Download effective nuclear charge

Chapter 7
Periodic Properties of the Elements
▶ POTASSIUM METAL
REACTING WITH WATER.
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Periodic Properties of the Elements
What’s Ahead
7.1 DEVELOPMENT OF THE PERIODIC TABLE
We begin our discussion with a brief history of the periodic table.
7.2 EFFECTIVE NUCLEAR CHARGE
We next explore the many properties of atoms that depend on the net attraction of the
outer electrons to the nucleus and on the average distance of those electrons from the
nucleus. The net positive charge of the nucleus experienced by the outer electrons is
called the effective nuclear charge.
7.3 SIZES OF ATOMS AND IONS
We explore the relative sizes of atoms and ions, both of which follow trends that are
related to their placement in the periodic table.
7.4 IONIZATION ENERGY
We next look at the trends in ionization energy, which is the energy required to remove
one or more electrons from an atom. The periodic trends in ionization energy depend on
variations in effective nuclear charge and atomic radii.
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Periodic Properties of the Elements
What’s Ahead
7.5 ELECTRON AFFINITIES
Next we examine periodic trends in the electron affinity, the energy released when an
electron is added to an atom.
7.6 METALS, NONMETALS, AND METALLOIDS
We learn that the physical and chemical properties of metals are different from those of
nonmetals. These properties arise from the fundamental characteristics of atoms,
particularly ionization energy. Metalloids display properties that are intermediate between
those of metals and those of nonmetals.
7.7 TRENDS FOR GROUP 1A AND GROUP 2A METALS
We examine some periodic trends in the chemistry of group 1A and group 2A metals.
7.8 TRENDS FOR SELECTED NONMETALS
Finally, we examine some periodic trends in the chemistry of hydrogen and the elements
in groups 6A, 7A, and 8A.
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Periodic Properties of the Elements
Periodic Properties of the Elements
• Elements in the same column of the periodic table contain the
same number of electrons in their valence orbitals (the
occupied orbitals that hold the electrons involved in bonding).
• Electron configurations can be used to explain differences as
well as similarities in the properties of elements.
– Ex) O ([He]2s22p4) vs. S ([Ne]3s23p4).
– Their similarities come from the same number of the valence electrons.
– They exhibits different properties due to the major difference in their
outermost electrons.
• Oxygen is a colorless gas but sulfur is a yellow solid at room temperature.
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7.1 DEVELOPMENT OF THE PERIODIC TABLE
Development of the Periodic Table
• Mendeleev and Meyer independently came to the same
conclusion about how elements should be grouped.
Figure 7.1 Discovering the elements.
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7.1 DEVELOPMENT OF THE PERIODIC TABLE
Development of the Periodic Table
• The identification of “holes” in the periodic table led to
the discovery of other unknown elements.
• Mendeleev, for instance, predicted the discovery of
germanium (which he called eka-silicon) as an element
with an atomic weight between that of zinc and arsenic,
but with chemical properties similar to those of silicon.
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7.1 DEVELOPMENT OF THE PERIODIC TABLE
Atomic Number
• Mendeleev’s table was based on atomic
masses. It was the most fundamental
property of elements known at the time.
• About 35 years later, the nuclear of the atom
was discovered by Ernest Rutherford.
• Henry Moseley developed the concept of
atomic number experimentally. The number
of protons was considered the basis for the
periodic property of elements.
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7.1 DEVELOPMENT OF THE PERIODIC TABLE
Periodicity
• Periodicity is the repetitive pattern of a property for
elements based on atomic number.
• The following properties are discussed in this chapter:
–
–
–
–
Sizes of atoms and ions
Ionization energy
Electron affinity
Some group chemical property trends
• First, we will discuss a fundamental property that
may leads to the trends, effective nuclear charge.
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7.2 EFFECTIVE NUCLEAR CHARGE
Effective Nuclear Charge
• Many of the properties of atoms
depend on their electron configurations
and on how strongly their outer
electrons are attracted to the nucleus.
• In a many-electron atom, electrons are
both attracted to the nucleus and
repelled by the other electrons.
