Transcript PPT - George Mason University

George Mason University
General Chemistry 212
Chapter 14
Main Group Element Patterns
Acknowledgements
Course Text: Chemistry: the Molecular Nature of Matter and
Change, 7th edition, 2011, McGraw-Hill
Martin S. Silberberg & Patricia Amateis
The Chemistry 211/212 General Chemistry courses taught at George
Mason are intended for those students enrolled in a science
/engineering oriented curricula, with particular emphasis on
chemistry, biochemistry, and biology The material on these slides is
taken primarily from the course text but the instructor has modified,
condensed, or otherwise reorganized selected material.
Additional material from other sources may also be included.
Interpretation of course material to clarify concepts and solutions to
problems is the sole responsibility of this instructor.
4/5/2017
1
Main-Group Elements

Chapter Overview

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Application of bonding, structure, and reactivity to
Main-Group Elements

Hydrogen

Period 2 Elements – Trends across Periodic Table

Group 1A – The Alkali Metals

Group 2A – The Alkaline Earth Elements

Group 3A – The Boron Family

Group 4A – The Carbon Family

Group 5A – The Nitrogen Family

Group 6A – The Oxygen Family

Group 7A – The Halogens

Group 8A – The Noble Gases
2
Main-Group Elements

In chemistry and atomic physics the periodic table divides
the elements into 4 groups
 Main Group Elements
 s-block: 1 (IA);
2 (IIA)
 p-block: 13 (IIIA); 14 (IVA); 15 (VA)
16 (VIA); 17 (VIIA); 18 (VIIIA)
 Transition (d-block) Elements
3
4
5
6
7
8
9
10
11 12
(IIIB IVB VB VIB VIIB VIIIB VIIIB VIIIB IB IIB)
 Lanthanides (f-block Elements)
 Elements in Period 6, group 3 (IIIB) whose
f-subshells are being filled
 Actinides
(f-block Elements)
 Elements in Period 7, group 3 (IIIB) whose
f-subshellls are being filled)
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3
Main Group Elements
Group→ 1
↓ Period IA
2
IIA
3
IIIB
1
4
IVB
5
VB
6
VIB
7
8
9
10
11
VIIB (VIIIB VIIIB VIIIB) IB
Main Group Elements
12
2B
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
18
VIIIA
p block
2
d-block
(Transition Metals)
3
4
5
6

7

s block


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 f-block - Lanthanoid (ide) series)
  f-block - Actinoid (ide) series)
4
Main Group Elements

Periodic Table Numbering System
 Old Systems
 Old IUPAC (Used in Europe)
 Used Roman Numerals I, II, III, IV, V, VI, VII,
VIII) and Letters (A &B) to indicate group
(columns)
 The numbers roughly indicated the highest
oxidation state of the elements, thus similar
chemical properties
 The letters A and B were designated to the left (A)
and right (B) part of the table
 CAS System (Used America)
 Similar to Old IUPAC except that the letter “A”
referred to the Main group Elements and the letter
“B” referred to the Transition Elements
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5
Main Group Elemetnts

Periodic Table Numbering System

Old systems confusing


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The use of the letters A & B in the old systems led to
a lot of confusion
New IUPAC System (Universally used)

Numbers the groups increasingly from 1 -18 left to
right on the standard periodic table incorporating the
10 Transition Elements groups

These group numbers correspond to the number
of s, p, and d orbital electrons added since the last
noble gas element (in column 18)
6
Main Group Elements

Main Group Elements

Elements that belong to the "s" and "p" blocks

Counting the columns (groups 1- 8) across the table
(ignoring the transition elements) gives 8 element
groups which match the filling of the eight spaces for
electrons in the ns and np subshells, ns2np6

One good aspect about the 1 to 8 group numbering
system is that the group number indicates the number
of valence (outer) electrons for atoms in the main
group elements
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Main Groupd Elemtns

The lightest Main Group members are represented by
Helium, Lithium, Beryllium, Boron, Carbon, Nitrogen,
Oxygen, And Fluorine

Main group elements (with some of the lighter
transition metals) are the most abundant elements on
the earth, in the solar system, and in the universe
They are sometimes called the representative elements
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Main-Group Elements

Hydrogen (1s1)
 90% of all atoms in Universe are
Hydrogen atoms
 Single Electron; Small Size


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No perfectly suitable position in
the periodic table
Depending on the property,
Hydrogen fits better in 1A, 4A, 7A
 +1 oxidation State (grp 1A?
 Relatively High Ionization Potential (grp 7?)
 Forms diatomic molecule (H2 - grp 7?)
 Shares electrons (grp 4?)
 Half-filled valence shell; ionization energy; electron
affinity, electronegativity, and bond energies most
similar to group 4
9
Hydrogen Chemistry

Hydrogen Bonding
Dipole-Dipole force between Hydrogen (H) and small,
highly electronegative atoms with lone electron pair:
Nitrogen (N); Oxygen (O); Fluorine (F)

Highly reactive, combining with nearly every element
Ionic (salt like) hydrides
 Group 1A & 2A metals
2Li(s) + H2(g)  2LiH(s) Lithium Hydride
Ca(s) + H2(g)  CaH2(s) Calcium Hydride



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In H2O, H- is a strong base that pulls H+ from water
Na+H-(s) + H2O  Na+(aq) + OH-(aq) + H2(g)
Hydride ion is also a strong reducing agent
Ti4+Cl4(l) + 4LiH(s) = Tio(s) + 4LiCl(s) + 2H2(g)
10
Covalent Hydrides

Hydrogen reacts with nonmetals to form covalent
hydrides
CH4

NH3
H 2O
HF
Conditions for forming Covalent Hydrides depend on
the reactivity of the nonmetal – the more stable, the
more temperature & pressure required for formation
Ex: Ammonia – 400oC & 250 atm
Catalyst
N2(g) + 3H2(g)  2NH3(g)
Horxn = -91.8 kJ
At low temperatures (-196oC) Hydrogen combines
readily with reactive Fluorine (F2)
F2(g) + H2(g)  2HF
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Horxn = -546 kJ
11
Metallic (Interstitial) Hydrides
Many Transition Elements form metallic
(interstitial) hydrides, where Hydrogen molecules
(H2) and Hydrogen atoms (H) occupy the holes in
the metal’s crystal structure.
 These are not compounds, but rather gas-solid
solutions
 They lack a Stoichiometric formula because metal
can incorporate a variable amount of hydrogen,
depending upon temperature and pressure

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Trends Across Periodic Table
 Electrons fill 1 ns – 3 np
orbitals according to Pauli
Exclusion Principle and
Hund’s Rule
 d orbitals in lower Periods can
be used to accommodate
additional oxidation states
 Atomic size generally
decreases
 1st ionization potential
increases
 Electronegativity increases
 Metallic character decreases
with increasing nuclear
charge
 Reactivity highest at right &
left sides, less in middle
 Bonding - metallic  covalent
 none (noble gas)
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Continued on next Slide
13
Trends Across Periodic Table
 Bonding between each element
and an active nonmetal changes
from ionic to polar covalent
 Bonding between each element
and an active metal changes
from metallic to polar covalent to
ionic
 Acid-Base behavior of common
element oxide in water changes
from basic to amphoteric (acts
as acid or base (H2O) to acidic as
bond between element and
oxygen becomes more covalent
 Reducing strength decreases
through the metals
 Oxidizing strength increases
through the nonmetals
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Group 1A - Alkali Metals (ns1)








