Download Chapter 8: Periodic Properties of the Elements

Chemistry 1A: Chapter 8
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Chapter 8: Periodic Properties of the Elements
Homework: Read Chapter 8. Work out sample/practice exercises
Suggested Chapter 8 Problems: 43, 45, 51, 55, 59, 63, 67, 71, 75, 79, 83, 99
Check for the MasteringChemistry.com assignment and complete before due date
The Periodic Table:
1869 Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) classified known
elements by organizing similar physical and chemical properties. Mendeleev’s periodic
table was an attempt to organize the known data at the time in a way that made sense.
Elements were arranged by increasing atomic mass and grouped together by chemical
reactivity. Where problems with mass positions occurred ( Te and I), he re-ordered by
other properties. Several holes led to predictions of elements and their properties that
were not yet discovered “eka-aluminum” (Ga) and “eka-silicon”(Ge).
1913 *Henry Moseley improved the periodic table by ordering the elements by
increasing atomic number. More “holes” were found, which led to the discovery of more
elements and the family of noble gases.
The periodic table gives a great amount of information in an organized manner. Vertical
columns are called groups or families. If you are aware of the properties of a couple
elements in a group, you can make a good guess at the properties of the other elements in
the same group. Periods are the horizontal rows in the periodic table. Many patterns can
be seen or predicted following periods and groups.
Electron Configurations and Orbital Diagrams:
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A. Energy Levels:
i. In the Bohr Atom (one electron systems) there is a one-to-one correspondence
between an orbit and its energy level (En = -B/n2). Example: the 3s, 3p, 3d
orbitals in Hydrogen all have the same energy.
ii. In the Quantum Mechanical version of the atom, the energy level of multielectron atoms depend on both the size (1, 2, 3, 4…) and shape (s, p, d, f).
B. Ground State: Filling orbitals of an atom where electrons go into the lowest
energy orbitals first.
i. Pauli Exclusion Principle, no two electrons in an atom may have all four
quantum numbers alike.
ii. Hund's Rule, when filling degenerate energy levels, each orbital fills one
electron, spin unpaired, before any orbital fills with two electrons.
C. Aufbau Principle: Electronic Configuration by energy
i. 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s2 5f146d107p6
ii. Isoelectronic  Same Electronic Configuration
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Watch for exceptions:
Electron Configurations (complete/short)
Core electrons
Valence electrons
Pseudocore electrons
Ground state
Excited state
By size
By energy
Pauli’s rule
Hund’s Rule
Paramagnetic
Diamagnetic
Quantum numbers (m, l ml, ms)
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Quantum Numbers:
A. Each electron in an atom is defined by four Quantum Numbers.
The principle quantum number, n, is related to the size (90% probability that the
electron is within a given radius)
Name
Symbol
Values
Significance
principle
n
1, 2, 3, ...
size
azimuthal

0, 1, ...(n-1)
shape
magnetic
ml
0, 1, 2,.... 
orientation
spin
ms
½
electron spin
B. The values of the azimuthal quantum number, , determines the shape of the
orbital. In the designation of an orbital, this quantum number is represented by a
letter.
Value
Letter
Shape
0
s
spherical
1
p
dumbbell
2
d
four-lobed
3
f
eight-lobed
C. The magnetic quantum number, ml, identifies the three dimensional orientation in
space. For an s orbital it is 0, for p it can be -1, 0, +1, d has 5 orientations, f has 7
Three quantum numbers are required to specify an orbital: principle, azimuthal and
magnetic.
An orbital is a place in an atom to hold electrons. An orbital may contain a
maximum of two electrons
D. The fourth quantum number is the electron spin, ms. Values are ½. Stern-Gerlach
experiment split a beam of silver atoms (47 electrons-odd number-paramagnetic) in
two by a magnetic field. As electrons spin they generate a magnetic field. About
have half the electrons point “north” (spin up) and others point “south” (spin down)
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Example 1:
Predict the ground state short electron configuration for each. Identify electrons
core, [valence], pseudocore. Draw the orbital diagram for the outershell electrons,
identify if species is diamagnetic or paramagnetic, write out the quantum numbers
for each of the outershell electrons.