• The nuclear charge that an electron
experiences depends on both factors
(the attraction and repulsion).
Figure 7.3 Effective
nuclear charge. The
effective nuclear
charge experienced
by the 3s electron in
a sodium atom
depends on the 11+
charge of the nucleus
and the 10– charge
of the core electrons.
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7.2 EFFECTIVE NUCLEAR CHARGE
Effective Nuclear Charge
• The effective nuclear
charge, Zeff acting on
an electron in an atom
is smaller than the
actual nuclear charge
(Zeff < Z) because the
effective nuclear
charge includes the
effect of the other
electrons in the atom.
Figure 7.2 An analogy for
effective nuclear charge. We
envision the nucleus as a light
bulb, and a valence electron
an observer. The amount of
light seen by the observer
depends on the screening by
the frosted glass lampshade.
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7.2 EFFECTIVE NUCLEAR CHARGE
Effective Nuclear Charge
• The effective nuclear charge,
Zeff, is found this way:
Zeff = Z − S
where Z is the atomic number
and S is a screening constant,
usually close to the number of
inner electrons.
• Ex) The magnitude of Zeff for
Na ([Ne]3s1) 3s electron.
– The nuclear charge: Z = 11+, the
screening constant: S = 10 (10
core electrons, 1s22s22p6).
– The effective nuclear charge for
the 3s electron: Zeff = 11 − 10 = 1+.
Figure 7.3 Effective nuclear
charge. The effective nuclear
charge experienced by the 3s
electron in a sodium atom
depends on the 11+ charge of
the nucleus and the 10–
charge of the core electrons.
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7.2 EFFECTIVE NUCLEAR CHARGE
Effective Nuclear Charge
• Calculation of Zeff with advanced methods: Zeff for
the 3s electron (valence electron) of Na atom is 2.5+.
– The 3s electron has a small probability of being closer to
the nucleus, in the region occupied by the core electrons.
– Thus, the electron experiences a greater attraction than our
simple S = 10 model suggests.
– The value of Zeff is increased from our expected Zeff = 1+ to
Zeff = 2.5+ and the value of S is changed from 10 to 8.5.
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7.2 EFFECTIVE NUCLEAR CHARGE
Effective Nuclear Charge
• The notion of Zeff explains this effect: For a manyelectron atom, the energies of orbitals with the same n
value increase with increasing l value.
– Ex) Carbon atom (1s22s22p2): The energy of the 2p orbital (l = 1)
is higher than that of the 2s (l = 0) orbital.
– This difference in energies is
due to the radial probability
functions for the orbitals:
The greater attraction
between the 2s electron and
the nucleus leads to a lower
energy for the 2s orbital than
for the 2p orbital.
– General trend in orbital
Figure 7.4 Comparison of 1s, 2s, and 2p
energies in many-electron
radial probability functions.
atom: ns < np < nd.
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7.2 EFFECTIVE NUCLEAR CHARGE
Effective Nuclear Charge
• The effective nuclear charge increases from left to right
across any period of the periodic table.
• Going down a column, the effective nuclear charge
experienced by valence electrons changes far less than
it does across a period.
• Effective nuclear charge increases slightly as we go
down a column because the more diffuse core electron
cloud is less able to screen the valence electrons from
the nuclear charge.
– Ex) Alkali metals: The Zeff for Li: 1.3+, Na: 2.5+, K: 3.5+.
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7.2 EFFECTIVE NUCLEAR CHARGE
EFFECTIVE NUCLEAR CHARGE
• The core electrons are much more effective than the valence
electrons at screening the charge of the nucleus.
• Effective nuclear charge is used to understand many properties,
such as ionization energy, atomic radii, and electron affinity.
• More accurate approach to calculate Zeff:
Slater’s rules (Zeff = Z − S).
–
–
–
–
–
Electrons with larger value of n: S = 0.
Electrons with the same value of n: S = 0.35.
Electrons with 1 less value of n: S = 0.85.
Electrons with even smaller value of n: S = 1.00.
Ex) For a valence electron in fluorine, S = (0.35 ×
6) + (0.85 × 2) = 3.8, Zeff = Z − S = 9 – 3.8 = 5.2+.
