Molecular Compounds
John W. Moore
Conrad L. Stanitski
Peter C. Jurs
• Contain 2 or more elements.
http://academic.cengage.com/chemistry/moore
• Form when nonnon-metals combine.
• Nanoscale: a discrete molecule.
Molecular formula
• shows the number and kind of elements used
Chapter 3
Chemical Compounds
e.g.
water
benzene
H2O
C6H6
ammonia NH3
Stephen C. Foster • Mississippi State University
Molecular Compounds
Molecular Formulas
Inorganic compounds
Ethanol has the formula C2H6O …
• Do not contain C or (C and H).
water
H2O
carbon dioxide
• Doesn’t show atom connections.
ammonia NH3
CO2
• A structural formula does.
C2H6O may not be ethanol.
Organic compounds
• always contain C, usually H
• may contain many other elements.
e.g.
benzene C6H6
ethanol
• most (but not all) are molecular.
Molecular Formulas
Condensed formula
Similar information in a compact form.
• Two C2H6O structural formulas:
C2H6O
H H
|
|
H–C–C–O–H
|
|
H H
ethanol
H
H
|
|
H–C–O–C–H
|
|
H
H
dimethyl ether
Molecular Formulas
More elaborate models:
C, what’s attached to it, C …
CH3CH2OH
ethanol
CH3OCH3
dimethyl ether
Groups of atoms attached to C (like OH) are
called functional groups
• –OH is the alcohol group.
Ball-and-stick
model
Space-filling
model
1
Naming Binary Inorganic Compounds
Naming Binary Inorganic Compounds
Binary compounds contain two different elements.
elements
Not monooxide, the
CO
NO2
N2O
P2O5
PBr3
PBr5
SF6
P4O10
• Name the elements in formula order
• Prefixes show the number of atoms present.
 but don’t use “mono” for the first element
1
mono
6
hexa
2
di
7
hepta
3
tri
8
octa
4
tetra
9
nona
5
penta
10
deca
extra vowel is
carbon monoxide
dropped to make it
sound better
nitrogen dioxide
dinitrogen monoxide
Not pentaoxide
diphosphorus pentoxide
phosphorus tribromide
phosphorus pentabromide
sulfur hexafluoride
Not decaoxide
tetraphosphorus decoxide
• Change the 2nd element’s name to end in “--ide”
ide
Naming Binary Inorganic Compounds
Hydrocarbons
Exceptions to these rules:
Binary molecules (C and H) are known
as hydrocarbons.
Hydrogen plus a group 6A or 7A element:
• All prefixes are omitted
HCl
H2S
HF
Alkanes are hydrocarbons with C-C
single bonds only.
hydrogen chloride
hydrogen sulfide
hydrogen fluoride
• Use an –ane ending.
• They exist as linear and branched
molecules
Common names in wide-spread use:
H2O water
NO
nitric oxide
NH3 ammonia
N2O nitrous oxide
PH3 phosphine
N2H4 hydrazine
 formula CnH2n+2
Butane, C4H10
Hydrocarbons
H
H
C
H
H
methane
H
H
H
C
C
H
H
H
ethane
Hydrocarbons
H
H
H
H
C
C
C
H
H
H
H
propane
H
H
H
H
C
C
C
H
C
H
H
H
H
H
# of C
prefix
1
meth
2
eth
3
prop
4
but
5
pent
6
hex
7
hept
8
oct
butane
Similar to molecular
compound prefixes
C8H18 = octane
C5H12 = pentane
Boiling points (°C)
-162
-89
-42
-1
Larger mass = higher b.p.
2
Hydrocarbons
Alkanes and Their Isomers
LineLine-angle structures
Branched alkanes occur.
• Lines represent C-C bonds.
Isomer:
Isomer same formula, different atom arrangement.
• Each junction and end is a C
• Each C needs 4 bonds.
• C–H bonds are omitted
ethane
H H
| |
H―C―C― H
| |
H H
propane
H H H
| | |
H―C―C―C― H
| | |
H H H
C with 3H
(missing 3 bonds)
becomes
C4H10
C with 2 H
(missing 2 bonds)
Alkanes and Their Isomers
Formula
methylpropane
|
CH3CHCH3
isopropyl
isomers
becomes
Alkyl groups
• An alkane with a H atom removed.
