Transcript Chapter 7 & 9.1-9.2 Powerpoint

Chapter 7
Periodic Properties of
Elements
7.1 Development of the
Periodic Table
Development of Periodic Table
► Elements
in the same
group generally have
similar chemical
properties.
► Physical properties are
not necessarily similar,
however.
Development of Periodic Table
Dmitri Mendeleev
and Lothar Meyer
independently
came to the same
conclusion about
how elements
should be grouped.
Development of Periodic Table
Mendeleev, for instance, predicted the
discovery of germanium (which he called ekasilicon) as an element with an atomic weight
between that of zinc and arsenic, but with
chemical properties similar to those of silicon.
7.2 Effective Nuclear
Charge
Periodic Trends
► In
this chapter, we will rationalize observed
trends in
 Sizes of atoms and ions.
 Ionization energy.
 Electron affinity.
Effective Nuclear Charge
► In
a many-electron
atom, electrons are
both attracted to the
nucleus and repelled by
other electrons.
► The nuclear charge
that an electron
experiences depends
on both factors.
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.
7.3 Sizes of Atoms and
Ions
What Is the Size of an Atom?
The bonding atomic
radius is defined as
one-half of the
distance between
covalently bonded
nuclei.
Practice
► 1)
Natural gas used in home heating and
cooking is odorless. Because natural gas
leaks pose the danger of explosion or
suffocation, various smelly substances are
added to detect leaks. One substance is
methyl mercaptan, CH3SH. Use Figure 7.6
to predict the lengths of the C-S, C-H, and
S-H bonds in the molecule.
► 2) Use Figure 7.6 to predict which will be
greater: the P-Br bond length in PBr3 or the
As-Cl bond length in AsCl3.
Sizes of Atoms
Bonding atomic
radius tends to…
…decrease from left to
right across a row
(due to increasing Zeff).
…increase from top to
bottom of a column
(due to increasing value
of n).
Practice
► 1)
Arrange the following atoms in order of
increasing size: 15P, 16S, 33As, and 34Se.
► 2) Arrange the following atoms in order of
increasing atomic radius: 11Na, 4Be, and
12Mg.
Sizes of Ions
► Ionic
size depends
upon:
 The nuclear
charge.
 The number of
electrons.
 The orbitals in
which electrons
reside.
Sizes of Ions
► Cations
are
smaller than their
parent atoms.
 The outermost
electron is
removed and
repulsions
between electrons
are reduced.
Sizes of Ions
► Anions
are larger
than their parent
atoms.
 Electrons are
added and
repulsions
between electrons
are increased.
Sizes of Ions
► Ions
increase in size
as you go down a
column.
 This is due to
increasing value of n.
Sizes of Ions
► In
an isoelectronic series, ions have the same
number of electrons.
► Ionic size decreases with an increasing nuclear
charge.
Sample Exercise 7.3 Atomic and Ionic Radii
Arrange these atoms and ions in order of
decreasing size: Mg2+, Ca2+, and Ca.
Practice Exercise
Which of the following atoms and ions is
largest: S2–, S, O2–?
Sample Exercise 7.4 Ionic Radii in an Isoelectronic Series
Arrange the ions K+, Cl–, Ca2+, and S2– in
order of decreasing size.
Practice Exercise
Which of the following ions is largest,
Rb+, Sr2+, or Y3+?
7.4 Ionization Energy
Ionization Energy
► The
ionization energy is the amount
of energy required to remove an
electron from the ground state of a
gaseous atom or ion.
 The first ionization energy is that energy
required to remove first electron.
 The second ionization energy is that
energy required to remove second
electron, etc.
Ionization Energy
► It
requires more energy to remove each
successive electron.
► When all valence electrons have been removed,
the ionization energy takes a quantum leap.
Trends in First Ionization Energies
► As
one goes down a
column, less energy is
required to remove
the first electron.
 For atoms in the same
group, Zeff is
essentially the same,
but the valence
electrons are farther
from the nucleus.
Trends in First Ionization Energies
► Generally,
as one goes
across a row, it gets
harder to remove an
electron.
 As you go from left to
right, Zeff increases.
Trends in First Ionization Energies
However, there are
two apparent
discontinuities in this
trend.
Trends in First Ionization Energies
► The
first occurs
between Groups IIA
and IIIA.
► In this case the
electron is removed
from a p-orbital rather
than an s-orbital.
 The electron removed is
farther from nucleus.
 There is also a small
amount of repulsion by
the s electrons.
Trends in First Ionization Energies
► The
second occurs
between Groups VA
and VIA.
 The electron removed
comes from doubly
occupied orbital.
 Repulsion from the
other electron in the
orbital aids in its
removal.
Sample Exercise 7.5 Trends in Ionization Energy
Three elements are indicated in the
periodic table in the margin (page 272).
Based on their locations, predict the one
with the largest second ionization energy.
Practice Exercise
Which will have the greater third
ionization energy, Ca or S?
Sample Exercise 7.6 Periodic Trends in Ionization Energy
Referring to a periodic table, arrange the
following atoms in order of increasing first
ionization energy: Ne, Na, P, Ar, K.
Practice Exercise
Which has the lowest first ionization energy, B,
Al, C, or Si? Which has the highest first
ionization energy?
Sample Exercise 7.7 Electron Configurations of Ions
Practice Exercise
Write the electron configuration for (a)
Ga3+, (b) Cr3+, and (c) Br–.
Write the electron configuration for (a)
Ga3+, (b) Cr3+, and (c) Br-.
7.5 Electron Affinities
Electron Affinity
Electron affinity is the energy change
accompanying the addition of an electron to
a gaseous atom:
Cl + e−  Cl−
Trends in Electron Affinity
In general, electron
affinity becomes
more exothermic as
you go from left to
right across a row.
