Salts & Metals

Report
Chapter 7– Ionic Compounds & Metals
7.0
7.1
7.2
7.3
7.4
Bonding Overview
Ion Formation
Ionic Bonds and Ionic Compounds
Names & Formulas for Ionic Compounds
Metallic Bonds & Properties of Metals
Section 7.0 Bonding Overview
There are 3 basic types of chemical bonding
which vary in how the valence electrons are
used. Electronegativity difference and the type
of element (metal or nonmetal) are the key
parameters in determining the type of
bonding that will occur.
• Name the three basic types of chemical bonding.
• Distinguish between metals and nonmetals using a periodic
table
• Explain the “octet rule”, the role it plays in chemical
bonding and its relationship to noble gas electron
configurations
Section 7.0 Bonding Overview
• Describe the role of electronegativity difference in
determining the type of bonding between two elements
• Distinguish between an ionic and a covalent (molecular)
compound and describe the basic difference in bond
formation between these two types of compounds
Section 7.0 Bonding Overview
Key Concepts
• A chemical bond is the force that holds two atoms together.
• There are three main types of chemical bonds – ionic,
covalent and metallic. All involve valence electrons in
some way.
• Both the electronegativity difference and the category of an
element (metal or nonmetal) determine the type of bond that
will form.
• Ionic bonds are formed between a metal and nonmetal with
a large electronegativity difference (> 1.7)
• Covalent bonds are generally formed between nonmetals or
between a metal and a nonmetal that have a small (< 1.7)
electronegativity difference.
• Metallic bonds are formed between metals and other metals.
Electron Categories
The electrons responsible for the
chemical properties of atoms are those
in the outer energy level
Valence electrons - The s and p
electrons in the outer energy level
• highest occupied energy level
Core electrons - those in energy levels
below the valence electrons
Chemical Bonds
Bond is force holding two atoms
together
When describing bond formation, focus
is on valence electrons
Elements tend to lose or gain electrons
to achieve an octet of electrons
• Extra stability associated with noble gas
configuration (filled outmost energy level)
• For low AN elements, have [He], not octet
Chemical Bonds & Valence Electrons (VEs)
Ionic Bonding
• Transfer of VEs from one atom to another
• Results in charged ions with opposite
sign that attract each other (e.g. Na+ Cl- )
Covalent Bonding
• Sharing of VEs between atoms
Metallic Bonding
• VEs become part of “sea” of electrons
Types of Chemical Bonds
Ionic
• Formed between metal and nonmetal
• Na (metal)
Cl (nonmetal)
NaCl sodium chloride
Covalent
• Generally formed between nonmetals
• C (nonmetal) O (nonmetal)
CO2 carbon dioxide
Metallic
• Formed between metals (same or similar)
Fe iron
Ionic vs Covalent Bonding
Key parameter in distinguishing types is
the electronegativity
Electronegativity
Relative ability to attract electrons in a
chemical bond
• Max value 3.98 - F
• Min value 0.7 - Fr
Elements with high EN tend to form
anions
Noble gases not tabulated
• Very few compounds to get info from
Electronegativity Ranges
(values slightly different than in book)
Below 1.0
1.0 – 1.4
1.5 – 1.9
2.0 – 2.4
2.5 – 2.9
3.0 – 4.0
Fig. 8.20 (p. 265) Electronegativity
Electronegativity & Bonding
[see p. 266 in section 8.1]
Difference in electronegativity (EN)
between the atoms involved in bond
formation determines type of bond
• = EN(atom 1) – EN (atom 2)
If difference large, electron transferred
 ionic bond
If difference small or zero, electron
shared  covalent bond
100
75
Ionic Bonds
50
25
% Ionic Character
EN Difference & Bond Character
0
Covalent
Bonds
1.0
2.0
3.0
Electronegativity Difference
Electronegativity (EN) & Bonding Type
Large EN difference (> 1.7) usually occurs
between 2 elements when one is a metal and
the other a nonmetal
Group 1 and 2 metals and highly
electronegative nonmetals form compounds
with high ionic character (>50% or DEN > 1.7)
Polar covalent bonds involve unequally shared
valence electrons (but not so unequal that ions
can form)
Most bonds have some mix of ionic & covalent
character
Electronegativity Difference
and Bond Character
(Table 8.7, p. 266)
Ionic Bonding
Subjects for Sections 7.1 and 7.2 are
formation of ions and formation of ionic
bonds
Chapter 7– Ionic Compounds & Metals
7.0
7.1
7.2
7.3
7.4
Bonding Overview
Ion Formation
Ionic Bonds and Ionic Compounds
Names & Formulas for Ionic Compounds
Metallic Bonds & Properties of Metals
Section 7.1 Ion Formation
Ions are formed when atoms gain or lose
valence electrons to achieve a stable octet
electron configuration.
• Define a chemical bond.
• Describe the formation of positive and negative ions from
the elements.
• Describe the size change that occurs when an atom becomes
an anion or a cation.
