Section 5.2

 If this rock were to tumble over, it would end up at a lower height. It would have less energy than before, but its position would be more stable. You will learn that energy and stability play an important role in determining how electrons are configured in an atom. 5.2

 In nature, change generally proceeds toward the lowest possible energy level  High energy systems are unstable and lose energy to become more stable  This is what you see when the crushed wintergreen mints emit light or when the samples of gases glowed- electrons are falling from a high energy level to a lower energy level  Electrons are arranged with lowest possible E level- we call this the electron configuration

 An electron configuration is the way that electrons are arranged in various orbitals around the nucleus  3 Rules must be followed to write an electron configuration  Aufbau Principle  Pauli Exclusion Principle  Hund’s Rule

 Electrons fill the orbitals with LOWEST ENERGY first  So in this chart, start with the 1s orbital and work your way upward

 An atomic orbital can hold at most 2 electrons  To occupy the same orbital, the electrons must have opposite spins  The opposite spins are indicated by arrows pointing in opposite directions (one up and one down)  Spinning electrons produce magnetic fields, which allow the electrons to attract (the attraction balances out the repulsion of like charges)

 Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.  In other words, electrons fill orbitals one at a time and have parallel spins.  After all orbitals in a sublevel have one electron, added electrons double up in orbitals and have opposite spins to electrons already there.

 For oxygen 8 electrons  The arrows represent the spin of the electron.  Notice that the electrons are spinning opposite ways when they are together and the same way in the orbitals with just one electron each  There are 6 paired electrons and 2 unpaired electrons

 Notice that all inner electrons are paired- there are 2 arrows in a box  Some outer electrons might not be paired- there may be only 1 arrow in a box

 Look at the periodic table to find the correct number of electrons  (look at the atomic number to find protons and for neutral atoms, the number of electrons matches the number of protons)  Use the Aufbau chart to figure out where the electrons go. Start with 1s and follow the arrows.  The numbers at the top of the chart indicate the maximum number of electrons that can be placed in that sublevel.  Fill up each sublevel until you have reached the correct number of electrons, which you found earlier on the periodic table

 Use the periodic table to help you check your work.  The last electrons written in your electron configuration should be the same as the principle energy level, sublevel, and number of electrons in the outermost energy level as indicated on your periodic table!

Valence electrons- outer shell electrons that determine the chemical properties of elements Since they are in the outer shell, you can find out how many are present by counting the electrons in the highest principle energy level This will ALWAYS include just the s and/or p orbitals Oxygen has 8 electrons 1s 2 2s 2 2p 4 The highest principle energy level is 2. It contains 6 electrons, so there are 6 valence electrons

 Shorthand electron configurations:  find the element you want on the P.T.  look up one row and find the noble gas in that row  Write the noble gas chemical symbol in brackets  Look back at the periodic table and find the information above the element you’re working with  Fill in the values for n and n-1 by looking at the row the element is in.  Make sure you pay attention to the presence of possible d and f orbitals if your element is in the p block.  You’ll need to look at the aufbau chart for this

 Example of Shorthand notation:  Chlorine: 17 electrons  Noble gas above chlorine: Ne  Neon has 10 electrons. The last 7 (to get to 17 electrons) must be written out.  Above Cl you will see ns 2 …np 5  Cl is in row 3, so replace n with 3.  There are no d orbitals after 3s, so write the answer like this:  [Ne]3s 2 3p 5

 According to the aufbau chart, copper should have this electron configuration:  Cu: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9  Instead, it has this electron configuration:  Cu: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10  Exceptions occur in groups 6B and 1B because:  Sublevels are most stable when they are full  Sublevels are fairly stable when they are ½ full  Sublevels lack stability when they are partly full  The d sublevel becomes more stable in groups 6B and 1B by stealing an electron from the previous s sublevel