Lithium (Li), Sodium (Na); Potassium (K); Rubidium
(Rb); Cesium (Cs); Francium (Fr)
Single electron relatively far from nucleus
weak metallic bonding - attraction between delocalized
electrons and metal-ion cores in crystalline structure
Low melting points, soft consistency
Reactive Metals
Powerful reducing agents – lose 1 electron becoming 1+
cations, donating the electron to other elements
ns1 configuration forms salts readily (+1 cations)
Low Heat of Atomization (oHatom ) – Recall Lattice
Energy)
 Energy to convert solid into individual gaseous atoms
oHatom (Li>Na>K>Rb>Cs)
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Group 1A - Alkali Metals (ns1)

Low Ionization Energy (IE) – Each alkali element
has the largest size and the lowest IE in its Period

Size of atom decreases considerably when
valence electron is lost

Lattice Energy – The atomic radius increases as
you move down a group. Since the square of the
distance is inversely proportional to the force of
attraction, lattice energy decreases as the atomic
radius increases

For a given anion, the Lattice Energy become
smaller as the cation becomes larger
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Group 1A - Alkali Metals (ns1)

Solubility – Despite strong ionic attractions, the
Group 1A salts are water soluble – attraction
between the ions and the polar Water molecule
creates highly Exothermic Heat of Hydration
(Hhydr)

Entropy – Entropy increases as ions disperse
going into solution overcoming the high lattice
energy

Magnitude of Hydration Energy decreases as ionic
size increases
H = -Hhydr
(Li+ > Na+ > K+ > Rb+ > Cs+
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Group 1A - Alkali Metals (ns1)

Anomalous Behavior of Lithium

Lithium ion (Li+) is small and highly positive

Dissociation of Lithium salts, such as LiF,
Li2CO3, LiOH, and Li3PO4, in water is much
more difficult than similar salts of sodium (Na)
and Potassium (K)

Only member of Alkali group that forms simple
Oxide and Nitride, Li2O & Li3N, on reaction with
O2 & N2 in air

Only Lithium forms organo-metalic molecular
compounds with hydrocarbon groups from
organic Halides
2Li(s) + CH3CH2Cl(g)  CH3CH2Li(s) + LiCl(s)
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Group 1A - Alkali Metals (ns1)

Reactions & Compounds of Alkali Metals

Alkali metals reduce Hydrogen in Water to form
Hydrogen gas
2E(s)
+ 2H2O 
2E+ + 2OH-(aq) + H2(g)
Where E = any alkali metal (Li, Na, K, Rb, Cs)
Reaction becomes more vigorous down group

Alkali metals reduce oxygen, but product depends
on the metal
4Li(s) + O2(g)
K(s) + O2(g)

 2Li2O(s) oxide
 KO2(s)
superoxide
Alkali metals reduce Hydrogen to form ionic hydrides
2E(s) + H2(g) 2EH(s)
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Group 1A - Alkali Metals (ns1)

Reactions & Compounds of Alkali Metals
 Alkali metals (E) reduce Halogens (X) to form Halides
2E(s) + X2  2EX(s)
X = F, Cl, Br, I)

Sodium Metal (Na) can be produced from Molten NaCL
and electricity
2NaCl(l)  2Na(l) + Cl2(g)
Sodium Hydroxide (Lye) can be produced from Salt
(NaCl), water (H2O) and electrolysis
2NaCl(s) + H2O(l)  2NaOH(aq) + H2(g) + Cl2(g)


In an ion-exchange process, water can be “softened”
by removal of dissolved hard-water cations to displace
Na+ ions from a “resin”
M2+(aq) + Na2Z(s)  MZ(s) + 2Na+(aq)
(M = Mg, Ca: Z = resin)
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Group 1A - Alkali Metals (ns1)
atomic properties
physical properties
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21
Group 2A - Alkaline Earth Metals (ns2)
Be, Mg, Ca, Sr, Ba, Ra (E2+ ions)
 Oxides (except Be) give basic (alkaline) solutions:
Ca(OH)2, Mg(OH)2
 High melting points (higher lattice energy than 1A)
 Atomic & Ionic sizes
 Smaller radii and higher ionization energy
 Increase in size down the group
 Combination of size, extra electron, and metallic
bonding result in stronger attractions between
delocalized electrons and the atom cores
 Thus, Melting Points and Boiling Points are much
higher than 1A alkali metals
 Harder & more dense than Alkali metals, but soft
and lightweight compared to transition metals (Fe,
Cr, etc)
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
22
Group 2A - Alkaline Earth Metals (ns2)

Even though the Alkaline Earth metals have higher
ionization potential, they still form ionic compounds (E2+),
but Beryllium (Be) is an exception forming covalent bonds

Like Alkali metals, Alkaline Earth metals are strong reducing
agents

Group 2A (Alkaline Earth) elements are reactive because
the higher lattice energy of their compounds more than
compensates for the large total Ionization Energy (IE)
needed to form the 2+ cations

The higher Lattice Energy (from the smaller cation size)
and higher Charge Density results in lower solubility

Ion-Dipole attraction is so strong that many slightly soluble
2A salts crystallize as “Hydrates”
Epsom salt – MgSO47H2O
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Gypsum – CaSO42H2O
23
Group 2A - Alkaline Earth Metals (ns2)

The anomalous behavior of Beryllium

Beryllium has smallest size; highest Ionization energy,
and highest Electronegativity of the Alkaline Earth
elements

Combined with the high charge density of the ion
(Be2+) it polarizes the nearby electron clouds very
strongly and causes extensive orbital overlap; this
results in covalent bonding

BeF2 is the most ionic of the Beryllium compounds, but
its melting point and electrical conductivity are
relatively low compared to the other alkaline earth
Fluorides

Unlike the other Alkaline Earth Metals, whose oxides
are basic, BeO is amphoteric and does not react with
water to form OH- ions
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Group 2A - Alkaline Earth Metals (ns2)

Diagonal relationships: Lithium and Magnesium
Certain Period 2 elements exhibit behaviors that are very
similar to those of the Period 3 elements immediately below
and to the right
3 relationships
1. Li, Mg
2. Be, Al
3. B, Si
 Lithium and Magnesium reflect similar atomic and ionic size
 Both elements form:
 Nitrides,
 Hydroxides and Carbonates (CO3) that decompose with
heat,
 Organic compounds with polar covalent metal-carbon
bonds
 Salts with similar solubilities
25
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
Group 2A - Alkaline Earth Metals (ns2)

Reactions & Compounds (E = Mg, Ca, Sr, Ba)

Metals reduce Oxygen (O2) to form Oxides
2E(s) + O2(g)  2EO(s)
Ba + O2  BaO2 (Barium Peroxide)

Larger metals reduce water to form hydrogen
gas
E(s) + 2H2O(l)  E2+aq) + 2OH- (aq) + H2(g)

Metals reduce Halogens to form ionic halides
E(s) + X2  EX2(s)

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(X = F, Cl, Br, I)
Most metals (Be exception) reduce Hydrogen to
form ionic hydrides
E(s) + H2(g)  EH2 (s)
(except Be)
26
Group 2A - Alkaline Earth Metals (ns2)