• Relate ion formation to electron configuration.
• Determine the oxidation state of metals and nonmetals
based on their position in the periodic table..
Section 7.1 Ion Formation
• Determine the electron configuration of an ion
• Explain why transition metals can have multiple oxidation
states.
• Name the transition metals that do not have multiple
oxidation states, list the oxidation number associated with
each, and explain using an argument involving the ion’s
electron configuration why this single oxidation state is
highly preferred.
Section 7.1 Ion Formation
Key Concepts
• A chemical bond is the force that holds two atoms together.
• Some atoms form ions to gain stability. This stable
configuration involves a complete outer energy level,
usually consisting of eight valence electrons.
• Ions are formed by the loss or gain of valence electrons.
• The number of protons remains unchanged during ion
formation.
• Transition metals can use electrons in d orbitals as valence
electrons to attain multiple oxidation states.
• The electron configuration of most ions is obtained by
removing electrons in reverse order of highest n from the
atom’s electrons configuration.
Bonding
Atom will try to form octet by gaining or
losing valence electrons
Bonding
Metals are reactive because they lose
valence electrons easily
Formation of Cations
Positive ions – called cations
• Ca+ions – write “t” as a plus sign
Energy equal to ionization energy (IE)
must be supplied to remove electron
• Expressed in kJ/mol (kilojoules per
mole)
Na + ionization energy  Na+ + electron
IE = 498 kJ/mol
Atom vs Cation Radius
Ionization of sodium
Na

Na+
+
Sodium atom Sodium cation
[Ne]3s1
[Ne]
e-
Formation of Cations
In most cases, lose enough electrons to
achieve noble gas configuration
Formation of Cations
For group 1 and 2 metals (and H), cation
charge = group number
Aluminum always has +3 charge
(Above metals lose all valence electrons)
• Na+1, Mg+2, Al+3
First two rows of table 7.7, page 218 lists
group 1 and 2 cations
Formation of Cations
For transition metals
• Generally have ns2 configuration, so
generally can form +2 cation
• Also have (n-1)dx configuration, so
some d electrons also lost
• Books says that “rule of thumb” –
generally can form +2 or +3

Not very good rule – wide range from +1
to +7 with +8 possible (osmium,
ruthenium)
Transition Metal Oxidation States
(uses old group designations)
Common
Less Common
Transition Metal Oxidation States
Special Case Transition Metals
Several transition metals only
commonly form a single type of ion:
Sc3+ (many sources don’t list)
Zn2+
Ag1+
Cd2+ (some sources don’t list)
Roman numerals are not used in
compound names involving these ions
Forming Transition Metal Cations
For atom, following standard Aufbau
order for transition metal means ns
sublevel fills before (n - 1)d sublevel
Example:
Ti [Ar]3d24s2 (filling order 4s then 3d)
To get electron configuration of ion,
remove valence electrons first
Ti2+ [Ar]3d2
Special Case Transition Metals
Pseudo-noble gas configuration
• Groups 3 - 14, periods 4 - 6
• Full sublevels have extra stability
• s2, p6, d10 – all with same
When an ion can be formed with this
configuration it is especially stable and
become the preferred (or only) ion
formed by this metal
Special Case Transition Metals
Zn = [Ne]3s23p63d104s2
Zn+2 = [Ne]3s23p63d10
Compare to Ar = [Ne]3s23p6
Consequence: Zn+2 only Zn ion formed
+2
Zn
– Pseudo-Noble Gas
Special Case Transition Metals
Pseudo-noble gas configuration
Ag [Kr]4d105s1
• Same situation as Cu (not 5s2)
Ag1+ [Kr]4d10 = [Ar] 3d104s24p64d10
Compare to Kr = [Ar] 3d104s24p6
Consequence: Ag1+ only Ag ion formed
Reasoning for Ag1+ applies to Cd2+
Cd [Kr]4d105s2 = [Ar]3d104s24p64d105s2
Cd2+ = [Ar]3d104s24p64d10
Special Case Transition Metals
Noble gas configuration
• Sc [Ar]3d14s2
• Sc3+ = [Ar] = Only common oxidation
state for Sc (+3)
Formation of Anions
Applies to nonmetals on upper RHS of
periodic table
Negative ions called anions
• aNions – N for negative
Some nonmetals can gain or lose
electrons to complete an octet
Formation of Anions
Gain enough electrons to form octet
• Halogen anions: –1 charge
• Oxygen (O), Sulfur (S): -2 charge
• Nitrogen (N), Phosphorus (P): -3 charge
• Carbon (C): - 4 charge
Charge = group # - 18 (groups 14 to 18)
Naming Monatomic Anions
For monatomic anions, name becomes
element root + “ide” ending
• Chlorine (Cl)  Chloride (Cl-)
• Oxygen (O)  Oxide (O2-)
• Sulfur (S)  Sulfide (S2-)
• Nitrogen  Nitride (N3-)
• Phosphorus  Phosphide (P3-)
• Carbon  Carbide (C4-)
Atom vs Anion Radius
Formation of chlorine ion
Cl + e- 
ClChlorine atom
Chloride anion
[Ne]3s23p5
[Ne]3s23p6 or [Ar]
Electron Dot diagrams
Way of keeping track of
valence electrons
How to write them?
