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Transcript
Periodic Properties of the Elements
Key Concepts:
1.
Understand and be able to predict and explain trends in effective nuclear charge, Zeff.
2.
Understand and be able to predict and explain the periodic trends in:
2.1.
2.2.
2.3.
2.4.
2.5.
atom size
ion size
ionization energy
electron affinity
Properties of elements: Metals, nonmetals, groups (descriptive chemistry)
Periodic Table: first proposed in 1869 separately
by Dmitri Mendeleev in Russia and Lothar Meyer
in Germany. The Periodic Table proposed by
Mendeleev and Meyer was arranged in order of
increasing atomic weight. Some elements seemed
“out of order” though. The modern period table is
arranged by rows and columns in order of
increasing ATOMIC NUMBER. The properties of
the elements tend to repeat, are periodic, from row
to row.
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Effective Nuclear Charge, Zeff
The splitting of the principle energy level into the s, p, d, and f energy sublevels is best
explained by using the concept of “effective” nuclear charge, Zeff.
An electron in a higher energy level is “screened” from seeing 100% (all the protons) of the
nuclear charge by the electrons in lower energy levels. We usually talk about the valence
electrons and how they are screened from experiencing the complete nuclear charge. This
screening depends on the sublevel (orbital type) occupied by the electron being screened.
The Effective Nuclear Charge is the NET
NUCLEAR charge an electron experiences
when other electrons “screen” the nuclear
charge. An analogy is looking at a lightbulb that
is covered by a frosted-glass lamp shade. The
lampshade “screens” our eyes from the full
brightness of the lightbulb.
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Zeff - Effective Nuclear Charge
Sodium valence electron Zeff at 100% screening: Zeff = 11-10 = +1
Lower energy (inner) electrons “shield” higher
energy (outer) electrons from seeing a full
nuclear charge. This screening is not 100%.
Actual Na atom 3s valence electron
screening by core electrons:
Zeff = 11-8.49 = +2.51
Zeff = Z - S
where Z is the atomic number (number of
protons) and S is the screening constant.
S is a positive number with a value
that is dependent on the energy
subshell.
Electrons in the same valence shell screen each
other very little, but do have a slight screening
effect. For valance electrons, the the core
electrons provide most of the shielding.
Screening electron density
from the core electrons:
1s, 2s, and 2p
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Trends in Effective Nuclear Charge
Equation 7.1:
Zeff = Z - S
Notice:
The Zeff experienced by the innermost electrons,
those in the 1s subshell (red circles), closely
tracks the increase in nuclear charge, Z (black
line).
The Zeff experienced by the outermost valence
electrons (blue squares) not only is significantly
smaller than Z, it does not evolve linearly with
increasing atomic number; it varies periodically.
Slater’s Rules: A Closer Look, page 253
1. Electrons for which the principle quantum
number n is larger than the value of n for
the electron of interest contribute zero to the
value of s.
2. Electrons with the same value of n as the
electron of interest contribute 0.35 to the
value of S. (Note: The electron does not
screen itself.)
3. Electrons for which n is one less than n for
the electron of interest contribute 0.85 to the
value of S, while those with even smaller
values of n contribute 1.00.
Graph showing the variations in effective nuclear
charge for period 2 and period 3 elements.
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Other Details: “Splitting” of Subshell Energies with the Same n Value
Remember that for many electron atoms, the energies of orbitals with the same n value
increase in the order ns < np < nd < nf.
This can be explained by the following:
• In general, for a given n value:
s electrons penetrate closer to the nucleus than p
p electrons penetrate closer to the nucleus than d
d electrons penetrate closer to the nucleus than f
•
Thus, for a given n value, the attraction
between the the electron and the
nucleus decreases in the order:
ns > np > nd > nf
The result is that the ns orbitals are
lower in energy then the (n-1)d orbitals.
This is why we fill the 4s before the 3d,
5s before 4d, etc.
Graphs showing the 2s and 2p
radial probability functions.
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Atomic Radius Trends (Outer Valence e–)
Atomic radii decrease along a row.
Why? Zeff increases as we add electrons to the same energy level. The increase in nuclear
charge as we move across a row is not completely screened by the additional valence
electrons so Zeff becomes larger for each valence electron.
(Atomic radii of transition metals decrease only slightly across a period.)
Atomic radii increase down a column.
Why? As we move down a column n increases for the valence electrons, hence the orbital
size also increases. Zeff also increases SLIGHTLY, but the valence electrons spend more
time further from the nucleus in the larger orbitals, 2s compared to 1s, etc.
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Atomic Radius - Predictions
When two atoms bond covalently, the bonding atomic
radius of the two atoms can be used to predict the covalent
bond length (the distance between the two nuclei). Bonding
atomic radii are shorter than nonbonding atomic radii due to
the attractive forces that lead to the bond.
Use the figure of bonding
atomic radii to predict:
1.
the largest diatomic
covalent molecule
bond distance.
2.
if a N-S bond is longer
or shorter than a P-O
bond.
