Download Chapter 8 (Hill/Petrucci/McCreary/Perry Electron Configurations and

Chapter 8 (Hill/Petrucci/McCreary/Perry
Electron Configurations and Atomic Properties
This chapter deals with atoms that have more than one electron … we will look at the ways in
which electrons are arranged in the rest of the elements and introduce the concept of an orbital
energy diagram. These diagrams show the electron occupancy of energy levels for shells, subshells and orbitals in various atoms … their electron “configurations.”
“orbitals are not actual regions in an atom … they are mathematical expressions used to
determine the probabilities of finding an electron in various regions of an atom.” (Hill, p.301)
An Introduction to Electron Configurations
] Read Sections 8.1 through 8.3, pp. 299-303 (Hill)
electron configuration: a description of the way in which electrons in a particular type of atom
occupy the various orbitals in that atom
spdf notation: a notation used to designate a particular shell and subshell in an atom … the shell
number is followed by a sub-shell letter, e.g. 3p
subshell occupancy: indicated by a superscript that tells us how many electrons there are in that
sub-shell, e.g. 3d7 (7 electrons in the 3d sub-shell)
Convention for Writing Electron Configurations
Electron configuration for a helium (He) atom with 2 electrons …1s2
n = 1 (shell)
l = 0 (s subshell)
2 = 2 electrons
An Orbital Energy Diagram
Each “line” represents an orbital that can hold up to a maximum of 2 electrons …
Generalizations About Subshell Energies
In a given shell, s < p < d < f
2s < 2p; 3s < 3p < 3d; 4s < 4p <4d <4f
For atoms that are not in a magnetic field, the energies of all of the orbitals in a given subshell
are equal … the orbitals are said to be degenerate (of equal energy). … we depict this on an
orbital energy diagram by showing them all on the “same line”
An Orbital Diagram for a Nitrogen Atom …Nitrogen atoms have 7 electrons …1s2 2s2 2p3
An occupancy diagram in which a p-subshell is not first half- filled before electrons are paired is
incorrect, because electrons half- fill all orbitals in any subshell before filling the subshell.
“Filling order” for Electrons in Subshells
Note how filling order works on orbital energy diagram
1s"2s"2p"3s"3p"4s"3d"4p"5s"4d"5p"6s"4f"5d"6p"7s"5f"6d"7p
The “Aufbau Principle”
Now let’s build our own atoms, starting with a hydrogen atom, H.
H
1 proton, 1 electron
Z=1
The idea: the electron goes in the lowest energy state, water filling a glass from the bottom up
Where does the electron “go”? Into the 1s sub-shell ! Electron configuration: 1s1
.... More Elements
Next, we’ll “build” a helium atom, He.
He
2 protons, 2 electrons Z = 2
Again, we want to put the electron in the lowest energy state so, where does the electron “go”?
Again, into the 1s sub-shell ! Electron configuration: 1s2 (now this sub-shell is full)
Orbital Diagrams for Elements 6-10
C
1s2 2s2 2p2
F
2 2
3
N
1s 2s 2p
Ne
O
1s2 2s2 2p4
1s2 2s2 2p5
1s2 2s2 2p6
Results of Aufbau for First 12 Elements
Subshells in italics are filled and are called the “core” and subshells in bold are in the valence
shell and contain the valence electrons.
H 1s 1
N 1s2 2s 2 2p3
2
He 1s
O 1s2 2s 2 2p4
Li 1s2 2s 1
F 1s2 2s2 2p5
2 2
Be 1s 2s
Ne 1s2 2s 2 2p6
B 1s2 2s 2 2p1
Na 1s2 2s2 2p6 3s 1
2
C 1s 2s2 2p2
Mg 1s2 2s2 2p6 3s2
Electron Configuration and the Periodic Table
Group 1 (I-A)
alkali metals
ns1
Group 2 (II-A)
alkaline earth metals ns2
These two groups or families are called the s-block elements. Groups 13-18 (III-A through VIIIA) are p-block elements. Together these elements make up the main-group elements.
Groups 3-12 (the B- metals) are d-block elements
Lanthanides and actinides are the f-block elements
Periodic Table with Block- and Group-Designations (see figure 8.6, p. 310, Hill)
Learn block locations: s-block; p-block; d-block; f-block
Electron Configurations of Ions
Metals lose electrons from valence shells to form cations, and nonmetals gain electrons in their
valence shells to form anions ….