• Write the symbol
• Put one dot for each
valence electron
• Don’t pair up until you
have to (Hund’s rule)
X
Electron Dot diagram for Nitrogen
5 valence electrons
First write chemical symbol
Add 1 electron at a time to each side
Until they are forced to pair up.
N
Electron Dots For Cations
Metals generally have few valence
electrons
These will come off
Forming positive ions
2+
Ca
Electron Dots For Anions
Nonmetals have many valence
electrons (usually 5 or more)
Gain electrons to fill outer energy level
P
P
3-
Chapter 7– Ionic Compounds & Metals
7.0
7.1
7.2
7.3
7.4
Bonding Overview
Ion Formation
Ionic Bonds and Ionic Compounds
Names & Formulas for Ionic Compounds
Metallic Bonds & Properties of Metals
Section 7.2 Ionic Bonds and Ionic Compounds
Oppositely charged ions attract each other,
forming electrically neutral ionic compounds.
• Describe the formation of ionic bonds using multiple
methods including the use of electron configurations, orbital
diagrams and dot diagrams.
• Describe the structure of ionic compounds.
• Generalize about the strength of ionic bonds based on the
physical properties of ionic compounds.
• Explain why ionic compounds are brittle.
Section 7.2 Ionic Bonds and Ionic Compounds
• Identify the conditions under which ionic compounds do or
don’t conduct electricity and explain why.
• Categorize ionic bond formation (starting with the
component ions) as exothermic or endothermic.
• Describe lattice energy and the sign convention used to
report it.
• Relate the strength of ionic bonds and the lattice energy to
the size of the ions and the charge on the ions using
reasoning based on Coulomb’s law.
Section 7.2 Ionic Bonds and Ionic Compounds
Key Concepts
• Ionic compounds contain ionic bonds formed by the
attraction of oppositely charged ions. Ions in an ionic
compound are arranged in a repeating pattern known as a
crystal lattice.
• Ionic compound properties are related to ionic bond
strength.
• Ionic compounds are electrolytes; they conduct an electric
current in the liquid phase and in aqueous solution.
• Lattice energy is the energy needed to remove 1 mol of ions
from its crystal lattice.
Ionic Bonding
Bond formed through transfer of electrons
to form anion and cation
In most cases, electrons transferred to
achieve noble gas configuration in each
ion
No net loss or gain of electrons
• Total number lost = total number gained
• NaCl Sodium loses 1, chlorine gains 1
• AlCl3 Al loses 3, three Cl each gain 1
Ionic Bonding
Elements start out electrically neutral
(no charge) and ionic compound must
also be neutral
Ionic Compounds
Ionic compounds called salts
Salts involving oxygen called oxides
• CuO Copper (II) Oxide
Simplest ratio called formula unit (for ionic
compounds, there are no molecules)
• NaCl (1:1 ratio)
• Al2O3 (2:3 ratio)
• Be4Al2Si6O18 formula unit for beryl
Simple binary compounds = 2 different
elements
• MgO, KCl, FeBr , Al O
NaCl Formation
Electron Configuration Picture
Transferred
Electron
NaCl Formation
Orbital Notation Picture
Transferred
Electron
Na
Cl
Na+
Cl-
NaCl Formation
Electron Dot Picture
Na Cl
NaCl Formation
Electron Dot Picture
+
Na
Cl
Note that after electron transfer, the
atoms are charged
Sodium Chloride Formation
No single isolated unit of + and – charge
(no molecule
as
such)
11 e
10 e-
Electron loss
Sodium Ion
Neutral Na
Atom
18 e-
17 eElectron gain
Neutral Cl
Atom
Chloride Ion
NaCl Crystalline structure
Ionic Bonding
All the electrons must be accounted for!
