1. General Chemistry: An Integrated Approach Hill, Petrucci, 4th Edition
Chapter 8 Electron Configurations, Atomic Properties, and the Periodic Table
Mark P. Heitz
State University of New York at Brockport
© 2005, Prentice Hall, Inc.

2. Multielectron Atoms
Electrons are attracted to the nucleus while simultaneously repelling one another
EOS
In the hydrogen atom, all subshells of a principal shell are at the same energy level
recall En = –B/n2
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

3. Multielectron Atoms
In a multielectron atom the various subshells of a principal shell are at different energy levels, but all orbitals within a subshell are at the same energy level
The increasing energy order of subshells is generally:
s < p < d < f
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

4. Multielectron Atoms
Orbital energies are lower in multielectron atoms than in the hydrogen atom
In higher numbered principal shells of a multielectron atom, some subshells of different principal shells have nearly identical energies
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

5. Electron Configurations
Electron configuration describes the distribution of electrons among the various orbitals in the atom
The spdf notation uses numbers to designate a principal shell and the letters to identify a subshell; a superscript number indicates the number of electrons in a designated subshell
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

6. Formula Mass
An orbital diagram uses boxes to represent orbitals within subshells and arrows to represent electrons:
EOS
Each box has arrows representing electron spins; opposing spins are paired together
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

7. Rules for Electron Configurations
Electrons occupy the lowest available energy orbitals
Pauli exclusion principle – no two electrons in the same atom may have the same four quantum numbers
EOS
Orbitals hold a maximum of two electrons
spins must be opposed
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

8. Rules for Electron Configurations
For orbitals of identical energy, electrons enter empty orbitals whenever possible – Hund’s rule
Electrons in half-filled orbitals have parallel spins
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

9. Rules for Electron Configurations
Capacities of shells (n) and subshells (l)
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

10. Rules for Electron Configurations
Subshell filling order ...
Each subshell must be filled before moving to the next level
EOS
1s22s22p63s23p6 ...
Illustration
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

11. The Aufbau Principle
A hypothetical building up of an atom from the one that precedes it in atomic number
(Z = 1) H 1s1
(Z = 2) He 1s2
(Z = 3) Li 1s22s1
EOS
(Z = 3) Li 1s22s1  [He]2s1
Abbreviated electron configuration
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

12. The Aufbau Principle ... [He]2p2 [He]2p3 [He]2p4 [He]2p5 [He]2p6
EOS
[He]2p6
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

13. Main Group and Transition Elements
Elements in which the orbitals being filled in the aufbau process are either s or p orbitals of the outermost shell are called main group elements
“A” group designation on the periodic table
The first 20 elements are all main group elements
In transition elements, the subshell being filled in the aufbau process is in an inner principal shell
EOS
Fourth period transition elements have n = 4 as their outermost shell as the 3d subshell fills
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

14. Main Group and Transition Elements
Completely filled and half-filled sublevels are more energetically favorable configurations
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

15. Periodic Relationships
The valence shell is the outermost occupied shell
The period number = principal quantum number, n, of the electrons in the valence shell
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

16. Periodic Relationships
For main group elements the number of valence shell electrons is the same as the periodic table “A” group number
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

17. Periodic Relationships
We can deduce the general form of electron configurations directly from the periodic table
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

18. Valence Electrons and Core Electrons
Valence electrons are those with the highest principal quantum number
EOS
Sulfur has six valence electrons
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

19. Valence Electrons and Core Electrons
Electrons in inner shells are called core electrons
EOS
Sulfur has 10 core electrons
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

20. Electron Configurations of Ions
Anions: gain e– to complete the valence shell
Example:
EOS -
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

21. Electron Configurations of Ions
Cations: lose e– to attain a complete valence shell
Example:
(Z = 11) Na
EOS
(Z = 11) Na+
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

22. Electron Configurations of Ions
Cations formed from transition metals lose e– from the highest principal energy level (n)
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

23. Magnetic Properties
Diamagnetism is the weak repulsion associated with paired electrons
Paramagnetism is the attraction associated with unpaired electrons
EOS
Ferromagnetism is the exceptionally strong attractions of a magnetic field for iron and a few other substances
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