• Named by replacing “-ane” with “-yl”
methyl
ethyl
propyl
methylpropane
H H H
| | |
H ―C―C―C― H
Ι
Ι
H
H
H C H
|
H
C4H10
Alkanes and Their Isomers
-CH3
-CH2CH3
-CH2CH2CH3
butane
H H H H
| | | |
H ―C―C―C―C― H
| | | |
H H H H
H H H
| |
|
H―C―C―C― H
Ι
Ι
H
H
H C H
|
H
a methyl
group
Ions and Ionic Compounds
Ions - charged units
- formed by transfer of e- between elements.
Cation = positive ion. Metals form cations
Na
Na+ + e-
Anion = negative ion. Nonmetals form anions
S + 2 eS2-
Isomers
Formula
Isomers
C6H14
CH4
1
C9H20
35
C2H6
1
C10H22
75
C3H8
1
C4H10
2
C12H26
355
C5H12
3
C15H32
4,347
C6H14
5
C20H42
366,319
C7H16
9
C30H62
4.1 x 109
C8H18
18
C40H82
6.3 x 1013
Monatomic Ions
Main group elements
Add/lose enough e- to “get to” the nearest noble gas.
• Charge on ion = group A# or (8-grpA#)
• Number of e- transferred = charge
S
Na
P
Sr
→
→
→
→
S2Na+
P3Sr2+
16 e11 e15 e38 e-
→
→
→
→
18 e10 e18 e36 e-
(like Ar)
(like Ne)
(like Ar)
(like Kr)
3
Monatomic Ions
Monatomic Ions
Transition metals:
metals
• lose varying number of e-.
• old (and new) group number not very helpful.
Ti2+
Cr2+
Fe2+
Cu+
Mn2+
Ti
Cr
Fe
Cu
Mn
(grp 4B)
(grp 6B)
(grp 8B)
(grp 1B)
(grp 7B)
Cr3+
or
or Fe3+
or Cu2+
Mn5+ or Mn7+
Polyatomic Ions
Ionic Compounds
Multiple atom “units” with a net electrical charge.
NH4+
ammonium ion
OH-
hydroxide ion
SO42-
sulfate ion
Common monatomic ions:
Ions are held together by electrostatic forces.
• Cations (+) and anions (–) attract each other.
• Larger charges = larger attraction.
• Larger separation = lower attraction.
• Coulomb’s law:
CN-
cyanide ion
Force between
ions
QQ
F = k 12 2
d
constant
Distance between
ions
Memorize all the ions in table 3.7!
Ionic Compounds
Q1 = charge on ion
1, Q2 = …
Naming Ions and Ionic Compounds
Charges are always balanced.
Positive ions
Most are metal ions (exception: ammonium NH4+ ).
Ionic compounds are always neutral!
• metal ion with only one charge state?
• Use metal name + ion.
ion
Ions
Compound
Mg2+ and FMg2+
Mg2+
Charges
MgF2
(2+) + 2(1-) = 0
and SO4
2-
MgSO4
(2+) + (2-) = 0
and PO4
3-
Mg3(PO4)2
3(2+) + 2(3-) = 0
• metal ion with multiple charge states?
• Use metal name + (Roman numeral) to show charge.
Na+ sodium ion
Fe2+ iron(II) ion
Ca2+ calcium ion
Fe3+ iron(III) ion
4
Naming Ions and Ionic Compounds
Two forms exist: –ate and –ite endings used.
 More oxygen = “-ate”
 Less oxygen = “-ite”
Negative ions
• Monatomic ion?
Increase O
• Add “--ide
ide” to the name stem.
• Polyatomic ion?
• Memorize these.
P
S
phosphorus
sulfur
Oxoanions
P3S2SO32-
phosphide ion
sulfide ion
sulfite ion
Oxoanions
SO42- sulfate ion
NO3- nitrate ion
SO32- sulfite ion
NO2- nitrite ion
If they contain H, add a prefix “hydrogen”
HSO4- hydrogen sulfate ion
(common name=bisulfate ion)
HCO3- hydrogen carbonate ion
(common name=bicarbonate ion)
Naming Ionic Compounds
When four forms exist
 Add “per_____ate
per_____ate” and “hypo____ite
hypo____ite” names
Name the ions and add together…
… cation then anion (drop “ion” from both)
Increase O
Single charge metalmetal-ion examples
FO4
-
FO3
-
perfluorate
fluorate
-
ClO4 perchlorate
-
ClO3 chlorate
FO2- fluorite
ClO2- chlorite
FO-
ClO- hypochlorite
hypofluorite
Naming Ionic Compounds
iron(II) chloride
FeCl3
iron(III) chloride
sodium chloride
KHSO4
potassium hydrogen sulfate
SrO
Mg(OH)2
strontium oxide
KMnO4
potassium permanganate
magnesium carbonate
magnesium hydroxide
Naming Ionic Compounds
Multiple charge examples
FeCl2
NaCl
MgCO3
When are Roman numerals used?