Figure 7.5 Variations in effective nuclear charge for period 2 and
period 3 elements. Moving from one element to the next in the periodic
table, the increase in Zeff felt by the innermost (1s) electrons (red circles)
closely tracks the increase in nuclear charge Z (black line) because
these electrons are not screened. The results of several methods to
calculate Zeff for valence electrons are shown in other colors.
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7.3 SIZES OF ATOMS AND IONS
Sizes of Atoms
• The nonbonding atomic
radius or the van der Waals
radius is half of the shortest
distance separating the two
nuclei during a collision of
atoms.
• The bonding atomic radius
is half of the distance between
covalently bonded nuclei.
– The C–Cl bond length in CCl4:
1.77 Å, very close to the sum
(0.77 + 0.99 Å) of the bonding
atomic radii of C and Cl.
Figure 7.6 Distinction between nonbonding
and bonding atomic radii within a molecule.
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7.3 SIZES OF ATOMS AND IONS
Periodic Trends in Atomic Radii
• The bonding atomic radius tends to
– Decrease from left to right across a row (due to increasing Zeff).
– Increase from top to bottom of a group (due to the increasing
value of n).
Figure 7.7 Trends in bonding atomic radii for periods 1 through 5.
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7.3 SIZES OF ATOMS AND IONS
Periodic Trends in Ionic Radii
• Sizes of ions are
determined by
interatomic distances
in ionic compounds.
• Ionic size depends
up the nuclear
charge, the number
of electrons, and the
orbitals in which
electrons reside.
Figure 7.8 Cation and anion size. Radii,
in angstroms, of atoms and their ions for
five groups of representative elements.
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7.3 SIZES OF ATOMS AND IONS
Periodic Trends in Ionic Radii
• Cations are smaller than
their parent atoms.
– The outermost electron is
removed and repulsions
between electrons are reduced.
• Anions are larger than their
parent atoms.
– Electrons are added and
repulsions between electrons
are increased.
Figure 7.8 Cation and anion size. Radii,
in angstroms, of atoms and their ions for
five groups of representative elements.
• Ions increase in size as you
go down a column due to
the increasing value of n.
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7.3 SIZES OF ATOMS AND IONS
Periodic Trends in Ionic Radii
• In an isoelectronic series, ions have the same number
of electrons.
• Ionic size decreases with an increasing nuclear charge.
– Ex) Each ion in the isoelectronic series O2–, F–, Na+, Mg2+, Al3+
has 10 electrons.
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Ionization Energy
7.4 IONIZATION ENERGY
• The ionization energy is the minimum energy required to
remove an electron from the ground state of a gaseous
atom or ion.
– The first ionization energy, I1, is the energy required to remove the
first electron from a neutral atom.
• Ex) I1 for the sodium atom is the energy required for the process.
– The second ionization energy, I2, is the energy required to remove
the second electron, etc.
• Ex) I2 for the sodium atom is the energy associated with the process.
• The greater the ionization energy, the more difficult it is to
remove an electron.
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7.4 IONIZATION ENERGY
Variations in Successive Ionization Energy
• Ionization energies for a given element increase as
successive electrons are removed: I1 < I2 < I3, and so forth.
– With each successive removal, an electron is being pulled away from
an increasingly more positive ion, requiring increasingly more energy.
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7.4 IONIZATION ENERGY
Variations in Successive Ionization Energy
• When all valence electrons have been removed, a sharp
increase in ionization energy occurs.
– Ex) Si (1s22s22p63s23p2): Removal of the fifth electron (from the 2p
subshell) requires a great deal more energy, 16,091 kJ/mol. The 2p
electron experiences a much greater effective nuclear charge than
do the 3s and 3p electrons.
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7.4 IONIZATION ENERGY
Periodic Trends in First Ionization Energies
• As one goes down a column, less energy is required to
remove the first electron.
– As we move down a column, the atomic radius increases while
the effective nuclear charge increases rather gradually. Thus,
the attraction between the nucleus and the electron decreases.