24. Periodic Trends
Use must justify the trend across the period, you cannot simply state the trend.
A trend is an observation, not an explanation!
You should state the trend in your answer, but you must also go further by explaining what causes the observed trend!
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

25. Periodic Atomic Properties of the Elements
Periodic law states that certain sets of physical and chemical properties recur at regular intervals when the elements are arranged according to increasing atomic number
EOS
Consider atomic radii: distance between the nuclei of two atoms
The distance between the nucleus and the outer edge of the electron cloud
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

26. Atomic Radii
Atomic Radii decrease as atomic numbers increase in an given period (going across).
A proton and electron are added so the effective nuclear charge increases because each proton has more of an effect than each additional electron
As that attraction between the nucleus and electrons increases, and the atomic radius decreases
Atomic Radii increase Sgoing down
In going from top to bottom of a group, the valence electrons are assigned to orbitals with increasingly higher values of n (prin. Quantum number)
The underlying electrons requires some space, so the electrons of the outer shell must be further (Your are adding energy levels)

27. Atomic Radii
Zeff effective nuclear charge: the nuclear charge experienced by a particular electron in a multielectron atom
Increases the attraction of the nucleus and pulls the electron cloud closer to the nucleus resulting in a smaller atomic radius
Atomic radii of transtion metals trend a little differently
Exceptions in atomic radii also exist in the lanthanide and actinide series because of how the f subshells are uniquely filled by electrons

28. Zeff & Shielding
The order in which electrons are assigned to subshells in an atom, as well as other properties are because of Zeff
Shielding: electrons closer to the nucleus screen or shield the effect of nuclear charge on valence electrons
the number of shielding electrons increases when you reach the end of the periodic table and go on to the next period.
Shielding increases in steps as you start a new period or go down a group
Video
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

29. Transition Metal Atomic Trends
From left to right across a period, the radii initially decrease, then size remains almost the same, then slightly increases toward the end.
The small increase in atomic radii is because of the d subshell is filled with electrons and thus the ele-eletron repulsions cause the size to increase

30. Atomic Radii Properties
The increased number of energy levels (n) increases the distance over which the nucleus must pull and therefore reduces the attraction for electrons
Full energy levels provide shielding between the nucleus and valence electrons, so you see an increase in shielding as the level gets full
Illustration
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

31. Ionic Radii
The ionic radius of each ion is the portion of the distance between the nuclei occupied by that ion
If the size of an atom is determined by the outermost electrons, what happens if you remove or add an electron?
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

32. Ionic Radii
Cations are smaller than the atoms from which they are formed
– the nucleus attracts the remaining electrons more strongly
EOS
Anions are larger than the atoms from which they are formed
– the greater number of electrons repel more strongly
Think of the proton/electron ratio, -as electrons are lost, the ratio of p+/e- increases and so the electrons are held closer vv.

33. Isoelectronic Configurations
Isoelectronic species are elements that all have the same number of electrons
For isoelectronic species, the greater the nuclear charge, the smaller the species
Effective nuclear charge
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

34. Atomic and Ionic Radii
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

35. Ionization Energy
Ionization energy is the energy required to remove an electron from a ground state atom in the gaseous state
to remove an electron, energy must be supplied to overcome the attraction of the nuclear charge (endothermic, always +)
Continual removal of electrons increases ionization energy greatly
B  B+ + e– I = 801 kJ mol–1
B+  B+2 + e– I = 2427 kJ mol–1
B+2  B+3 + e– I = 3660 kJ mol–1
B+3  B+4 + e– I = 25,025 kJ mol–1
Illustration
EOS
B+4  B+5 + e– I = 32,822 kJ mol–1
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

36. Ionization energy
First ionization energy- energy is increased with each successive removal because the electron is being removed from an increasingly positive ion
The remaining electrons are held more tightly
Notice the large jump at the 3rd level for Mg.
There is a large increase as you remove electrons from lower (inner) energy subshells

37. First Ionization Energies
Illustration
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

38. Ionization energy
Ionization energy increases as atomic number increases in any given period
Zeff increases the attraction of the nucleus and holds the electrons more tightly
Exceptions: group II to III, IE drops because the p electrons do not penetrate the nuclear region as well as s electrons so aren’t as tightly held
Drop in IE also occurs between V & VI because of increased repulsion created by the first pairing of electrons, that is stronger than the increase in Zeff, lowering the energy required to remove the electron
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

