Transcript Group IA, Alkali Metals

Periodic Relationships Among
the Elements
Chapter 8
Periodic Table - In the Beginning
• A necessary prerequisite to the construction of the 1
periodic table was the discovery of the individual
elements.
• Although elements such as gold, silver, tin, copper,
lead, and mercury have been known since antiquity, the
first scientific discovery of an element occurred in
1649 when Hennig Brand discovered phosphorous.
• By 1869, a total of 63 elements had been discovered.
As the number of known elements grew, scientists
began to recognize patterns in properties and began to
develop classification schemes.
Periodic Table – Law of Triads
• In 1817 Johann Dobereiner noticed that the atomic weight
1
of strontium fell midway between the weights of calcium
and barium, elements possessing similar chemical
properties.
• In 1829, after discovering the halogen triad composed of
chlorine, bromine, and iodine and the alkali metal triad of
lithium, sodium and potassium he proposed that nature
contained triads of elements the middle element had
properties that were an average of the other two members
when ordered by the atomic weight (the Law of Triads).
Periodic Table – Law of Triads
• This new idea of triads became a popular area of study. 1
Between 1829 and 1858 a number of scientists (Jean
Baptiste Dumas, Leopold Gmelin, Ernst Lenssen, Max von
Pettenkofer, and J.P. Cooke) found that these types of
chemical relationships extended beyond the triad.
• During this time fluorine was added to the halogen group;
oxygen, sulfur, selenium and tellurium were grouped into a
family while nitrogen, phosphorus, arsenic, antimony, and
bismuth were classified as another.
• Unfortunately, research in this area was hampered by the
fact that accurate values of were not always available.
First Attempts At Designing a
Periodic Table
1
• in 1862 French geologist, A.E.Beguyer de Chancourtois
transcribed a list of the elements positioned on a cylinder in
terms of increasing atomic weight.
• When the cylinder was constructed so that 16 mass units
could be written on the cylinder per turn, closely related
elements were lined up vertically.
• De Chancourtois was first to recognize that elemental
properties reoccur every seven elements, and using this
chart, he was able to predict the stoichiometry of several
metallic oxides. Unfortunately, his chart included some
ions and compounds in addition to elements.
Law of Octaves
1
• John Newlands, an English chemist, wrote a paper in 1863
which classified the 56 established elements into 11 groups
based on similar physical properties, noting that many pairs
of similar elements existed which differed by some multiple
of eight in atomic weight.
• In 1864 Newlands published his version of the periodic
table and proposed the Law of Octaves (by analogy with the
seven intervals of the musical scale). This law stated that
any given element will exhibit analogous behavior to the
eighth element following it in the table.
Father of Periodic Table
1
• There has been some disagreement about who deserves
credit for being the "father" of the periodic table, the
German Lothar Meyer or the Russian Dmitri Mendeleev.
• Both chemists produced remarkably similar results at the
same time working independently of one another.
Father of Periodic Table
1
• Meyer's 1864 textbook included a rather abbreviated version
of a periodic table used to classify the elements. This
consisted of about half of the known elements listed in order
of their atomic weight and demonstrated periodic valence
changes as a function of atomic weight.
• In 1868, Meyer constructed an extended table which he
gave to a colleague for evaluation.
• Unfortunately for Meyer, Mendeleev's table became
available to the scientific community via publication (1869)
before Meyer's appeared (1870).
Periodic Table – Noble Gases
1
• In 1895 Lord Rayleigh reported the discovery of a new
gaseous element named argon which proved to be
chemically inert. This element did not fit any of the known
periodic groups.
• In 1898, William Ramsey suggested that argon be placed
into the periodic table between chlorine and potassium in a
family with helium, despite the fact that argon's atomic
weight was greater than that of potassium. This group was
termed the "zero" group due to the zero valiancy of the
elements. Ramsey accurately predicted the future discovery
and properties neon.
The Modern Periodic Table
1
• The last major changes to the periodic table resulted from
Glenn Seaborg's work in the middle of the 20th Century.
Starting with his discovery of plutonium in 1940, he
discovered all the transuranic elements from 94 to 102 using
the particle accelerator at University of California,
Berkeley.
• He reconfigured the periodic table by placing the actinide
series below the lanthanide series.