Reactions & Compounds (E = Ca, Mg, Sr, Ba)
 Most elements reduce Nitrogen to form ionic Nitrides
3E(s) + N2(g)  E3N2(s)
(except Be)
 Element Oxides are Basic (except for amphoteric BeO)
EO(s) + H2O(l)  E2+(aq) + 2OH-(aq)
 All Carbonates undergo thermal decomposition to the
oxide
heat
ECO3(s)  EO(s) + CO2(g)
(CaO – Lime)
 Beryl (Be3Al2Si6O18) - Gemstone, source of Be
 Magnesium oxide (MgO) – Refractory material for
furnace bricks
 Alkyl Magnesium Halides – RMgX
(R=Hydrocarbon)
Grignard Reagents – organic compound synthesis
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Group 2A - Alkaline Earth Metals (ns2)
atomic properties
physical properties
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28
Group 3A – Boron Family (ns2np1)
B
Al
Ga
In
Tl

Boron heads family, but other elements in group
3A exhibit diverse properties

Boron & Aluminum, especially Aluminum, are
much more abundant than the others, but still
quite rare

Group 3A elements include “p” orbitals for first
time

In Period 4 (transition elements) the “d” orbitals
are present

Physical Properties are influenced by type of
bonding
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Group 3A – Boron Family (ns2np1)

Boron is a network covalent metalloid - Black,
hard, very high melting point

A network solid or covalent network solid is
a chemical compound in which the atoms are
bonded by covalent bonds in a continuous
network

In a network solid there are no individual
molecules and the entire crystal may be
considered a macromolecule

Boron (metalloid) is much less reactive than the
others members of the 3A group because it forms
covalent bonds
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Group 3A – Boron Family (ns2np1)






Other group members are metals – shiny, relatively
soft with low melting points
Aluminum is more ionic; its low density and 3 valence
electrons make it a good electrical conductor
Although Aluminum is a metal, its halides exist in the
gaseous state as covalent dimers - AL2Cl6 (contrast
salts of group 1 & 2 metals)
Aluminum Oxide, Al2O3, is amphoteric (can act as an
acid or base) rather than basic like the Group 1A &
2A metals
Although the other Group 3A elements are basically
ionic they exhibit more Covalent character than
similar 2A compounds.
3A cations are smaller with more charge density than
2A cations and they polarize an anion’s electron cloud
more effectively
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Group 3A – Boron Family (ns2np1)

Oxidation-Reduction (REDOX) behavior in Group 3A

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Presence of Multiple Oxidation States

In Groups 3A – 6A many of the larger elements
(down the group) exhibit an oxidation state “two
lower” than the A-Group number

This lower state occurs when the atoms lose their
np electrons, not the ns electrons.

The lower oxidation state is the result of lower
bond energies

Bond energies decrease as the size of the atom
and the bond length increase for elements lower
in the group
32
Group 3A – Boron Family (ns2np1)

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Increasing prominence of the low oxidation
state

When a group exhibits more than one
oxidation state, the lower state becomes more
prominent going down the Group

All members of the 3A group exhibit the +3
state, but the +1 state appears first with
some compounds of Gallium (Period 4)

The +1 state becomes the most important
state of Thallium (Period 6)
33
Group 3A – Boron Family (ns2np1)

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Relative Basicity of Group 3 oxides
 Recall: A1 oxides (ionic charge +1 and more
metallic) are more basic than A2 oxides (ionic
charge +2 and less metallic)
 In general, oxides with the element in a lower
oxidation state (less positive) are more basic
than oxides with the element in a higher
oxidation state
 For Indium oxides in Group A3, In+12O acts more
like a metal and is more basic than In+32O3
 The lower charge density of In+1 does not
polarize the O-2 ion as much as the In+3 ion
 Thus, in E2O compounds, the E-O bonding is
more ionic than in E2O3 compounds, thus; the O-2
ion is more available to act as a base – donate
electron pair or accept a proton
34
Group 3A – Boron Family (ns2np1)

Boron Chemistry
 Boron compounds are covalent (unique within
group)
 Forms network covalent compounds or large
molecules with metals, H, O, N
 Electron deficient; uses two approaches to
complete octet
 Accepting a Bonding Pair from Electron-Rich
atom
BF3(g) + NH3(g)  F3B-NH3(g)
(BF3 acts as acts as Lewis acid in accepting the
electron pair from the Nitrogen in NH3)
B(OH)3 + H2O(l)  B(OH)4-(aq) + H+(aq)
(Acts as acid by accepting electron pair from H2O)
Note: Water is acting as the base
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Group 3A – Boron Family (ns2np1)

Boron Chemistry
 Two approaches to filling octet (con’t)
 Accepting electron pair from Electron-Rich atom
(con’t)
 Boron-Nitrogen compounds are similar in
structure to elemental Carbon and some of its
organic compounds
 Size, Ionization Energy, Electronegativity of
Carbon is between Boron & Nitrogen
 Ethane & Amine – Borane have the same
number & electron configuration
 
 C – C
 
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 
B – N
 
36
Group 3A – Boron Family (ns2np1)

Boron Chemistry

Two approaches to filling octet (con’t)

Forming Bridge Bonds with Electron-Poor Atoms

Boron Hydrides - Boranes

2 types of B – H bonds
 Normal electron-pair bond
o sp3 orbital of B overlaps 1s orbital of H in
each of the four terminal B-H bonds
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Group 3A – Boron Family (ns2np1)
 Hydride Bridge Bond (3-center, 2 electron bond)
o Each B – H – B grouping is held together by
only two electrons
o Two sp3 orbitals, one from each B, overlap an
H 1s orbital between them
o Two electrons move through this extended
bonding orbital – one from one of the B atoms
and the other form the H atom – and join the
2 B atoms via the H atom bridge
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38
Group 3A – Boron Family (ns2np1)
atomic properties
1/19/2015
physical properties
39
Group 3A – Boron Family (ns2np1)

Reactions & Compounds
Elements react sluggishly, if at all, with water (H2O)
2 Ga(s) + 6H2O(hot)  2Ga3+(aq) 6OH-(aq) + 3H2(g)

2Tl(s) + 2H2O(steam)  2Tl+(aq) +2OH-(ag) + H2(g)
Note different oxidation numbers for Ga3+ & Tl+

All members form oxides when heated in pure O2
4E(s) + 3O2(g)  2E2O3(s)
(E = B, Al, Ga, In)
4Tl(s) + O2  2Tl2O3(s)

Oxide acidity decreases down the group:
B2O3 > Al2O3 > Ga2O3 > In2O3 > TlO2
(weakly acidic)
(strongly basic)
The +1 oxide (TlO2) is more basic than the +3 oxide
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40
Group 3A – Boron Family (ns2np1)

Reactions & Compounds

All members reduce Halogens
2E(s) + 3X2

2EX3
(E = B, Al, Ga, In)
2Tl(s) + X2  2TlX(s)


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Trihalides of AL, Ga, In are mostly ionic but
exist as dimers in the gas phase
Acid (H2SO4) treatment of Al2O3 produces Al2SO4, a
colloid (coagulant) used in water purification
Al2O3 + 3H2SO4  Al2SO4(s) + 3H2O(l)
41
Group 4A – Carbon Family (ns2np2)






The whole range of elemental behavior occurs
within the 4A group
 Non metalic Carbon (C)
 Metalloids (Silicon (Si) & Germanium (Ge)
 Metallic (Tin (Sn) & Lead (Pb)
 Newly synthesized element at bottom of group
Carbon forms the basis of “Organic Chemistry”
 20,000,000 compounds
Polymer Chemistry
Biochemistry based on Carbon
Geochemistry
Electronic technologies bases on Si
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42
Group 4A – Carbon Family (ns2np2)

1/19/2015
Bonding effects on Physical Properties

Silicon has a much lower melting point than Carbon
because of the longer, weaker bonds.

