CHEMISTRY
The Central Science
9th Edition
Chapter 7
Periodic Properties of the
Elements
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Chapter 7
7.1: Development of the
Periodic Table
• Dimitri Mendeleev and Lothar Meyer arranged the elements
in order of increasing atomic weight
• Certain elements were missing from this scheme:
• Mendeleev noted:
• As properly belonged underneath P and not Si
• A missing element underneath Si
• He predicted a number of properties for this element
• In 1886 Ge was discovered
• Properties of Ge match Mendeleev’s predictions
• Modern periodic table: arrange elements in order of
increasing
atomic number Chapter 7
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Effective Nuclear Charge
• Effective nuclear charge is the charge experienced by an
electron on a many-electron atom
• The effective nuclear charge is not the same as the
charge on the nucleus
• Due to the effect of the inner electrons
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Chapter 7
7.2: Effective Nuclear
Charge
• Electrons are attracted to the nucleus, but repelled by the
electrons that screen them from the nuclear charge
• The nuclear charge experienced by an electron depends
on
• its distance from the nucleus
• the number of core electrons
• As the average number of screening electrons (S)
increases, the effective nuclear charge (Zeff) decreases
• OR…as the distance from the nucleus increases, S
increases and Zeff decreases
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Chapter 7
• The ns orbitals all have the same shape, but have different
sizes and different numbers of nodes (places where the
probability function = 0…see text, P. 215)
Text, P. 240
Zeff = Z – S
or
Protons – Electrons
The radial electron density is the
probability
of finding an
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electron at a given distance
Chapter 7
• Zeff increases across a period
– Added electrons don’t increase shielding
• Zeff increases down a group
– Added electrons are in increasingly higher energy levels
but they are less able to shield the outer electrons from
the nucleus
The period trend is of greater significance than the group trend
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Chapter 7
7.3: Sizes of Atoms and Ions
• Consider a simple
diatomic molecule:
• The distance between the
two nuclei is called the
bond distance
• Half of the bond
distance is called the
covalent radius of the
atom
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Chapter 7
Periodic Trends in Atomic Radii
• The properties of elements vary periodically
• Atomic size varies consistently through the periodic table
• Atomic radius increases down a group (more noticeable)
• Atomic radius decreases across a period
• There are two factors at work:
• Principal quantum number
• Electrons fill higher energy levels
• Effective nuclear charge
• Electrons fill the same energy level, no increase in
shielding
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Chapter 7
Text, P. 242
Trends in the Sizes of Ions
• Ion size is the distance between ions in an ionic compound
• Cations vacate the most spatially extended orbital and are
smaller than the parent atom
• Anions add electrons to the most spatially extended orbital
and are larger than the parent atom
Remember: electrons are gained or lost to achieve a noble gas
configuration
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Chapter 7
Text
P. 244
• For ions of the same charge, ion size increases down a
group
• Electrons in higher energy levels
• All the members of an isoelectronic series have the same
number of electrons (O2-, F-, Na+, Mg2+, Al3+)
• As nuclear charge increases in an isoelectronic series, the
ions become smaller:
O2- > F- > Na+ > Mg2+ > Al3+
• Added electrons are in the same energy level, so no
increase in shielding
• Removed electrons result in an increasingly positive
nucleus which pulls electrons toward it
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Chapter 7
7.4: Ionization Energy
• The first ionization energy, I1, is the amount of energy
required to remove an electron from a gaseous atom:
Na(g)  Na+(g) + e• The second ionization energy, I2, is the energy required
to remove an electron from a gaseous ion:
Na+(g)  Na2+(g) + e• The larger ionization energy, the more difficult it is to
remove the electron
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Chapter 7
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Chapter 7
Variations in Successive Ionization
Energies
• There is a sharp increase in ionization energy when a
core electron is removed:
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Chapter 7
Text, 246
Periodic Trends in Ionization Energies
Think: does it want to lose an electron???
• Ionization energy decreases down a group
• The outermost electron is more readily removed as
electrons fill higher energy levels
• The s electrons are more effective at shielding than p
electrons
• Irregularity in group 2: sharp increase in I1
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Chapter 7
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Chapter 7
Text, P. 248
• Ionization energy generally increases across a period
• Across a period, Zeff increases
• It becomes more difficult to remove an electron
• Irregularity in group 15: sharp increase in I1 (half-filled
sublevel)
• Transition metals and f-block elements show slow
variations in I1 due to shielding
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Chapter 7
Text, P. 248
Text, P. 247
Electron Configuration of Ions
• Cations: electrons removed from orbital with highest
principle quantum number, n, first:
Li (1s2 2s1)  Li+ (1s2)
Fe ([Ar]3d6 4s2)  Fe3+ ([Ar]3d5)
• Anions: electrons added to the orbital with highest n:
F (1s2 2s2 2p5)  F- (1s2 2s2 2p6)
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Chapter 7
7.5: Electron Affinities
Think: does it want to gain an electron???