Figure 7.10 Trends in first ionization
energies of the elements.
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7.4 IONIZATION ENERGY
Periodic Trends in First Ionization Energies
• Generally, as one goes across a row, it gets harder to
remove an electron.
– As we move across a period, there is both an increase in Zeff
and a decrease in atomic radius, causing the ionization energy
to increase.
Figure 7.10 Trends in first ionization
energies of the elements.
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7.4 IONIZATION ENERGY
Periodic Trends in First Ionization Energies
• The s- and p-block elements show a larger range of I1
values than do the transition-metal elements.
– Generally, the ionization energies of the transition metals increase
slowly from left to right in a period.
– The f-block metals also show only a small variation in the values of I1.
Figure 7.10 Trends in first ionization
energies of the elements.
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7.4 IONIZATION ENERGY
Periodic Trends in First Ionization Energies
• Two apparent discontinuities in this trend.
– Between Groups IIA (Be: [He]2s2) and IIIA (B: [He]2s22p1).
• The electron is removed from a p orbital rather than an s orbital.
• The 2p subshell is at a higher energy than the 2s subshell.
• The electron removed is farther from the nucleus.
– Between Groups VA (N: [He]2s22p3) and VIA (O: [He]2s22p4).
• The electron removed comes from a doubly occupied orbital.
• Repulsion of paired electrons in the p4 configuration aids in the removal.
Figure 7.10 Trends in first ionization energies of the elements.
Figure 7.11 2p orbital filling
in nitrogen and oxygen.
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7.4 IONIZATION ENERGY
Electron Configurations of Ions
• When electrons are removed from an atom to form a cation,
they are removed from the occupied orbitals having the
largest principal quantum number, n.
– Ex) The 2s1 electron from Li (1s22s1):
The 4s2 electrons from Fe ([Ar]3d64s2):
An additional electron from a 3d orbital of Fe2+:
• Electrons added to an atom to form an anion are added to
the empty or partially filled orbital having the lowest value of n.
– Ex) An electron added to a fluorine atom to form the ion goes into the
one remaining vacancy in the 2p subshell:
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Electron Affinity
7.5 ELECTRON AFFINITIES
• Electron affinity is the energy change accompanying the
addition of an electron to a gaseous atom:
• It is typically exothermic, so, for most elements, it is negative!
– The greater the attraction between a given atom and an added
electron, the more negative the atom’s electron affinity.
– For some elements, such as the noble gases, the electron affinity has
a positive value, meaning that the anion is higher in energy than are
the separated atom and electron (X– is unstable):
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Electron Affinity
7.5 ELECTRON AFFINITIES
• In general, electron affinity becomes more exothermic
as you go from left to right across a row.
• There are three notable exceptions in this trend.
– The group IIA and VA elements have low electron affinity.
– The group VIIIA (noble gas) elements have positive electron
affinity.
Figure 7.12 Electron affinity in kJ/mol for selected s- and p-block elements.
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Electron Affinity
7.5 ELECTRON AFFINITIES
– Between Groups IA (Li: [He]2s1) and IIA (Be: [He]2s2).
• The added electron must go in a previously empty p subshell
that is higher in energy.
– Between Groups IVA (C: [He]2s22p2) and VA (N: [He]2s22p3).
• Group VA has no empty orbitals.
• The extra electron must go into an already occupied orbital,
resulting in larger electron–electron repulsions.
Figure 7.12 Electron affinity in kJ/mol for selected s- and p-block elements.
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Electron Affinity
7.5 ELECTRON AFFINITIES
• Electron affinities do not change greatly in the same group.
– As we move down a group, the distance between the added electron
and the nucleus increases: The electron–nucleus attraction decreases.
– As we move down a group, the orbital that holds the outermost
electron is increasingly spread out: The electron–electron repulsions
are also reduced.
– The reduction in the electron–nucleus attraction is counterbalanced by
the reduction in electron–electron repulsions.
Figure 7.12 Electron affinity in kJ/mol for selected s- and p-block elements.