Ca
P
Ionic Bonding
Ca
P
Ionic Bonding
2+
Ca
P
2-
Ionic Bonding
2+
Ca
Ca
P
2-
Ionic Bonding
2+
Ca
1+
Ca
P
3-
Ionic Bonding
2+
Ca
P
1+
Ca
P
3-
Ionic Bonding
2+
Ca
2+
Ca
P
P
3-
1-
Ionic Bonding
Ca
2+
Ca
2+
Ca
P
P
3-
1-
Ionic Bonding
Ca
2+
Ca
2+
Ca
P
P
3-
1-
Ionic Bonding
2+
Ca
2+
Ca
2+
Ca
P
P
33-
Ionic Bonding – Formula Unit
= Ca3P2
Formula Unit for
Calcium Phosphide
Simplest ratio of ions in compound is
called the formula unit
Properties of Ionic Compounds
Crystalline structure, usually solids
A regular repeating 3D arrangement of
ions in the solid – crystal lattice
Crystals vary in shape due to variation
in relative number and sizes of ions
No single isolated unit of + and –
charge (no molecule as such)
Crystalline structure
Ionic Bonding
Anions and cations are held together by
force between opposite charges
• Strong electrostatic force
• Magnitude of force (F) given by
Coulomb’s Law
• F proportional to q+q- /r2



q+ = magnitude of positive (cation) charge
q- = magnitude of negative (anion) charge
r = distance between charge centers
Force vs distance for 1/r2
0
0
F o rce
0
0
0
0
0
0
0
5
15
25
35
r (d ista n ce )
45
Coulomb’s Law & Ionic Bonds
F  q+q-/r2
The larger the ionic charges (q+, q-), the
stronger the force between them
The closer together the ions are, the
stronger the force
• Smaller ions can get closer together
than larger ions
Properties of Ionic Compounds
Ions are strongly bonded together
Structure is hard, rigid, brittle
High melting and boiling points
Strength of bond depends on relative
size and charge of ions
Ionic Solids – Melting, Boiling Points
Compound
MP ( C)
BP ( C)
NaI
KBr
NaBr
CaCl2
CaI2
NaCl
MgO
660
734
747
782
784
801
2852
1304
1435
1390
>1600
1100
1413
3600
Ionic solids are brittle
+
+
-
+
+
+
+
-
+
+
Ionic solids are brittle
Strong Repulsion breaks crystal apart.
- + - +
+ - + - + - +
Conductivity of Ionic Solids
Substance that conducts electricity is
allowing charges to move
In a solid, ions locked in place
• Ionic solids are insulators
When melted, ions can move around
• Melted ionic compounds conduct
• NaCl: must get to about 800 ºC
Conductivity of Ionic Solids
Dissolved in water (aqueous) they conduct
• Solution called electrolyte
• Each individual ion, surrounded by H2O
molecules, free to move about and carry
charge from one place to another
Conductivity of Ionic Solids
Conduct when
dissolved in water
(electrolyte)
Energy and the Ionic Bond
Energy absorbed or released during a
chemical reaction
• Released – Exothermic
• Absorbed – Endothermic
Formation of ionic compounds from
cations & anions always exothermic
Energy and the Ionic Bond
Ions sit in crystal lattice
Takes energy to separate them into
individual ions
Energy required is called lattice energy
Like ionization energy, lattice energy
expressed in kJ/mol (kilojoules/mole)
The more negative the value, the
stronger the forces of attraction
• Can correlate LE with melting point
Lattice Energy (LE)
Related to size and charge of ions
Smaller ions generally have more
negative values (stronger forces)
• Electrons approach closer to + nucleus
• Li (period 2) compound more negative
LE than K (period 4) compound

Follows ionic radius trend – Li ion < K ion
Lattice Energy (LE)
Related to size and charge of ions
Bonds formed from ions with larger
charges have more negative LE
• MgO (+2, -2 72, 140 pm)
4 X more negative LE than
NaF (+1, -1 102, 133 pm)
Lattice Energy (LE)
See next 2 slides for explanation of arrows
Comp
LE
ound (kJ/mol)
KI
RbF
KF
AgCl
SrCl2
MgO
-632
-774
-808
-910
-2142
-3795
Cation
r (pm)
138
152
138
126
118
72
Anion
r (pm)
220
133
133
181
181
140
Lattice Energy (LE) – Size Effects
Comp
LE
ound (kJ/mol)
KI
RbF
KF
-632
-774
-808
Cation
r (pm)
138
152
138
Anion
r (pm)
220
133
133
I vs F – anion size decrease
Rb vs K – cation size decrease
Lattice Energy (LE) – Charge Effect
Comp
LE
ound (kJ/mol)
AgCl
SrCl2
-910
-2142
Cation
r (pm)
126
118
Anion
r (pm)
181
181
Ag vs Sr – cation charge increase
from +1 to +2
Chapter 7– Ionic Compounds & Metals
7.0
7.1
7.2
7.3
7.4
Bonding Overview
Ion Formation
Ionic Bonds and Ionic Compounds
Names & Formulas for Ionic Compounds
Metallic Bonds & Properties of Metals
Section 7.3 Names and Formulas for Ionic
Compounds
In written names and formulas for ionic
compounds, the cation appears first,
followed by the anion.
• Determine the oxidation state of metals and nonmetals
based on their position in the periodic table
• Relate a formula unit of an ionic compound to its
composition.
• Know the formulas and charges of common polyatomic
ions
• Name ionic compounds and produce the formula of a
compound given its name
Section 7.3 Names and Formulas for Ionic
Compounds
• Know how to transform an oxyanion name to adjust for
increased or decreased oxygen, addition of hydrogen, or
change in halogen
• Manipulate subscripts (including use of parentheses for
polyatomic ions) in the chemical formula of an ionic
compound to produce a neutral compound
• Know when to use roman numerals in the names for ionic
compounds
Section 7.3 Names and Formulas for Ionic
Compounds
Key Concepts
• A formula unit gives the ratio of cations to anions in the
ionic compound.
