Docstoc

CHAPTER 8_ ELECTRON CONFIGURATIONS AND CHEMICAL PERIODICITY

Document Sample
CHAPTER 8_ ELECTRON CONFIGURATIONS AND CHEMICAL PERIODICITY Powered By Docstoc
					CHAPTER 8: ELECTRON CONFIGURATIONS AND CHEMICAL PERIODICITY Electron Configuration - shows the distribution of electrons into the ________ ________ Three rules govern the placement of electrons into the orbitals: 1. The ________ principle The process of building up atoms by adding one electron at a time, beginning with the lowest energy orbitals, in order to obtain the electron configuration for the element in question 2. ________ Rule Whenever orbitals of equal energy are available, electrons always arrange themselves as to have the highest number of ________ ________ 3. The ________ ________ Principle No two electrons may have the same set of quantum numbers (another way of saying that an orbital can hold a maximum of two electrons)

The Energies of Orbitals 1. The lower the ______ ______ ______, the lower the energy ______ orbital has lower energy than ______ ______ orbital has lower energy than ______

1

2. Within a shell, the lower the value of ________ the lower the energy of the orbital ______ orbital has lower energy than ______ ______ orbital has lower energy than ______ 3. Orbitals than have the same quantum number ______ but different values of ______ all have the same energy all three 2p orbitals have the same energy all three 3p orbitals have the same energy all five 3d orbitals have the same energy

Orbital Filling Order The orbitals are therefore filled in the following order: (memorize this) 1s > 2s > 2p > 3s > 3p > 4s > 3d > 4p > 5s Ways of Representing the Orbitals 1. Orbital Diagrams – Boxes or (or lines or circles) represent ________ Arrows represent the ________ Direction of arrow represents ________ ______ Orbitals with same ________ are grouped together  1s  2s  2p 3s

2

2. Orbital diagrams can also be shown vertically 3s 2p 2s 1s   

3. Shorthand notation – Superscripts represent the number of electrons in each orbital 1s22s22p1 The Electron Configurations Of The Elements 1. Hydrogen – 1s1 Electron must go into the 1s orbital (Aufbau Principle)  1s

2s

2p

3s

2. Helium – 1s2 Second electron also goes into the 1s orbital 2s doesn’t begin to fill until 1s is filled  1s

2s

2p
3

3s

3. Lithium – 1s22s1 1s orbital is filled 1s can’t take a third electron (Pauli exclusion principle) Begin filling second shell; third electron goes into the 2s orbital  1s  2s

2p

3s

4. Beryllium – 1s22s2 2s not completely filled Fourth electron goes into the 2s orbital  1s  2s

2p

3s

5. Boron – 1s22s22p1 1s completely filled; 2s completely filled Fifth electron goes into the 2p orbital Doesn’t matter which 2p orbital – we make no distinction  1s  2s  2p 3s

6. Carbon – 1s22s22p2 Sixth electron goes into vacant 2p orbital Would be incorrect to pair the two 2p electrons (Hund’s Rule) Carbon has two unpaired electrons  1s  2s   2p

3

4

7. Nitrogen – 1s22s22p3 Seventh electron goes into remaining vacant 2p orbital Has three unpaired electrons  1s  2s   2p  3s

8. Oxygen – 1s22s22p4 All three 2p orbitals are half-filled 3s cannot be filled until all 2p orbitals are completely filled Eighth electron must pair with one of the three 2p electrons  1s  2s   2p  3s

9. Fluorine – 1s22s22p5 Ninth electron completes the filling of a second 2p orbital  1s  2s   2p  3s

10. Neon – 1s22s22p6 Tenth electron goes into remaining half-filled 2p orbital Second energy shell is completely filled  1s  2s   2p  3s

5

The Elements of the Third Shell Same pattern all over again 11. Sodium – 1s22s22p63s1  1s  2s    2p  3s

3p

12. Magnesium – 1s22s22p63s2  1s  2s    2p  3s

3p

13. Aluminum – 1s22s22p63s23p1  1s  2s    2p  3s  3p

14. Silicon – 1s22s22p63s23p2  1s  2s    2p  3s   3p

15. Phosphorous – 1s22s22p63s23p3  1s  2s    2p  3s   3p 

16. Sulfur – 1s22s22p63s23p4  1s  2s    2p
6

 3s

  3p



17. Chlorine – 1s22s22p63s23p5  1s  2s    2p  3s    3p

18. Argon – 1s22s22p63s23p6  1s  2s    2p  3s    3p

The Elements of the Fourth Shell 19. Potassium - 1s22s22p63s23p64s1  4s

3d

4p

20. Calcium - 1s22s22p63s23p64s2  4s

3d

4p

21. Scandium - 1s22s22p63s23p64s23d1  4s  3d 4p

22. Titanium - 1s22s22p63s23p64s23d2  4s   3d 4p

7

23. Vanadium - 1s22s22p63s23p64s23d3  4s    3d

4p

24. Chromium – one of two exceptions What we expect is 1s22s22p63s23p64s23d4 What is actually observed is 1s22s22p63s23p64s13d5 An electron is “borrowed” from 4s to half-fill the 3d  4s    3d   4p