Trends in Electron Affinity
There are again,
however, two
discontinuities in
this trend.
Trends in Electron Affinity
► The
first occurs
between Groups IA
and IIA.
 The added electron
must go in a p-orbital,
not an s-orbital.
 The electron is farther
from nucleus and
feels repulsion from
the s-electrons.
Trends in Electron Affinity
► The
second occurs
between Groups IVA
and VA.
 Group VA has no
empty orbitals.
 The extra electron
must go into an
already occupied
orbital, creating
repulsion.
7.6 Metals, Nonmetals,
and Metalloids
Properties of Metal, Nonmetals,
and Metalloids
Metals versus Nonmetals
Differences between metals and nonmetals tend to
revolve around these properties.
Metals versus Nonmetals
► Metals
tend to form cations.
► Nonmetals tend to form anions.
Metals
They tend to be
lustrous, malleable,
ductile, and good
conductors of heat and
electricity.
Metals
► Compounds
formed
between metals and
nonmetals tend to be
ionic and solid.
► Metal oxides tend to
be basic.
Sample Exercise 7.8 Metal Oxides
(a) Would you expect scandium oxide to
be a solid, liquid, or gas at room
temperature? (b) Write the balanced
chemical equation for the reaction of
scandium oxide with nitric acid.
Practice Exercise
Write the balanced chemical equation for
the reaction between copper(II) oxide and
sulfuric acid.
Nonmetals
► These
are dull, brittle
substances that are
poor conductors of
heat and electricity.
► They tend to gain
electrons in reactions
with metals to acquire
a noble gas
configuration.
Nonmetals
► Substances
containing
only nonmetals are
molecular compounds.
► Most nonmetal oxides
are acidic.
Sample Exercise 7.9 Nonmetal Oxides
Write the balanced chemical equations for
the reactions of solid selenium dioxide
with (a) water, (b) aqueous sodium
hydroxide.
Practice Exercise
Write the balanced chemical equation for
the reaction of solid tetraphosphorus hexoxide with water.
Metalloids
► These
have some
characteristics of
metals and some of
nonmetals.
► For instance, silicon
looks shiny, but is
brittle and fairly poor
conductor.
7.7 Group Trends for the
Active Metals
Alkali Metals
► Alkali
metals are soft,
metallic solids.
► The name comes from
the Arabic word for
ashes.
Alkali Metals
► They
are found only in compounds in nature, not
in their elemental forms.
► They have low densities and melting points.
► They also have low ionization energies.
Alkali Metals
Their reactions with water are famously exothermic.
Alkali Metals
► Alkali
metals (except Li) react with oxygen to
form peroxides.
► K, Rb, and Cs also form superoxides:
K + O2  KO2
► They produce bright colors when placed in a
flame.
Practice
► 1)
Write a balanced equation that predicts
the reaction of cesium metal with:
 A) Cl2(g)
 B) H2O(l)
 C) H2(g)
► 2)
Write a balanced equation for the
reaction between potassium metal and
elemental sulfur.
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.
Alkaline Earth Metals
► Beryllium
does not
react with water and
magnesium reacts only
with steam, but the
others react readily
with water.
► Reactivity tends to
increase as you go
down the group.
7.8 Group Trends for
Selected Nonmetals
Group 6A
► Oxygen,
sulfur, and selenium are nonmetals.
► Tellurium is a metalloid.
► The radioactive polonium is a metal.
Oxygen
► There
are two allotropes of
oxygen:
 O2
 O3, ozone
► There
can be three anions:
 O2−, oxide
 O22−, peroxide
 O21−, superoxide
► It
tends to take electrons
from other elements
(oxidation).
Sulfur
► Sulfur
is a weaker
oxidizer than
oxygen.
► The most stable
allotrope is S8, a
ringed molecule.
Group VIIA: Halogens
► The
halogens are prototypical nonmetals.
► The name comes from the Greek words halos
and gennao: “salt formers”.
Group VIIA: Halogens
► They
have large, negative
electron affinities.
 Therefore, they tend to
oxidize other elements
easily.
► They
react directly with
metals to form metal
halides.
► Chlorine is added to water
supplies to serve as a
disinfectant
Group VIIIA: Noble Gases
► The
noble gases have astronomical ionization
energies.
► Their electron affinities are positive.
 Therefore, they are relatively unreactive.
► They
are found as monatomic gases.
Group VIIIA: Noble Gases
► Xe
forms three
compounds:
 XeF2
 XeF4 (at right)
 XeF6
► Kr
forms only one stable
compound:
 KrF2
► The
unstable HArF was
synthesized in 2000.
Integrative Exercise
►
Bismuth is the heaviest member of Group 5A and is used in
Pepto Bismol.
 A) The covalent atomic radii of thallium and lead are 1.48 Å and 1.47 Å
respectively. Using this and Figure 7.6, predict the radius of bismuth.
 B) What accounts for the general increase in atomic radius going down
group 5A?
 C) Bismuth is also used in low-melting alloys such as sprinkler systems.
The element itself is a brittle, white crystalline solid. How do these
characteristics fit with the fact that bismuth is in the same group as N
and P?
 D) Bi2O3 is a basic oxide. Write a balanced equation for its reaction with
nitric acid. If 6.77 g of Bi2O3 is dissolved in dilute acidic solution to make
up to 500 mL of solution, what is the molarity of the solution of Bi3+ion?
 E) 209Bi is the heaviest stable isotope of any element. How many
protons and neutrons are present in the nucleus?
 F) The density of Bi at 25ºC is 9.808 g/cm3. How many Bi atoms are
present in a cube of the element that is 5.00 cm on each edge? How
many moles of the element are present?