Bonding Atomic Radii in Angstroms
1 Å = 10–10 m
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Electron Configurations of Ions
Main Group Elements: electrons are lost or gained so that the electron configuration of the
ion matches that of the nearest Noble Gas.
Metals lose electrons to become cations. Nonmetals gain electrons to become anions. We
can use spdf notation to show this.
1. Al —> Al3+ + 3e–
[Ne]3s23p1 —> [Ne] + 3e– (loses the 3s and 3p electrons)
2. Ca —> Ca2+ + 2e–
[Ar]4s2 —> [Ar] + 2e– (loses the 4s electrons)
3. O + 2e– —> O2[He]2s22p4 + 2e– —> [Ne] (gains two e– to fill the 2p shell)
In general, what type of orbitals are filled when nonmetals gain electrons?
Transition metals (d-block) lose the (n+1)s electrons first!
1. Fe —> Fe2+ + 2e–
[Ar]4s23d6 —> [Ar]3d6 + 2e– (loses the 4s electrons)
2. Fe —> Fe3+ + 3e–
[Ar]4s23d6 —> [Ar]3d5 + 3e– (loses the 4s & a 3d electron)
Write electron configurations for Li+, Zn2+, Mn4+, P3–, Sn2+ and Sn4+.
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Ionic Size Trends
Ions show a trend in ionic size as well.
Cations are smaller than the atoms they come from because they have lost outer electrons. Also,
electron-electron repulsions are reduced.
Anions are larger than the atoms they come from because of increased electron-electron repulsions.
Also, Zeff decreases for added valence electrons.
about 2x size
about 1/2 size
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Period Trends - Comparison of Atomic and Ionic Radii
(Units are angstroms: 1 Å = 10–10 m)
Grey: Neutral radius
Pink: Cation radius
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Grey: Neutral radius
Blue: Anion radius
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Isoelectronic Series - Same Valence Shell Electron Configuration
Isoelectronic
with [He]
Isoelectronic
with [Ne]
Isoelectronic
with [Ar]
Isoelectronic
with [Kr]
Isoelectronic
with [Xe]
For isoelectronic series, what is the trend in size?
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Period Trends - Ionization Energy, IE
Ionization Energy, IE, is the energy needed to remove an outer electron from an atom in the
gas phase to make a positive ion.
Each atom can have a series of ionizations to produce a multi-charged cation.
For example consider the ionization of Mg(g):
1.
2.
3.
First:
Mg(g) —> Mg+(g) + e– IE1 = +738 kJ/mol
Second: Mg+(g)
IE2 = +1451 kJ/mol
Third: Mg2+(g)
—>
Mg2+(g)
—>
+
Mg3+(g)
e–
+
e–
1.
Why the increase from IE1 to IE2?
2.
Why the HUGE increase from IE2 to IE3?
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IE3 = +7733 kJ/mol
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Periodic Trends - First Ionization Energy, IE1
In general:
first ionization energy increases across a row.
Zeff increases across a row. As Zeff increases there is more
attraction of the electrons to the nucleus thus more
difficult to remove.
In general
first ionization energy decreases down a column.
The outer electrons are in higher principle quantum shells
and are further from the nucleus. Less attraction to the
nucleus thus easier to remove.
We see some exceptions however. For example, IE1 of N is
greater than IE1 of O. Why?
Half-filled p-sublevel for N is more stable than the
partially filled p-sublevel for O. In N, we have no e– - e–
repulsive pairing energy since all p-orbitals have only 1
e–. In O we have a p-orbital with two electrons, the
pairing energy in this p-orbital leads to a slightly less
stable electron configuration and thus lower ionization
energy.
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Periodic Trends - Electron Affinity, EA
Electron affinity is the energy change when an electron is added to a neutral atom in the gas
phase.
For example:
F(g) + e– —> F–(g)
Na(g) + e– —> Na–(g)
N(g) + e– —> N–(g)
EA = -328 kJ/mol
EA = -53 kJ/mol
EA > 0 kJ/mol so N–(g) is unstable
All second electron affinities are positive. For example:
O-(g) + e– —> O2–(g) EA2 = +744 kJ/mol
Does this make sense?
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Periodic Trends - Electron Affinity, EA
The trends are not as “regular” as for ionization energies.
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Summary of Atomic Trends
Increase
Increase
Zeff
The individual atomic properties of atoms can be related to the observed macroscopic behavior of the elements.
The trends we observe across the periodic table help explain chemical behavior:
• Nonmetals have high electron affinities and tend to form (-) ions.
• Metals tend to have low ionization energies and form (+) ions.
• Compounds formed by a metal and a nonmetal tend to be ionic substances.
• Compounds formed by two nonmetals tend to be molecular substances.