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7.6 METALS, NONMETALS, AND METALLOIDS
Metals, Nonmetals, and Metalloids
• The elements can be broadly grouped as metals, nonmetals,
and metalloids.
• The close relationships exist between electron configurations
and the properties of metals, nonmetals, and metalloids.
Figure 7.13 Metals, metalloids, and nonmetals.
40
7.6 METALS, NONMETALS, AND METALLOIDS
Metallic Character
• Metallic character generally increases as we proceed
down a group and left across a period of the periodic table.
Figure 7.13 Metals, metalloids, and nonmetals.
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7.6 METALS, NONMETALS, AND METALLOIDS
Metals
• Most of the elements in nature are metals.
• Properties of metals:
–
–
–
–
–
Shiny luster
Conduct heat and electricity
Malleable and ductile
Solids at room temperature (except mercury)
Low ionization energies/form cations easily
Figure 7.14 Metals
are shiny, malleable,
and ductile.
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7.6 METALS, NONMETALS, AND METALLOIDS
Metals
• Metals tend to have low ionization energies and therefore
form cations relatively easily.
– Metals are easily oxidized when they undergo chemical reactions.
– Cf) Nonmetals tend to form anions.
Figure 7.15 Representative oxidation states of the elements. Note that
hydrogen has both positive and negative oxidation numbers, +1 and –1.
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7.6 METALS, NONMETALS, AND METALLOIDS
Metals
• Representative oxidation states of metals.
– The charge on any alkali metal ion is always 1+, and that on any
alkaline earth metal is always 2+ in their compounds.
– For metals belonging to groups with partially occupied p orbitals
(groups 3A–7A), cations are formed either by losing only the outer p
electrons (such as Sn2+) or the outer s and p electrons (such as Sn4+).
– Transition metals have various oxidation states (Fe2+ and Fe3+).
Figure 7.14 Representative oxidation states of the elements. Note that
hydrogen has both positive and negative oxidation numbers, +1 and –1.
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7.6 METALS, NONMETALS, AND METALLOIDS
Metals
• Compounds made up of a metal and a nonmetal tend to
be ionic substances.
– Most metal oxides and halides are ionic solids.
– Ex) The reaction between nickel metal and oxygen:
• Most metal oxides are basic.
– Those that dissolve in water react to form metal hydroxides:
– The basicity of metal oxides is due to the oxide ion, which reacts
with water:
45
7.6 METALS, NONMETALS, AND METALLOIDS
Metals
– Even metal oxides that are insoluble in water react with acids:
Figure 7.16 Metal oxides react with acids. NiO does not dissolve in water
but does react with nitric acid (HNO3) to give a green solution of Ni(NO3)2.
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7.6 METALS, NONMETALS, AND METALLOIDS
Nonmetals
• Nonmetals are found on the right
hand side of the periodic table.
• Properties of nonmetals include the
following:
– Solid, liquid, or gas (depends on element)
– Solids are dull, brittle, poor conductors.
– Because of their relatively large, negative
electron affinities, nonmetals readily form
anions when they react with metals.
Figure 7.17 Sulfur, known
to the medieval world as
“brimstone,” is a nonmetal.
• Ex) The reaction of aluminum with bromine:
48
7.6 METALS, NONMETALS, AND METALLOIDS
Nonmetals
• Substances containing only nonmetals are molecular
compounds.
– Ex) The common hydrocarbons (CH4, C8H18), the gases HCl,
NH3, and H2S, and many nonmetal drugs.
• Most nonmetal oxides are acidic.
– Ex) Nonmetal oxides that dissolve in water form acids:
SO2 and SO3 dissolve in water to produce acid rain.
Figure 7.18 The reaction of CO2 with water containing
a bromthymol blue indicator. Initially, the blue color
tells us the water is slightly basic. When a piece of solid
carbon dioxide (“dry ice”) is added, the color changes to
yellow, indicating an acidic solution. The mist is water
droplets condensed from the air by the cold CO2 gas.
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7.6 METALS, NONMETALS, AND METALLOIDS
Metalloids
• Metalloids have some characteristics of metals and
some of nonmetals.
– Ex) Silicon looks shiny, but is brittle and a fairly poor conductor.