• Electron affinity is the energy change when a gaseous
atom gains an electron to form a gaseous ion:
Cl(g) + e-  Cl-(g)
• Electron affinity can either be exothermic (above) or
endothermic:
Ar(g) + e-  Ar-(g)
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Chapter 7
• EA becomes more negative across a period
• A greater attraction between the atom and the added electron
•Groups 2, 15 and 18
•EA has no real group trend
Text, P. 251
7.6: Metals, Nonmetals, and
Metalloids
Text, P. 252
Metals
• Metallic character refers to the properties of metals
• shiny or lustrous
• malleable and ductile
• oxides form basic ionic solids
• form cations in aqueous solution
• Metallic character increases down a group
• Metallic character decreases across a period
• Metals have low ionization energies
• Most neutral metals are oxidized rather than reduced
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2003
Chapter 7
•Prentice
form
cations
• Most metal oxides are basic:
Metal oxide + water  metal hydroxide
Nonmetals
• Nonmetals are more diverse in their behavior than metals
• Properties are opposite those of metals
• Compounds composed only of nonmetals are molecular
• When nonmetals react with metals, nonmetals tend to
gain electrons:
metal + nonmetal
 salt
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Chapter 7
• Most nonmetal oxides are acidic:
nonmetal oxide + water  acid
Nonmetal oxide + base  salt + water
• The greater the nonmetallic character of the central atom,
the stronger the acidic character of the oxide
Metalloids
• Metalloids have properties that are intermediate between
metals and nonmetals
• Si has a metallic luster but it is brittle
• Semiconductors
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Chapter 7
7.7: Group Trends for the
Active Metals
Group 1A: The Alkali Metals
• Alkali metals are all soft
• Chemistry dominated by the loss of their single s
electron:
M  M+ + e• Reactivity increases down the group
• Alkali metals react with water to form MOH and
hydrogen gas:
2M(s) + 2H2O(l)  2MOH(aq) + H2(g)
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Chapter 7
• Alkali metals produce different oxides when reacting with
O2:
4Li(s) + O2(g)  2Li2O(s)
(oxide)
2Na(s) + O2(g)  Na2O2(s)
(peroxide)
K(s) + O2(g)  KO2(s)
(superoxide)
(Rb and Cs, too)
• Alkali metals: Flame tests
• The s electron is excited by the flame and emits energy
when it returns to the ground state
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Chapter 7
Li line: 2p  2s
transition
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Na line (589 nm):
3p  3s transition
Chapter 7
K line: 4p  4s
transition
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Chapter 7
Text, P. 257
Group 2A: The Alkaline Earth Metals
• Alkaline earth metals are harder and more dense than the
alkali metals
• The chemistry is dominated by the loss of two s electrons:
M  M2+ + 2eMg(s) + Cl2(g)  MgCl2(s)
• Increasing activity down the group:
• Be does not react with water
• Mg will only react with steam
• Ca, Sr and Ba:
M(s) + 2H2O(l)  M(OH)2(aq) + H2(g)
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Chapter 7
Group 2A: The Alkaline Earth Metals
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Chapter 7
Text, P. 260
7.8: Group Trends for
Selected Nonmetals
Hydrogen
• Colorless diatomic gas, H2
• High IE because there are no inner electrons
• It can either gain another electron to form the hydride
ion, H-, or lose its electron to become H+:
2Na(s) + H2(g)  2NaH(s)
2H2(g) + O2(g)  2H2O(g)
• H+ is a proton
• The aqueous chemistry of hydrogen is dominated by
+(aq)
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Chapter 7
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Group 16: The Oxygen Group
• Metallic character increases down a group (O2 is a gas,
Te is a metalloid, Po is a metal)
• Two important forms of oxygen: O2 and ozone, O3
• Ozone can be prepared from oxygen:
3O2(g)  2O3(g) H = +284.6 kJ.
• Ozone is pungent and toxic
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Chapter 7
Group 16: The Oxygen Group
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Chapter 7
Text, P. 261
• Oxygen (O2) is a potent oxidizing agent
• O2- ion has a noble gas configuration
• There are two oxidation states for oxygen: 2- and 1• Sulfur is another important member of this group
• Most common form of sulfur is yellow S8
• Sulfur tends to form S2- in compounds (sulfides)
• Allotropes: different forms of the same element in the
same state
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Chapter 7
Group 17: The Halogens
• The chemistry of the halogens is dominated by gaining an
electron to form an anion:
X2 + 2e-  2X• Fluorine is one of the most reactive substances known
• All halogens consists of diatomic molecules, X2
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Chapter 7
Group 17: The Halogens
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Chapter 7
Text, P. 262
• Chlorine is the most industrially useful halogen
• Produced by the electrolysis of brine (NaCl):
2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) + Cl2(g)
• Hydrogen compounds of the halogens are all strong acids
with the exception of HF
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Chapter 7
Group 18: The Noble Gases
• These are all nonmetals and monatomic
• Unreactive because they have completely filled s and p subshells
• In 1962 the first compound of the noble gases was prepared:
XeF2, XeF4, and XeF6
• Xe has a low enough IE to react with substances that
readily remove electrons
• To date the only other noble gas compounds known are
KrF2 and HArF
• Prentice
“Argon
hydrofluoride” Chapter 7
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Group 18: The Noble Gases
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Chapter 7
Text, P. 263