• Main block metals form one type of ion:
 omit Roman numerals.
 exceptions:
exceptions lead (Pb2+, Pb4+), tin (Sn2+, Sn4+)…
Cu2O
copper(I) oxide
CuO
copper(II) oxide
• Transition metals form multiple ions
 use Roman numerals.
 exceptions:
exceptions silver (Ag+), zinc (Zn2+), cadmium
Cu2O
CuO
(Cd2+)…
5
Naming Compounds
Naming Compounds
Generalizations
Metallic element in a formula?
… the compound is usually ionic.
All non-metal formula?
… the compound is usually molecular.
Metalloid in a formula?
… no easy way to tell if ionic or molecular.
Naming Compounds
sulfur trioxide
CuSO4
copper(II) sulfate
AlCl3
aluminum chloride
AgF
silver fluoride
SF6
sulfur hexafluoride
PbO2
lead(IV) oxide
Bonding and Properties of Ionic Compounds
sodium hypochlorite
NaClO
dinitrogen pentoxide
N2O5
potassium dichromate
K2Cr2O7
ammonium perchlorate
NH4ClO4
hydrogen chloride
HCl
Ionic compounds
Not individual molecules. Crystal lattices
• Each ion is surrounded by many others
NaCl
sodium chloride
Formula unit = smallest ratio of anions to cations
Bonding and Properties of Ionic Compounds
Electrostatic forces hold ionic compounds together:
F=k
SO3
Bonding and Properties of Ionic Compounds
Ionic crystals can be cleaved:
Q1Q2
d2
External force
displaces layers
High melting points
 strong forces.
 high charge = high m.p.
Repulsion
occurs
Na+
Cl-
ions
Similar sized ions:
m.p. (°C)
NaF
+1 -1
993
CaO
+2 -2
2572
6
Bonding and Properties of Ionic Compounds
Ionic Compounds in Aqueous Solution: Electrolytes
Ionic compounds are electrical
insulators when SOLID.
If an ionic compound dissolves in
water:
water
• will conduct if molten.
• It dissociates
 breaks apart into its ions.
• It is a strong electrolyte
Many are soluble in water.
 the solution is a good
electrical conductor.
Molecular & Ionic Compounds
Property
Formation
Molecular
Non-metal
combinations
Ionic
Metal/non-metal
combinations
Physical
state
Gases, liquids &
solids. Brittle & weak
or soft & waxy
Crystalline solids
Hard & brittle
mp & bp
Low
High
Conductivity
Poor heat & electrical Poor heat & electrical.
conductors
Good electrical if molten
Solubility
Few soluble in water
Many soluble in water
In solution
Remain molecular
Dissociate
Moles of Compounds
A mole of XmYn contains:
m moles of atom X and n moles of atom Y
1 mol of H2O contains:
2 mol of H atoms and 1 mol of O atoms
Molar mass = sum of the atomic masses
Molar Mass of Ionic Compounds
Mass of 1 water molecule:
= 2(1.008 amu) + 1(15.999 amu) = 18.015 amu
Molar mass of water:
= 2(1.008 g/mol) + 1(15.999 g/mol) = 18.015 g/mol
Gram-Mole Conversions
How many moles of Ca3(PO4)2 are in 10.0 g of the
compound?
Ionic compounds do not contain molecules.
Don’t use “molecular weight” to describe mass.
Formula mass = 3 40.08
Formula weight (or molar mass) should be used
g
mol
+ 8 16.00
Compound
atomic wts
amu
NaCl
22.99 + 35.45
Ca(NO3)2 40.08+2(14.01)+6(16.00)
Formula wt.
amu
Molar mass
g/mol
58.44
58.44
164.10
164.10
= 310.18
+ 2 30.97
g
mol
g
mol
g
mol
Moles of Ca3(PO4)2 = 10.0 g
1 mol
= 0.0322 mol
310.2 g
7
Gram-Mole Conversions
Moles of Ionic Hydrates
Find the mass of cobalt in 3.49 g of cobalt(II) sulfate.