• In 1951, Seaborg was awarded the Nobel Prize in chemistry
for his work.
• Element 106 has been named seaborgium (Sg) in his honor.
The Periodic Law and the
Periodic Table
• Periodic Law - the physical and chemical
properties of the elements are periodic
functions of their atomic numbers.
1
When the Elements Were Discovered
4f
5f
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
Classification of the Elements
Electron Arrangement and the
Periodic Table
• Electron configuration - describes the
arrangement of electrons in atoms.
• The electron arrangement is the primary
factor in understanding how atoms join
together to form compounds.
• Valance electrons - the outermost electrons.
– These are the electrons involved in chemical
bonding.
3
Valance Electrons
• For the representative elements:
– The number of valance electrons is the
group number.
– The period number gives the energy
level (n) of the valance shell.
• For an atom of fluorine, how many valance
electrons does it have and what is the energy
level of these electrons?
• Fluorine has 7 electrons in the n=2 level
• Let’s look at fluorine more closely.
• What is the total number of electrons in
fluorine?
– The atomic number is 9. It therefore
has 9 protons and 9 electrons.
• If there are 7 electrons in the valance shell,
(with n = 2 energy level) where are the other
two electrons?
– In the n = 1 energy level. This level
holds two and only two electrons.
• Isoelectronic - they have the same electron
configuration (same number of electrons)
• Nonmetallic elements tend to form negatively
charged ions called anions.
• Nonmetals tend to gain electrons so they become
isoelectronic with its nearest noble gas neighbor.
O + 2e[He]2s22p4
O2[He]2s22p6 or [Ne]
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
Atoms gain electrons
so that anion has a
noble-gas outer
electron configuration.
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
Na+: [Ne]
Al3+: [Ne]
O2-: 1s22s22p6 or [Ne]
F-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He
Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal,
electrons are always removed first from the ns orbital and
then from the (n – 1)d orbitals.
Fe:
[Ar]4s23d6
Fe2+: [Ar]4s03d6 or [Ar]3d6
Fe3+: [Ar]4s03d5 or [Ar]3d5
Mn:
[Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
Trends in the Periodic Table
• We will look at the following trends
–
–
–
–
–
In effective nuclear charge
in atomic size
in ion size
in ionization energy
in electron affinity
Periodic Properties
• Two factors determine the size of an atom.
One factor is the principal quantum
number, n. The larger is “n”, the larger the
size of the orbital.
The other factor is the effective nuclear
charge, which is the positive charge an
electron experiences from the nucleus
minus any “shielding effects” from
intervening electrons.
Effective nuclear charge (Zeff) is the “positive charge” felt
by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Z
Core
Zeff
Radius (pm)
Na
11
10
1
186
Mg
12
10
2
160
Al
13
10
3
143
Si
14
10
4
132
Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
Atomic Size
7
1. The size of the atoms increases from top to
bottom.
• This is due to the valance shell being
higher in energy and farther from the
nucleus.
2. The size of the atoms decreases from left to
right.
• This is due to the increase in
magnitude of positive charge in
nucleus. The nuclear charge pulls the
electrons closer to the nucleus.
Atomic Radii
Variation in Size of Atoms
Comparison of Atomic Radii with Ionic Radii
Cation is always smaller than atom from
which it is formed.
Anion is always larger than atom from
which it is formed.
Ion Size
7
• Cations are always smaller than their parent
atom.
– This is due to more protons than
electrons. The extra protons pulls the
remaining electrons closer.
– Which would be smaller, Fe2+ or
Fe3+?
– Fe3+
– This size trend is also due to the fact
that it is the outer shell that is lost.
• Anions are always larger than their parent
atom.
– This is due to the fact that anions
have more electrons than protons.
The Radii (in pm) of Ions of Familiar Elements
Ionization Energy
• Ionization energy - The energy required to
remove an electron from an isolated atom.
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8
ionization energy + Na  Na+ + e• The magnitude of ionization energy correlates with
the strength of the attractive force between the
nucleus and the outermost electron.
• The lower the ionization energy, the easier to form
a cation.
Ionization energy is the minimum energy (kJ/mol) required
to remove an electron from a gaseous atom in its ground
state.