25. Manganese - 1s22s22p63s23p64s23d5  4s    3d   4p

26. Iron - 1s22s22p63s23p64s23d6  4s    3d   4p

27. Cobalt - 1s22s22p63s23p64s23d7  4s    3d   4p

8

28. Nickel - 1s22s22p63s23p64s23d8  4s    3d   4p

29. Copper - second exception What we expect is 1s22s22p63s23p64s23d9 What we observe is 1s22s22p63s23p64s13d10 An electron is borrowed from 4s to fill the 3d  4s    3d   4p

30. Zinc - 1s22s22p63s23p64s23d10  4s    3d   4p

9

Inner (core) Electrons – Those in the previous ______ ______ and any ______ ______ ______ Outer Electrons – those electrons in the ______ ______ ______ (higest n value); they spend most of their time farthest from the nucleus Valence Electrons – those electrons involved in ________ ________. For the main group-electrons, the outer electrons are the same as the valence electrons Example: Sodium – 1s22s22p63s1 1s, 2s, 2p electrons = ______ electrons 3s = the ______ electron Abbreviated Electron Configurations Use the ______ ______ ______ to represent the core electrons Example: Sodium = 1s22s22p63s1 Neon = 1s22s22p6 So we can abbreviate sodium as [Ne] + 3s1 Second Example: Calcium = 1s22s22p63s23p64s2 Argon = 1s22s22p63s23p6 So we can abbreviate calcium as [Ar] + 4s2

10

The Basis for Chemical Periodicity All elements within a group have ______ ______ ______ All elements within a group have ______ ______ ______ – the same number of ______ ______ Example: All alkali metals have just ______ valence electron Hydrogen = 1s1 Lithium = 1s22s1 = [He] + 2s1 Sodium = 1s22s22p63s1 = [Ne] + 3s1 Potassium = 1s22s22p63s23p64s2 = [Ar] + 4s1 All form a +1 ion because they have just the one valence electron; my losing an electron they revert to a _____ ______ ______ Example: All halogens have seven ______ electrons (s2p5) Fluorine = 1s22s22p5 = [Ne] + 2s22p5 Chlorine = 1s22s22p63s23p5 = [Ne] + 3s23p5 Bromine = 1s22s22p63s23p64s23d104p5 = [Ar] + 4s24p5

11

Trends in Some Key Periodic Atomic Properties 1. Trends in atomic size Measured in terms of how close one atom lies to another ______ Radius : one-half the difference between nuclei of adjacent atoms in a crystal of the element ______ Radius : one-half the distance between nuclei of identical covalently bonded atoms  Moving down a ______, the atomic radius ______ Each successive element has the same number of electrons but in the next highest shell Atomic radii (in picometers) H 37 Li 152 Na 186 K 227 Rb 248

 Moving across a ______ the atomic radius ______ Even though we are moving towards atoms with greater numbers of electrons and greater numbers of protons, we are building up within the ______ ______ The greater ______ ______ in the nucleus and the greater ______ ______ in the orbitals leads overall to smaller atoms Atomic radii (in picometers) Li 152 Be 112 B 85 C 77
12

N 75

O 73

F 72

Ne 71

2. Trends in Ionization Energy Defined as the energy required to ______ ______ the outermost valence electron from an atom  Moving down a ______, the ionization energy ______ Moving down a group we encounter atoms with increasingly larger ______ _____ electrons spend more time further from the nucleus and are therefore less tightly held Electrons in outermost orbitals are also said to be ________from the nucleus due to all the electrons in the inner orbitals; the ______ ______ felt by these electrons is diminished Li Na K Rb 520 kJ/mol 496 kJ/mol 419 kJ/mol 376 kJ/mol

Moving across a ______, the ionization energy ______ Moving towards smaller atoms, where outermost electrons will be more tightly held and therefore more difficult to remove Li Be B C N O F Ne 520 kj/mol 900 kj/mol 800 kj/mol 1090 kj/mol 1400 kj/mol 1310 kj/mol 1680 kj/mol 2080 kj/mol

