really doesn’t orbit (like a little planet) around the nucleus.

S Quantum mechanics describes “electron clouds” and where they are in relation to the nucleus.

S Again, the electron can move from one energy level to another, IF it absorbs a quantum of energy.

Energy Levels & Quantum Numbers

http://www.chem4kids.com/files/art/elem_pertable2.gif

S Quantum mechanics has a principal quantum number. S It is represented by a little n. It represents the “energy level” similar to Bohr’s model.

S S n=1 describes the first energy level n=2 describes the second energy level. Etc.

S Each energy level represents a period or row on the periodic table. S Isn’t it amazing how all this stuff just “fits” together?

Atomic Orbitals

http://milesmathis.com/bohr2.jpg

S The energy levels in quantum mechanics describe locations where you are likely to find an electron cloud. S Schroedinger used calculus to calculate the PROBABILITY of finding an electron in a particular location.

S These locations are called ORBITALS. S S Orbitals are “geometric shapes” around the nucleus where electrons are found.

There must be at least a 90% probability of finding an electron there. S The 4 different types of orbitals are s, p, d, and f.

Atomic Orbitals

http://courses.chem.psu.edu/chem210/quantum/quantum.html

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Think of orbitals as sort of a "border” for spaces around the nucleus inside which electrons are allowed.

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No more than 2 electrons can ever be in 1 orbital.

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The orbital just defines an “area” where you can find 1 or 2 electrons.

No more than 2 can fit into any one orbital.

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What is the chance of finding an electron in the nucleus?

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Yes, of course, it’s zero.

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There aren’t any electrons in the nucleus.

A node = a location where the probability of finding an electron there = 0.

Atomic Orbitals define an area where electrons are moving

http://www-hep.phys.unm.edu/~gold/phys492/orbitals.gif

3s 2s 1s S Quantum mechanics doesn’t predict a SPECIFIC orbit, like the Bohr model does.

S We don’t really know how the electron is moving, or if it follows any particular path as it moves.

Energy Sub-level = Specific Atomic Orbital

Blue = s block (0) S levels” which describe the specific S n = 1 has 1 sub-level (the “s” orbital) S S S n = 2 has 2 sub-levels (“s” and “p”) n = 3 has 3 sub-levels (“s”, “p” and “d”) and “f ”) S Periodic Table where those orbitals S A second quantum number identifies the specific orbital.

Shapes of These Orbitals

(the nucleus is ALWAYS at the center of the orbital) S The s orbital looks like a ball or sphere. S The p orbital looks like a dumb-bell.

S These orbitals are all perpendicular to each other.

S The d orbitals have two shapes.

S 4 of the 5 look like “4-leaf clovers.” S The 5 th one looks like a “big dumb-bell” with a “hula-hoop” around the middle.

S The shapes of the f orbitals are complex. S We have a slide showing them, but you don’t need to remember them, nor will they be on the test. But s, p and d will be.

Shapes of s, p, and d Orbitals

http://media-2.web.britannica.com/eb-media/54/3254-004-AEC1FB42.gif

http://upload.wikimedia.org/wikipedia/commons/thumb/e/e1/D_orbitals.svg/744px-D_orbitals.svg.png

S In the s block, electrons are going into s orbitals.

S In the p block, the s orbitals are full. S New electrons are going into the p orbitals.

S In the d block, the s is full but the p orbitals are not full. S New electrons are going into the d orbitals, because we are in the transition metals. THIS is characteristic of the d block.

f orbitals

http://antoine.frostburg.edu/chem/senese/101/electrons/faq/f-orbital-shapes.shtml

g orbitals = Science Fiction?

2,8,18,32…50?

http://antoine.frostburg.edu/chem/senese/101/electrons/faq/f-orbital-shapes.shtml

S Dr. Seaborg predicted the g orbitals would start with element number 121, which has not been invented yet. The g block will have 18 elements.

S Will his hypothesis be proven true?

To Summarize

Energy Level Sub levels Total Orbitals Total Electrons Total Electrons per Level

n = 1 n = 2 s s p 1 (1s orbital) 1 (2s orbital) 3 (2p orbitals) 2 2 6 2 8 n = 3 S S n = 4 S f s Complete the chart in your notes as we discuss this.

d s 1 (3s orbital) 3 (3p orbitals) 1 (4s orbital) We call this orbital the 1s orbital.

7 (4f orbitals) 2 10 2 6 10 14 18 32

Island of Stability

http://www.nytimes.com/1999/02/27/us/glenn-seaborg-leader-of-team-that-found-plutonium-dies-at-86.html

S This is another hypothesis from Dr. Seaborg. His thought was that element 114 would be an “island of stability,” especially if it also had 184 neutrons. It would aehv a mass number of 298.

S However, other “islands” might be 120 or 126. Detailed and complicated math calculations are necessary to figure out these numbers.

S Most synthesized elements only last for fractions of seconds. However, in 1998 researchers synthesized element 114 and it lasted for 30 seconds. Perhaps this is the “shore” of the Island of Stability that Dr. Seaborg hypothesized.

S S The element 114 was made using some of the original Pu-244 that Dr. Seaborg himself made in the early 1940s. They bombarded plutonium with Ca-48 atoms to form some of the new element 114.

Element 114 is now know as Flerovium (symbol Fl); it was named in 2012. S It took 14 years to agree on the name. S All of the atoms so far have had mass numbers of 285-289. Therefore, the “island” still remains undiscovered.