• Several metalloids are electrical semiconductors.
– Very pure silicon is an electrical insulator, but its conductivity
can be dramatically increased with the addition of specific
impurities called dopants.
– Silicon is the principal element used in integrated circuits and
computer chips.
Figure 7.19 Elemental silicon.
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7.7 TRENDS FOR GROUP 1A AND GROUP 2A METALS
Group Trends
• Elements in a group have similar properties.
• Trends also exist within groups.
• Groups Compared:
–
–
–
–
–
Group 1A: The Alkali Metals
Group 2A: The Alkaline Earth Metals
Group 6A: The Oxygen Group
Group 7A: The Halogens
Group 8A: The Noble Gases
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7.7 TRENDS FOR GROUP 1A AND GROUP 2A METALS
Group 1A: The Alkali Metals
• Alkali metals are soft, metallic solids.
• The name alkali comes from the Arabic
word for ashes.
• Typical metallic properties (luster,
conductivity) are seen in them.
• They have low densities and melting points.
• The outer s electron can be easily removed:
Low I1 value.
Figure 7.20 Sodium,
like the other alkali
metals, is soft enough
to be cut with a knife.
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7.7 TRENDS FOR GROUP 1A AND GROUP 2A METALS
Group 1A: The Alkali Metals
• They are found only in compounds in nature, not
in their elemental forms.
– All alkali metals combine directly with most nonmetals.
– Ex) They react with hydrogen to form hydrides and with
sulfur to form sulfides:
In hydrides of the alkali metals
(LiH, NaH, and so forth), hydrogen
is present as H–, the hydride ion.
• The alkali metals react vigorously with water.
– These reactions are very exothermic and they produce
hydrogen gas and a solution of an alkali metal hydroxide:
Figure 7.21 The alkali metals
react vigorously with water.
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7.7 TRENDS FOR GROUP 1A AND GROUP 2A METALS
Differences in Alkali Metal Chemistry
• The reaction between the alkali metals and oxygen are
complex.
– Li form a metal oxide (containing the O2– ion):
4 Li(s) + O2(g) → 2 Li2O (s)
Lithium oxide
– Other alkali metals (except Li) form metal peroxides (O22– ion):
2 Na(s) + O2(g) → Na2O2(s)
Sodium peroxide
– K, Rb, and Cs also form superoxides (O2– ion):
K(s) + O2(g) → KO2(s)
Potassium superoxide
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7.7 TRENDS FOR GROUP 1A AND GROUP 2A METALS
Group 1A: The Alkali Metals
• Each emits a characteristic color when placed in a flame.
– The ions are reduced to gaseous metal atoms in the flame.
– The light emitted is at a specific wavelength for each element.
Figure 7.22 Placed in a flame,
ions of each alkali metal emit light
of a characteristic wavelength.
– The characteristic yellow
emission of sodium at 589 nm is
the basis for sodium vapor lamps.
Figure 7.23 The characteristic yellow light in a sodium
lamp results from excited electrons in the high-energy
3p orbital falling back to the lower-energy 3s orbital.
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7.7 TRENDS FOR GROUP 1A AND GROUP 2A METALS
Group 2A: The Alkaline Earth Metals
• Alkaline earth metals have higher densities and melting
points than alkali metals.
• Their ionization energies are low, but not as low as those
of alkali metals.
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7.7 TRENDS FOR GROUP 1A AND GROUP 2A METALS
Group 2A: The Alkaline Earth Metals
• Reactivity tends to increase as you go down the group.
– Beryllium does not react with water, and magnesium reacts only with
steam, but the other alkaline earth metals react readily with water:
Figure 7.25 Elemental calcium reacts with water.
• The heavier alkaline earth ions give off characteristic colors
when heated in a hot flame.
– Ex) Sr: The brilliant red color, Ba: The green color.
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7.8 TRENDS FOR SELECTED NONMETALS
Hydrogen
• Electron configuration of hydrogen: 1s1.
– Its position in the periodic table is above the alkali metals;
however, it does not truly belong to any particular group.
