Chemistry Matter and Change:
Chapter 7
Atoms in ionic compounds
are held together by chemical
bonds formed by the
attraction of oppositely
charged ions.
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
Relate ion formation to electron
configuration.
Ion
Valence electron
Octet
Electron configuration
Lewis-dot diagrams
Electron affinity
Chemical
bond
Ionic bond
Cation
Anion
The
force that holds two atoms together
Three types
Ionic bonds
Metallic bonds
Covalent bonds
*Chap 7
*Chap 7
*Chap 8
Each valence electron
is represented as a dot
around the nuclear core of the element.
The
most stable electron configuration for
an element is the nearest noble gas.
ns2np6
Octet
Ions gain or
lose electrons to achieve noble
gas configurations
Cation:
a positively charged ion
Results when electrons are lost
Group
1 loses 1 electron
+1 charge
Group
2 loses 2 electrons
+2 charge
Group
13 loses 3 electrons
+3 charge
Groups
3-12 usually lose 2 electrons
Most have +2 charge (range from +1 to +3)
Anion:
negatively charged ion
Formed when electrons are gained
Non-metals
Group
15 gains 3 electrons
3- charge
Group
16 gains 2 electrons
2- charge
Group
17 gains 1 electron
1- charge
A
chemical bond is the force that holds two
atoms together
Some atoms gain or lose electrons to gain a
stable configuration; these are called ions
Most stable configurations end: ns2np6.
•
•
•
Define a chemical bond.
Describe the formation of positive and
negative ions
Relate ion formation to electron
configuration.
Oppositely charged ions
attract each other forming
electrically neutral ionic
compounds.
Describe
the formation of ionic bonds and
the structure of ionic compounds
Generalize about the strength of ionic
bonds based on the physical properties of
ionic bonds
Categorize ionic bond formation as
exothermic or endothermic
Compound
Chemical
bond
Physical property
Chemical property
Electronegativity
Ionic
bond
Ionic compound
Crystal lattice
Binary compound
Electrolyte
Electrons are exchanged
between atoms
Increases stability of both
Ions are
charges
held together by the opposite
Bond
formed between two elements with an
electronegativity difference > 1.7
Crystallize as sharply defined particles
Formed
from a metal and a non-metal
Contain only two elements
Examoples
NaCl
MgO
CaCl2
Fe2O3
Net charge on all
ions in a compound must
be zero (0)!
More on this later!!!
Crystal
Lattice: Highly organized crystal of
cations and anions
Anion
Cation
Crystalline shape depends on
involved
the ions
Physical
properties
Very strong
Solid at normal temperatures
Very high melting point and boiling point
Many have brilliant colors due to transition metals
Hard, rigid
Brittle
Conductivity
(ability for electric charge to
move through a substance
Solids have electrons locked in place
Non conductive
Aqueous solutions have easily moveable electrons
Electrolytes
Good conductors
Dissolve
in water
May have radically different properties than
the elements that compose them
Polar
Dissolution
Formation of
lattice is always exothermic.
Ionic compounds contain
ionic bonds
formed by the attraction of oppositely
charged ions.
Ions in an ionic compound are arranged in a
repeating pattern called a lattice.
Ionic compounds are electrolytes; they
conduct electricity in liquid and aqueous
states.
Describe
the formation of ionic bonds and
the structure of ionic compounds
Generalize about the strength of ionic
bonds based on the physical properties of
ionic bonds
Categorize ionic bond formation as
exothermic or endothermic
In written names and
formulas for ionic
compounds, the cation
appears first, followed
by the anion.
•
•
•
Relate a formula unit of an ionic
compound to its composition
Write formulas for ionic
compounds and oxyanions.
Apply naming conventions to ionic
compounds and oxyanions.
Anion
Cation
Metal
Non-metal
Formula unit
Monatomic
ion
Polyatomic ion
Oxidation number
Oxyanion
Formula unit-
simplest way to indicate the
composition of an ionic substance
NaCl
MgCl2
Cl-
Mg2+
Cl-
Ions
in which only one element is present
• Na+, Cl-, Mg2+, P3-
Fancy word
for “charge”
aka oxidation state
Transition metals may have multiple
oxidation states
Must tell the oxidation state
Ex: Iron 2+ is Iron II; Iron 3+ is Iron III
CxAy
C is cation
A is anion
x number of cations in one unit
y is number of anions in one unit
CxAy
Cation is always first
Anion is always second
Net oxidation
MUST BE ZERO
1.
2.
3.
4.
Write out each ion.
Place oxidation number under each ion
Cross multiply
Reduce to simplest form
Sodium
and chlorine
Na
Cl
Sodium
and chlorine
Na
Cl
+1
-1
Sodium
and chlorine
Na
Cl
+1
-1
Na 1 Cl 1
Remove any
 NaCl
Reduce
“1”s
if needed
NaCl
Iron
III and oxygen
Fe
O
Iron
III and oxygen
Fe
O
3+
2-
Iron
III and oxygen
Fe
O
3+
2-
Fe2 O 3
Fe2O3
Cannot be reduced
Magnesium
Mg
O
and oxygen
Magnesium
Mg
2+
O
2-
and oxygen
Magnesium
and oxygen
Mg
O
2+
2Mg2 O 2
Mg2O2
Both subscripts can
formula is
MgO
be divided by 2 so the final
Silver
I and chlorine
Magnesium and fluorine
Iron II and oxygen
Calcium and phosphorus
Why did we
specify some
oxidation
numbers, but
not others?