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Metals
Properties
1. Low ionization energies - oxidized easily
2. Metallic bonding in elemental form
Chemistry
1. Metals and nonmetals react to form ionic
compounds (salts):
Metals + nonmetals —> salts
2 Fe(s) + 3 Cl2(g) —> 2 FeCl3(s)
2. Metal Oxides are basic since they contain a basic
oxide ion:
Soluble Metal oxide + water —> metal hydroxide
Na2O(s) + H2O(l) —> 2 NaOH(aq)
The oxide ion is basic in water:
O2-(aq) + H2O(l) —> 2 OH–(aq)
3. Metal oxides react with acids:
Metal oxide + acid —> salt + water
Al2O3(s) + 6 HNO3(aq) —> 2 Al(NO3)3(aq) + 3 H2O(l)
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Nonmetals
Properties
1. Vary greatly in appearance
2. High electron affinities - tend to be reduced
3. Compounds of nonmetals are typically
molecular substances (covalent bonding)
Chemistry
1. Nonmetal Oxides are acidic in solution:
Nonmetal oxide + water —> acid
CO2(g) + H2O(l) —> H2CO3(aq)
2. Nonmetal oxides react with bases:
Nonmetal oxide + base —> salt + water
SO3(g) + 2 NaOH(aq) —> Na2SO4(aq) + H2O(l)
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Questions: Trends in Properties of the Elements
•
Compare B, Al, and C
Which has the largest atomic radii?
Which has the highest electron affinity?
Rank them in order of INCREASING first ionization energy.
Which has the most metallic Character?
•
Which experiences the greatest effective nuclear charge, a 2p electron in F−, a 2p electron
in Ne, or a 2p electron in Na+?
•
Text Question 7.36 Consider S, Cl and K and their most common ions.
(a) List the atoms in order of increasing size.
(b) List the ions in order of increasing size.
(c) Explain any differences in the orders of the atomic and ionic sizes.
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Problems From Text
•
7.55 Consider the first ionization energy of neon and the electron affinity of fluorine.
(a) Write equations, including electron configurations, for each process.
(b) These two quantities will have opposite signs. Which will be positive and which will
be negative?
(c) Would you expect the magnitudes of these two quantities to be the same. If not,
which one would you expect to be larger and why?
•
7.53 While the electron affinity of bromine is a negative quantity, it is positive for Kr. Use
the electron configurations of the two elements to account for this observation.
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Problems From Text
•
7.95
(a) Use orbital diagrams to illustrate what happens when an oxygen atom gains two
electrons.
(b) Why does O3– not exist in nature?
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Problems From Text
•
Predict whether each of the following oxides is ionic or molecular:
CO2 BaO
SO3
Fe2O3
Li2O
H2O
•
7.67
Write balanced chemical equations for the following reactions:
(a) barium oxide with water
(b) iron(III) oxide with perchloric acid
(c) sulfur trioxide gas with water
(d) carbon dioxide gas with aqueous sodium hydroxide.
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Problems From Text
•
7.106 (Time permitting)
(a) Write the electron configuration for Li, and estimate the effective nuclear charge
experienced by the valence electron.
(b) The energy of an electron in a one-electron atom or ion equals
where Z is the nuclear charge and n is the principal quantum number of the electron.
Estimate the first ionization energy of Li.
(c) Compare the result of your calculation with the value reported in Table 7.4, and
explain the difference.
(d) What value of the effective nuclear charge gives the proper value for the ionization
energy? Does this agree with your explanation in (c)?
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Problems From Text
•
7.109 (Time permitting)
Consider the gas-phase transfer of an electron from a sodium atom to a chlorine atom:
(a) Write this reaction as the sum of two reactions, one that relates to an ionization
energy and one that relates to an electron affinity.
(b) Use the result from part (a), data in this chapter, and Hess’s law to calculate the
enthalpy of the above reaction. Is the reaction exothermic or endothermic?
(c) The reaction between sodium metal and chlorine gas is highly exothermic and
produces NaCl(s), whose structure was discussed in Section 2.6. Comment on this
observation relative to the calculated enthalpy for the aforementioned gas-phase
reaction.
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Alkali Metals
Very reactive because of low ionization energy. Easily oxidized. Very good reducing agents.
Found only as compounds in nature. Low melting points and densities when pure.
All will form metal hydrides: 2 Na(s) + H2(g) —> 2 NaH(s)
All will react with water to form hydroxides and H2:
2 Na(s) + 2 H2O(l) —> 2 NaOH(aq) + H2(g)
Reactivity increases down the column, why?
(a) The reaction of lithium is evidenced by the bubbling
of escaping hydrogen gas.
(b) The reaction of sodium is more rapid and is so
exothermic that the hydrogen gas produced burns in air.
(c) Potassium reacts almost explosively.
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Alkaline Earth Metals
Harder and more dense than alkali metals. Not as reactive.
Generally don’t form metal hydrides.
All form stable oxides, their most common form in nature.
Only Ca, Sr and Ba react with water at room temperature to form hydroxides:
Ca(s) + 2 H2O(l) —> Ca(OH)2(aq) + H2(g)
Reactivity again increases down the column. Why?
Calcium metal reacts
with water to form
hydrogen gas and
aqueous calcium
hydroxide, Ca(OH)2(aq).
Strontium oxide, SrO(s)
The colors of firework displays
originate from the characteristic
emissions of elements, including
the alkaline earths.
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