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EMR and the atom: Part Deux http://imagers.gsfc.nasa.gov/ems/waves3.html Electron Configurations

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First, a review… 1.What do scientists mean when they say light is quantized? 2.Compare a ground state and an excited state electron. 3.Explain the origin of the atomic emission spectrum of an element. 4.What do we know so far about the arrangement of electrons in atoms?

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

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Which is really not true- why? Because orbitals- the “ electron cloud ” are – 3-D, not flat – are not round in most cases – e - s are spread out as much as possible – (e - s are moving very rapidly)

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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: (#, #, #, ±#)

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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

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Quantum numbers

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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 1 st period, n=1 – For the 2 nd period, n=2 – And so on….

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The shape of things Is the shape of the orbital There are 4 shapes (although we only deal with the first three) – s = sphere – p = peanut – d = double peanut – f = flower

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The s orbital http://www.sfu.ca/~nbranda/28xweb/images/s_orbital.gif

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p orbitals

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d orbitals

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f orbitals

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General tutorials for electron configuration stuff some slides in this PowerPoint are from this site already http://www.wwnorton.com/college/chemistry /gilbert/tutorials/ch3.htm http://www.wwnorton.com/college/chemistry /gilbert/tutorials/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)

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

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“ Flavors ” of m l s orbitals have one sublevel; a sphere has one orientation in space

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“ Flavors ” of m l p orbitals have three sublevels; 3 orientations in space

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“ Flavors ” of m l d orbitals have five sublevels; 5 orientations in space

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“ Flavors ” of m l f orbitals have seven sublevels; 7 orientations in space

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Spin It is the last number in the description (#,#,#,±½) Spin is +½ or -½ – Up or down

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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)

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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 sublevels before placing electrons

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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

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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 ↑↓

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

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Next… 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 l )

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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)

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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 1s 2 2s 2 2p 2 Notice the superscripts add up to 6

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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 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 – Cr is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 – 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

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Electron Arrangement in Atoms. Electron configurations Aufbau principle Pauli exclusion principle Hund’s rule.

Electron Arrangement in Atoms. Electron configurations Aufbau principle Pauli exclusion principle Hund’s rule.

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