State the name of the cation.
1.

2.
(If using a transition metal, you must state
the oxidation number if there is more than
one possibility.)
State the name of the anion, but change
the ending to “ide.”
NaCl
Sodium chlorine chloride
MgO
Magnesium oxide
K2S
Potassium sulfide
Fe2O3
Iron is a transition metal so we need to figure out
the charge before we can name the compound.
We know oxygen is always -2, so there is an overall
charge of -6 from the oxygen
That means Iron must supply an overall charge of
+6
This indicates that iron must have an oxidation
number of +3 in this case
Fe2O3
Iron III oxide
CuS
AgCl*
H2O
(trick!)
State the name of the cation.
1.

2.
(If using a transition metal, you must state
the oxidation number if there is more than
one possibility.)
Name the anion
AgNO3
Silver nitrate
CaCO3
Calcium carbonate
NH4Cl
Ammonium chloride
FeSO4
Iron II sulfate
1.
2.
3.
4.
Write out each ion.
Place oxidation number under each ion
Cross multiply
Reduce to simplest form
Oxyanions are any
polyatomic anions that
contain oxygen
Your book likes to sound fancy!
You
may not, under any conditions, change
the subscripts within the polyatomic ion
when balancing the charge. You may only
adjust the number of units of each
polyatomic ion!!
Use parentheses to remind yourself that the
units go together and cannot be changed.
Potassium
K
permanganate
(MnO4)
Potassium
K
+1
permanganate
(MnO4)
-1
Potassium
K
+1
permanganate
(MnO4)
-1
K 1(MnO4)1
Potassium permanganate is
K(MnO4)
Calcium
Ca
hydroxide
(OH)
Calcium
Ca
+2
hydroxide
(OH)
-1
Calcium
Ca
+2
hydroxide
(OH)
-1
Ca(OH)2
Ca(OH)2
Ammonium
(NH4)
phosphate
(PO4)
Ammonium
(NH4)
+1
phosphate
(PO4)
-3
Ammonium
(NH4)
+1
phosphate
(PO4)
-3
(NH4)3(PO4)
(NH4)3(PO4)
Hydrogen
peroxide
Sodium nitrite
Calcium sulfate
Aluminum hydroxide
A
formula unit gives the ration of cations to
anions in the ionic compound.
A monatomic ion is formed from one atom.
Roman numerals indicate the oxidation
numbers of any element with more than one
oxidation number.
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 outside the parentheses.
Relate
a formula unit of an ionic compound
to its composition
Write formulas for ionic compounds and
oxyanions.
Apply naming conventions to ionic
compounds and oxyanions.
Metals
form crystal lattices
and can be modeled as
cations surrounded by a “sea”
of freely moving valence
electrons.
Describe
a metallic bond
Relate the electron sea model the physical
properties of metals
Define alloys and categorize them into two
basic types.
Physical
property
Metal
Malleable
Electron
sea model
Delocalized electron
Metallic bond
Alloy
Lattice structures with
freely moving
electrons
Electrons are not firmly attached to any one
nucleus, but instead “visit” many nuclei
Attraction of
a metallic cation for
delocalized electrons
Freely
moving electrons are referred to as
“delocalized” (lacking a location)
Video
-
-
+
+
+
-
-
-
+
+
-
+
-
+
-
-
+
+
-
+
-
-
+
-
-
+
-
+
+
-
+
-
-
+
+
+
+
+
-
-
-
-
-
+
+
+
-
+
+
+
-
+
-
+
-
-
+
-
+
Melting
and Boiling points
Vary greatly
Most are moderately high melting points and very
high boiling points
Malleability,
ductility and durability
Nuclei move relatively free of each other due to the
sea of electrons
Thermal
and electrical conductivity
Delocalized electrons quickly move heat from one
part of the metal to other parts
Delocalized electrons can move in one direction
and create a “current.”
Hardness
and strength
The number of delocalized electrons plays a role in
the hardness of the metal
More delocalized electrons means a harder metal
Sometimes d level electrons are delocalized as well as the
s resulting in very hard metals.
Alloy-
a mixture of elements that has
metallic properties
Characteristics may differ from the “parent” metals
Include brass, bronze, 14-carat gold, stainless steel,
etc.
Substitutional
Some of the atoms from one metal are replaced by
atoms of the other metal
Ex: brass
Interstitial
Small holes in the lattice are filled by atoms of
another element
Example: Steel
A
metallic bond forms when metal cations
attract freely moving, delocalized valence
electrons.
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.
Describe
a metallic bond
Relate the electron sea model the physical
properties of metals
Define alloys and categorize them into two
basic types.