Chemical Bonding
When two atoms are joined to make a chemical compound, the force of attraction between the two atoms is the chemical bond .
• Ionic bonding , electrons are transferred before bond formation, forming an ion pair.
• Covalent bonding , electrons are shared between atoms in the bonding process.
• Polar covalent bonding , like covalent bonding, is based on the concept of electron sharing; however, the sharing is unequal and based on the electronegativity difference between joined atoms
Lewis Symbols
The Lewis symbol, showing only valence electrons, is a convenient way of representing atoms to show their chemical bonding pattern. It shows how many electrons will be shared , or lost/gain by an atom. If too many electrons are to be lost/ gained or shared atoms tend to share electrons forming covalent bonding.
Principal Types of Chemical Bonds: Ionic and
Covalent
Ionic bonds - electrostatic forces that exist between ions of opposite charge
• typically involves a metal with a nonmetal
Sodium metal reacts with chlorine (non-metal) gas in a violently exothermic reaction to produce NaCl (composed of Na + and Cl ions):
2Na( s ) + Cl
2
( g ) -> 2NaCl( s )
Problem : Predict number and kid of ions, and the formulas of the compounds formed from the combination of ions of the following elements a.
Potassium and Chlorine b.
Magnesium and Bromine c.
Manesium and Nitrogen
Answer : Metal loses electrons and non-metal gain electrons. From the Lewis symbol one can tell how many electrons are lost or gained. To get the formula of ionic compound one need to balance the opposite charges. If they are equal already formula has 1:1 anions and cation like in K+ and Cl-, therefore formula become KCl. If charges are different like in Mg2+ and N3- to get the formula usually cross multiply with charges to obtain 3 Mg2+ and 2 N3-) and drop the charges and write formula Mg
3
N
2
. a.
(one K+ and one Cl-) KCl

b.
(one Mg2+ and two Br-) MgBr
2 c.
(three Mg2+ and two N3-) Mg
3
N
2
Covalent bonds - results from the sharing of electrons between two atoms
• typically involves two nonmetallic elements
Homo-nuclear covalent bonds- Covalent bond between identical atoms are Non-polar covalent bonds there is no charge separation.
The diatomic hydrogen molecule (H
2
) is the simplest model of a covalent bond, and is represented in Lewis structures as:
The shared pair of electrons provides each hydrogen atom with two electrons in its valence shell (the 1 s ) orbital.
In a sense, it has the electron configuration of the noble gas helium
When two chlorine atoms covalently bond to form Cl
2
, the following sharing of electrons occurs:
Each chlorine atom shared the bonding pair of electrons and achieves the electron configuration of the noble gas argon.
In Lewis structures the bonding pair of electrons is usually displayed as a line , and the unshared electrons as dots:
The shared electrons are not located in a fixed position between the nuclei. In the case of the H
2
compound, the electron
Polar Covalent Bonding and Electronegativity
Hetero-nuclear covalent bonds- Covalent bond between two different atoms leads to polar covalent bonds with a charge separation depending of the electronegativities of atoms.
Electronegativity

" Electronegativity is the power of an atom when in a molecule to attract electrons to itself." The electronegativity will depend upon a number of factors including other atoms in the molecule. There are a number of ways to produce a set of numbers which represent electronegativity scales. The Pauling scale is perhaps the most famous .
Non-polar covalent bonds : Electron negativity difference between 0.2 - 0.5 to indicate non-polar covalent.
CS
2
∆ EN 2.5-2.5= 0, non-polar CH
4
∆ EN 2.5-2.1 = 0.4, non-polar
Polar Covalent bonds: Electron negativity difference between 0.5 - 1.6 to indicate polar covalent. This type of bond occurs when there is unequal sharing (between the two atoms) of the electrons in the bond. Molecules such as NH
3
and H
2
O are the usual examples.
CS
2
∆ EN -2.1= 1.4, polar covalent CH
4
∆ EN 3.0-2.1 = 0.9, polar covalent