I1 + X (g)
X+(g) + e-
I1 first ionization energy
I2 + X+(g)
X2+(g) + e-
I2 second ionization energy
I3 + X2+(g)
X3+(g) + e-
I3 third ionization energy
I1 < I2 < I3
Variation of the First Ionization Energy with Atomic Number
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
• Ionization
decreases down a
family because the
outermost
electrons are
farther from the
nucleus.
• Ionization increases across a period because
the outermost electrons are more tightly held.
• Why do you think that the noble gases
would be so unreactive?
General Trend in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
Electron Affinity
7
8
• Electron Affinity - The energy change when a
single electron is added to an isolated atom.
Br + e-  Br- + energy
• Electron affinity gives information about the
ease of anion formation.
– Large electron affinity indicates an
atom becomes more stable as it
forms an anion.
Electron affinity is the negative of the energy change that
occurs when an electron is accepted by an atom in the
gaseous state to form an anion.
X (g) + e-
X-(g)
F (g) + e-
X-(g)
DH = -328 kJ/mol
EA = +328 kJ/mol
O (g) + e-
O-(g)
DH = -141 kJ/mol
EA = +141 kJ/mol
Periodic Properties
• Electron Affinity
The electron affinity is the energy
change for the process of adding
an electron to a neutral atom in the
gaseous state to form a negative
ion.
For a chlorine atom, the first electron
affinity is illustrated by:


Cl([Ne]3s 3p )  e  Cl ([Ne]3s 3p )
2
5
2
6
Electron Affinity = -349 kJ/mol
Periodic Properties
• Electron Affinity
The more negative the electron affinity, the
more stable the negative ion that is formed.
Broadly speaking, the general trend goes
from lower left to upper right as electron
affinities become more negative.
Variation of Electron Affinity With Atomic Number (H – Ba)
• E.A. generally
decreases
down a group.
• E.A. generally increases across a period.
General Observations
• Several general observations can be made
about the main-group elements.
First, the metallic characteristics of these
elements generally decrease across a
period from left to right in the periodic
table.
Second, metallic characteristics of the
main-group elements become more
pronounced going down any column
(group).
General Observations
• Several general observations can be made
about the main-group elements.
Finally, a second-period element is
usually rather different from the other
elements in its group.
Hydrogen
• Hydrogen is the most abundant element in
the universe and is the third most abundant
element of the surface of the earth.
Most of the hydrogen on earth is found in
water.
Hydrogen has three isotopes: protium,
deuterium, and tritium.
Hydrogen
• Hydrogen
Protium is the most abundant, with less
that 0.02% being deuterium and only a
trace being radioactive tritium.
Deuterium and tritium isotopes can be
substituted for protium in chemical
compounds in order to provide markers
that can be followed during a chemical
reaction, or to change the chemical and
physical properties of the compound.
Hydrogen
• Hydrogen
Elemental hydrogen is produced on an
industrial scale by the steam-reforming
process in which a hydrocarbon is reacted
with water in the presence of a catalyst at
high temperature.
The bulk of the hydrogen produced in this
manner is used to make organic
compounds including methanol.
Hydrogen
• Hydrogen
Hydrogen forms three classes of binary
compounds called binary hydrides: ionic
hydrides, covalent hydrides, and
metallic hydrides.
The ionic hydrides are reactive solids
formed either by the reaction of hydrogen
with an alkali metal to form compounds with
the formula MH, or with larger alkaline earth
metals to form MH2.
Hydrogen
• Hydrogen
The covalent hydrides are compounds in
which hydrogen is covalently bonded to
another element.
Hydrogen
• Hydrogen
The metallic hydrides contain a transition
metal element and hydrogen.
In these compounds, the lattice of metal
atoms forms a porous structure that allows
hydrogen atoms to enter and bond.
Metallic hydrides are often
nonstoichiometric, meaning that the ration
of hydrogen atoms to metal atoms is not a
whole number.
An abridged periodic table showing the main-group elements.
Group IA, Alkali Metals
Largest atomic radii
React violently with water to form H2
Readily ionized to 1+
Metallic character, oxidized in air
R2O in most cases
Group 1A Elements (ns1, n  2)
M+1 + 1e-
2M(s) + 2H2O(l)
4M(s) + O2(g)
2MOH(aq) + H2(g)
2M2O(s)
Increasing reactivity
M
Group 1A Elements (ns1, n  2)
Group IA: The Alkali Metals
• The Group IA metals (alkali metals) are soft,
chemically reactive elements.