Ag
Ag+1
Ga
Ga+3
Fe
Fe+2
Fe+3
Sn
Sn+2
Sn+4
P
P-3
P+3
P+5
S
S-2
Ce
Ce +2
Ce+3
Ce+4
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Coulombs Law: F = kQ1Q2/d2 Since Energy is force times distance, E = kQ1Q2/d
Coulombs Law describes attractions and repulsions between charged particles.
Attraction is stronger as atomic sizes decrease and charge differences increase.
Effective Nuclear Charge:
Negatively charged electrons are attracted to the positively charged nucleus and
repelled by other electrons in the atom.
The force of attraction depends on the magnitude of the net nuclear charge
acting on an electron and the average distance between the nucleus and the
electron.
We can estimate the net attraction and average environment of a single electron
through Zeff (effective nuclear charge), which is always smaller than the total
charge of the nucleus. Zeff = Ztotal - Sscreening constant
S is generally close to the # of core electrons, (i.e. for Na, 10 core electrons, 1
valence electron. The Zeff = 11-10 = +1 for the valence electron 3s1)
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Penetration:
• The closer an electron is to the
nucleus, the more attractions it
experiences.
• The degree of penetration is related
to the orbitals radial distribution
function
• The radial distribution function
shows that the 2s orbital penetrates
more than the 2p
• The weaker penetration of the 2p
sublevel means that electrons in
the 2p sublevel 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
• Penetration causes the energies of
sublevels in the same principal
level to not be degenerate (2s and 2p are different energies)
• In the 4th and 5th principle levels, the effects of penetration cause the s orbital to be
lower in energy than d orbitals of the previous principal level (4s is lower than 3d)
• The energy separations between one set of orbitals and the next become smaller
beyond the 4s so the ordering can vary among elements causing variations
(exceptions) in the electron configurations of the transition metals and their ions
Periodic Trends:
Properties of the elements follow a periodic pattern: Same column have similar
properties and in a period a pattern repeats. This is explained with the
quantum-mechanical model because the number of valence electrons and
types of orbitals they occupy are periodic.
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Size (atomic and ionic radii):
Atomic radii increase from right to left; top to bottom of periodic table.
Zeff is the effective nuclear charge. Zeff = Zactual – electron shielding
This is the charge felt by the outer electrons that are shielded from the full power
of the positive nuclear charge
Along a period the Zeff increases left to right pulling in electrons closer to the nucleus
and causing the atoms to decrease in size.
Transition metals in the same d block are roughly the same size
Vertically, the size of the orbitals (quantum number n) increases from top to bottom
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Ionic Radii:
Cations lose electrons and are therefore smaller than the original atom
Anions gain electrons and are larger than the original atom
Isoelectronic series (all have the same number of electrons, same electron
configuration) the size increases as the charge of the nuclei decreases.
(smallest Sr+2, Rb+1, Kr, Br-1, Se-2 largest)
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First Ionization Energy, Ei:
 Energy required to remove the outermost ground state electron, endothermic
 Ionization Energy decreases from right to left; top to bottom of periodic table.
The small nonmetals require the highest ionization energy, they do not want to
lose electrons
Large metals have lowest ionization energy, they want to lose electrons and
become positively charged cations.
Minor irregularities occur
a. Ei of Be is larger than B and that of Mg is larger than Al. An
explanation is that Be and Mg lose an s2 electron while B and Al are
losing the p1 electron. An s electron spends more time closer to the
nucleus and is therefore harder to remove. Additionally, the p electrons
are shielded somewhat by the s electrons and feel a smaller Zeff.
b. Ei of N is larger than O. The explanation lies in the difference between
losing an electron from a half filled orbital verses a filled orbital.