• A monatomic ion is formed from one atom. The charge of a
monatomic ion is its oxidation number.
• Roman numerals indicate the oxidation number of cations
having multiple possible oxidation states.
• Polyatomic ions consist of more than one atom and act as a
single unit. To indicate more than one polyatomic ion in a
chemical formula, place parentheses around the polyatomic
ion and use a subscript.
Formulas
formula unit = simplest ratio of ions in
compound
• Overall charge is zero
• KBr (K+ Br-)
• MgCl2 (Mg2+ 2 x Cl-)
Monatomic ion = one atom ion
• Na+1, O2-
Oxidation Number
Charge on monatomic ion = oxidation
number or oxidation state
• Na+1 has oxidation number of +1
• Equals number of electrons transferred
from atom to form the ion


+ sign transferred from
- sign transferred to
• Oxidation numbers of elements in ionic
compound must sum to zero
Rules for Naming Ionic Compounds
(see page 223)
1. Name cation first, anion second
Cation is always written first in formula
2. Monatomic cations use element
name
sodium magnesium lead
3. Monatomic anions use element root
name plus suffix –ide
chloride oxide sulfide nitride
Naming Monatomic Anions
For monatomic anions, name becomes
element root + “ide” ending
• Chlorine (Cl)  Chloride (Cl-)
• Oxygen (O)  Oxide (O2-)
• Sulfur (S)  Sulfide (S2-)
• Nitrogen  Nitride (N3-)
• Phosphorus  Phosphide (P3-)
• Carbon  Carbide (C4-)
Ionic Compound Names
Examples for Monatomic Ions
Cations Either Al or Groups 1A, 2A
NaCl – Sodium chloride
MgO – Magnesium oxide
Al2O3 – Aluminum oxide
Ca3P2 – Calcium phosphide
Li3N – Lithium nitride
Naming Ionic Compounds
Binary Compounds
Write formulas from names
Practice problems 19-23, page 221
Write names from formulas
Practice problems 28-30, page 223
Polyatomic Ions
Ions made up of more than one atom
• OH- (hydroxide)
• SO42- (sulfate)
• NH4+ (ammonium)
Note: atoms within the polyatomic ion
are covalently bonded to each other
Never change the subscripts of
polyatomic ions
Polyatomic Ions
p 221
Polyatomic Ions
Table 7.9, page 221 lists common ones
You must know:
• Ammonium
NH4+
• Nitrite, Nitrate
NO2- NO3• Hydroxide
OH• Bicarbonate*, Carbonate HCO3- CO32• Phosphate
PO43• Peroxide
O22*
AKA
• Sulfate
SO42hydrogen
• Chlorate
ClO3carbonate
Polyatomic Ions
Table R-5, page 970 has more
comprehensive list sorted by net
charge
See also Wikipedia
http://en.wikipedia.org/wiki/Polyatomic_ion
Naming Ionic Compounds
With Polyatomic Ions
Write formulas from names
Practice problems 24-27, page 222
Write names from formulas
Practice problems 30-33, page 223
Naming Ions - Oxyanions
Oxyanion = Polyatomic anions
• Same nonmetal element
• Differing number of oxygen atoms
NO3SO42-
NO2SO32-
Nonmetal N 3 vs 2 O
Nonmetal S 4 vs 3 O
Naming Ions - Oxyanions
Rules:
• Greater O atoms, use nonmetal root +
suffix –ate
• Fewer O atoms, use nonmetal root +
suffix –ite
NO3- nitrate NO2- nitrite
SO42- sulfate SO32- sulfite
Naming Ions
Oxyanions with Halogen or P
Rules: More complicated because can
have three or four different anions
ClO4ClO3- ClO2ClOperchlorate chlorate chlorite hypochlorite
PO43Phosphate
PO33-
PO23-
Phosphite Hypophosphite
Oxyanion Naming Conventions
for Chlorine (table 7.11, p 223)
Naming Oxyanions with Br and I
All rules followed by chlorine are
followed by bromine and iodine as well
Change root name
Chlorate Bromate Iodate
ClO3BrO3IO3IO4- ?