The alkali metals usually react by losing an
electron to become +1 cations.
Because of their reactivity, they never
occur as free metals.
They do occur extensively in silicate
minerals.
Group IA: The Alkali Metals
• The Group IA metals (alkali metals) are soft,
chemically reactive elements.
Lithium, sodium, and potassium are
important alkali metals.
In recent years, the commercial uses of
lithium chloride and is used in the
production of low-density alloys and as
a battery anode.
Group IA: The Alkali Metals
• Lithium
The commercial source of lithium is the ore
spudomene, LiAl(SiO3)2.
Lithium metal is obtained by the
electrolysis of of the chloride salt.
Group IA: The Alkali Metals
• Lithium
Lithium, like other alkali metals reacts with
water to produce lithium hydroxide and
hydrogen gas.
Lithium burns in air to produce lithium
oxide, Li2O, a white powder.
Group IA: The Alkali Metals
• Lithium
LiOH is used to make lithium soap for
lubricating greases.
LiNH2 is used in the preparation of
antihistamines.
LiH is used as a reducing agent in organic
synthesis.
A roll of lithium metal for
batteries.
Lithium battery.
Group IA: The Alkali Metals
• Sodium
Sodium metal is prepared in large
quantities.
It is used as a reducing agent in the
preparation of other metals, such as
titanium and zirconium, and in the
preparation of dyes and pharmaceuticals.
Sodium compounds available at the grocery store.
Group IA: The Alkali Metals
• Sodium
Sodium compounds are of enormous
economic importance.
Sodium chloride is the source of sodium
and most of its compounds.
Group IA: The Alkali Metals
• Sodium
Sodium hydroxide is prepared by the
electrolysis of aqueous sodium chloride;
as a strong base, it has many useful
commercial applications, including
aluminum production and petroleum
refining.
Group IA: The Alkali Metals
• Sodium
Sodium carbonate is obtained from the
mineral trona, which contains sodium
carbonate and sodium hydrogen
carbonate, and by the Solvay process
from salt (NaCl) and limestone (CaCO3).
Sodium carbonate is used to make glass.
Group IA: The Alkali Metals
• Potassium
Potassium metal is produced in relatively
small quantities, but potassium compounds
are important.
Large quantities of potassium chloride are
used as a plant fertilizer.
Group IIA, Alkali Earth
Metals
Readily ionized to 2+
React with water to form H2
Closed s shell configuration
Metallic
Group 2A Elements (ns2, n  2)
M+2 + 2e-
Be(s) + 2H2O(l)
Mg(s) + 2H2O(g)
M(s) + 2H2O(l)
No Reaction
Mg(OH)2(aq) + H2(g)
M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba
Increasing reactivity
M
Group 2A Elements (ns2, n  2)
Group IIA: The Alkaline Earth
Metals
Magnesium and calcium are the most
important of the Group IIA (alkaline earth)
metals.
Magnesium and its alloys are important
structural metals.
Group IIA: The Alkaline Earth
Metals
Magnesium and calcium are the most
important of the Group IIA (alkaline earth)
metals.
Calcium is important primarily as its
compounds, which are prepared from
natural carbonates, such as limestone, and
the sulfates, such as gypsum.
Group IIA: The Alkaline Earth
Metals
Magnesium and calcium are the most
important of the Group IIA (alkaline earth)
metals.
When limestone is heated strongly, it
decomposes to calcium oxide (lime).
Enormous quantities of lime are used in
the production of iron from its ores.
Group III A
Metals (except for boron)
Several oxidation states (commonly 3+)
Group 3A Elements (ns2np1, n  2)
4Al(s) + 3O2(g)
2Al(s) + 6H+(aq)
2Al2O3(s)
2Al3+(aq) + 3H2(g)
Group 3A Elements (ns2np1, n  2)
Group IV A
Form the most covalent compounds
Oxidation numbers vary between 4+ and 4-
Group 4A Elements (ns2np2, n  2)
Sn(s) + 2H+(aq)
Sn2+(aq) + H2 (g)
Pb(s) + 2H+(aq)
Pb2+(aq) + H2 (g)
Group 4A Elements (ns2np2, n  2)
Group IIIA and Group IVA
Metals
• Of the Group IIIA and Group IVA metals,
aluminum, tin, and lead are especially
important.
