Chapter 6
Chemical & Physical Properties of the
Elements and the Periodic Table
Review Quiz Chapter 6
• Heats of (kJ/mol) conversion.
• ∆H summation formula.
Valence Electrons
• The valence electrons are the
electrons in the outer energy level
(valence level).
Alkali
Metals
Alkaline Earth
Metals
Transition
Elements
(Metals)
Halogens
Noble
Gases
Inner Transition
Elements (Metals)
The Representative Elements
Covalent radius
• Covalent radius is essentially the size of
an atom.
Covalent Radii (atomic radii)
Atomic Radius
Ionic Radius
• Ionic Radius is the size of an ion.
Isoelectronic Series
•
•
Substances are isoelectronic if they
have the same electron configuration.
Name two isoelectronic species.
Ionization Energy
• Ionization energy is the energy needed to
remove an electron from an atom or ion.
First Ionization Energy
• First Ionization energy is the energy
needed to remove the first electron from
an atom.
Multiple Ionization Energies
•
Second Ionization energy is the
energy needed to remove the second
electron from an atom.
•
Third Ionization energy is the energy
needed to remove the third electron from
an atom.
• Etc.
Ionization Energies in kJ/mol
1
H
2
3
4
5
6
7
1312
He 2372 5250
Li
520
7297 11810
Be 899
1757 14845 21000
B
800
2426 3659
C
1086 2352 4619
6221
37820 47260
N
1402 2855 4576
7473
9442
25020 32820
53250 64340
Write the equation representing the first
ionization energy of hydrogen.
First Ionization Energy of H
• H + 1312 kJ → H+ + e-
Electron Affinity
• The energy change that occurs when an
electron is added to an atom.
Write the equation representing the electron affinity of hydrogen.
Electron Affinity of H
• H + e- → H- + 72 kJ
Effective Nuclear Charge (Zeff)
• You will find many of the notes for
effective nuclear charge on a sheet in your
notebook titled “Effective Nuclear Charge”.
• The effective nuclear charge (Zeff) of an
atom is basically how well it is able to hold
on to its most loosely held electron.
Effective Nuclear Charge (Zeff)
•
We can estimate the effective nuclear
charge of an atom by using the following:
1. The nuclear charge (Z)
2. The shielding effect
3. Electron repulsions
The Nuclear Charge (Z)
• Based on the number of protons in the
nucleus.
– Example: Carbon vs. Nitrogen
The Nuclear Charge (Z)
The greater the number of protons in the nucleus
the greater the effective nuclear charge.
Nuclear Charge and Zeff
Shielding Effect.
• The shielding effect is when electrons
between the nucleus and the outermost
electrons in an atom shield or lessen the
hold of the nucleus on the outermost
electrons.
Shielding Effect.
Shielding Effect.
•
Shielding can be checked by writing the
electron configuration.
Example of the Shielding Effect
He atom in the excited state with
one electron in the 1s and one
electron in the 2p.
1s12p1
He+ ion in the excited state
with one electron in the 2p.
2p1
Shielding Effect
Energy Levels vs. Sublevels
•
•
Energy levels have the greatest effect on
shielding.
Sublevels increase shielding but to a far
lesser extent.
Ionization Energies in kJ/mol
1
H
2
3
4
5
6
7
1312
He 2372 5250
Li
520
7297 11810
Be 899
1757 14845 21000
B
800
2426 3659
C
1086 2352 4619
6221
37820 47260
N
1402 2855 4576
7473
9442
25020 32820
53250 64340
Zeff can help us
explain the ionization
energies.
Explain the first ionization energies of
Be and B
A
Explain the first ionization energies of
Be and Mg
Effective Nuclear Charge can be
used to help explain atomic radius.
Atomic Radius
Explain the difference in atomic radii
for Li and Be. Which are 1.52 and
1.11 angstroms respectively.
Explain the difference in atomic radii
for Li and Na. Which are 1.52 and
1.86 angstroms respectively.
Effective Nuclear Charge can be
used to help explain atomic radius.
• Based on nuclear charge and shielding.
Electron Repulsions:
Paired vs. Unpaired Electrons
• A paired electron has increased electron – electron
repulsion.
• It is easier (takes less energy) to remove a paired
electron than it does to remove an unpaired
electron.
• We check the pairing of electrons in the outer
sublevel by writing an orbital filling diagram.
Nitrogen vs. Oxygen
First Ionization Energy
Nitrogen vs. Oxygen
First Ionization Energy
It is much harder to remove an electron from helium
than it is Li. This is Illustrated by their respective
ionization energies given below. Explain.
• He = 2370 kJ/mol
• Li = 520 kJ/mol
Stability
Schmability
Penetration Effect
• Electrons in a higher energy level can
often penetrate (dive) through lower
energy levels because of the attraction
that the nucleus has on them.
• Smaller sublevels can penetrate closer
to the nucleus than larger sublevels.
Explain the relative energies of the
sublevels within the fourth energy level.
• The s sublevel penetrates closer to the
nucleus followed by the p, d and the f has
the least penetration. The closer to the
nucleus the lower the energy and therefore
the relative energies of the sublevels in the
fourth energy level is:
4s < 4p < 4d < 4f.
Explain why a 4s sublevel has a
lower energy than 3d.