Island of Stability

http://www.sciencecodex.com/files/Island%20of%20Stability%201.jpg

http://physicsworld.com/cws/article/print/19751 Famous picture of the “Island of Stability” showing the island off in the distance (top right) with 114 protons and 184 neutrons. An element with Z = 184 is also predicted to be another “island of stability.”

Timeline = Homework

S Check out the History timeline on page 133 in your book.

S Prepare a timeline listing the major developments listed up to 1932.

S Answer questions 1&2 at the bottom of the page (2 requires a 5 sentence paragraph as a minimum requirement).

S Add 3 things to your timeline that have happened in Chemistry SINCE 1935 that you think are significant. You might have to do research to answer this.

Electron Configurations Section 2

S What do I mean by “electron configuration?” S The electron configuration is the specific way in which the atomic orbitals are filled.

S Think of it as being similar to your address. The electron configuration tells me where all the electrons “live.”

Rules for Electon Configurations

https://teach.lanecc.edu/gaudias/scheme.gif

S In order to write an electron configuration, we need to know the RULES.

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3 rules govern electron configurations.

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Aufbau Principle

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Pauli Exclusion Principle Hund’s Rule

S Using the orbital filling diagram at the right will help you figure out HOW to write them S S S Start with the 1s orbital. Fill each orbital completely and then go to the next one, until all of the electrons have been accounted for.

FOLLOW the arrows!!!

Fill Lowest Energy Orbitals

Each line represents ONE orbital.

FIRST

http://www.meta-synthesis.com/webbook/34_qn/qn3.jpg

1 (s), 3 (p), 5 (d), 7 (f) S

The Aufbau Principle states that electrons enter the lowest energy orbitals first.

S The lower the principal quantum number (n) the lower the energy.

S Within an energy level S s orbitals have the lowest energy S S followed by p then d and then f. f orbitals are the highest energy for that level.

No more than 2 Electrons in Any Orbital…ever.

http://www.fnal.gov/pub/inquiring/timeline/images/pauli.jpg

S The next rule is the Pauli Exclusion Principal.

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The Pauli Exclusion Principle states that an atomic orbital may only have 1 or 2 electrons and then it is full.

S The spins have to be paired.

S We usually represent this with an up arrow and a down arrow.

Hund’s Rule (Dog’s Rule?)

http://intro.chem.okstate.edu/AP/2004Norman/Chapter7/Lec111000.html

Don’t pair up the 2p electrons until all 3 orbitals are half full.

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Hund’s Rule states that when you get to degenerate orbitals, you fill them all half way first, and then you start pairing up the electrons.

S Degenerate means they have the same energy.

S p orbitals are degenerate because there are 3 of them on EACH level.

S d and f orbitals are also degenerate.

Let’s Try Some…

S NOW that we know the rules, we can try to write some electron configurations.

S Remember to use your orbital filling guide to determine WHICH orbital comes next in the sequence (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc).

S Follow the arrows!!

S Lets write some electron configurations for the first few elements, and let’s start with hydrogen.

S There are also shorthand electron configurations, but we will look at those after Chapter 6.

Electron Configurations

Element

H Z=1 Li Z=3 B Z=5 N Z=7 F Z=9 Na Z=11 K Z=19 Fe Z=26

Configuration

1s 1 1s 2 2s 1 1s 2 2s 2 2p 1 1s 2 2s 2 2p 3 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6

Element Configuration

He Z=2 Be Z=4 C Z=6 O Z=8 1s 1s 1s 1s 2 2 2 2 (1s is now full) 2s 2s 2s 2 2 2 (2s is now full) 2p 2p 2 4 Ne Z=10 1s 2 2s 2 2p 6 (2p is now full) Cl Z=17 1s 2 2s 2 2p 6 3s 2 3p 5 Sc Z=21 Br Z=35 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5

Electron Configurations of Alkali Metals (and H)

Element

H Z=1 Li Z=3 Na Z=11 K Z=19

Configuration

1s 1 1s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1

Exceptions to the Rules for Electron Configurations

Element

Cr should be BUT Cr is Cu should be BUT Cu is

Configuration

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 (Z=24) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 (d half full) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 (d is full)

More HW…OMG!

S Chemistry: write full electron configurations for elements 1-36.

S Also write orbital diagrams for 3-10 S Advanced Chemistry: write full electron configurations for 1-36 + Rb, Sr, Y, Ag, I, Kr, Cs, Ba, La, Ce, Hf, Pb.

S Also write orbital diagrams for 11-18

Emission Spectra = Fingerprint of the Elements (Section 3)

http://www.cbu.edu/~jvarrian/252/emspex.jpg

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Atomic emission spectrum is sometimes called a line spectrum, to distinguish it from the continuous spectrum.

Emission Spectra = Fingerprint of the Elements (Section 3)

http://www.cbu.edu/~jvarrian/252/emspex.jpg

S The top 3 (H, Hg, Ne) are emission spectra.

S The bottom one is an absorption spectrum of H.

Emission Spectra = Fingerprint of the Elements

S the spectrum to identify the element (like a S Bohr’s model predicted and explained emission S one energy level to another.

they are heated.

S analyze it and determine what types of elements S Just by looking at the light!

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All the Rest of Section 3….

S …was covered in the Chemical History power point.

S Photoelectric effect S S S A photon is a quantum of light. It is light behaving as a particle. A photon has a certain wavelength, frequency and energy.

De Broglie equation S Showed that particles could also act as waves.

Heisenberg’s uncertainty principle S Principal = Dr. Gordon S Principle = a statement that explains how or why something works scientifically

The End

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