Aluminum is the third most abundant
element in the earth’s crust.
It is obtained commercially from bauxite;
through chemical processing bauxite yields
pure aluminum oxide.
Group IIIA and Group IVA
Metals
• Aluminum
Most of this aluminum oxide is used in the
production of aluminum by electrolysis.
Some aluminum oxide is used as a carrier
for heterogeneous catalysts and in
manufacturing industrial ceramic
materials.
Group IIIA and Group IVA
Metals
• Tin
Tin is normally a metal (called white tin)
but does undergo a low-temperature
conversion to a nonmetallic form (called
gray tin).
Tin is obtained by reduction of
cassiterite, a mineral form of SnO2.
Tin is used to make tin plate, bronze, and
solder.
Tin alloys.
Group IIIA and Group IVA
Metals
• Lead
Lead is obtained from galena, which is a
sulfide ore, PbS.
More than half of the lead produced is
used to make electrodes for lead storage
batteries.
Litharge, or lead(II) oxide, is an important
lead compound from which other lead
compounds are prepared.
Group IVA: The Carbon
Family
• Carbon is the least metallic of the Group
IVA elements.
Catenation is an important feature of
carbon chemistry and is responsible for the
enormous number of organic compounds.
Carbon has several allotropes, the principal
one being diamond and graphite, which are
covalent-network solids, and
buckminsterfullerene, which is molecular
(C60).
Group IVA: The Carbon
Family
• Carbon
The element has important industrial used,
including carbon black for rubber tires.
The principal oxides of carbon are CO and
CO2.
Mixtures of carbon monoxide and
hydrogen (synthesis gas) are used to
prepare various organic compounds.
Group IVA: The Carbon
Family
• Carbon
Liquid and solid carbon dioxide are used
as refrigerants, and the gas is used to
make carbonated beverages.
Group IVA: The Carbon
Family
• Silicon is the second most abundant
element in the crust of the earth.
The majority of silicon-containing compounds
consist of chains and networks of siliconoxygen bonds.
Sharing some of the structural attributes of
the network solids discussed in Chapter 13
are the silicones; materials that contain
chains or rings of Si—O bonds with organic
groups bonded to the silicon atoms.
Molecular model of a silicone.
Group IVA: The Carbon
Family
• Silicon
Depending on the amount and type of
bonding between the rings or chains, the
silicones can be formulated to be oils,
elastomers, or resins.
Due to their low reactivity and thermal
stability, the silicones find many
commercial and industrial applications
ranging from cosmetics to hydraulic fluids.
Products that contain silicone.
Group IVA: The Carbon
Family
• Silicon
The silicon hydrides (also known as
silanes) and silicon halides are siliconcontaining compounds that do not contain
Si—O bonds.
The silanes consist of straight chains or
branched chains of silicon atoms combined
with hydrogen atoms with the general
formula SinHn+2 or as rings with the formula
Si5H10 or Si6H12.
Group IVA: The Carbon
Family
• Silicon
The silicon halides can be formed by the
direct reaction of Si with a halogen.
Both the silanes and silicon halides are
very reactive.
Group V A
Form anions generally(1-, 2-, 3-), though
positive oxidation states are possible
Form metals, metalloids, and nonmetals
Group 5A Elements (ns2np3, n  2)
N2O5(s) + H2O(l)
P4O10(s) + 6H2O(l)
2HNO3(aq)
4H3PO4(aq)
Group 5A Elements (ns2np3, n  2)
Group VA: Nitrogen and the
Phosphorus Family
• Of the Group VA elements, nitrogen and
phosphorus are particularly important.
Nitrogen, N2, is obtained from liquid air by
fractional distillation; liquid nitrogen is used
as a refrigerant.
Ammonia, NH3, is the most important
compound of nitrogen.
It is prepared from the elements and is
used as a fertilizer.
Liquid nitrogen.