The melting point difference between Germanium
(Ge) and Tin (sn) is due to the change from network
covalent to metallic

Going from Group 3 to group 4 there are large
increases in melting point and the Hfus because of
the change from metallic to network covalent bonding
43
Group 4A – Carbon Family (ns2np2)

Allotropism: Different Forms of an Element
 Elemental Carbon – Graphite & Diamond
 Different crystalline & molecular forms with different
physical properties

Carbon Allotropes
 Graphite – Black, “greasy”, soft, more stable than
diamond



1/19/2015
Diamond – Colorless, electrical insulator,
extremely hard
Bucky-Balls (Buckminsterfullerene) – soccer ballshaped with the formula C60
Tin Allotropes
 -tin – stable at room temperature & above
 -tin – stable below 13oC
44
Group 4A – Carbon Family (ns2np2)


Bonding Changes in Group 4A

Carbon – Covalent (intermediate EN)

Si & Ge – strong polar bonds (silicate minerals)

Tin (Sb) & Lead(Pb) – Metallic (Ionic)
Multiple Oxidation States

Carbon (+4)

Silicon (+4 more stable than +2)

Lead (+2 more stable than +4)

Elements with lower oxidation states act more like
metals (more basic)
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45
Group 4A – Carbon Family (ns2np2)

Highlights of Carbon Chemistry

Carbon, like other elements in “Period 2, is the
anomalous element in the group

Carbon forms bonds with:
1/19/2015

Smaller Group 1A & 2A metals

Many transition metals

Halogens

Neighbors B, Si, N, O, P, S

Exhibits all possible oxidation states from +4
in CO2, and Halides to -4 in CH4
46
Group 4A – Carbon Family (ns2np2)

Highlights of Carbon Chemistry

Two main features of Carbon Chemistry

Catenation and the ability of Carbon to form
multiple bonds
Carbon can form chains, branches, and
rings (aromatic & aliphatic)
 Multiple bonds – sigma (), Pi (), Triple ()


1/19/2015
The C-C bond is short enough for side-toside overlap of two half-filled 2p orbitals to
form  bonds that give rise to many diverse
structures and reactivities of organic
compounds
47
Group 4A – Carbon Family (ns2np2)

The Other 4A elements

E-E bonds become longer going down the group,
with decreasing bond strength
C–C
> Si – Si > Ge – Ge

The empty d shell orbitals make these chains
susceptible to chemical attack – they are
reactive

The long bonds are not suitable for overlap of p
orbitals; thus, no  bonds
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48
Group 4A – Carbon Family (ns2np2)

Carbonates

Metal Carbonates are the main mineral
Marble, Limestone, Chalk, Coral, others

Antacids – Calcium Carbonate & Stomach Acid
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O(l)


Limestone (CaCO3) deposits help moderate the
effects of acid rain (H2SO4 & HNO3)
Carbon Dioxide (CO2)

Essential to all life as primary source of carbon in
plants & animals through photosynthsis

Atmospheric buildup from motor vehicles and fossil
fuel powerplant severely affect global climate
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49
Group 4A – Carbon Family (ns2np2)

Silicon Chemistry

Silicon Halides are more reactive than Carbon Halides
because Si (3s, 3p, 3d orbitals) has empty 3d orbitals
available for bond formation

The Si – X bond is long but stronger than
corresponding C – X bond

Si – X bond has some double bond character because
of the presence of a  bond and a different type of 
bond called a p,d- (side-to-side overlap of the Si d
orbital
and a Halogen
p p,d-
orbital
Trimethylamine
The impact of
bonding on the structure of trisilylamine
(CH3)3N
1/19/2015
trigonal
pyramidal
Trisilylamine
(SiH3)3N
trigonal planar
50
Group 4A – Carbon Family (ns2np2)

Silicon chemistry is dominated by the SiliconOxygen – (Si–O) bond

C – C bonds can repeat endlessly; similarly, the
Si–O bonds can also repeat forming large chains
in Silicate minerals in the earths crust and
Silicones, which are synthetic polymers with a
large number of industrial applications

Silicate Minerals

From common sand (SiO2) and clay to
semiprecious amethyst, Silicate minerals are
the dominant form of matter on the earth

Oxygen is the most common element on earth
and Silicon is the next most abundant
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51
Group 4A – Carbon Family (ns2np2)


The Orthosilicate (–SiO4 –) grouping is the building
unit for Silicate minerals
 Zircon ZrSiO4
1 Unit
 Hemimorphite
2 Units [Zn4(OH2Si2O7H2O]
 Beryl
6 Units [Be3Al2Si6O18]
Silicon Polymers
 Manufactured Substances
 Alternating Si & O atoms with two Organic groups
bonded to each Silicon atom
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52
Group 4A – Carbon Family (ns2np2)
atomic properties
physical properties
1/19/2015
53
Group 4A – Carbon Family (ns2np2)

Important Reactions

Group 4A elements are Oxidized by Halogens
E(s) + 2X2  EX4
(E = C, Si, Ge)
The +2 Halides are more stable for Tin (Sb) & Lead (Pb)
SnX2
PbX2
The Elements are oxidized by Oxygen (O2)
E(s) + O2(g)  EO2
(E = C, Si, Ge, Sn)
The oxides are more basic (metallic) going down the group
Lead (Pb) forms the +2 oxide (PbO) - basic
In natural streams, Carbon Dioxide (CO2) forms a weakly
“acidic” solution
CO + H O ⇄ H CO (aq) ⇄ H+(aq) + HCO -(aq)
2
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2
2
3
3
54
Group 4A – Carbon Family (ns2np2)

Air & Steam passed over hot coke (carbon) produce
gaseous fuel mixtures – producer gas & water gas
C(s) + H2O(g) + air(g)  CO(g) + CO2(g) + N2(g) + H2(g)
Note: This industrial reaction cannot be balanced

Hydrocarbons (C & H only) react with Oxygen (O2) to
form CO2 & Water (H2O), a source of heat to yield
steam (H2O) for electrical generation
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + Heat

Certain metal Carbides react with water to produce
Acetylene (H-CC-H), used in oxyacetylene torches
CaC2(s) + 2H2O(g)  Ca(OH)2(aq) + C2H2(g)
Acetylene is source material for organic compound
synthesis and a fuel for Welding
1/19/2015
55
Group 4A – Carbon Family (ns2np2)

Freon (chlorofluorocarbon) is formed from fluorinating
Carbon Tetrachloride
CCl4(l) + HF(g)  CFCl3(g) + HCL(g)
Production of Trichlorofluoromethane (Freon-11) is being
discontinued because it is an atmospheric pollutant

Silica (SiO2)is reduced to form elemental Silicon used in
the manufacture of computer chips
SiO2(s) + 2C(s)  Si(s) + CO2(g)
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56
Group 5A – Nitrogen Family (ns2np3)






Widest range of physical behavior in 1st 5 groups
Compounds of Nitrogen (gaseous nonmetal) and
Phosphorus (solid nonmetal) are important in
industrial and environmental processes
Arsenic (As) and Antimony (Sb) are network covalent
metalloids with highest melting points in group
Bismuth (Bi) exhibits metallic bonding
Nitrogen (N) exists as diatomic molecules, which
interact through very weak dispersion force
producing a boiling point 200 oC below room
temperature
Phosphorus (P) is heavier and more polarizable than
Nitrogen with stronger dispersion forces – higher
melting point 44oC
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57
Group 5A – Nitrogen Family (ns2np3)
atomic properties
physical properties
1/19/2015
58
Group 5A – Nitrogen Family (ns2np3)

