2
Chemical Formulas
and Composition
Stoichiometry
1
Chemical Formulas
• Show the chemical composition of the
substance containing 2 or more elements.
•
•
•
•
He, Au, Na – monatomic elements
O2, H2, Cl2 – diatomic elements
O3, P4, S8 - more complex elements
H2O, C12H22O11 – compounds
2
Chemical Formulas
Compound 1 Molecule Contains
HCl
H2O
NH3
C3H8
1 H atom & 1 Cl atom
2 H atoms & 1 O atom
1 N atom & 3 H atoms
3 C atoms & 8 H atoms
3
Ions and Ionic Compounds
• Ions are atoms or groups of atoms that
possess an electric charge.
• Two basic types of ions:
– Positive ions or cations
• one or more electrons less than neutral
• Na+, Ca2+, Al3+
• NH4+ - polyatomic cation
– Negative ions or anions
• one or more electrons more than neutral
• F-, O2-, N3• SO42-, PO43- - polyatomic anions
4
Ions and Ionic Compounds
• Sodium chloride
– table salt is an ionic compound
5
Names and Formulas of
Some Ionic Compounds
• Common Ion Sheet: Quiz Wednesday 9/9
• Some examples are:
– Anions - Cl1-, OH1-, SO42-, PO43– Cations - Na1+, NH41+, Ca2+, Al3+
• Some examples are:
– H2SO4 - sulfuric acid
– FeBr2 - iron(II) bromide
– C2H5OH - ethanol
6
Names and Formulas of
Some Ionic Compounds
• Ionic formulas are determined by the charges of
the ions.
– Charge on the cations=charge on the anions.
– The compound must be neutral.
•
•
•
•
NaCl
KOH
CaSO4
Al(OH)3
sodium chloride
(Na1+ & Cl1-)
potassium hydroxide(K1+ & OH1-)
calcium sulfate
(Ca2+ & SO42-)
aluminum hydroxide (Al3+ & 3 OH1-)
7
Names and Formulas of
Some Ionic Compounds
•
•
•
•
•
•
You do it!
What is the formula of nitric acid?
HNO3
What is the formula of sulfur trioxide?
SO3
What is the name of FeBr3?
iron(III) bromide
8
Names and Formulas of
Some Ionic Compounds
•
•
•
•
•
•
You do it!
What is the name of K2SO3?
potassium sulfite
What is charge on sulfite ion?
SO32- is sulfite ion
What is the formula of ammonium sulfide?
(NH4)2S
9
Names and Formulas of
Some Ionic Compounds
•
•
•
•
•
•
You do it!
What is the charge on ammonium ion?
NH41+
What is the formula of aluminum sulfate?
Al2(SO4)3
What are the charges on both ions?
Al3+ and SO4210
7
Chemical Bonding
Introduction
• Attractive forces that hold atoms together
in compounds are called chemical bonds.
• The electrons involved in bonding are
usually those in the outermost (valence)
shell.
12
Introduction
• Chemical bonds are classified into two types:
o Ionic bonding electrostatic attractions among
ions, formed by the transfer of one or more
electrons from one atom to another.
o Covalent bonding results from sharing one or
more electron pairs between two atoms.
13
Comparison of Ionic and
Covalent Compounds
• Melting point comparison
– Ionic compounds usually have high MP
• Typically > 400oC
– Covalent compounds usually have low MP
• Typically < 300oC
• Solubility in polar solvents
– Ionic compounds are generally soluble
– Covalent compounds are generally insoluble
14
Comparison of Ionic and
Covalent Compounds
• Solubility in polar solvents
– Ionic compounds are generally soluble
– Covalent compounds are generally insoluble
• Solubility in nonpolar solvents
– Ionic compounds are generally insoluble
– Covalent compounds are generally soluble
15
Comparison of Ionic and
Covalent Compounds
• Conductivity
– Ionic compounds generally conduct electricity
• They contain mobile ions
– Covalent compounds generally do not
conduct electricity
16
Comparison of Ionic and
Covalent Compounds
• Formation of Compounds
– Ionic compounds: large EN differences
• Often a metal and a nonmetal
– Covalent compounds: small differences
in EN
• Usually two or more nonmetals
17
Lewis Dot Formulas
of Atoms
• Lewis dot formulas or Lewis dot
representations are used for tracking
valence electrons.
– Valence electrons are those electrons that are
transferred or involved in chemical bonding.
• They are chemically important.
18
Lewis Dot Formulas
of Atoms
....
H
H
H
H
....
Li
Li
Li
Li
....
..
Be
Be
Be
Be
....
.. ..
B
B
B.
B
....
.. ..
...C
C
C.
C
....
.. N
..
N
N
.
.N
....
.. ..
O
..O
O
.
.O
....
.. ..
F
....
F
F
F
..
....
....
He
He
He
He
..
.
.
Ne
. Ne
Ne.
Ne
..
19
Lewis Dot Formulas
of Atoms
• Elements that are in the same periodic
group have the same Lewis dot structures.
.
.
Li & Na
..
..
. N. & .P .
.
.
..
..
. .
. .
. ..F & . Cl
..
20
Ionic Bonding
•
An ion is an atom or a group of atoms
possessing a net electrical charge.
• Ions come in two basic types:
1. positive (+) ions or cations
•
These atoms have lost 1 or more electrons.
2. negative (-) ions or anions
•
These atoms have gained 1 or more electrons.
21
Formation of
Ionic Compounds
• Monatomic ions consist of one atom.
• Examples:
– Na+, Ca2+, Al3+ - cations
– Cl-, O2-, N3- -anions
• Polyatomic ions contain more than one
atom.
– NH4+ - cation
– NO2-,CO32-, SO42- - anions
22
Formation of
Ionic Compounds
• Ionic bonds are formed by the attraction
between cations and anions, usually, to
form solids.
• Commonly, metals react with nonmetals to
form ionic compounds.
• The formation of NaCl is one example of
an ionic compound formation.
23
Formation of
Ionic Compounds
• Reaction of Group IA Metals with Group
VIIA Nonmetals
IA metal VIIA nonmetal
2 Li (s)  F2(g)
silver
yellow
solid
gas
24
Formation of
Ionic Compounds
• Reaction of Group IA Metals with Group
VIIA Nonmetals
IA metal VIIA nometal
2 Li (s)  F2(g)  2 LiF(s)
silver
solid
yellow
gas
white solid
o
with an 842 C
melting point
25
Formation of
Ionic Compounds
• 1s 2s
2p
Li  
F   
These atoms form ions with these configurations.
Li+ 
same configuration as [He]
F-    same configuration as [Ne]
26
Formation of
Ionic Compounds
• We can also use Lewis dot formulas to
represent the neutral atoms and the ions
they form.
Li .
+
..
.. .
F
..
Li
+
..
.. ..
F
..
[ ]
27
Formation of
Ionic Compounds
• The Li+ ion contains two electrons, same
as the helium atom.
– Li+ ions are isoelectronic with helium.
• The F- ion contains ten electrons, same as
the neon atom.
– F- ions are isoelectronic with neon.
28
Formation of
Ionic Compounds
• The reaction of potassium with bromine is
another example of a group IA metal with
a Group VIIA non metal.
– Write the reaction equation.
You do it!
IA metal VIIA nonmetal
2 K (s)