– It is a nonmetal that occurs as a colorless diatomic gas, H2(g).
– High ionization energy due to the absence of nuclear shielding.
• Reactions with another nonmetal can be quite exothermic.
– Ex) The combustion reaction with oxygen to form water:
• It reacts with active metals to form solid metal hydrides.
• It loses an electron to form a cation.
– It is the most important characteristic in the aqueous chemistry of
hydrogen.
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7.8 TRENDS FOR SELECTED NONMETALS
Group 6A: The Oxygen Group
• Group 6A contains nonmetals (O, S, and Se), a metalloid
(Te) and a radioactive metal (Po).
• Oxygen is a colorless gas at room temperature; all of the
other members of group 6A are solids.
• The thermal stability of group 6A compounds with hydrogen
decreases down the column: H2O > H2S > H2Se > H2Te.
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7.8 TRENDS FOR SELECTED NONMETALS
Group 6A: The Oxygen Group
• There are two allotropes of oxygen: O2 and O3 (ozone).
– Allotropes are different forms of the same element in the same state.
– O3 is less stable than O2.
– They are colorless and, therefore, do not absorb visible light.
– Ozone absorbs UV light: Ozone in the upper atmosphere filter out
harmful UV light.
– Ozone is a powerful oxidant: It is added to water to kill bacteria.
• Oxygen has a great tendency to attract electrons from other
elements (oxidation).
– There can be three anions of oxygen: The oxide (O2–) ion, the
peroxide (O22–) ion, and the superoxide (O2–) ion.
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7.8 TRENDS FOR SELECTED NONMETALS
Group 6A: The Oxygen Group
• The most stable allotrope of sulfur
is S8, a ringed molecule.
Figure 7.27 Elemental sulfur exists as the
S8 molecule. At room temperature, this is
the most common allotropic form of sulfur.
• Sulfur has a tendency to gain electrons from other elements
to form sulfides (Contain the S2– ion).
– Most sulfur in nature is present as metal sulfides.
– Sulfur is a weaker oxidizer than oxygen; sulfur can be burned in
oxygen to form sulfur dioxide (a major air pollutant):
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7.8 TRENDS FOR SELECTED NONMETALS
Group 7A: The Halogens
• The halogens are nonmetals.
• Each element consists of diatomic
molecules (F2, Cl2, Br2, I2).
• They have large, negative electron
affinities.
– They tend to oxidize other elements
very exothermically:
Figure 7.28 The elemental halogens
exist as diatomic molecules.
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7.8 TRENDS FOR SELECTED NONMETALS
Group 7A: The Halogens
• Chlorine is the most industrially useful of the halogens.
– Chlorine is added to drinking water and swimming pools, where the
generated HOCl(aq) serves as a disinfectant:
• The halogens directly react with metals to form metal halides.
• The halogens also react with hydrogen to form gaseous
hydrogen halide compounds:
– These compounds are very soluble in water and dissolve to form the
hydrohalic acids.
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7.8 TRENDS FOR SELECTED NONMETALS
Group 8A: The Noble Gases
• The noble gases are all monatomic gases.
• They have completely filled s and p subshells (stable
electron configurations).
– The noble gases have very large ionization energies.
– Their electron affinities are positive (can’t form stable anions).
– Therefore, they are relatively unreactive.
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7.8 TRENDS FOR SELECTED NONMETALS
Group 8A: The Noble Gases
• The elements were called the inert gases.
– Until the early 1960s, they were thought to be incapable of
forming chemical compounds.
– Bartlett synthesized the first noble-gas compound (compound
containing both XeFPtF6 and XeFPt2F11).
– Direct reaction of Xe with F2(g) to form XeF2, XeF4, and XeF6.
• Krypton has a higher I1 value than xenon and is therefore
less reactive.
– Only a single stable compound of krypton is known, KrF2.
• In 2000, the first neutral molecule, that contains argon
(HArF), was reported.
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68
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Chapter 7. Homework
Exercises 7.6
7.7
7.14
7.16
7.19
7.22
7.34
7.42
7.53
7.55
7.65
7.70
7.73
7.99
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