Chemical Behavior in Group 5A Patterns
 Nitrogen forms a maximum of 4 covalent bonds
 The other elements in the group can expand the
valence shell by using empty ‘d’ orbitals

The noble gas configuration is attained by group 5
elements gaining 3 electrons – the first is Exothermic
and the last two are Endothermic (requiring input of
energy from surroundings
As in groups 3A & 4A, fewer oxidation states occur
moving down the group with the lower oxidation state
becoming prominent
 Oxidation states



1/19/2015

Nitrogen
P, As, Sb
Bi
– from +5 to -3
– +5 & +3
– +3
59
Group 5A – Nitrogen Family (ns2np3)

Nitrogen Chemistry
 Nitrogen Oxides – 6 stable forms
1/19/2015
60
Group 5A – Nitrogen Family (ns2np3)

Nitrogen Oxoacids & Oxoanions
Oxoacid
Oxoanion
1/19/2015
61
Group 5A – Nitrogen Family (ns2np3)

Nitric Acid (oxidizing agent) Reactions

Active Metal (Al in dilute HNO3 solution)
8Al(s) + 30 HNO3(aq)  8Al(NO3)3 + 3NH4NO3 + 9H2O(l)
Al(s) + 30H+(aq) + 3(N+5O3)-(aq)  8Al+3(aq) + 3(N+3H4)+(aq) +
9H2O(l)
(net ionic equation)

Less reactive metal (Cu), more conc HNO3
(N2+ forms)
3Cu(s) + 8HNO3(aq)  3Cu(NO3)2 + 4H2O(l) + 2N2+O(g)
3 Cu(s) + 8H+(aq) + 2(N+5O3)-  3Cu2+(aq) + 4H2O(l) +
2N+2O(g)
(net ionic equation)

Copper with still more concentrated HNO3
(N4+ forms)
Cu(s) + 4HN+5O3(aq)  Cu+2(NO3)2(aq) + 2H2O(l) + 2N4+O2(g)
Cu(s) + 4H+(aq) + 2NO3-(aq)  Cu+2(aq) + 2H2O(l) + 2NO2(g)
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(net ionic equation)
62
Group 5A – Nitrogen Family (ns2np3)

Nitrates form when Nitric Acid (HNO3) reacts with the
hydroxides, oxides, or carbonates of metals
HNO3 + NaOH  NaNO3 + H2O

Nitrous Acid (HNO2), a much weaker acid, is formed
from the reaction of a strong acid (HCl) and metal
Nitrites
NaNO2(aq) + HCl(aq)  HNO2(aq) + NaCl(aq)

Strong acid vs Weak acid – The more oxygen atoms
bonded to the central nonmetal, the stronger the acid
Oxygen atom pulls electron density from the Nitrogen
atom, which in turn pulls electron density from the
Oxygen of the O-H bond, facilitating the release of
the H+ ion – The more Protons (H+) in solution, the
stronger the acid.

1/19/2015
63
Group 5A – Nitrogen Family (ns2np3)

Phosphorus Chemistry
 Phosphorus forms two important Oxides
 Tetraphosphorus Hexaoxide, P4O6
 P (+3) has tetrahedral orientation with an
Oxygen between pair of P atoms
 Reacts with water to form Phosphorus acid,
H3PO3
P4(s) + 3O2(g)  P4O6(s) + 6H2O(l)  4H3PO3(l)
 Only two of the H atoms are acidic, the third is
bonded to the central P and does not dissociate
 Dissociation is complete in strong base solution
1/19/2015
64
Group 5A – Nitrogen Family (ns2np3)

Tetraphosphorus Decaoxide (con’t)

Phosphorus (+5) oxidation state
P4(s) + excess 5O2  P4O10(s)
P4O10(s) + 6H2O  4H3PO4(l)

1/19/2015
Phosphoric Acid (H3PO4) is a weak triprotic
acid

In water it loses one proton to form H2PO4-

In excess strong base all three protons
dissociate to form the Phosphate ion, PO43+ 3H+
65
Group 5A – Nitrogen Family (ns2np3)

Diphosphate & Polyphosphates
 Polyphosphates are formed by heating
Hydrogen Phosphates (ex. Na2HPO4)

2Na2HPO4(s)  Na4P2O7(s) + H2O(g)
 The Diphosphate ion, P2O74-, is the smallest of
the polyphosphates consisting of tetrahedral
PO4 units linked through a common Oxygen
1/19/2015
66
Group 5A – Nitrogen Family (ns2np3)

Reactions & Compounds
 Converting Nitrogen to other forms (fixing) is quite
difficult because of the strength of the triple bond
(NN) between the Nitrogen atoms. It can be fixed
industrially by the Haber process
N2(g) + 3H2(g) ⇄ 2NH3(g)

Non Nitrogen Hydrides from metal Phosphides
Ca3P2(s) + 6H2O(l)  2PH3(g) + 3Ca(OH)2

Halides formed by direct combination of elements
2E(s) + 3X2  2EX3 (E= P, As, Sb not N)
EX3 + X2  EX5 (all except N & Bi with X = F & Cl)

P4 in basic solution increases & decreases oxidation
number
P4(s) + 3OH-(aq) + 3H2O  P3+H3(g) + 3H2P1+O2-
1/19/2015
67
Group 6 – Oxygen Family (ns2np4)

The Oxygen Family

First two members of group – gaseous
nonmetallic oxygen (O) & solid nonmetallic
sulfur (S) are among most important elements
in industry, the environment and living
organisms

Selenium (Se) & Tellurium (Te) are metalloids

Polonium (Po) is radioactive and only metal in
the group
1/19/2015
68
Group 6 – Oxygen Family (ns2np4)
atomic properties
1/19/2015
physical properties
69
Group 6 – Oxygen Family (ns2np4)

Oxygen Family vs Nitrogen Family

Groups 5A & 6A have similar Physical & Chemical
Properties

Oxygen & Nitrogen are low-boiling Diatomic gases

Phosphorus (P) & Sulfur (S) occur as polyatomic
molecules – P4 & S8

Arsenic (Ar) & Selenium (Se) occur as gray metalloids

Antimony (Sb) & Tellurium more more metallic than
preceding group members, but display network
convalent bonding

Bismuth & Polonium are metallic crystals

Electrical conductivity increases down group as bonding
changes from individual molecules (insulators) to
metalloid networks to metallic solids (conductors)
1/19/2015
70
Group 6 – Oxygen Family (ns2np4)

Oxygen Family vs Nitrogen Family
 Allotropism (two or more crystalline or molecular forms
of an element) is more common in Group 6A than in
Group 5A




1/19/2015
Oxygen has 2 allotropes
Dioxygen O2 & Ozone O3
Oxygen (O2) gas is colorless, odorless, paramagnetic
(unpaired electrons attracted by outside magnetic
field), and thermally stable
Ozone (O3) gas is bluish, has pungent odor, is
diamagnetic (paired electrons not affected by
external magnetic field), decomposes in heat and
Ultraviolet light.
Ozone in upper atmosphere protects living organisms
from Ultraviolet radiation
71
Group 6 – Oxygen Family (ns2np4)

Oxygen Family vs Nitrogen Family

1/19/2015
Sulfur Allotropes

10 forms

Sulfur bonds to other Sulfur atoms creating rings
and chains

The most stable form is orthorhombic, S8,
a crown-shaped ring of 8 atoms
Top View
side View
72
Group 6 – Oxygen Family (ns2np4)