Br2( )  2 KBr(s)
ionic solid
29
Formation of
Ionic Compounds
4s
4p
K [Ar] 
Br [Ar]     and the 3d electrons
The atoms form ions with these electronic structures.
4s
K+
Br-
4p
same configuration as [Ar]

  
same configuration as [Kr]
30
Formation of
Ionic Compounds
• Write the Lewis dot formula representation for
the reaction of K and Br.
You do it!
K.
+
..
.. .
Br
..
K
+
..
.. Br
..
..
[ ]
31
Formation of
Ionic Compounds
• Cations become isoelectronic with the
preceding noble gas.
• Anions become isoelectronic with the
following noble gas.
32
Formation of
Ionic Compounds
• In general for the reaction of IA metals and
VIIA nonmetals, the reaction equation is:
2 M(s) + X2  2 M+ X-(s)
– where M is the metals
– X is the nonmetals
ns
np
M
X


  
ns
 M+
 X- 
np
  
33
Formation of
Ionic Compounds
• IIA metals with VIIA nonmetals.
• Forms mostly ionic compounds.
– Notable exceptions are BeCl2, BeBr2, and
BeI2 which are covalent compounds.
• One example is the reaction of Be and
F 2.
Be(s) + F2(g) BeF2(g)
34
Formation of
Ionic Compounds
2s
2p
2s
2p
Be [He] 
 Be2+ [He]
F [He]      F- [He]   
Next, draw the Lewis dot formula representation
of this reaction.
You do it!
35
Formation of
Ionic Compounds
..
. F ..
.. .
2+
..
.
.
Be
Be .
2 .F.
..
..
. F ..
..
• Symbolically this can be represented as:
M(s) + X2  M2+ X2M can be any of the metals
X can be any of the nonmetals
36
Formation of
Ionic Compounds
• IA metals with VIA nonmetals
• The reaction equation is:

2
4 Li (s)  O 2(g)  2 Li O
2s 
37
Formation of
Ionic Compounds
• Draw the electronic configurations for Li, O, and
their appropriate ions.
You do it!
2s
2p
2s
2p
Li [He] 
 Li+1 [He]
O [He]    
 O-2 [He]    
Draw the Lewis dot formula representation of
this reaction.
You do it!
38
Formation of
Ionic Compounds
Li .
Li .
+
..
. O ..
.
Li
Li
+
+
. .. 2-.
.O
.. .
• Symbolically this can be represented as:
2 M (s) + X  M21+ XM can be any of the metals
X can be any of the nonmetals
39
Formation of
Ionic Compounds
Simple Binary Ionic Compounds Table
• Reacting Groups
Compound General Formula
IA + VIIA
IIA + VIIA
IIIA + VIIA
IA + VIA
IIA + VIA
IIIA + VIA
MX
MX2
MX3
M2X
MX
M2X3
Example
NaF
BaCl2
AlF3
Na2O
BaO
Al2S3
40
Formation of
Ionic Compounds
• Reacting Groups Compound General Formula
IA + VA
IIA + VA
IIIA + VA
M3X
M3X2
MX
Example
Na3N
Mg3P2
AlN
H, a nonmetal, forms ionic compounds with IA and IIA
metals for example, LiH, KH, CaH2, and BaH2.
Other hydrogen compounds are covalent.
41
Formation of
Ionic Compounds
• Attraction of positive ions for negative ions
due to the opposite charges.

q q 
F


2
d
where
F  force of attraction between ions
q  magnitude of charge on ions
d  distance between center of ions
42
Formation of
Ionic Compounds
• Small ions with high ionic charges have
large Coulombic forces of attraction.
• Large ions with small ionic charges have
small Coulombic forces of attraction.
43
Covalent Bonding
• Covalent bonds are formed when atoms
share electrons.
• 2 electrons=single covalent bond
• 4 electrons=double covalent bond
• 6 electrons=triple covalent bond
The atoms have a lower potential energy when bound.
44
Formation of
Covalent Bonds
• This figure shows the potential
energy of an H2 molecule as a
function of the distance between
the two H atoms.
45
Formation of
Covalent Bonds
• Representation of
the formation of an
H2 molecule from
H atoms.
46
Formation of
Covalent Bonds
• We can use Lewis dot formulas to show
covalent bond formation.
1. H molecule formation representation.
H.
+
H.
H .. H or H2
2. HCl molecule formation
H.
+
..
. Cl ..
..
..
.
.
.
. or HCl
H Cl
..
47
Bond Lengths and Bond Energies
• For any covalent bond there is an
internuclear distance where the attractive
and repulsive forces balance
• This distance is the bond length
48
Bond Lengths and Bond Energies
• More stable than the separated atoms by
an amount of energy called bond energy.
49
Writing Lewis Formulas:
The Octet Rule
• N - A = S rule
– Simple mathematical relationship to help us write Lewis dot
formulas.
• N = number of electrons needed
– N usually has a value of 8 for representative elements.
– N has a value of 2 for H atoms.
• A = number of electrons available
– A is equal to the periodic group number for each element.
– A is equal to 8 for the noble gases.
• S = number of electrons shared in bonds.
• A-S = number of electrons in unshared, lone, pairs.
50
Lewis Formulas for Molecules
and Polyatomic Ions
• First, we explore Lewis dot formulas of
homonuclear diatomic molecules.
– Two atoms of the same element.
1. Hydrogen molecule, H2.
.
H. H
H H
or
2. Fluorine, F2.
.. ..
.
.
.
. F . F .
.. ..
or
..
.
.F
..
..
.
.
F
..
3. Nitrogen, N2.
·· N ·· ·· ·· N ··
or
·· N N ··
51
Lewis Formulas for Molecules and
Polyatomic Ions
•
Next, look at heteronuclear diatomic
molecules.
–
Two atoms of different elements.
•
1.
2.
Hydrogen halides are good examples.
hydrogen fluoride, HF
. ·· ·
H. F ·
··
hydrogen chloride, HCl
. ·· ·
H . Cl ·
··
3.
or
or
··
H F ··
··
··
H Cl··
··
hydrogen bromide, HBr
··
. ·· ·
H . Br· or H Br··
··
··
52
Lewis Formulas for Molecules
and Polyatomic Ions
• Now we will look at a series of slightly
more complicated heteronuclear
molecules.
• Water, H2O
··
H ·· O ··
··
H
53
Lewis Formulas for Molecules
and Polyatomic Ions
• Ammonia molecule , NH3
··
H ·· N ·· H
··
H
54
Lewis Formulas for Molecules
and Polyatomic Ions
• Lewis formulas can also be drawn for
polyatomic ions.
• One example is the ammonium ion , NH4+.
H +
··
H ·· N ·· H
··
H
•Notice that the atoms other than H
in these molecules have eight
electrons around them.
55
Writing Lewis Formulas:
The Octet Rule
• The octet rule states that representative
elements usually attain stable noble gas
electron configurations in most of their
compounds.
• Lewis dot formulas
• We need to distinguish between bonding
(or shared) electrons and nonbonding (or
unshared or lone pairs) of electrons.
56
Writing Lewis Formulas:
The Octet Rule
• For ions we must adjust the number of electrons
available, A.
– Add/subtract one e- for each charge.
• The central atom in a molecule or polyatomic ion
is determined by:
– Largest number of electrons needed.
– Same periodic group, the less electronegative.
57
Writing Lewis Formulas:
The Octet Rule
Example 7-2: Write Lewis dot and dash formulas
for hydrogen cyanide, HCN.
58
Writing Lewis Formulas:
Limitations of the Octet Rule
1.
2.
3.
4.
The covalent compounds of Be.
The covalent compounds of the IIIA Group.
Odd number of electrons.
Central element must have a share of more than 8
valence electrons.
5. d- and f-transition metals.
59
Formal Charge
• The formal charge is the hypothetical
charge.
• The best Lewis structures will have formal
charges on the atoms that are zero or
nearly zero.
60
Formal Charge
Rules for Assigning Formal Charge
1. Formal Charge = grp number – (number of
bonds + number of unshared e-)
2. Same number of bonds as periodic group =
formal charge of 0.
3. a. Molecules must have a formal charge of 0.
b. The formal charges of elements must sum
to the ion’s charge for polyatomics.
61
Formal Charge
Consider nitrosyl chloride, NOCl
Cl N O
Cl
N
O
7 – (2+4) = +1
5 – (3+2) = 0
6 – (1+6) = -1
Cl N O
Cl
N
O
7 – (1+6) = 0
5 – (3+2) = 0
6 – (2+4) = 0
62
Resonance
·· O S
··
·· O ·
·· ·
·· ·
O·
··
··
·· O
··
S
·· O ··
··
O ··
··
··
·· O
S O ··
··
··
·· O ··
··
When two or more Lewis formulas are necessary, we must
use equivalent resonance structures to show structure.
Double-headed arrows are used to indicate resonance
formulas.
63
Polar and Nonpolar Covalent
Bonds
• Covalent bonds in which the electrons are
shared equally are nonpolar.
• To be nonpolar the two atoms involved in
the bond must be the same element to
share equally.
64
Polar and Nonpolar Covalent
Bonds
• Some examples of nonpolar covalent
bonds.
• H2 H .. H or H H
• N2
·· N ·· ·· ·· N ··
or
·· N N ··
65
Polar and Nonpolar Covalent
Bonds
• Covalent bonds in which the electrons are
not shared equally are polar
– Have an asymmetrical charge distribution
• To be a polar covalent bond the two atoms
involved in the bond must have different
electronegativities.
66
Polar and Nonpolar Covalent
Bonds
• Some examples of polar covalent
bonds.
H
F
• HF
Electroneg ativities 2.1
4.0