Problem: Classify following bonds in compounds (NaCl , CO, ICl and H
2 covalent .
)as ionic or
Answer: a. Ionic (NaCl is formed from a metal and a nonmetal.) b. Covalent (CO is formed from two nonmetals.) c. Covalent (ICl is formed from two nonmetals.) d. Covalent (H2) is formed from two identical atoms.)
Problem: Predict the whether following bond between atoms will be non-polar, polar or ionic.
Answer : a. Cl and Cl non-polar; since there is no electronegativity difference between the atoms sharing the electrons, the electrons are shared equally. b. H and H non-polar; there is no electronegativity difference between the atoms sharing the electrons, the electrons are shared equally. c. C and H non-polar or only very slightly polar; the electronegativity difference between carbon and hydrogen (0.4) is negligibly small, thus the electrons are shared essentially equally. d. Li and F ionic; the electronegativity difference between fluorine and lithium
(3.0) is large and also the bond is between a metal and a nonmetal. e. O and O non-polar; since there is no electronegativity difference between theatoms sharing the electrons, the electrons are shared equally.
Ionic bond : Electron negativity difference between 1.6 or higher indicate ionic bond.
If the ∆ EN is between 1.6 and 2.0 and if a metal is involved, then the bond is considered ionic. If only nonmetals are involved, the bond is considered polar covalent. This type of bond occurs when there is complete transfer (between the two atoms) of the electrons in the bond. Substances such as NaCl and MgCl
2
are the usual examples.
Naming Compounds and Writing Formulas of Compounds
A substance is given systematic name of substance according to certain rules. Before the rules are made common names was given without following systematic rules. The
"shorthand" symbol for a compound is its formula. Formula gives types atoms and number each one in the Chemical compound.
Naming Ionic Compounds
Names of ionic compounds are based on names of the ions making them. Ions are classified as monatomic or polyatomic.
In the Stock system a Roman numeral indicates the charge of the cation. This system is preferred over the older "common nomenclature" system.
Charges on monatomic ions of metals and nonmetals
Symbols and Names of monoatomic ions:
"Representative = Fixed Charge " ions:

Symbol
H +
Li +
Na +
K +
Be 2+
Mg 2+
Ca 2+
Ba 2+
Zn 2+
Name
Hydrogen ion
Lithium ion
Sodium ion
Potassium ion
Beryllium ion
Magnesium ion calcium ion barium ion zinc ion
Symbol Name
H Hydride ion
F Fluoride ion
Cl -
Br -
I -
O 2-
S 2-
N 3-
P 3-
Chloride ion
Bromide ion
Iodide ion
Oxide ion
Sulfide ion
Nitride ion
Phosphide ion
"Variable Charge" Cations
Symbol (Stock system) Common Symbol
(Stock system) Common
Cu +
Cu 2+
Fe 2+
Fe 3+ copper(I) copper(II) iron(II) iron(III) cuprous Hg
2
2+ cupric ferrous ferric
Hg
Pb
Pb
2+
2+
4+ mercury(I) mercurous mercury(II) mercuric lead(II) lead(IV) plumbous plumbic
Sn 2+
Sn 4+
Cr 2+
Cr 3+ tin(II) tin(IV) stannous stannic
Co
Co
2+
3+ chromium(II) chromous Ni 2+ cobalt(II) cobaltous cobalt(III) cobaltic nickel(II) nickelous nickel(IV) nickelic
Mn 2+
Mn 3+ chromium(III) chromic Ni 4+ manganese(II) manganese(III) manganou s manganic
Au
Au
+
3+ gold(I) gold(III)
Symbols and Charges for Polyatomic Anions
Formul a Name Formula Name aurous auric
NO
3
-
NO
2
-
CN nitrate nitrite cyanide
MnO
4
permanganate
CO
3
2-
SO
4
2-
SO
3
2-
PO
4
3carbonate sulfate sulfite phosphate
OH -
O
2
2hydroxide peroxide
HCO
3
hydrogen carbonate(bi
PO
3
3-
ClO
4
ClO
3
-
phosphite perchlorate chlorate

carbonate)
HSO
4
hydrogen sulfate (bisulfate)
HSO
3
hydrogen sulfite (bisulfite)
HPO
4
2hydrogen phosphate
H
2
PO
4
dihydrogen phosphate
ClO
2
ClO
CrO
4
-
Cr
2
O
7
-
2-
2chlorite hypochlorite chromate dichromate
C
2
H
3
O
2
- acetate
Writing Formulas of Ionic Compounds
For ionic compounds, the name of the positive ion (cation) is given first, followed by the name of the negative ion (anion). There for conversion of name to formula is easy if you know the metal and nonmetal ion symbols and charges. Use the periodic table to decide the charge on both the cation and anion (or the tables) and determine the formula of the compound(s) formed in each case. For transition metals the common ionic charges are given in after the metal name in parenthesis.
Writing basic ionic compound formulas.
Examples: lithium sulfide; lithium =Li +1 ; sulfide =S -2
Write ions on a line: Li +1 S -2
Then remove cation and anion charges and exchange them without charge as subscripts on the metal and nonmetal
Li +1 S -2 becomes Li
2
S
1
Remember we omit "1" from the subscript formula becomes Li
2
S
Problem : What is the formula of the following compounds given their names? a.
Potassium chloride b.
Magnesium bromide c.
Magnesium nitirde
Answer : First get the formula of ions in the compound. Potassium consists of cation K + and chloride Cl . Look in the table to get charges on the ions and one need to balance the opposite charges. If charges are equal already formula has 1:1 anions and cation like in K+ and Cl-, therefore formula become KCl. If charges are different like in Mg2+ and N3- to get the formula usually cross multiply with charges to obtain 3 Mg2+ and
2 N3-) and drop the charges and write formula Mg
3
N
2. a.
Potassium chloride (one K+ and one Cl-) KCl b.
Magnesium bromide (one Mg2+ and two Br-) MgBr
2 c.
Magnesium nitride (three Mg2+ and two N3-) Mg
3
N
2
Problem: Give formula of following ionic compounds a) sodium chloride b) aluminum phosphate c) magnesium fluoride d) potassium nitrate e) calcium sulfate f) mercury(II) chloride g) iron(II) chloride h) cobalt(III) nitrate i) potassium chromate