Group VA: Nitrogen and the
Phosphorus Family
• Nitrogen
Ammonia is also the starting compound for
the manufacture of other nitrogen
compounds.
For example, in the Ostwald process for
the preparation of nitric acid, ammonia is
burned in the presence of a catalyst to
nitric oxide, NO.
The nitric oxide reacts with oxygen to give
nitrogen dioxide, which reacts with water to
give nitric acid.
Industrial preparation of ammonia.
Group VA: Nitrogen and the
Phosphorus Family
• Phosphorus
Phosphorus has two common allotropes,
white phosphorus (P4) and red
phosphorus (chain structure).
White phosphorus is obtained by heating a
phosphate mineral with sand and coke in
an electric furnace.
When phosphorus burns in air, it forms
phosphorus(V) oxide, P4O10.
Group VA: Nitrogen and the
Phosphorus Family
• Phosphorus
This oxide reacts with water to give
orthophosphoric acid, H3PO4.
Phosphorus has many oxoacids; most are
obtained by condensation reactions with
orthophosphoric acid.
Group VA: Nitrogen and the
Phosphorus Family
• Phosphorus
One series is called the polyphosphoric
acids; they have the general formula
Hn+2PnO3n+1.
Triphosphoric acid is an example; sodium
triphosphate, Na5P3O10, is used in
detergents.
The metaphosphoric acids have the
general formula (HPO3)n.
Allotropes of phosphorus.
Group VI A
Form 2- anions generally, though positive
oxidation states are possible
React vigorously with alkali and alkali
earth metals
Nonmetals
Group 6A Elements (ns2np4, n  2)
SO3(g) + H2O(l)
H2SO4(aq)
Group 6A Elements (ns2np4, n  2)
Group VIA: Oxygen and the
Sulfur Family
• Oxygen, a Group VIA element, occurs in the
atmosphere (as O2), but mostly it is present on
earth as oxide and oxoanion minerals.
Oxygen has two allotropes: dioxygen, O2,
and ozone, O3.
Dioxygen, usually called simply oxygen, is
obtained commercially from liquid air.
Oxygen reacts with almost all elements to
give oxides or in some cases, peroxides or
super-oxides.
Group VIA: Oxygen and the
Sulfur Family
• Sulfur, another Group VIA element, occurs
in sulfate and sulfide minerals.
Free sulfur, S8 , occurring in deep underground deposits is mined by the Frasch
process.
Sulfur is also produced by the Claus
process, in which hydrogen sulfide
(obtained from natural gas and petroleum)
is partially burned.
Sulfur obtained from underground deposits.
The Frasch
process for
mining sulfur.
Group VIA: Oxygen and the
Sulfur Family
• Sulfur
Most of the sulfur is used to prepare
sulfuric acid by the contact process.
In this process, sulfur is burned to sulfur
dioxide, SO2, which in the presence of a
catalyst and oxygen forms sulfur trioxide,
SO3.
Group VIA: Oxygen and the
Sulfur Family
• Sulfur
This oxide dissolves in concentrated
sulfuric acid, which when diluted with water
gives additional sulfuric acid.
Sulfuric acid is the most important
compound of sulfur.
Halogens
Form monoanions
High electronegativity (electron affinity)
Diatomic gases
Most reactive nonmetals (F)
Group 7A Elements (ns2np5, n  2)
X2(g) + H2(g)
X-1
2HX(g)
Increasing reactivity
X + 1e-
Group 7A Elements (ns2np5, n  2)
Group VIIA: The Halogens
• The Group VIIA elements, or halogens, are
reactive.
Chlorine (Cl2), a pale greenish-yellow gas,
is prepared commercially by the
electrolysis of aqueous sodium chloride.
Its principal uses are in the preparation of
chlorinated hydrocarbons and as a
bleaching agent and disinfectant.
Group VIIA: The Halogens
• The Group VIIA elements, or halogens, are
reactive.
Hydrogen chloride, HCl, is one of the
most important compounds of chlorine;
aqueous solutions of HCl are known as
hydrochloric acid.
Group VIIA: The Halogens
• The Group VIIA elements, or halogens, are
reactive.
The halogens form a variety of oxoacids.
Noble Gases
Minimal reactivity
Monatomic gases
Closed shell
Group 8A Elements (ns2np6, n  2)
Completely filled ns and np subshells.