Oxygen Chemistry vs Nitrogen Chemistry

Oxygen & Sulfur occur as anions more often than
Nitrogen & Phosphorus

Oxygen & Sulfur bond covalently with almost all
nonmetals

Selenium & Tellurium do some covalent bonding,
whereas Polonium behaves like a metal

Oxygen has few oxidation states (O2- most
common)

The other elements in the family exhibit +6. +4,
-2 oxidation states, with the +4 state most
common in Tellurium and Polonium
1/19/2015
73
Group 6 – Oxygen Family (ns2np4)

Oxygen Chemistry vs Nitrogen Chemistry

Oxidizing strength of Oxygen is 2nd only to
Fluorine

The other members of the group behave very
little like oxygen being less electronegative and
forming anions less often

Except for Oxygen, all elements of group 6A
form foul smelling, poisonous, gaseous
hydrides (H2E) upon treatment of the metal
Sulfide, Selenide, etc., with an acid
FeSe(s) + HCl (aq)  H2Se(g) + FeCl2(aq)
1/19/2015
74
Group 6 – Oxygen Family (ns2np4)

Oxygen Chemistry vs Nitrogen Chemistry
 Bonding & Thermal stability of Group 6A elements
have several features in common
 Only Water (H2O) forms Hydrogen bonds, so it
melts & boils at higher temperatures than the
group 6A H2E compounds (E = S, Se, Te, Po)
 Bond angles drop from the nearly tetrahedral for
H2O (104.5o) to around 90o for the Group 6A
element hydrides (unhybridized p orbitals)
 E-H bond length increases (bond energy
decreases) down group
 Thus, H2Te is stable above 0oC, but H2Po is only
stable at extremely cold temperatures; it even
decomposes from the heat generated by the
radioactivity of the Polonium.
1/19/2015
75
Group 6 – Oxygen Family (ns2np4)

1/19/2015
Bonding & Thermal Stability (con’t)
 Except for Oxygen, Group 6A elements form a
wide range of Halides
 The Halide structure and reactivity patterns
depend on the sizes of the central atom and the
surrounding Halogens
Radius of S < Se < Te < Po
 Sulfur – Many Fluorides, Few Chlorides, one
Bromide
 As the central atom becomes larger, the
Halides become more stable
 Tetrachlorides & Tetrabromides of Se, Te,
Po are known
 Tetraiodides of Te & Po are known
 Hexafluorides of S, Se, Te are known
76
Group 6 – Oxygen Family (ns2np4)

1/19/2015
Halide Structure (con’t)

Inverse relationship between bond length
& bond strength does not explain pattern
of Group 6 Halide formation

Crowding of lone electron pairs and
Halogen (X) atoms around the central
atom

With S (small central atom) the larger
Halides further down group 7 would be too
crowded, which explains why Sulfur
Iodides don’t occur
77
Group 6 – Oxygen Family (ns2np4)

1/19/2015
Halide activity (con’t)

Sulfur Tetrafluoride vs Sulfur Hexafluoride

SF4 has a lone pair of unshared electrons
and empty d orbitals which can be
involved in bonding – Highly reactive

SF6 uses all of the bonds allowable for S
and is tightly packed – chemically inert
78
Group 6 – Oxygen Family (ns2np4)

Highlights of Oxygen Chemistry
 Most abundant element of the Earth’s surface
 Many oxides – water, silicates, carbonates, phosphates
 Most free Oxygen (O2) has biological origin from
photosynthesis in algae and multicellular plants
 Although much more complicated, the basic reaction
between oxygen, CO2 and light to form carbohydrates
can be represented as:
nH2O(l) + nCO2(g)  nO2(g) + (CH2O)n
 The reverse process of combustion and respiration
produce CO2
 Every element, except He, Ne, Ar (noble gases) form at
least one oxide
 Some oxides have Endothermic heats of reaction, while
others have Exothermic ones
1/19/2015
79
Group 6 – Oxygen Family (ns2np4)

Highlights of Sulfur Chemistry
 Two important oxides – SO2 & SO3
 Sulfur Dioxide (+4 oxidation state) is a colorless
choking gas that forms when S, H2S, or a metal
sulfide burns in air (oxygen)
2H2S (g) + 3O2(g)  2H2O(g) + 2SO2(g)
FeS2(s) + 11O2(g)  2Fe2O3(s) + 8SO2(g)
Sulfur Dioxide dissolves in water to form Sulfurous
acid (H2SO3) – weak acid - which dissociates into an
equilibrium solution of hydrated SO2, H+ ions, &
Bisulfite (HSO3-) ions
SO2(g) + H2O(l) ⇄ [H2SO3(aq)] ⇄ H+(aq) + HSO3-(aq)


1/19/2015
Neither H2SO3 or H2CO3 (both weak acids) can exist
as isolated molecules – they dissociate immediately
in water
80
Group 6 – Oxygen Family (ns2np4)
The S in the Sulfite ion (SO32-) is in the 4+ state and
can be easily oxidized to 6+ state.
 Thus, Sulfites are good Reducing Agents
 SO3 (Sulfur Trioxide) is produced by oxidizing SO2
V O /K O
SO2(g) + 1/2O2 ⇄ SO3(g)

2




1/19/2015
5
2
Sulfuric Acid (H2SO4) is a strong acid and the most
common industrial chemical
It is prepared from SO3, H2O, & conc H2SO4
SO3(g) + conc H2SO4 + H2O  H2SO4(l)
Like other strong acids, sulfuric acid dissociates
completely in water forming the Bisulate (HS6+O4)ion
Conc Sulfuric acid is an excellent dehydrating agent
The loosely held proton transfer to water in an
exothermic formation of Hydronium ion (H3O+)
81
Group 7 – Halogen Family (ns2np5)





Trends in properties down the group are just the opposite
of those in Group 1A
For Halogens, boiling point, melting point, heats of fusion
& vaporization increase down the group
The reason for the opposite trends is the different type of
bonding:
 Alkali metals exhibit metallic bonding, which decreases
in strength as atoms become larger down group
 Halogens exist as diatomic molecules that interact
through “Dispersion” forces
Halogens are quite reactive reacting with metals and
nonmetals to form ionic and covalent compounds – metal
& nonmetal halide oxides, and oxoacids
The halides must gain a single electron to attain the noble
gas configuration as a negatively charged anion
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82
Group 7 – Halogen Family (ns2np5)
atomic properties
1/19/2015
physical properties
83
Group 7 – Halogen Family (ns2np5)
Halogen redox behavior is based
on electron affinity, ionic charge
density, and electronegativity
Halogen higher in group can
oxidize halide ion lower in group
1/19/2015
84
Group 7 – Halogen Family (ns2np5)

The reactivity of Halogens decreases down group
because of the decrease in electronegativity
Note: Fluorine is the most electronegative element

The F-F bond is the weakest, despite it short length,
because the lone pairs of electrons around the first
Fluorine atom repel those on the other Fluorine atom
weakening the bond

Because of the weak bond, F2 reacts with every
element, except the noble gases, in many cases
explosively
The Halogens display the largest range of
electronegativity, but all are electronegative enough to
behave as nonmetals
85
1/19/2015

Group 7 – Halogen Family (ns2np5)


Halogens act as oxidizing agents (they are reduced,
gaining electrons) in the majority of their reactions
Halogens higher in group oxidize Halide ions lower in the
group – Oxidizing ability of X decreases down group
F2(g) + 2X-(aq)  2F-(aq) + X2 (X = Cl, Br, I)

Reaction of Chlorine with water produces Hypochlorous
acid
Cl2 + 2H2O(l) ⇄ Cl-(aq) + HClO(aq) + H+ + H2O(l)