Answers: a) NaCl b) AlPO
4 c) MgF
2 d) KNO
3 e) CaSO
4 f) Hg Cl
2 g) FeCl
3 h) Co(NO
3
)
3 i) KmnO
4
Problem: Give names of following ionic compounds a) iron(II) bromide b) copper(II) sulfate c) Sodium phospate d) Sodium sulfite e) Iron (II) nitrate f) lithium carbonate g) Gold (II) chloride h) calcium bisulfate i) potassium bicarbonate
Answers : a) FeBr
2
CuSO
4
PO
4 d) Na
2
SO
3
Fe(NO
3
)
2 f) Li
2
CO
3 g) AuCl
2
Ca(HSO
4
)
2
KHCO
3
Naming Bases
Most bases have a formula that ends with OH. Naming is similar to naming salts or ionic compounds: Name the metal and then the OH, hydroxide. Subscripts are obtained by cation charge and charge on the hydroxide ion, OH .
NaOH
Ba(OH)
2 sodium hydroxide barium hydroxide
KOH
NH
4
Ca (OH)
2 potassium hydroxide hydroxide calcium hydroxide
Covalent Compounds
Most covalent compounds are formed by the reactions of non-metals with another non-metal. Covalent compounds exist as molecules and are named using prefixes that denote the number of each element present in the compound.
Prefixes used:
1 mono
2 di
6
7 hexa hepta
3 tri
4 tetra
8
9 octa nona
5 penta 10 deca
Many familiar covalent compounds have common names. It is useful to correlate both systematic and common names with the corresponding molecular formula.
Problem : Give the names of following formulas from the names of following covalent compunds names:

a.
b.
d.
H
P
Answer :
2
2
CS
S c.
PCl
2
O
5
5 hydrogen
b. diphosphorus
Naming Binary Molecular Compounds
H
H
2
2
Binary molecular compounds are composed of only two elements. Examples are
O, NO, SF
6
etc. .Sometimes these compounds have generic or common names (e.g.,
O is "water") and they also have systematic names (e.g., H
2
O, dihydrogenmonoxide).
The common name must be memorized. The systematic nameis more complicated but it has the advantage that the formula of the compound can be deduced from the name.
Common name
Compound Systematic name
(if it has one)
NF
3
NO nitrogen trifluoride
nitrogen monoxide note: for first element we don't use mono- prefix
nitrogen dioxide
nitr ic oxide higher oxidation #
NO
N
2
2
O dinitrogen monoxide laughing gas nitr ous oxide lower oxidaton #
N
2
O
4
PCl
5
SF
6
S
2
F
10
H
2
O
H
2
S
NH
3
N
2
H
4
dinitrogen tetraoxide
phosphorous pentachloride
sulfur hexafluoride
disulfur decafluoride
dihydrogen monoxide
dihydrogen monosulfide
nitrogen trihydride
dinitrogen tetrahydride water
hydrogen sulfide
ammonia
hydrazine
PH
3
phosphorous trihydride phosphine
Problem : Give the formula of following covalent compounds: a.
Nitrogen trifluoride b.
Carbon monoxide
Answer :