periodate
Polyatomic Ions
In series with varying oxygen (only), charge
fixed
• Nitrate, Nitrite
NO3- NO2• Sulfate, Sulfite
SO42-, SO32• Phosphate, Phosphite, Hypophosphite
PO43-, PO33- , PO23• Perchlorate, Chlorate, Chlorite, Hypochlorite
ClO4- ClO3- ClO2- ClO-
Polyatomic Ions
For series differing by an H, charge
increases by +1 for each added H
• Carbonate, Hydrogen carbonate
CO32- HCO3• Sulfate, Hydrogen Sulfate
SO42- HSO4• Phosphate, Hydrogen Phosphate
Dihydrogen Phosphate
PO43- HPO42- H2PO4-
Rules for Naming Ionic Compounds
4. A. Group 1, 2 metals or Al – no
additional work necessary
B. Some group 13 to 15 metals and
most transition metals (see following
slide) have multiple oxidation states
- use Roman numeral in
parentheses to indicate which one
d-block
transition metals
13 14 15
Group 13
to 15
metals
Zn Ga
with
multiple
Ag Cd In Sn
oxidation
Tl Pb Bi
states
Lanthanides
Actinides
f-block transition metals
Rules for Naming Ionic Compounds
4. Group 13 to 15 metals + transition
metals
Iron(II) Iron(III)
Copper(I) Copper(II)
FeO – Iron(II) oxide
CuCl – Copper(I) chloride
Fe2O3 – Iron(III) oxide CuCl2 – Copper(II) chloride
Outdated but still commonly used
naming system uses –ous, -ic suffixes
(not responsible for these)
Ferrous Ferric
Cuprous Cupric
Table 7.8, page 219: values for some
transition and group 3A / 4A (13/14)
metals
Cation Oxidation Numbers
There are certain common cations
(beyond group 2) with fixed oxidation
numbers that do not get roman
numerals in their compound names
For table 7.8, these are:
Ag+, Zn2+, Cd2+, Al3+ (silver, zinc,
cadmium, aluminum)
You must know the four ions above
ZnCl2 zinc chloride, not zinc(II) chloride
Sometimes Sc3+ included in this list
Figure top of p 224
Rules for Naming Ionic Compounds
5. If compound contains a polyatomic
ion, use the ion name
NH4Cl
ammonium chloride
NaOH
sodium hydroxide
(NH4)2SO4
ammonium sulfate [note
use of parentheses and
subscript in this
compound to obtain
neutral compound]
Section 7.3 Assessment
Which subscripts would you most likely use
for an ionic compound containing a group 1
metal and a group 17 nonmetal?
(Remember, 1 = no written subscript)
A. 1 and 2
B. 2 and 1
C. 2 and 3
D. 1 and 1
???
Section 7.3 Assessment
Name of the compound Ca(OH)2?
A. calcium oxide
B. calcium (II) hydroxide
C. calcium hydroxide
D. calcium oxyhydride
???
Ionic Compounds Practice
Name the following:
Ba(NO3)2
barium nitrate
CaSO4
calcium sulfate
Mg3(PO4)2
magnesium phosphate
SrSO3
strontium sulfite
Ionic Compounds Practice
Name the following:
BaO
barium oxide
CaF2
calcium fluoride
Mg3N2
magnesium nitride
SrS
strontium sulfide
Mn(H2PO4)3 manganese (III)
dihydrogen phosphate
Ionic Compounds Practice
Name the following:
iron(II) bromide
FeBr2
FeBr3
iron(III) bromide
SnCl2
tin(II) chloride
SnCl4
tin(IV) chloride
Ionic Compounds Practice
Name the following:
AgCl
silver chloride
Na2Se
sodium selenide
Fe2O3
iron(III) oxide
CrI3
chromium(III) iodide
Ionic Compounds Practice
Name the following:
BaBr2
barium bromide
K2S
potassium sulfide
AlN
aluminum nitride
SnF4
tin(IV) fluoride
Cd(OH)2
cadmium hydroxide
Ionic Compounds Practice
Name the following:
Fe(OH)3
iron(III) hydroxide
NH4I
ammonium iodide
Na2O2
sodium peroxide
Ca(ClO)2 calcium hypochlorite
Ionic Compounds Practice
Name following compounds:
Fe(NO3)2
iron(II) nitrate
Fe(NO2)3
iron(III) nitrite
Sn(ClO)2
tin(II) hypochlorite
Sn(ClO2)4
tin(IV) chlorite
Ionic Compounds Practice
Name following compounds:
AgHSO4
Silver hydrogen sulfate
(NH4)2CO3
Ammonium carbonate
Naming Ionic Compounds
Problem 81 page 233 (get formula)
Problem 82 page 233 (get name)
Systematic vs Common Names
Elaborate rules exist for assigning
names to chemical substances on
basis of their structures
Called systematic names; they uniquely
identify given substance
Rules for these names are defined by
international body (IUPAC)
http://www.chem.qmul.ac.uk/iupac/
http://www.acdlabs.com/iupac/nomenclature/
Chapter 7– Ionic Compounds & Metals
7.0
7.1
7.2
7.3
7.4
Bonding Overview
Ion Formation
Ionic Bonds and Ionic Compounds
Names & Formulas for Ionic Compounds
Metallic Bonds & Properties of Metals
Section 7.4 Metallic Bonds and the Properties
of Metals
Metals form crystal lattices and can be
modeled as cations surrounded by a “sea” of
freely moving (delocalized) valence electrons.
• Describe a metallic bond.
• Describe the meaning of the words/terms “delocalized
electron”, malleable, and ductile.