Ex. Na0 " Na+ + e1Na0 1s2 2s2 2p6 3s1
Na+ 1s2 2s22 p6 3s0
+
0
Na is “isoelectronic” with Ne (a very stable noble gas configuration)
Ex. S0 + 2 e1- " S2S0 1s2 2s2 2p6 3s2 3p4
S21s2 2s2 2p6 3s2 3p6
20
S is “isoelectronic” with Ar (another very stable noble gas configuration)
Charges on Metal Cations
Group IA (+1) Group IIA (+2); Transition metal ions (most have +2)
Atomic Radii
atomic radius: ½ the distance between two bonded atoms
covalent radius: ½ the distance between two identical bonded atoms, e.g. I-I (I2) shown at left
metallic radius: ½ the distance between two metal atoms in solid metal
Trends in Atomic Radii (in pm) of Elements 1-82
What group of elements has the largest radii in the series shown? Group IA!!
Why? “effective nuclear charge” increases from left to right in a period (row)
An Explanation of “Effective” Nuclear Charge
Inner shell electrons (below the valence shell) “shield” the valence electrons (-) in the valence
shell from the attraction due to the + charge on the nucleus.
As we add more and more electrons in the valence shell going across a given period, the
effective nuclear charge increases as shielding remains about the same.
Summary of Trends in Atomic Size : Size increases down and to the left
Ionic Radii: The Radii for Cations and Anions
ionic radius: that part of the inter- ionic distance (d) due to a cation or an anio n
d = r(cation) + r(anion)
Cations Have Smaller Radii than Their Respective Atoms
Ex. r(Na+) << r(Na0 ) …. 99 pm << 186 pm
Anions Have Larger Radii than Their Respective Atoms
Ex. r(Cl1-) >> r(Cl0 ) …. 181 pm >> 99 pm
See Example 8.6, p. 318, Hill
Some Representative Atomic and Ionic Radii
See Figure 8.14, p. 319, Hill
Ionization Energy: Removal of Electrons from Atoms and Ions
ionization energy: the energy (kJ/mol) required to remove an electron from a neutral atom or
from an ion in the gaseous state
Principle. It is successively more difficult to remove an electron (-) from an atom or ion as
the positive charge on the species increases (electrostatic attraction increases with
increasing + charge)
B(g) " B+ (g) + eB+(g) " B2+(g) + eI2 > I1
I.E. = 801 kJ/mol
I.E. = 2427 kJ/mol
1st I.E. (I1 )
2nd I.E. (I2 )
Trends in Ionization Energy for IA and IIA Elements (see Table 8.4, p. 320, Hill)
What two trends are obvious in this table? 1. I.E. decreases down a group and 2. I2 > I1 Why?
I.E. decreases with increasing size (down a group) … electron farther from nucleus and
I2 > I1 because electron is removed form a more positive charge
Trends in Ionization Energies
Nonmetals have very high I.E. Metals tend to have relatively low I.E. values
I.E. increases up and to the right.
Change in Ionization Energy with Increasing Atomic Numbe r
Note which elements have the highest I.E. values. Why? Noble gases have very stable electron
configurations.
Trends in “Metallic Character”
Metallic character is related to the ease of removal of electrons from an atom or an ion …
Metals tend to form cations (n+). Cs and Fr (most metallic). F (least metallic)
Electron Affinity: Addition of Electrons to Atoms and Ions
electron affinity: the energy (kJ/mol) change that occurs when and electron is added to a
gaseous atom or io n
Principle. electrons approaching nuclei are attracted by the positive charge on the nucleus and
repelled by the electrons already around the nucleus … when an electron is gained by the atom
or ion, energy is usually (but not always) given off
O(g) + e- " O1-(g)
O1-(g) + e- " O2-(g)
E.A. = -141 kJ/mol
E.A. = +744 kJ/mol
1st E.A. (EA1 )
2nd E.A. (EA2)
EA2 < EA1 (less energy given off with each new e1-)
Trends in Electron Affinities
Nonmetals have very high E.A.. Metals tend to have relatively low E.A. values
E.A. increases up and to the right
Metals and Nonmetals on the Periodic Table
Most of the known elements are metals. Metalloids include: B, Si, Ge, As, Sb and Te
Trends in Atomic and Elemental Properties (see Figure 8.17, p. 325, Hill)
The Noble Gases Comprise Group VIIIA (or Group 18)
Include: He, Ne, Ar, Kr, Xe and Rn
Flame Tests for Selected Alkali Metals and Alkaline Earths
See Figure 18.9, p. 327, Hill