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
• In forming compounds, atoms tend to achieve the
electron configuration of a noble gas
•ns2np6
•Atoms of metals tend to lose their valence electrons
leaving a complete octet in the next lowest energy
level
•Atoms of non-metals tend to gain electrons or to
share electrons with another non metal to achieve a
complete octet
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.
http://www.youtube.com/watch?v=lODqdh
xDtHM&feature=related
 Electron loss or ionization of sodium atom
 Na 1s22s22p63s1 → Na+ 1s22s22p6
Chlorine
 A gain of one electron gives chlorine an octet and
converts a chlorine atom into a chloride ion.
 It has the same electron configuration as argon
Gain of valence
electrons
1. How many valence electrons does each of
the following atoms have?
a. gallium
b. fluorine
c. selenium
2. For each element below, state (i) the
number of valence electrons in the atom,
(ii) the electron dot structure, and (iii) the
chemical symbol(s) for the most stable
ion.
a. Ba
b.
I
c.
K
3. Write the electron configuration for each of
the following atoms and ions.
a. K atom
b. K ion
c. Na atom
d. Na ion
e. Phosphorous atom
f. Phosphide ion
4. How many electrons will each element gain or
lose in forming an ion? State whether the resulting
ion is a cation or an anion.
a. strontium (Sr)
b. tellurium (Te)
c. Bromine (Br)
d. aluminum (Al)
e. rubidium (Rb)
f. Phosphorus (P)
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
Examples
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
Example
1:
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
Example
2:
Iron III and oxygen
Fe
O
3+
2-
Iron
III and oxygen
Fe
O
3+
2-
Fe2 O 3
Fe2O3
Cannot be reduced
Example
3:
Magnesium and oxygen
Mg
O
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
Antimony (V) oxide
Aluminum sulfide
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!)
Monatomic
Ex. Mg 2+
ion: a one atom ion
Polyatomic
ion: ions made up of more than
atom
Ex. PO4 3-
Oxyanions-
a polyatomic ion composed of
an element, usually a nonmetal, bonded to
one or more oxygen atoms.
The
ion with the greater number of oxygen
atoms ends in –ate.
The ion with the fewer number of oxygen
atoms ends in –ite.
Ex. NO3NO2Nitrate
Nitrite
Directions:
Draw 5 by 5 chart.
Write 5 metals ions in the left margin
Write 5 nonmetal ions on the top
Trade charts with the person next to you
Fill in the charts by writing the correct formulas for
the ionic compounds formed in the square.
Hand the chart back to your partner when your
done and check each others responses.
Once finished, raise your hand and ask Mrs. Chebib
to come over and stamp your work for credit
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)
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
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+
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+
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+
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.