• Describe how the properties of conductivity, reflectivity
malleability and ductility are related to the presence of
delocalized electrons (electron sea model).
• Describe the similarities and differences between ionic and
metallic bonding.
Section 7.4 Metallic Bonds and the Properties
of Metals
Metals form crystal lattices and can be
modeled as cations surrounded by a “sea” of
freely moving (delocalized) valence electrons.
• Define alloys, categorize them into two basic types, list the
two types of solution alloys, and give examples of each.
• List possible advantages of using an alloy over using a pure
metal
• Explain the role that carbon plays in steel alloys.
• Describe the roles that imperfections play in the properties
of metals and list various physical methods that are used to
alter these imperfections.
Section 7.4 Metallic Bonds and the Properties
of Metals
Key Concepts
• A metallic bond forms when metal cations attract freely
moving, delocalized valence electrons.
• In the electron sea model, electrons move through the
metallic crystal and are not held by any particular atom.
• The electron sea model explains the physical properties of
metallic solids.
• Metal alloys are formed when a metal is mixed with one or
more other elements.
Metallic Bonds
Metals don’t form ionic bonds
Do form solid state lattices
• Lattice similar to ionic crystal lattice
Have valance electrons, but these are
free to roam in a “sea” of other
electrons
Electrons are “delocalized” – not
confined to any particular location
Metallic Bonds – “Electron Sea Model”
Metal
ion (+)
Metal
lattice
structure
Free
electron
(-)
Metallic Bonds – “Electron Sea Model”
Metal
ion (+)
Free
electron
“sea”
Metallic Bonds
Metallic bond is the attraction of a
metallic cation for the delocalized
electrons
Not very directional, so metal atoms
can be rearranged without problem
• Gives ductility and malleability
Metallic Bonds – MP & BP
Indicate strength
of metallic bond
BP more extreme
than MP – large
energy required to
separate atoms
from soup of
cations and
electrons
Metals – Malleable & Ductile
Malleable
+
+ +
+
+
+
+
+
+ +
+
+
Malleable
Electrons allow atoms to slide by
+
+ +
+
+
+
+
+ +
+
+
+
Mobile Electrons
Impart good electrical conductivity
Interact with light, absorbing &
releasing photons
• Redirected light gives luster
Delocalized Electrons & Properties
As number of delocalized electrons
increases, so does hardness and
strength
Alkali metals soft (1 valence electron)
In transition metals, unpaired d
electrons are delocalized, so transition
metals in the middle of the d block tend
to be harder and stronger and also to
have higher MPs
Melting Points (C)
# U n p aire d
d e le c s
6
5
4
3
2
1
0
-1 2 0
25
A tom ic N um ber
U npa ire d d
M P (K )
30
2500
2250
2000
1750
1500
1250
1000
750
500
250
M P (K )
T rend of U npaired d E lec trons and
M elting P oint vs A N - P eriod 4
Period 6 – s block & TM Melting Points
(4 unpaired d
electrons)
Melting Point (K)
Peak occurs at W
Atomic Number
Alloys
Alloys have more than one element
(one a metal)
• Alloy has metal characteristics
Pure metals and alloys have different
physical and chemical properties
• Strength, hardness, corrosion resistance
In jewelry, alloy of gold & copper used
• alloy harder (& cheaper) than pure gold
Alloys - Types
Solution alloys are homogeneous
Heterogeneous alloys: components
are not dispersed uniformly
• Steel with >1.4% C has 2 phases:
almost pure Fe and cementite, Fe3C
(iron carbide)
• Fe3C is white, hard, brittle – makes
steel less ductile but much stronger
Alloys
Two types of solution alloy
• Substitutional alloys - some atoms in the
original metallic solid are replaced by
other metals of similar atomic structure
• Interstitial alloys - formed when small
holes in a metallic crystal are filled with
smaller atoms (solute occupies
interstitial sites in metallic lattice)
Alloys
Substitutional
Interstitial
Alloys
Alloys
Substitutional alloys
• atoms must have similar atomic radii
• elements must have similar bonding
characteristics

Sterling silver – Ag 92.5% Cu 7.5%
Interstitial alloys
• one element must have a significantly
smaller radius than the other (must fit
into interstitial site)

e.g. a nonmetal – Carbon Steel
Metal Properties
The chemical composition (alloying
elements) of a metal is only one factor that
determines metal properties
Properties such as hardness and strength
also depend on any mechanical and heat
treatments that may be applied
These treatments effect how the alloying
elements are distributed within the alloy, the
crystal size, and the number and type of
crystal defects within the material
Classification of Commerically
Important Metals
Ferrous Metals
Iron
Low Carbon Steel
Medium Carbon Steel
High Carbon Steel
Cast Iron
Alloy Steel
Stainless Steel
Others
Non- Ferrous Metals
Aluminum
Copper
Brass
Bronze
Zinc
Lead
Tin
Others
Steel
0.001% to 1.5% carbon
Wide range of properties due to
• Variation in carbon content
• Cold working (work hardening)
• Heat treatment
• Addition of alloying elements
Steel and Carbon