Ionic hydrate: ionic compound with water trapped in the
crystal.
• the water of hydration.
• use “hydrate” with a Greek prefix for the number.
• heat can remove some, or all, of this water.
Formula wt CoSO4 = 58.93 + 32.07 + 4(16.00)
= 155.00 g
3.49 g CoSO4
1 mol CoSO4
1 Co
155.0 g CoSO4 1 CoSO4
= 0.02252 mol Co
0.02252 mol Co
58.93 g Co
1 mol Co
= 1.33 g Co
Percent Composition
Examples
MgSO4•7H2O magnesium sulfate heptahydrate
(Epsom salt).
CuSO4•5H2O
copper(II) sulfate pentahydrate.
Percent Composition
Two names used:
• percent composition by mass,
mass or
%Na =
• mass percent of the compound.
Example
What is the mass percent of each element in sodium
chlorite, NaClO2?
molar mass = (22.990 g) + (35.453 g) + 2(15.999 g)
= 90.441 g
Percent Composition
%Cl =
mass of Cl …
x 100 %
mass of NaClO2 …
= 35.453 g x 100 % = 39.20%
90.441 g
mass of Na in 1 mol NaClO2
x 100 %
mass of NaClO2 in 1 mol NaClO2
= 22.990 g x 100 % = 25.42%
90.441 g
%O =
mass of O …
x 100 %
mass of NaClO2 …
= 2(15.999) g x 100 % = 35.38%
90.441 g
Determining Empirical and Molecular Formulas
Last example:
molecular formula
percent composition
The process can be reversed:
percent composition
empirical formula
Not molecular
formula
Check your work:
%Na + %O + %Cl = 25.42 + 35.38 + 39.20 = 100%
Empirical formula = the simplest ratio of atoms in a
compound.
8
Determining Empirical and Molecular Formulas
Examples
Compound
mol. formula
emp. formula
Determining Empirical and Molecular Formulas
Example
An orange compound was 26.6% K, 35.4% Cr and
38.0% O. Determine its empirical formula.
hydrogen peroxide
H2 O 2
HO
borane (boron trihydride)
BH3
BH3
diborane (diboron hexahydride)
B2H6
BH3
octene
C8H16
CH2
• Divide each mass by its atomic mass.
 Gives the number of moles of each (in 100 g).
butene
C4 H8
CH2
• Divide each by the smallest answer found.
• Assume a 100.0 g sample.
 % becomes mass in grams
 The smallest integer ratio = empirical formula.
Determining Empirical and Molecular Formulas
Unknown: 26.6% K
35.4% Cr
38.0% O
Determining Empirical and Molecular Formulas
Empirical formula = smallest integer ratio.
Divide by the smallest value
(ratios stay the same!)
In 100.0 g
26.6 g K
35.4 g Cr
38.0 g O
1 mol K
39.10 g K
= 0.6803 mol K
1 mol Cr
52.00 g Cr
= 0.6808 mol Cr
1 mol O
16.00 g O
= 2.375 mol O
K
0.6803 mol = 1.000
0.6803 mol
x2
2
Cr
0.6808 mol = 1.001
0.6803 mol
x2
2
O
2.375 mol = 3.491
0.6803 mol
x2
7
Choose a multiplier to make integer
Determining Empirical and Molecular Formulas
The molecular formula can be determined if the
molecular mass is known.
Example
Vitamin C has the empirical formula C3H4O3 and
molecular mass = 175 g/mol.
The empirical formula is K2Cr2O7
The Biological Periodic Table
Element in the
body
Symbol
Abundance
atoms/106 atoms
Hydrogen
H
630,000
Oxygen
O
255,000
Carbon
C
94,500
Nitrogen
N
13,500
Calcium
Ca
3,100
Phosphorus
P
2,200
Empirical mass:
3(12.01) + 4(1.008) + 3(15.99) = 88.03 g/mol
Chlorine
Cl
570
Sulfur
S
490
Na
410
Empirical mass ≈ ½(molecular mass)
Mol. formula = 2(emp. formula) = C6H8O6
Potassium
K
260
Magnesium
Mg
130
Sodium
98.0%
99.3%
0.7%
9
The Biological Periodic Table
10