13

Trends In Successive Ionization Energies ______ ionization energy – energy required to remove one electron from an atom (always the smallest) ______ ionization energy – energy required to remove a second electron (always larger than the second) ______ ionization energy – energy required to remove a third electron (always larger than the third) There is always a big jump in ionization energy once all the ________ electrons are removed; any further electrons would be ________ electrons Na = 1s22s22p63s1 Na+ = 1s22s22p6 Na2+ = 2s22s22p5 First ionization energy = 496 kJ/mol Second ionization energy = 4560 kJ/mol Mg = 1s22s22p63s2 Mg+ = 1s22s22p63s1 Mg2+ = 1s22s22p6 First ionization energy = 737 kJ/mol Second ionization energy = 1450 kJ/mol Third ionization energy = 7731 kJ/mol * Sodium never forms Na2+ * Magnesium never forms Mg3+

14

3. Trends In Electron Affinities Electron affinity – the energy change the addition of 1 mole of ______ to 1 mole of ______ ______ * Usually has a ______ value for H (energy released rather than energy required) * Moving down a ________, the general trend is towards ______ ______ values * Moving across a ______, the general trend is towards ______ ______ values The alkali metals (H, Li, Na, K, Rb, Cs) all have ______ ______ electron attachment enthalpies H Li Na K -72.8 kJ/mol -59.6 kJ/mol -52.9 kJ/mol -48.4 kJ/mol

The halogens all have ______ ______values (F, Cl, Br, I) all have large negative values – the tendency to form the negative ion is very strong F Cl Br I -328 kJ/mol -349 kJ/mol -325 kJ/mol -295 kJ/mol

The noble gases have small negative or positive electron attachment enthalpies (have little tendency to form a negative ion)

15

Summary of Periodic Properties 1. Atomic Radii: size of atoms increases moving down a group decreases moving across a period 2. Ionization energy – energy required to remove an electron decreases moving down a group increases moving across a period 3. Electron Affinity – enthalpy change associated with the addition of an electron generally less negative values moving down a group generally more negative values moving across a period

Alkali Metals, Alkaline Earth metals Lose electrons readily Attract electrons weakly Form postitive ions Oxygen and Halogen families Much higher ionization energies Attract electrons much more readily Form negative ions Noble Gases Very high ionization energies (do not lose electrons) Very little tendency to attract electrons
16

Trends in Metallic Behavior Properties of Metals * tendency to give up an electron and form a ______ ______ * ______ solids * good conductors of ______ and ______ * react with ______ to form ______ ______ Metallic behavior ________ moving left to right across a period Related to increasing ______ ______ moving left to right across a group; decreasing tendency to lose an electron means decreasing metallic behavior Na, Mg, Al – metals Si – metalloid P,S,Cl - nonmetals (do not form positive ions at all )

Metallic behavior increases moving top to bottom down a group: N – nonmetal P – nonmetal As, Sb – metalloids Bi – definite metallic properties (forms a +3 cation)

17

Acid-Base Behavior of the Element Oxides Metal Oxides – dissolve in water to form ______ solutions Nonmetal Oxides – dissolve in water to form ______ solutions

Acid-Base trends moving across a group: Na2O, MgO – react with water to form basic solutions of NaOH Mg(OH)2 SO3, Cl2O7 – react with water to form acidic solutions of H2SO3/H2SO4 and HClO4 Acid-base trends moving down a group: Since metallic properties ________ moving down a group, the oxides of the heavier members of a group should be more ______ N2O5 – reacts with water to form an ______ solution of HNO3 Bi2O3 – reacts with water to form a ______ solution of Bi(OH)3 Electron Configurations of Transition Metal Ions Transition metal ions rarely lose enough electrons to obtain noblegas electron configurations – for most this would involve the loss of too many electrons Remember successive ionization energies – each successive electrons is harder to remove - +3 ions can form, a few +4 ions are known, +5 ions are not formed

18

Electrons are lost out of the 4s orbital first, and any additional electrons are removed from the 3d Ti = 1s22s22p63s23p64s23d2  4s   3d 4p

Ti+2 = 1s22s22p63s23p63d2  4s  3d 4p

Fe = 1s22s22p63s23p64s23d6  4s    3d   4p

Fe3+ = 1s22s22p63s23p63d5  4s   3d   4p

Cu = 1s22s22p63s23p64s13d10  4s    3d   4p

Cu+ = 1s22s22p63s23p63d10  4s   3d   4p

19

________ – the tendency for a species with unpaired electrons to be attracted to an external magnetic field ________ – the tendency for a species with no unpaired electrons to be unaffected by a magnetic field Cu – paramagnetic Cu+ - diamagnetic Fe, Fe3+ = both paramagnetic Can determine the number of unpaired electrons in a material from magnetic measurements

The Relative Sizes of Ions Cations are ________ than their parent atoms Removing an electron results in a positive charge; the remaining electrons are pulled in closer to the nucleus Anions are ________ than their parent atoms Electrons are added to an outer level without any increase in the positive charge of the nucleus

20


				
DOCUMENT INFO