Chlorination of drinking water (disinfectant)
Household bleach is a dilute 5.25% solution of
Sodium Hypochlorite (NaClO)


Hydrogen Chloride (HCL) is extremely water soluble
forming H+ & Cl- ions in solution (Hydrochloric Acid)
 Hydrochloric Acid occurs in animal stomach fluids and
has many industrial uses
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86
Group 7 – Halogen Family (ns2np5)

Highlights of Halogen Chemistry

The Hydrogen Halides (HX) are formed from the
reaction of metal halides and a concentrated acid
CaF2(s) + H2SO4(l)  CaSO4(s) + 2HF(g)
2NaBr(s) + H3PO4(l)  Na3PO4(s) + 3HBr(g)

HCl is formed as a byproduct in the chlorination of
Hydrocarbons for plastics production
CH2=CH2(g) + Cl2(g)  ClCH2CH2Cl(g) 
Ethylene

1/19/2015
1,2 Dichloroethane
CH2=CHCl(g) + HCl(g)
1-chloroethene
Hydrogen Fluoride, with its short, strong bond
forms a weak acid (Hydrofluoric acid) with water
HF(g) + H2O(l)  H3O+ + F-
87
Group 7 – Halogen Family (ns2np5)
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
The other Hydrogen Halides (Cl ,Br, I) dissociate
completely to form stoichiolmetric amounts of
H3O+ ions – strong acids
HBr(g) + H2O(l)  H3O+(aq) + Br-(aq)

Halogens react Exothermically with one another
to form many “Interhalogen” compounds
 Diatomic Molecules (ClF, BrCl, IF)
The more electronegative atom has the 1charge; the other less electronegative atom has
1+ charge
 The XYn interhalogens (n = 3,4,5) form when
the larger members of the group (X) use “d”
orbitals to expand the valence shell
The central atom in these molecules has the
lower electronegativity and positive charge
88
Group 7 – Halogen Family (ns2np5)

Commercially useful interhalogens include Fluorine
compounds used as powerful fluorinating agents,
reacting with metals, nonmetals, oxides, and even
wood and asbestos
Sb(s) + ClF3(l)  SnF2(s) + ClF(g)
P4(s) + 5ClF3(L)  4PF3(g) + 3ClF(g)

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The Fluoro Interhalogens react very actively
(explosively) with water yielding Hydrogen Fluoride
(HF) and the oxoacid.
There is no oxidation-reduction reaction and the
central atom of the oxoacid retains the same
oxidation state
3H2O(l) + Br5+F5(l)  5HF(g) + HBr5+O3
89
Group 7 – Halogen Family (ns2np5)

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Oddness & Eveness of Oxidation States
 Odd numbered groups exhibit odd-numbered
oxidation states (Na+1, P+5, Cl-1)
 Even numbered groups exhibit even-numbered
oxidation states (Ca+2, C+4, O-2)
 Reason: Almost all stable molecules have “Paired”
electrons either as bonded or lone pairs
 When bonds form or break, two electrons are
involved and the oxidation state changes by “2”
Ex. Consider Interhalogens – XY, XY3, XY5, XY7
With Y in the -1 state, the X atoms must be in the
+1, +3, +5, +7 state, respectively
The X+1 state arises when Y fills its valence shell
The X+7 state arises when X is completely oxidized
(all electrons shifted to more electronegative Y atom
90
Group 7 – Halogen Family (ns2np5)

Odd Numbered Oxidation States

When I2 reacts with F2, Iodine Fluoride (IF) forms
I2 + F2  2I+1F-1
Each of the two shared electrons in I2 are used to filled
the valence shell of each Fluorine

In IF3, Iodine uses two more valence electrons to
form two more bonds
I+1F- + F2  I+3F3

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If only 1 electron changed, then an unstable loneelectron species containing 2 Fluorines would form
91
Group 7 – Halogen Family (ns2np5)

Even Numbered Oxidation States

An element in an even-numbered group, such as Sulfur in Group
6A(16), shows the same tendency to have paired electrons in its
compounds

Elemental Sulfur (Ox No = 0) gains or shares 2 electrons to
complete its shell

The Sulfur atom loses 2 electrons to react with Fluorine to form:
SF2 (Ox No = + 2)
SF2 forms a bent compound. There are a total of 20 valence
electrons: 6 by Sulfur, 7 from each Fluorine = 20 total. Eight are
placed around the sulfur. Six are placed around each Fluorine. A
Fluorine is placed on two sides of the Sulfur. The two unshared
electron pairs take up more space than the shared pairs and so the
shared pairs move closer together approximately 105 degrees apart.
AX2E2 = Tetrahedral Bent, just like water.
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92
Group 7 – Halogen Family (ns2np5)
Molecular shapes of the main types
of interhalogen compounds
ClF3
linear, XY
Tshaped
XY3
ClF
IF7
BrF5
900
Square Pyramidal
XY5
Pentagonal
Bipyramidal,
XY7
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93
Group 7 – Halogen Family (ns2np5)

Halogen Oxides

Group 7A Halogens form many Oxides that are
powerful oxidizing agents (they are reduced by
gaining the electrons lost by the oxidized species)

The Oxides form Acids with water

Dichlorine Monoxide (Cl2O) & Chlorine Dioxide (ClO2)
are used to bleach paper
2NaClO3(s) + SO2(g) + H2SO4(aq)  2ClO2(g) + 2NaHSO4(aq)
ClO2 has an unpaired electron and
the Chlorine (Cl) in the unusual
+4 oxidation state
(4 electrons are shared with the 2 oxygen
atoms, leaving 3 unshaired)
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chlorine dioxide
ClO2
94
Group 7 – Halogen Family (ns2np5)

Halogen Oxoacids and oxoanions

Oxoacids and Oxoanions are formed by reacting the
Halogens and their Oxides with Water

Most Oxoacids are stable only in solution

There are four Oxoacids & Oxoanions
The known Halogen Oxoacids
Acid
Salt
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Hypochlorous,
Chlorous
Chloric
Perchloric
Sodium Hypochlorite, Sodium Chlorite, Sodium Chlorate, Sodium Perchlorate
95
Group 7 – Halogen Family (ns2np5)

Electronegativity of the Halogen
 Relative strengths of the Halogen Oxoacids depend
on two factors
 Electronegativity of the Halogen
 The more electronegative the halogen, the
more electron density it removes from the
O-H bond, and the more easily the proton is
lost
 Among Oxoacids with the oxidation state of
the Halogen in each Halogen the same, the
Acidity (acid strength) decreases as the
Halogen’s Electronegativity (EN) decreases
Electronegativity –
Cl
>
Br
>
I
Acidity
– HOClO2 > HOBrO2 > HOIO2
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96
Group 7 – Halogen Family (ns2np5)

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Oxidation state of the Halogen
 The oxidation number is a number identical
with the valency but with a sign, expressing
the nature of the charge of the species in
question when formed from the neutral atom
 The oxidation number of Chlorine in Chlorine
Oxoacids
 Hydrochloric Acid (HCl)
- 1
 Hypochlorous acid (HOCl)
+1
 Chlorous Acid (HOCLO or HClO2) + 3
 Chloric acid (HClO3 or HOCLO2)
+5
 Perchloric acid (HClO4 or HOCLO3) + 7
97
Group 7 – Halogen Family (ns2np5)

Oxidation State of the Halogen

Among Oxoacids of a given Halogen, such
as Chlorine, acid strength decreases as the
oxidation state of Halogen decreases