Oxygen’s last filled electron is 2p4 . Electrons repel each other and
when electrons are forced to share space in a filled orbital they are
slightly higher in energy, so it is slightly easier to remove one giving O
a smaller ionization energy compared to N.
Chemistry 1A: Chapter 8
Higher Ionization Energies:
 Energy required removing the second, third
of even more electrons from an atom
 Larger amounts of energy are required to
remove each successive electron. It is
relatively easier to remove from partially
filled valence shell and harder to remove
from filled d shells or core electrons.
Electron Affinity, Eea:
 Energy given away when adding an
electron, exothermic:
 The greatest negative value (most
preferred) electron affinity is for F.
Small nonmetals.
 Ignore noble gases.
 -Eea generally decreases from right to
left; top to bottom of periodic table.
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Chemistry 1A: Chapter 8
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Electronegativity: (Chapter 9: page 394-395)
 The ability to attract electrons toward the atom
 Electronegativity is important in the covalent bonding or ionic transfer of
electrons in molecules and compounds. (Ch 9: Lewis Structure, Valence Shell
Electron Pair Repulsion, determining polarity of a molecule)
 Increases from left to right; bottom to top of periodic table. Small nonmetals are
much better at attracting electrons.
 Ignore the noble gases as most do not attract electrons (except xenon which may
make a few compounds such as XeOF4, XeCl2 or XeF4)
http://www.green-planet-solar-energy.com/electronegativity-values.html
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Metallic Character:
Increases from right to left; top to bottom of periodic table
Octet Rule:
Main group elements tend to undergo reactions that leave them with 8 outer shell
electrons, obtaining the noble gas configuration.
Chemistry by Group:
Alkali Metals:
Metallic, soft enough to cut with a knife, silver color, low melting points,
malleable, conductive, reactive and must be stored under oil to prevent reaction
with air and moisture.
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Reaction of alkali metal (M) with a halogen (X)  MX
Reactions:
2 M + X2 
2 M + H2 
6 M + N2 
4 M + O2 
2 M + 2 H2O 
Alkaline Earth Metals:
Metallic, silver color, malleable, conductive, can lose 2 electrons easily
causing them to be powerful reducing agents.
Reactions:
M + X2 
(X is a halogen)
M + H2 
2 M + O2 
M + 2 H2O 
Halogens:
Nonmetals, diatomic molecules, high electron affinities (tendency to gain
electrons), powerful oxidizing agents
Reactions: (X is a halogen)
2 M + n X2  2MXn
H2 + X2 
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Noble Gases:
Unreactive nonmetals, low melting and boiling points, colorless, odorless,
filled core electron configuration.
He and Ne undergo no known reactions. Argon is known to form HArF. Kr
and Xe may react with fluorine (XeF2, XeF4, XeF6, XeOF4)
Problems:
1. Use the concepts of effective nuclear charge, shielding, and n value of the valence orbital to
explain the trends in atomic radius as you (a) move across the periodic table, (b) move down the
periodic table.
2. Is the order of electron removal upon ionization simply the reverse of electron addition upon
filling? Why or why not. Complete the short electron configurations for thalium…Tl, Tl+1, Tl+3
3.
Arrange the elements as follows:
C, Mg, He, Sr, O, Fr
a) Increasing metallic character
b) Increasing atomic size
c) Increasing Ionization energy
d) Increasing electronegativity
4. Identify 2 cations and 2 anions that are isoelectronic with Xe. Place them in order of increasing
atomic radii.
5. Estimate the normal melting point of Br from the given melting points of atoms in the same
group… F = -219°C, Cl = -101°C, I = 114°C
6. Estimate the density of Kr from the given density of atoms in the same group at STP…
Ne = 0.90 g/L, Ar= 1.78 g/L, Xe = 5.86 g/L
7. Use Coulomb’s Law to arrange the ionic compounds by expected increasing normal melting
points. BaO, CsI, NaCl, Fe2O3