Transcript Chapter 8 Periodic Properties of the Elements

Chemistry: A Molecular Approach, 1st Ed.
Nivaldo Tro
Roy Kennedy
Massachusetts Bay Community College
Wellesley Hills, MA
2007, Prentice Hall
Mendeleev
 order elements by atomic mass
 saw a repeating pattern of properties
 Periodic Law – When the elements are arranged in
order of increasing atomic mass, certain sets of
properties recur periodically
 put elements with similar properties in the same column
 used pattern to predict properties of undiscovered
elements
 where atomic mass order did not fit other properties, he
re-ordered by other properties
 Te & I
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Periodic
Pattern
nm H O
2
H
1
a/b
H2
m Li2O m/nm BeOnm B2O3 nm CO2 nm N2O5nm
O2 nm
Li b
Be a/b
B a
C
O
F
a N a
7 LiH 9 BeH2 11 ( BH3)n 12 CH4 14 NH3 16 H2O 19 HF
m Na2O m MgO m Al2O3 nm/m SiO2nm P4O10nm SO3 nm Cl2O7
Na b
Mg b
Al a/b
Si a
P
S a
Cl
a
a
23 NaH24 MgH2 27 (AlH3) 28 SiH4 31 PH3 32 H2S 35.5 HCl
m = metal, nm = nonmetal, m/nm = metalloid
a = acidic oxide, b = basic oxide, a/b = amphoteric oxide
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Mendeleev's Predictions
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What vs. Why
 Mendeleev’s Periodic Law allows us to predict what the
properties of an element will be based on its position on the
table
 it doesn’t explain why the pattern exists
 Quantum Mechanics is a theory that explains why the
periodic trends in the properties exist
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Electron
Spin
 experiments by Stern and Gerlach showed a beam of silver
atoms is split in two by a magnetic field
 the experiment reveals that the electrons spin on their axis
 as they spin, they generate a magnetic field
 spinning charged particles generate a magnetic field
 if there is an even number of electrons, about half the atoms
will have a net magnetic field pointing “North” and the other
half will have a net magnetic field pointing “South”
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Electron Spin Experiment
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Spin Quantum Number, ms
 spin quantum number describes how the electron spins on its
axis
 clockwise or counterclockwise
 spin up or spin down
 spins must cancel in an orbital
 paired
 ms can have values of ±½
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Pauli Exclusion Principle
 no two electrons in an atom may have the same set of 4
quantum numbers
 therefore no orbital may have more than 2 electrons, and
they must have with opposite spins
 knowing the number orbitals in a sublevel allows us to
determine the maximum number of electrons in the
sublevel




s sublevel has 1 orbital, therefore it can hold 2 electrons
p sublevel has 3 orbitals, therefore it can hold 6 electrons
d sublevel has 5 orbitals, therefore it can hold 10 electrons
f sublevel has 7 orbitals, therefore it can hold 14 electrons
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Allowed Quantum Numbers
Quantum
Number
Principal, n
Number of Significance
Values
1, 2, 3, ...
distance from
nucleus
Azimuthal, l 0, 1, 2, ..., n-1
n
shape of
orbital
Magnetic, ml -l,...,0,...+l
2l + 1
orientation of
orbital
Spin, ms
-½, +½
2
direction of
electron spin
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Values
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Quantum Numbers of
Helium’s Electrons
 helium has two electrons
 both electrons are in the first energy level
 both electrons are in the s orbital of the first energy level
 since they are in the same orbital, they must have opposite spins
first
electron
second
electron
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n
l
ml
ms
1
0
0
+½
1
0
0
-½
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Electron Configurations
 the ground state of the electron is the lowest energy
orbital it can occupy
 the distribution of electrons into the various orbitals
in an atom in its ground state is called its electron
configuration
 the number designates the principal energy level
 the letter designates the sublevel and type of orbital
 the superscript designates the number of electrons in
that sublevel
 He = 1s2
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Orbital Diagrams
 we often represent an orbital as a square and the
electrons in that orbital as arrows
 the direction of the arrow represents the spin of the
electron
unoccupied
orbital
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orbital with
1 electron
orbital with
2 electrons
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Sublevel Splitting in
Multielectron Atoms
 the sublevels in each principal energy level of Hydrogen
all have the same energy – we call orbitals with the same
energy degenerate
 or other single electron systems
 for multielectron atoms, the energies of the sublevels
are split
 caused by electron-electron repulsion
 the lower the value of the l quantum number, the less
energy the sublevel has
 s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3)
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Penetrating and Shielding
 the radial distribution function shows that
the 2s orbital penetrates more deeply into
the 1s orbital than does the 2p
 the weaker penetration of the 2p sublevel
means that electrons in the 2p sublevel
experience more repulsive force, they are
more shielded from the attractive force of
the nucleus
 the deeper penetration of the 2s electrons
means electrons in the 2s sublevel
experience a greater attractive force to the
nucleus and are not shielded as effectively
 the result is that the electrons in the 2s
sublevel are lower in energy than the
electrons in the 2p
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Penetration & Shielding
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7s
6s
Energy
5s
4s
6d
6p
5p
4d
3d
3p
2p
1s
4f
4p
3s
2s
5d
5f
Notice the following:
1. because of penetration, sublevels within an energy level
are not degenerate
2. penetration of the 4th and higher energy levels is so strong
that their s sublevel is lower in energy than the d sublevel
of the previous energy level
3. the energy difference between levels becomes smaller for
higher energy levels
Order of Subshell Filling
in Ground State Electron Configurations
start by drawing a diagram
putting each energy shell on
a row and listing the subshells,
(s, p, d, f), for that shell in
order of energy, (left-to-right)
next, draw arrows through
the diagonals, looping back
to the next diagonal
each time
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1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
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Filling the Orbitals with
Electrons
 energy shells fill from lowest energy to high
 subshells fill from lowest energy to high
 s→p→d→f
 Aufbau Principle
 orbitals that are in the same subshell have the same
energy
 no more than 2 electrons per orbital
 Pauli Exclusion Principle
 when filling orbitals that have the same energy, place one
electron in each before completing pairs
 Hund’s Rule
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Example 8.1 – Write the Ground State
Electron Configuration and Orbital
Diagram and of Magnesium.
Determine the atomic number of the element from the
Periodic Table
1.