1.9
Difference  1.9 very polar bond
67
Polar and Nonpolar Covalent
Bonds
• Shown below is an electron density map of
HF.
– Blue areas indicate low electron density.
– Red areas indicate high electron density.
68
Polar and Nonpolar Covalent
Bonds
• Compare HF to HI.
H
Electroneg ativities
I
2.1
2.5


0.4
Difference  0.4 slightly polar bond
69
Polar and Nonpolar Covalent
Bonds
– Notice that the charge separation decreases as
we move from HF to HI.
70
Polar and Nonpolar Covalent
Bonds
• Polar molecules can be attracted by
magnetic and electric fields.
71
Dipole Moments
• Centers of positive and negative charge do not
coincide
• Have an asymmetric charge distribution
• Polar.
• The dipole moment has the symbol .
•  is the product of the distance,d, separating charges
of equal magnitude and opposite sign, and the
magnitude of the charge, q.
72
Dipole Moments
• Molecules that have a small separation
of charge have a small .
• Molecules that have a large separation
of charge have a large .
• For example, HF and HI:

  H - F 1.91 Debye units

  H -I 0.38 Debye units
73
Dipole Moments
•
•
There are some nonpolar molecules that have
polar bonds.
There are two conditions that must be true for a
molecule to be polar.
1. Must be at least one polar bond or one lone pair of
electrons.
2. The polar bonds and lone pairs must be arranged so
that their dipole moments cannot cancel out one
another.
74
The Continuous Range of
Bonding Types
•
Covalent and ionic bonding represent two
extremes.
1.
2.
•
•
Covalent bonds: e- are equally shared
In pure ionic bonds: e- are completely lost or
gained by one of the atoms
Most compounds fall somewhere between
these two extremes.
The greater the electronegativity
differences the more polar the bond.
75