39. Ionization energy
Ionization energy decreases as atomic number increases down a column or group
The increased number of energy levels (n) increases the distance over which the nucleus must pull, reducing the attraction for electrons
A full energy level provides some shielding between the nucleus and valence electrons
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

40. Electron Affinity
Electron affinity is the energy change that occurs when an electron is added to a gaseous atom
How much an atom ‘likes’
Electrons (+ or -)
-the more negative it is the higher the EA
(energy is flowing out of the system)
EOS
Electron affinities are expressed as negative because the process is exothermic
Illustration
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

41. Electronegativity
A measure of the attraction of an atom for the pair of outer shell electrons in a covalent bond with another atom
Pattern is same as electron affinity for same reasons
Both are attraction nucleus has for electrons, one in forming an ion (EA) and one in forming a molecule (EN)
Fluorine is the most electronegative. The closer it is to fluorine, the more electronegative it is.

42. Metals, Nonmetals, and Metalloids
Metals have a small number of electrons in their valence shells and tend to form positive ions
Except for hydrogen and helium, all s-block elements are metals
EOS
All d- and f-block elements are metals
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

43. Metals, Nonmetals, and Metalloids
Atoms of a nonmetal generally have larger numbers of electrons in their valence shell than do metals, and many tend to form negative ions
Nonmetals are all p-block elements and include hydrogen and helium
EOS
Metalloids have properties of both metals and nonmetals
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

44. Metals Metals react by losing electrons
A loosely held electron will result in a more reactive metal
This is tied directly to ionization energy
With an increased # of energy levels (n), comes increased distance from the nuclear attraction and thus a more loosely held electron available for reactions
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

45. Non-metals
Non-metals tend to gain electrons, a strong nuclear attraction will result in a more reactive non-metal
This means that an atom with the highest Zeff and the least number of energy levels should be the most reactive nonmetal (F) because its nucleus exerts the strong pull
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

46. A Summary of Periodic Trends
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

47. The Inert (Noble) Gases
The six noble gases, He, Ne, Ar, Kr, Xe, and Rn, rarely enter into chemical reactions
All have complete octets ...
= stability!
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

48. “Periodic” Behavior of Elements
Flame tests: elements with low first ionization energies are excited in a flame
EOS
Atoms emit energy when electrons fall from higher to lower energy states
FlameTests
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

49. “Periodic” Behavior of Elements
halogens (Group 7A) are good oxidizing agents
EOS
When Cl2 is bubbled in a solution containing iodide ions, chlorine oxidizes I– to I2
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

50. The s-Block Metals as Reducing Agents
2 K + 2 H2O  2 K+ + 2 OH– + H2
EOS
Recall activity series ...
H+ is reduced by these metals
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

51. Acidic, Basic, and Amphoteric Oxides
Acidic oxides are oxides that produce acids by reacting the oxide with water
e.g., SO3 + H2O  H2SO4
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

52. Acidic, Basic, and Amphoteric Oxides
Basic oxides are oxides that produce bases by reacting with water
e.g., MgO + H2O  Mg(OH)2
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

53. Acidic, Basic, and Amphoteric Oxides
Oxides that can react with either acids or bases are amphoteric oxides
e.g., Al2O3
Behavior of Oxides
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

54. Summary of Concepts
The wave-mechanical treatment of the hydrogen atom can be extended to multielectron atoms, but with two differences
Electron configuration is the distribution of electrons in orbitals among the subshells and principal subshells
EOS
There are two types of electron configuration notation:
spdf and orbital
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

55. Summary of Concepts
The aufbau principle describes a process of hypothetically building up an atom from the atom of the preceding atomic number
Elements in similar electron configurations fall in the same group of the periodic table
An atom with all the electrons paired is diamagnetic; an atom with one or more unpaired electrons is paramagnetic
EOS
Certain atomic properties, such as atomic radius, ionic radius, ionization energy, and electron affinity, vary periodically with increasing atomic number
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table

56. Summary of Concepts
The regions of the periodic table ascribed to metals, nonmetals, metalloids, and the noble gases are related to the value of atomic properties
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table