The higher the oxidation state (also stated
as the number of attached O atoms) of the
Halogen, the more electron density it pulls
from the O-H bond
HOCL+7O3 > HOCL+5O2 > HOCl+3O >
Perchloric
Acid
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Chloric
Acid
Chlorous
Acid
HOCl+1
Hypochlorous
Acid
98
Group 8 – Noble Gas Family (ns2np6)








The Noble gases have completed outer s & p shells
Noble gases are generally not reactive – nearly inert
Behave more like “Ideal Gases” than any other gases
The smallest radii in their period
Condense and solidify only at very low temperatures
Helium solidifies (with pressure) at -272.2oC
(absolute zero is -273.15oC) and boils only 3 degrees
higher
A few Noble gas compounds have been prepared
PtF6 + Xe  XePtF6
Other Xenon compounds
Xe+2F2
Xe+4F4
Xe+6F6
Xe+8O4
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Group 8 – Noble Gas Family (ns2np6)
atomic properties
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physical properties
100
Practice Problem

What trends exist for Zeff (Effective Nuclear Charge)
across a period and down a group
Ans: Zeff, the effective nuclear charge
In multi-electron atoms an electron feels the
attraction from the positively charged nucleus
(protons) and the repulsion of like-charged electrons
Electron repulsion shields the electron from the
nuclear attraction making the electron easier to remove
Shielding reduces the “full nuclear charge” to an
effective nuclear charge (Zeff)
Zeff increases across a period and decreases down a
group
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101
Practice Problem

How does the Effective Nuclear Charge, Zeff , influence
atomic size, Ionization Energy (IE), and Electronegativity
(EN)?
Ans:
As you move to the right across a Period:
The Atomic Size decreases
The Ionization Energy increases
The Electronegativity increases
All because of the increased Zeff
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102
Practice Problem

How are covalent and metallic bonding similar?
Ans:
Covalent and Metallic bonding involve sharing of
electrons between atoms
Covalent bonding includes sharing between a small
number of atoms (usually two)
Metallic bonding involves essentially all the atoms in
a given sample
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103
Practice Problem

Which of following pairs react to form covalent
compounds, ionic compounds?
a. Be & C
Ans:
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b. Sr & O
c. Ca & Cl
d. P & F
a. Covalent
(2 non-metals)
b. Ionic
(Oxygen is most ionic in group 6)
c. Ionic
Metal & non-metal)
d. Covalent
(2 non-metals)
104
Practice Problem

Which member of each pair gives the more acidic
solution?
a. CO2 or SrO
b. SnO or SnO2
c. Cl2O or Na2O
d. SO2 or MgO
Ans:
a. Carbon dioxide will form a more acidic solution
2 CO2(g) + H2O(l) → H2CO3(aq)
(Weak acid)
SrO(s) + H2O(l)
→ Sr(OH)2(aq) (Sr oxides form basic solutions)
b. Tin(IV) oxide (SnO)
An element with more than 1 oxidation state
exhibits less metallic behavior in its higher
state – more acidic
c. Dichlorine Oxide (Cl2O) Non-metal oxides form acidic solution;
metallic oxides form basic solution
d. Sulfur Dioxide (SO2)
Non-metal oxides for acidic solution;
metallic oxides form basic solution
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105
Practice Problem

Each of the following properties shows a regular trend in
Group 1A. Predict whether each increases or decreases
up the group
a. Melting Point
b. E – E bond length
d. Molar Volume
e. Lattice Energy of E-Br
Ans:
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c. Hardness
a. Increases
Increased Lattice Energy
b. Decreases
Decreasing radius of atom
c. Increases
Increased Lattice Energy
d. Decreases
Decreasing radius of atom
e. Increases
The atomic radius decreases
as you move up a group
increasing the lattice energy.
106
Practice Problem

The melting points of Alkaline Earth metals (group 2A) are
many times higher than the Alkali metals (group 1A)
Explain this difference on the basis of atomic properties
Ans: Metal atoms are held together by metallic bonding, a
sharing of valence electrons.
Alkaline earth metal atoms have one more valence electron
than alkali metal atoms, so the number of electrons shared
is greater
Thus, metallic bonds in alkaline earth metals are stronger
than in alkali metals.
Melting requires overcoming the metallic bonds
To overcome the stronger alkaline earth metal bonds
requires more energy (higher temperature) than to
overcome the alkali earth metal bonds.
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107
Practice Problem

Many compounds of group 3A elements have chemical
behavior that reflects an electron deficiency
Explain electron deficiency with illustrative reactions

Ans: Compounds of Group 3A(13) elements (ns2np1), like Boron (B),
have only six electrons in their valence shell when combined
with Halogens to form three bonds
Having six electrons, rather than an octet, results in an
“electron deficiency,” i.e., violates octet rule
As an electron deficient central atom, Born (B) is
trigonal planar (AX3). Upon accepting an electron pair to form a
bond, the shape changes to tetrahedral (AX4)
BF3(g) + NH3(g) → F3B–NH3(g)
F
AX3
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B
F
F
AX4
108
Practice Problem
Nearly every compound of Silicon (Group 4A) has the element in the +4
oxidation state. In contrast, most compounds of Lead have the element
in the +2 state
a. What general observation does this fact illustrate?
Ans: The increased stability of the lower oxidation state as one goes
down a group
b. Explain in terms of atomic structure and molecular properties
Ans: As the atoms become larger (Pb > Si), the strength of the
bonds to other elements becomes weaker, and insufficient
energy is gained in forming the bonds to offset the additional
ionization or promotion energy
c. Give an analogous example from Group 3A
Ans: Thallium(Tl+) is more stable than Tl3+, but Al3+ is the only
table oxidation state for Al
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109
Practice Problem
Based on the relative sizes of Fluorine (F) and Chlorine(Cl),
predict the structure of PF2Cl3
Ans: From the Lewis structure, the Phosphorus (Central
atom) has 5 electron groups for a trigonal bipyramidal
molecular shape
In this shape, the three groups in the equatorial plane
have greater bond angles (120°) than the two groups
above and below this plane (90°)
The Chlorine atoms (larger than Fluorine atoms)
would occupy the planar sites where there is more
space for the larger atoms
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110
Practice Probem
A Halogen (X2) disproportionates (one X is reduced and
one is oxidized) in base in several steps to X- and XO-3.
Write the overall reaction for the disproportionation of Br2
to Br- and BrO3-1
Ans: A substance that disproportionates serves as both an
oxidizing and reducing agent.
Assume that OH– serves as the base

Write the reactants and products of the reaction, and
balance like a redox reaction
Br2(l) + OH– (aq) → Br– (aq) + BrO3-1 (aq) + H2O(l)
Br2(l) + 6 OH– (aq) → Br– (aq) + BrO3-1 (aq) + 3 H2O(l)
Balance e- on each side using coefficients
0
5 x -1 = -5
+5
3 Br2(l) + 6 OH– (aq) → 5 Br– (aq) + BrO3-1 (aq) + 3 H2O(l)
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111
Practice Problem

The main reason Alkali metal dihydrides (MX2) do not form
is the Ionization Energy (IE) of the metal
Why is the IE so high for alkali metals
Ans:
Alkali metals have an ns1 outer electron configuration
The first electron lost by the metal is the ns
electron, giving the metal a noble gas configuration
Second ionization energies for alkali metals are high
because the electron being removed is from the next
lower energy level and electrons in a lower level are
more tightly held by the nucleus.
The metal would also lose its noble gas configuration
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