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CHAPTER 7
The Structure of
Atoms and Periodic
Trends
Chapter Outline
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7-1 The Pauli Exclusion Principle
7-2 Atomic Subshell Energies and Electron
Assignments
7-3 Electron Configuration of Atoms
7-4 Electron Configuration of Ions
7-5 Atomic Properties and Periodic Trends
7-6 Periodic Trends and Chemical Properties
Arrangement of Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
Chapter Goal: Use our knowledge of atomic structure
to rationalize chemical behavior.
The Pauli Exclusion Principle •  The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers n, l, ml and ms. –  For a given orbital the values of n, l, and ml are fixed. –  If we want to put more than one electron in an orbital they must have different ms values. –  Thus, an orbital can hold a maximum of two electrons, and they must have opposite spins. Wolfgang Pauli 1900-­‐1958, Austrian physicist, 1945 Nobel Prize in physics "for the discovery of the Exclusion Principle, also called the Pauli Principle" 4 Orbital energies In hydrogen electron energy is only dictated by n Because of electron interac5ons, larger atoms have energy dictated by n AND l 1.) The energy of an electron is ordered relaPve to n+l 2.) For electrons with same n+l, the one with a lower n is lower in energy 5 PenetraPon and Orbital Energy The 2s orbital distribuPon has a small probability contribuPon closer to nucleus (orbital penetra,on) This region lowers the screening of a 2s electron by the 1s electron 2p d6 oesn’t have this penetraPon, so it is higher in energy (less stable) Factors AffecPng Atomic Orbital Energies The Effect of Nuclear Charge (Zeffective or Z*)
Higher nuclear charge lowers orbital energy (stabilizes the
system) by increasing nucleus-electron attractions.
The Effect of Electron Repulsions (Shielding or Screening)
Additional electron in the same orbital
An additional electron raises the orbital energy through
electron-electron repulsions.
Additional electrons in inner orbitals
Inner electrons shield outer electrons more effectively than
do electrons in the same sublevel.
Problem Determine the relaPve energies of all the orbitals in n=1, 2 and 3 8 Orbital Filling: The Aufbau Principle & Hund’s Rule
•  Aufbau Principle:
•  Lower energy orbitals fill first to give the ground state.
• 
Hund’s Rule:
•  Degenerate orbitals (those of the same energy) are filled with
electrons until all are half filled before pairing up electrons.
•  Pauli exclusion principle:
•  Individual orbitals only hold two electrons, and each should have
different spin.
Orbital Filling: Notation
Orbital diagram: represent the orbital with a box
and the electron with an arrow
spdf notation: write the orbital with a superscript
indicating the number of electrons in that orbital
Noble gas notation: represent the core electrons with a noble
gas symbol [X] and write out the valence electrons
Problems
Write out (i) full spdf notation, (ii) orbital diagrams and (iii) noble gas
notation for the following:
(a) Li (b) B (c) N (d) F (e) Ne (f) Sc and (g) Ti
Transition Metals
Transition elements have partially filled d-orbitals
Electron filling follows the same principles as we have already
discussed.
Two notable exceptions involve having a ground state
configuration with an unfilled s-orbital in order to:
1.) maximize the electron spin in the d-orbitals (Hund)
2.) Completely filling out the d-orbitals
Problems
Write out (i) full spdf notation, (ii) orbital diagrams and (iii) noble gas
notation for the following:
(a) Cr (b) Cu
Electron Configuration of Ions
Start with electron configurations of the neutral element
For:
Anions: add electrons to obtain a noble gas configuration
Cations: remove electrons from orbital with highest n until
a particularly stable configuration is obtained (filled subshell,
high spin)
Problems
Write out (i) full spdf notation, (ii) orbital diagrams and (iii) noble gas
notation for the following:
(a) Ca2+ (b) Ti2+ (c) Cr+ (d) Ag+
Problems
We have already encountered transition metal ions that can take on
more than one charge state.
Predict the (i) full spdf notation, (ii) orbital diagrams and (iii) noble gas
notation the following:
(a) Fe2+ and Fe3+ (b) Cu+ and Cu2+ (c) Co2+ and Co3+
Magnetic Properties of Atoms and Ions
Paired electrons – diamagnetic
Ferromagnetic- metals with magnetic properties
Unpaired electrons - paramagnetic
attracted by a magnetic field
attraction proportional to number of unpaired e–
Problem
Suppose an element has the electronic
configuration:
1s2 2s2 2p6 3s2 3p6 4s13d3
Is the element in the ground state?