Carbon even at relatively low levels has an
impact on steel properties
Because iron and carbon form an interstitial
alloy, carbon acts as a “stiffener” to prevent
the layers of iron ions from moving freely
relative to each other
Result is a harder, stronger but more brittle
alloy as the carbon content increases
Steel - Effect of Increasing Carbon
Decreases ductility
Decreases machinability
Lowers melting point
Increases tensile strength
Increases hardness
Makes steel easier to harden with heat treatments
Lowers temperature required to heat treat steel
Increases difficulty of welding
Steel Composition - % by Weight
Balance is Fe
Type C
SAE
1010 0.080.13
1040 0.370.44
1552 0.470.55
Mn
P
S
Si
Remarks
0.30- 0.030 0.035 -- Common
0.60
0.60- 0.030 0.035 0.35
Tools
0.90
1.20- 0.030 0.035 0.35 Tempered
1.55
Parts
Nonmetals
Metalloid
Metal Properties – Other Factors
Although chemical composition (% Fe, % C,
etc) plays important role, other factors
strongly influence metal’s properties
(hardness, toughness, etc)
• Mechanical treatment (working)
• Heat treatment (tempering, quenching)
• Distribution of elements within metal
(often not homogeneous)
All of above can interact – study is field of
metallurgy
Cooling Rate and Crystal Size
The way metal prepared can have large
impact on how it behaves
Many metals prepared in liquid state &
cooled; rate of cooling can have significant
effect on properties of solid because it
controls crystal size/grain structure
Grain Structure & Imperfections (NIB)
Structure not continuous throughout
As metal cools, have l  s phase change,
atoms come together to form grains
Crystal structure not continuous
Steel
paper clip
Fe crystal
structure
Grain
Grain
Boundary
Formation of Grain Structure
Solidification of molten material
Two steps starting with molten material (all liquid)
1) Nuclei form
2) Nuclei grow to form crystals
Crystals grow until they meet each other to form
grain structure
nuclei
liquid
crystals growing
grain structure
Metal Crystal Size
Small crystals make metal
harder because ions less able
to move; also means there is
more disruption between
crystals making them brittle
(easy to break)
Larger crystals make metal soft
Imperfections and Alloys
Many imperfections within each crystal
Flaws produce weak points in bonds
between atoms
Adding other elements to produce an alloy
can counteract effects of imperfections and
make metal harder and stronger
Heat and mechanical treatment also effect
these imperfections
Area Defects: Grain Boundaries
Grain boundaries
•
•
•
•
boundaries between crystals
produced by solidification process
have change in crystal orientation across them
impede dislocation motion
grain
boundaries
Imperfections in Solids
Schematic drawing of poly-crystal with
many defects
Grain Structure & Imperfections (NIB)
Micrograph of metal that
has undergone
intergranular corrosion
Grain
Grain Boundary
Heat Treatment of Metals
3 ways of treating a metal with heat:
Annealing Quenching Tempering
Steel is alloy most commonly treated
Used to:
Soften part that is too hard
Harden part that is not hard enough
Put hard skin on part that is soft
Make good magnets out of ordinary material
Make selective property changes within parts
Heat Treatment
Metal on striking face of hammer heattreated differently than that on rest of head
Hardness on front traded for toughness at
back
Treatment
Process
Effect on metal
properties
Effect on metal
structure
Annealing
A metal is heated to a
moderate temperature
and allowed to cool
slowly
A metal is heated to a
moderate temperature
and cooled quickly
(sometimes by plunging
into water
The metal is
softer with
improved
ductility
The metal is
harder and
brittle.
Larger metal
crystals form
A quenched metal is
The metal is
heated (to a lower
harder but
temperature than is
less brittle.
used for quenching and
allowed to cool
Crystals of
intermediate
size form.
Quenching
Tempering
Tiny metals
crystals form.
Mechanical & Heat Treatment of Metals
Blacksmith creates objects from wrought iron or
steel by forging the metal (using tools to
hammer, bend, and cut) and in the process can
also change the characteristics of the metal
Arvind Thekdi - E3M, Inc.
Also uses heat treatment
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Cold Working
Increases strength at the expense of
ductility
Arvind Thekdi - E3M, Inc.
Sales
Mechanical Treatment of Metals
Work hardening (aka strain hardening or cold
working) is strengthening of metal by plastic
deformation. Strengthening occurs because of
dislocation movements and dislocation
generation within crystal structure of the material
Most non-brittle metals with a reasonably high
melting point as well as several polymers can be
strengthened in this fashion
Alloys not amenable to heat treatment, including
low-carbon steel, are often work-hardened
Arvind Thekdi - E3M, Inc.
Sales
Mechanical Treatment of Metals
Cold rolling increases strength via strain
hardening – metal grains become elongated
Cold
rolled
steel

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