This gives the number of protons and electrons in the
atom
Mg Z = 12, so Mg has 12 protons and 12 electrons
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Example
 Write the Ground State Electron Configuration and
Orbital Diagram and of the following element
 Lithium
 Magnesium
 Sulfur
Valence Electrons
 the electrons in all the subshells with the highest
principal energy shell are called the valence electrons
 electrons in lower energy shells are called core electrons
 chemists have observed that one of the most important
factors in the way an atom behaves, both chemically and
physically, is the number of valence electrons
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Electron Configuration of Atoms in
their Ground State
 Kr = 36 electrons
1s22s22p63s23p64s23d104p6
 there are 28 core electrons and 8 valence electrons
 Rb = 37 electrons
1s22s22p63s23p64s23d104p65s1
[Kr]5s1
 for the 5s1 electron in Rb the set of quantum numbers is n
= 5, l = 0, ml = 0, ms = +½
 for an electron in the 2p sublevel, the set of quantum
numbers is n = 2, l = 1, ml = -1 or (0,+1), and ms = - ½ or
(+½)
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Examples
 Write an electron configuration for phosphorous and
sodium then identify the valence electrons and core
electrons
Electron Configurations
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Electron Configuration & the
Periodic Table
 the Group number corresponds to the number of valence
electrons
 the length of each “block” is the maximum number of
electrons the sublevel can hold
 the Period number corresponds to the principal energy level
of the valence electrons
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s1
1
2
3
4
5
6
7
s2
p1 p2 p3 p4 p5 s2
p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 f14d1
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Transition Elements
 for the d block metals, the principal energy level is one less than
valence shell
 one less than the Period number
 sometimes s electron “promoted” to d sublevel
Zn
Z = 30, Period 4, Group 2B
[Ar]4s23d10
4s
3d
• for the f block metals, the principal energy level is two less
than valence shell
 two less than the Period number they really belong to
 sometimes d electron in configuration
Eu
Z = 63, Period 6
[Xe]6s24f 7
6s
4f
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Practice – Use the Periodic Table to write the short
electron configuration and orbital diagram for each
of the following
 Na (at. no. 11)
 Te (at. no. 52)
 Tc (at. no. 43)
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Properties & Electron
Configuration
 elements in the same column
have similar chemical and
physical properties because
they have the same number of
valence electrons in the same
kinds of orbitals
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Electron Configuration &
Element Properties
 the number of valence electrons largely determines the behavior of
an element
 chemical and some physical
 since the number of valence electrons follows a Periodic pattern,
the properties of the elements should also be periodic
 quantum mechanical calculations show that 8 valence electrons
should result in a very unreactive atom, an atom that is very stable
– and the noble gases, that have 8 valence electrons are all very
stable and unreactive
 conversely, elements that have either one more or one less electron
should be very reactive – and the halogens are the most reactive
nonmetals and alkali metals the most reactive metals
 as a group
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Electron Configuration &
Ion Charge
 we have seen that many metals and nonmetals form one ion,
and that the charge on that ion is predictable based on its
position on the Periodic Table
 Group 1A = +1, Group 2A = +2, Group 7A = -1, Group 6A
= -2, etc.
 these atoms form ions that will result in an electron
configuration that is the same as the nearest noble gas
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Anions in their Ground
State
 anions are formed when atoms gain enough electrons to have
8 valence electrons
 filling the s and p sublevels of the valence shell
 the sulfur atom has 6 valence electrons
S atom = 1s22s22p63s23p4
 in order to have 8 valence electrons, it must gain 2 more
S2- anion = 1s22s22p63s23p6
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Cations in their Ground
State
 cations are formed when an atom loses all its valence electrons
 resulting in a new lower energy level valence shell
 however the process is always endothermic
 the magnesium atom has 2 valence electrons
Mg atom = 1s22s22p63s2
 when it forms a cation, it loses its valence electrons
Mg2+ cation = 1s22s22p6
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Examples
 Write the electron configuration and orbital diagram for
each of the following ions and predict whether the ion will
be paramagnetic or diamagnetic
 Al3+
 N3 Fe2+