1. AP Chemistry Chapter 6 Electronic Structure and the Periodic Table

2. Nature of Light  Wavelengths and frequencies  Wavelength – distance between “troughs” Measured in meters or nanometer Measured in meters or nanometer 1nm = 10 -9 m 1nm = 10 -9 m  Frequency – number of of wave cycles that pass a given point in unit time Hertz – represents cycles per second Hertz – represents cycles per second Ex. If 10 10 cycles pass a particular point in one second v = 10 10 Hz Ex. If 10 10 cycles pass a particular point in one second v = 10 10 Hz

3.  λv = c  c = speed of light in vacuum  2.998 x 10 8 m/s  λ expressed in meters  v in hertz

4. Photon energies  E = hc λ Use to determine energy in joules of a photon emitted by an excited atom Also use to determine energy, in joules, of a mole of photons multiple by 6.02 x 10 23

5.  Remember to convert to kilojoules if necessary!!!  10 3 J = 1 kJ

6. Bohr model  Based on the hydrogen atom  Why the electrons that were circling the nucleus did not release their energy and spiral into the nucleus  He calculated energies associated with each allowed orbit

7.  E n = -2.178 x 10 -18 J n 2 n 2 E n = energy of the electron n = principal energy level This formula is included on AP constants sheet

8. Quantum Mechanical Model of electron placement FFFFirst quantum number, n designates Principal Quantum Level must be an integer Important for determining energy of the electron

9. l = Second Quantum Number  Determines shape of electron cloud  l = 0, 1,2,3……(n-1)  n = 1 l = 0  n = 2 l = 0, 1  n = 3 l = 0,1,2  n = 4 l = 0,1,2,3

10. m l = Third Quantum Number  Determines electrons orientation in space Corresponds to number of orbitals allowed in that sublevel ( l ) Corresponds to number of orbitals allowed in that sublevel ( l ) l = 0 m l = 0(1 orbital) l = 1 m l = 1,0,-1(3 orbitals) l = 2 m l = 2,1,0,-1,-2(5 orbitals) l = 3 m l = 3,2,1,0,-1,-2,-3(7 orbitals)

11. M s = Fourth Quantum Number; electron spin  Electron has magnetic properties like that of charged particles spinning on an axis  Either of two spins is possible – clockwise or counterclockwise  +1/2 -1/2

12.  Electrons with different m s values (one +1/2 and the other -1/2) (one +1/2 and the other -1/2) Said to have “opposed” spins Said to have “opposed” spins  Electrons with same value for m s (both +1/2 or -1/2) (both +1/2 or -1/2) Said to have “parallel” spins Said to have “parallel” spins

13. Pauli Exclusion Principle  No two electrons in an atom can have the same set of four quantum numbers  If they occupy the same orbital, must have opposing spins  Pg. 140 example 6.4, 6.5

14.  Review shape of sublevels s - sphere, p - figure-8 s - sphere, p - figure-8  Pg. 412 shape of d sublevel orbitals

15. Hund’s Rule  When several orbitals of equal energy are available, as in a given sublevel, electrons enter singly with parallel spins

16.  Solids – is possible to determine number of unpaired electrons in an atom by their behavior in a magnetic field  If unpaired electrons are present, the solid will be attracted into the field That substance called “paramagnetic” That substance called “paramagnetic”

17.  If the atoms in the solid contain only paired electrons, is slightly repelled by the field Called “ diamagnetic ” Called “ diamagnetic ”

18. Review electron configurations and orbital notation  Aufbau Principle Sublevels are filled in order of increasing energy Sublevels are filled in order of increasing energy Stability exceptions: Cr, Mo, W, Cu, Ag, Au Stability exceptions: Cr, Mo, W, Cu, Ag, Au break from strict Aufbau Principlebreak from strict Aufbau Principle

19.  Review “blocks” on Periodic Table  Lanthanides – filling 4 f  Actinides – filling 5 f (all of these elements are radioactive, only thorium and uranium are found in nature Stability decreases with increasing atomic number Stability decreases with increasing atomic number

20. Monatomic ions  Electrons are added to or removed from sublevels in the highest principal energy level  Want to achieve configuration like a noble gas – more stable  Na+1 (1s 2 2s 2 2p 6 ) + e -

21.  Species (whether ion or not) with same electron configuration called “isoelectronic”

22. Transition Metal Cations  When transition metals from positive ions, the outer s electrons are lost first  25 Mn [Ar] 4s 2 3d 5 to  Mn +2 [Ar]3d 5

23.  After the outer s electrons are lost, then the d can be lost  26 Fe [Ar]4s 2 3d 6 forms Fe +3  [Ar]3d 5  “first in, first out” rule

24. Periodic Trends  Atomic Radius – one half the distance of closest approach between atoms in an elemental substance (pg. 151 drawing) Decreases across a period from left to right Decreases across a period from left to right Increases down a group as atomic number increases Increases down a group as atomic number increases

25.  In a group, are increasing one whole pel as you go down The inner electrons “shield” the outer electrons from the positive nucleus The inner electrons “shield” the outer electrons from the positive nucleus  In periods, inner electrons are a poor “shield” because they are at about the same distance from the nucleus

26. Period trend continued  Effective nuclear charge (charge felt by an outer electron) increases steadily with atomic number  As effective nuclear charge increases, the outermost electrons are pulled in more tightly, and atomic radius decreases

27. Ionic Radius  Increases moving down a group  Both cations and anions decrease from left to right across a period  Positive ions smaller that their atoms  Negative ions larger than their atoms  Pg. 152 Figure 6.13

28.  Cation – excess of protons draws the outer electrons closer  Anion – extra electron adds to the repulsion between outer electrons (makes the negative ion larger that the corresponding atom)  Pg. 153 example 6.10

29. Ionization Energy  Measure of how difficult it is to remove an electron from a gaseous atom  Energy must always be absorbed to remove an electron, so always a positive quantity  First ionization energy – removal of outermost electron  X (g)  X + + e - ΔE 1 = first ionization energy ΔE 1 = first ionization energy

30. Trends  Increases across a period (left to right)  Decreases down a group (increasing a.n.)  Indirect relationship between atomic radius and ionization energy Large atom, electron far from the nucleus, easier to remove Large atom, electron far from the nucleus, easier to remove Smaller atom, electrons closer to the nucleus, held tighter, so harder to remove Smaller atom, electrons closer to the nucleus, held tighter, so harder to remove

31.  Pg. 153 Figure 6.15  First ionization energies in kJ/mol  Pg. 154 Example 6.11

32. Electronegativity  Measure the ability of an atom in a molecule to attract electrons to itself  The greater the electronegativity, the greater is its ability to attract electrons to itself  Dependent on ionization energy and electron affinity

33.  Electron affinity – tendency to form anions  EA = energy required to add an electron  Z + e - + energy  Z -  EA = energy released on removing e- from anion anion  Z -  Z + e - + energy

34.  If EA is large and negative, the atom “wants” to add an electron and form an anion  Atom with a very negative electron affinity and a high ionization energy will attract e - and resist any e - being removed from it.  Is highly electronegative

35. EN Trends  Period – generally a steady increase (metal to nonmetal)  Group – decrease within a group (are some exceptions)  Scale of electronegativities pg. 154  Important scale when we get to bonding!