Problem
(a) Depict the electron configuration of manganese, Mn,
and its 4+ cation, Mn4+, using noble gas configuration
and orbital diagrams.
(b) Determine the magnetic properties of MnO2.
Will this substance be more or less magnetic than solid
manganese Mn (s)
General Periodic Trends
Elements were arranged in the periodic table historically based on their
chemical properties.
The modern understanding of atomic structure allows us to rationalize this
arrangement
Will look at three properties:
Atomic and ionic size
- radius
Ionization energy
- complete removal of an electron
Electron affinity
- addition of an electron
General Periodic Trends
All three properties can be rationalized by considering effective nuclear
charge. Higher Z* yields more tightly bound electrons and:
Smaller atomic and ionic size
- electrons orbit closer to nucleus
Larger Ionization energy
- more tightly bound = more energy required to remove
More negative electron affinity
- addition of an electron more favorable
Atomic Radii of Elements
IonizaPon Energy •  IonizaPon energy – energy (in kJ) required for the removal of 1 mol of electrons from 1 mol of atoms or ions in the gas phase. •  First ionizaPon energy, IE1, is the energy needed to remove the first electron from a neutral atom. Na(g) → Na+(g) + e-­‐ •  Second ionizaPon energy, IE2, is the energy needed to remove the second electron. Na+(g) → Na2+(g) + e-­‐ –  IonizaPon energies are always posiPve, i.e., removing an electron is endothermic. 24 Problem Predict the periodic trend for first ionizaPon energy: (a)  Across a period (b)  Down a group Removing More Than 1 Electron
The general trend: IE1 < IE2 < IE3 < IE4…
Notice that the jump from IE1 to IE2 for Be is not as great as that for Li.
Use orbital diagrams to rationalize
Problem: Determine the group that element ‘X’ belongs to:
50,000
45,000
E, kJ
40,000
35,000
X
30,000
25,000
20,000
15,000
10,000
5,000
0
1st
2nd
3rd
4th
5th
6th
Ionization Energies
7th
8th
Electron Affinity •  Electron Affinity (EA) – energy change accompanying the addiPon of 1 mol of electrons to 1 mol of gaseous atoms or ions. –  In most cases, energy is released when the first electron is added. X(g) + e-­‐ → X-­‐(g) ΔE = EA1 28 Electron Affinity EA1 show numerous excepPons to expected periodic trends. Group 6A & 7A – Highly negaPve EA1 values, tend to form anions. Group 1A & 2A – Only slightly negaPve EA1 values, tend to form caPons. Noble Gases – Slightly posiPve EA1 and very high IE1, tend not to form ions. In general, groups 2 and 5 have less electron affinity than expected. Can this be ra5onalized with electron configura5ons? 29 Electron Affinity - Problem
Would you expect the following reaction to occur:
Br2 + I− → I2 + Br −
Problem: ion sizes
Rank the following ions in order of decreasing size.
Na+,
N3-,
Mg2+,
F–
O2–
Notice that they all have 10 electrons: They are
isoelectronic (same electron configuration) as Ne.
Ion size
Group Trends for the
Active Metals
Group 1A:
The Alkali Metals
Form +1 ions because
removing core electrons
energetically unfavorable
Alkali Metals + H2O
2 M (s) + 2 H2O (l) → 2 MOH (aq) + H2 (g)
This reaction typically more favorable for metals with
lower ionization energies
Group Trends for the
Active Metals
Group 7A:
The Halogens
Form -1 ions because this give
a noble gas configuration which
fills out p subshell.
Halogens
Group 7A: The Halogens
The chemistry of the halogens is dominated by gaining an
electron to form an anion:
X2 + 2e− → 2X−
With water:
2F2(g) + 2H2O(l) → 4HF(aq) + O2(g) ΔH = −758.7 kJ
This reaction typically more favorable for halogens
with higher electron affinities
Suggested problems:
from the end of chapter 7 of the text book: Electron configura5ons: 1, 3, 5, 7, 9 Quantum numbers: 11, 13, 15 Ions: 17, 19, 21 Periodic proper5es: 23, 25, 27, 29, 31 General: 37, 39, 41, 43, 45, 47, 49, 51, 53, 55, 57, 59, 63, 65, 67, 69, 71, 75, 77