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.
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1
Main-Group Elements

Chapter Overview
 Application of bonding, structure, and reactivity to MainGroup 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
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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


 f-block - Lanthanoid (ide) series)
  f-block - Actinoid (ide) series)
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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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Main Group Elemetnts

Periodic Table Numbering System
 Old systems confusing
● 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)
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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
 No perfectly suitable position in
the periodic table
 Depending on the property,
Hydrogen fits better in 1A, 4A, 7A





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+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
 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)
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Covalent Hydrides

Hydrogen reacts with nonmetals to form covalent hydrides
CH4

NH3
H2O
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





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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
14
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)  2Li2O(s) oxide
K(s) + O2(g)  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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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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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
Gypsum – CaSO42H2O

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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
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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)
(X = F, Cl, Br, I)
 Most metals (Be exception) reduce Hydrogen to form ionic
hydrides
E(s) + H2(g)  EH2 (s)
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(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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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
 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
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Group 3A – Boron Family (ns2np1)
 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)
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Group 3A – Boron Family (ns2np1)
 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
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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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Group 3A – Boron Family (ns2np1)
atomic properties
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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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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)
 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)
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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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Group 4A – Carbon Family (ns2np2)

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
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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
 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
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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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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:
● 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
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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 ()
 The C-C bond is short enough for side-to-side
overlap of two half-filled 2p orbitals to form 
bonds that give rise to many diverse structures
and reactivities of organic compounds
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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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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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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 orbital
Trimethylamine
(CH3)3N
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trigonal
pyramidal
The impact of p,d- bonding on the structure of trisilylamine
Trisilylamine
(SiH3)3N
trigonal planar
50
Group 4A – Carbon Family (ns2np2)



Silicon chemistry is dominated by the Silicon-Oxygen –
(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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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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Group 4A – Carbon Family (ns2np2)
atomic properties
physical properties
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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
CO2 + H2O ⇄ H2CO3(aq) ⇄ H+(aq) + HCO3-(aq)

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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

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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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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
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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
● Nitrogen
– from +5 to -3
● P, As, Sb
– +5 & +3
● Bi
– +3
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Group 5A – Nitrogen Family (ns2np3)

Nitrogen Chemistry
 Nitrogen Oxides – 6 stable forms
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Group 5A – Nitrogen Family (ns2np3)

Nitrogen Oxoacids & Oxoanions
Oxoacid
Oxoanion
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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.
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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
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Group 5A – Nitrogen Family (ns2np3)
● Tetraphosphorus Decaoxide (con’t)
 Phosphorus (+5) oxidation state
P4(s) + excess 5O2  P4O10(s)
P4O10(s) + 6H2O  4H3PO4(l)
● 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+
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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
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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-
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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
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68
Group 6 – Oxygen Family (ns2np4)
atomic properties
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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)
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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
● 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
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71
Group 6 – Oxygen Family (ns2np4)

Oxygen Family vs Nitrogen Family
 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
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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
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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)
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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.
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75
Group 6 – Oxygen Family (ns2np4)
 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
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76
Group 6 – Oxygen Family (ns2np4)
● 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
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77
Group 6 – Oxygen Family (ns2np4)
● 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
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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
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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)
● Neither H2SO3 or H2CO3 (both weak acids) can exist
as isolated molecules – they dissociate immediately
in water
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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
SO2(g) + 1/2O2 V O⇄/K OSO3(g)
● 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+)
2
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5
2
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 anion26
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82
Group 7 – Halogen Family (ns2np5)
atomic properties
physical properties
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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
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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
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85
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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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)  CH2=CHCl(g) + HCl(g)
Ethylene
1,2 Dichloroethane
1-chloroethene
 Hydrogen Fluoride, with its short, strong bond forms a
weak acid (Hydrofluoric acid) with water
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HF(g) + H2O(l)  H3O+ + F-
87
Group 7 – Halogen Family (ns2np5)
 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 1- charge;
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
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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)
 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
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Group 7 – Halogen Family (ns2np5)
 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
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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
 If only 1 electron changed, then an unstable loneelectron species containing 2 Fluorines would form
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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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Group 7 – Halogen Family (ns2np5)

Molecular shapes of the main types of interhalogen
ClF3
compounds
linear, XY
Tshaped
XY3
ClF
IF7
BrF5
900
Square Pyramidal
XY5
Pentagonal
Bipyramidal,
XY7
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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
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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
Hypochlorous,
Chlorous
Chloric
Perchloric
Salt Sodium Hypochlorite, Sodium Chlorite, Sodium Chlorate, Sodium Perchlorate
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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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Group 7 – Halogen Family (ns2np5)
● 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
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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 > HOCl+1
Perchloric
Chloric
Chlorous Hypochlorous
Acid
Acid
Acid
Acid
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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
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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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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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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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Practice Problem

Which of following pairs react to form covalent
compounds, ionic compounds?
a. Be & C
b. Sr & O
Ans: a. Covalent
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c. Ca & Cl
d. P & F
(2 non-metals)
b. Ionic
(Oxygen is most ionic in group 6)
c. Ionic
Metal & non-metal)
d. Covalent
(2 non-metals)
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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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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
c. Hardness
d. Molar Volume e. Lattice Energy of E-Br
Ans:
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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.
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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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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
F
B
F
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AX4
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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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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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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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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 outer electron configuration of
ns1
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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