• A 4s sublevel penetrates closer towards the
nucleus than does a 3d so even though the
3d is part of the third energy level the 4s on
average is closer to the nucleus and is
therefore lower in energy than the 3d.
Periodicity
Periodicity
• What would you expect the properties of
element 118 to be?
Periodicity
• How do you expect the reactivity of Rb to
compare to K?
Periodicity
• How do you expect the reactivity of Rb to
compare to K? Explain your answer in
terms of effective nuclear charge.
Physical Properties of Metals
• Metals are ductile (they can be drawn into
wires) and metals are malleable (they can be
pounded into thin sheets)
• Metals have relatively high melting points and
boiling points.
• Metals conduct electricity and heat.
• Metals are usually shiny. They have luster.
• These properties can be explained by
understanding the metallic bond.
The Metallic Bond
• The metal is held together by the strong
forces of attraction between the positive
nuclei and the delocalized electrons in
what is termed a metallic bond.
• The metallic bond is sometimes described
as "an array of positive ions in a sea of
electrons".
• If you are going to use this view, beware!
Is a metal made up of atoms or ions?
Sodium
• Atoms
Strength of Metallic Bonding
• Metallic Bonds are formed when atomic
orbitals overlap to form molecular orbitals
which stretch across the entire metal.
• Bonding electrons then become
delocalized and move throughout these
molecular orbitals.
• In general the more delocalized electrons
the stronger the metallic bond.
Strength of Metallic Bonding
• Based on the number of delocalized electrons.
Example:
• Explain the variation in boiling points for
hydrogen (-257ºC); sodium (883ºC);
magnesium (1107ºC); and iron (2750ºC); in
terms of metallic bonding.
Hydrogen boiling point (-257ºC)
• Need to overcome van der Waals forces
which are much weaker than metallic
bonds.
Boiling point sodium (883ºC)
• The 3s orbitals overlap on all the atoms to
form metallic bonds.
• [Ne]3s1
• Na+
Boiling point magnesium (1107ºC)
• The 3s orbitals overlap on all the atoms to
form metallic bonds.
• [Ne]3s2
• Mg2+
• There are twice as many delocalized
electrons as in sodium creating a stronger
metallic bond and a more positive ionic
charge.
Boiling points iron (2750ºC)
• The 3s orbitals overlap on all the atoms to
form metallic bonds.
• [Ne]4s23d6
• In transition metals the d electrons can
become involved in metallic bonding
allowing for a larger number of delocalized
electrons and a stronger metallic bond.
World of Chemistry “Metals”
Periodic Variations
• Metallic behavior increases and nonmetallic
behavior decreases as we move down a
group in the periodic table.
Metals, Nonmetals, Metalloids
Periodic Variations
• Metallic behavior decreases and nonmetallic
behavior increases as we move from left to
right across a period in the periodic table.
Chemical Properties of Metals
Metals “reduce” elemental nonmetals
Na + Cl2 → NaCl
Nonmetals “oxidize” elemental
metals and other nonmetals of
lower electronegativity
Na + Cl2 → NaCl
C + O2 → CO2
Oxidation – Reduction (Redox)
• Reduction is the decrease in oxidation number
caused by the gain of electrons.
• Oxidation is the increase in oxidation number
caused by the loss of electrons.
• A reducing agent is the substance in a redox
reaction that gives up electrons and is oxidized.
• A oxidizing agent is the substance in a redox
reaction that gains electrons and is reduced.
LEO the lion says GER
Lose Electrons Oxidation
Gain Electrons Reduction
Oxidation – Reduction (Redox)
0
0
+
_
Na + Cl2 → NaCl
Cl2 is reduced and is therefore an oxidizing agent
Na is oxidized and is therefore an reducing agent
Chemical Properties of Metals
Review the rest of table 6.7 and the supporting
reactions on pp. 169-171 and have questions ready
for me tomorrow.
Standard Reduction Potentials Table
(in notebook)
Use the five substances listed below to answer
questions 1 – 3.
Cl2
Al3+
Hg
Cr2+
I-
1. Which of the substances are oxidizing
agents?
2. Which substance is the strongest reducing
agent?
3. Which substance(s) could oxidize Ag?
4. A piece of gold (Au) is dropped into a
solution of Pb(NO3)2. Does a reaction
occur? Support your answer.
5. A piece of cobalt (Co) is dropped into a
solution of Pb(NO3)2. Does a reaction
occur? Support your answer.
Cl2
Al3+
Hg
Cr2+
I‾
Which of the substances are oxidizing
agents?
Cl2
Al3+
Hg
Cr2+
Which substance is the strongest
reducing agents?
I‾
Cl2
Al3+
Hg
Cr2+
I‾
Which substance(s) could oxidize Ag?
A piece of gold (Au) is dropped in
a solution of Pb(NO3)2. Does the
reaction occur? Support your
answer.
A piece of cobalt (Co) is dropped
in a solution of Pb(NO3)2. Does
the reaction occur? Support your
answer.
Homework
• There is another practice quiz in your
notebook with a corresponding redox table
on the back. Do this quiz as homework.
• Should be able to do all homework both in
the book and on the worksheets.
• Activity Series pre-lab and lab write up for
tomorrow.