Highest ionization energy of all elements.
No tendency to accept extra electrons.
Group VIIIA: The Noble Gases
• The Group VIIIA elements, the noble
gases, were discovered at the end of the
nineteenth century.
Although the noble gases were at first
thought to be unreactive, compounds of
xenon, krypton, radon, and argon are now
known.
Chemistry in Action: Discovery of the Noble Gases
Sir William Ramsay
Crystals of xenon tetrafluoride.
The Main-Group Elements
• The physical and chemical properties of the maingroup elements clearly display periodic behavior.
Variations of metallic-nonmetallic character.
Basic-acidic behavior of the oxides.
Properties of Oxides Across a Period
basic
acidic
Transition Elements and
Coordination Compounds
• The transition elements are defined as those
metallic elements that have an incompletely
filled d subshell or easily give rise to ions with
incompletely filled d subshells.
Sometimes chemists include the two rows
of elements at the bottom of the periodic
table; the lanthanides and the actinides.
Classification of the transition elements.
Properties of the Transition
Elements
• In addition to their commercial uses, many
transition elements have biological
importance.
Transition Elements and
Coordination Compounds
• The transition elements are defined as those
metallic elements that have an incompletely
filled d subshell or easily give rise to ions with
incompletely filled d subshells.
They have a number of characteristics,
including high melting points and a
multiplicity of oxidation states.
Transition Elements and
Coordination Compounds
• The transition elements are defined as those
metallic elements that have an incompletely
filled d subshell or easily give rise to ions with
incompletely filled d subshells.
Compounds of transition elements are
frequently colored and many are
paramagnetic.
These properties are due to the
participation of d orbitals in bonding.
Periodic Trends in the
Transition Elements
• In this section we will examine trends in
melting point, boiling point, covalent
radii, and ionization energies.
Recall that the strength of the metallic
bond is roughly proportional to the number
of unpaired electrons a metal has.
Melting Points and Boiling
Points
• As you read across a row of transition
elements, melting points and boiling
points increase, reaching a maximum at the
group VB or VIB elements, after which, the
melting point decreases.
These properties depend on the strength of
metal bonding, which in turn depends
roughly on the number of unpaired
electrons.
Atomic Radii
• Atomic size decreases across a row in the
transition elements, similar to the trend seen
in the main-group elements.
It is due to an increase in effective
nuclear charge that acts on the outer
electrons.
Ionization Energies
• The first ionization energies of the fourth
period transition elements tend to increase
from left to right.
As with the main-group elements, this
trend seems to mirror a decrease in atomic
radius.
Oxidation States
• Most of the transition elements have a
doubly filled s subshell making a +2
oxidation state relatively common.
In addition, d electrons can be lost,
producing many polyvalent transition
metal ions.
The maximum oxidation state equals the
number of ns and (n-1)d electrons the
metal has.
The Chemistry of Two
Transition Elements
• In the following section we will describe the
chemical properties of two transition
elements: Cr and Cu.
Both elements are well known and
commercially important; in addition, they
provide examples of colorful compounds.
Chromium
• Chromium
Chromium metal reacts with acids to give
Cr2+ ion, which is readily oxidized to Cr3+.
The +3 state is the most common oxidation
state; chromium(III) oxide is a green
pigment.
The +6 oxidation state is represented by
chromates and dichromates.
Aqueous chromium ion.
Chromium
• Chromium
The dichromate ion, Cr2O72-, in acid
solution is a strong oxidizing agent.
Chromium metal is obtained by reduction
of the ore chromite, FeCr2O4.
Chromate-dichromate equilibrium.
Copper
• Copper
Copper metals reacts only with acids
having strongly oxidizing anion, such as
HNO3; it give Cu2+ ion.
Copper(I) oxide, Cu2O, occurs naturally as a
copper mineral.
Copper
• Copper
Cu2O also forms as a brick-red precipitate
when copper(II) ion is reduced in basic
solution. The reaction is used as a test for
glucose.
Benedict’s test for glucose.
Pg 346
WORKED
EXAMPLES
Worked Example 8.1
Worked Example 8.2
Worked Example 8.4
Worked Example 8.5
Worked Example 8.6