Names of Acids and Bases
Binary acids: made up of only two elements - hydrogen and one other element.
Naming binary acids:
Begin with the prefix hydro.
Determine the "stem" - part of the name of the element that combines with hydrogen.
Add the suffix ic.
Examples:
HF -hydro fluor ic - hydrofluor ic acid
HCl - hydro chlor ic - hydrochloric acid
HBr - hydro brom ic - hydrobromic acid
HI -hydro iod ic - hydroiodic acid
Ternary acids : made up of three elements - hydrogen, oxygen, and another element.
Naming ternary acids:
Acids made up of three elements including hydrogen
Determine the "stem" - part of the name of the third element.
The most common acid is given the suffix ic.
Add the prefix per for the acid with one more oxygen.
The suffix ous is given to the acid with one less oxygen.
Add the prefix hypo for the acid with two less oxygen atoms.
Examples :
HClO
HClO
3
4
- per chlor ic - perchloric acid - one more oxygen atom.
- chlor ic - chloric acid - the most common form of the acid.
HClO
2
- chlor ous - chlorous acid - one less oxygen atom.
HClO - hypo chlor ous - hypochlorous acid - two less oxygen atoms.
HNO
3
-nitric acid
HNO
2
-nitrous acid
H
2
SO
4
- sulfuric acid
H
H
3
2
SO
PO
4
3
-sulfurous acid
-phosphoric acid
H
3
PO
3
-phosphorous acid
H
3
BO
3
-boric acid
Properties of Ionic and Covalent Compounds
A compound consists of two or more atoms or ions bonded together.There are two major classes of compounds defined by chemical bonding:
• Ionic compounds are formed by the attractions between oppositely charged ions as described above.
• Covalent compounds are formed when nonmetallic atoms share attractionsfor each other's electrons.Each class of compounds has unique distinguishing properties.
Imagine that you are cleaning up tiny white crystals that have been spilled on your kitchen counter. How can you tell whether the crystals are table salt (an ionic compound), sugar (a molecular compound), or something else?
Compounds may be classified as ionic or covalent by performing some simple diagnostic tests. These include state of matter, melting point, solubility in water, and electrical conductivity tests.

Electrical conductivity test: A simple conductivity tester can be made using some wire, a battery, a light bulb, and light bulb socket. Ionic compounds dissolve in water to form electrically conductive solutions. Dissolved molecular substances do not conduct electricity.
All solutions that contain dissolved ionic compounds will conduct electrical current,
But what if the substance you want to test is not soluble in water? In such a case
You have to melt it. If the molten (liquid) substance conducts electricity, then it is ionic; if it doesn’t, then it is molecular.
Solutions of Ionic and Covalent Compounds
In the solid state, covalently bonded molecules are discrete units and have fewer tendencies to form an extended structure, while ionic compounds form a crystal lattice.
Some aqueous solutions conduct electricity due to the presence of freely moving ions.
As an example, ionic compounds consisting of anions and cations dissolved in water. In the presence of an electric field the ions in the solution begin to move, and it is the moving charges that form an electric current. Water itself is a poor conductor of electricity due to each water molecule being electrically neutral. Covalent compounds do not conduct electricity unless they break up into ions in the water solution
Electrolyte
An electrolyte is a substance that dissolves in water to give an electrically conducting solution. There are two different types of electrolytes (strong and weak electrolytes) summarized below.
Strong-Electrolytes, Weak-electrolytes and Nonelectrolytes
Strong electrolyte
Most ionic compounds (salts) e.g. NaCl(aq) Æ Na + (aq) + Cl (aq)
The single arrow indicates that the Na + form NaCl.
and Cl ions have no tendency to recombine to
Some molecular compounds such as strong acids e.g. HCl(aq) Æ H +
(aq)
+ Cl (aq)
The single arrow indicates that the H + and Cl ions have no tendency to recombine to form HCl.
Weak electrolytes
Weak electrolytes incompletely dissociate to form ions in solution. These substances exist as a mixture of molecules and ions in solution. The double arrow indicates that the reaction is significant in both directions. Both forward and reverse reactions occur constantly and simultaneously. As a result only a small amount of reactant ionizes to form the products. This balance between the forward and reverse reactions produces a state of chemical equilibrium.
Weak acids (molecular compounds) e.g. acetic acid (HC
2
H
3
O
2) weak bases (molecular compounds)

e.g. ammonia (NH
3
)
A non-electrolyte is a substance that dissolves in water to give a non-conducting or poorly conducting solution. Molecules of the substance mix with water molecules and dissolve but do not dissociate to form ions in solution. Molecules of a non-electrolyte are not charged and therefore do not carry an electric current.
Covalent compounds other than strong acids, weak acids and weak bases. e.g. sugar (C
12
H
22
O
11
)
C
12
H
22
O
11
(s) Æ C
12
H
22
O
11
(aq)
The melting and boiling temperatures for ionic compounds are generally higher than those of covalent compounds. Ionic solids are crystalline, whereas covalent solids may be either crystalline or amorphous. Many ionic solids dissolve in water, dissociating into positive and negative ions (an electrolytic solution). Because these ions can carry
(conduct) a current of electricity, they are called electrolytes. Covalent solids in solution usually retain their neutral character and are nonelectrolytes.
State of matter : Ionic compounds are solids at room temperature whereas molecular substances as a group are variable in their states of matter - some are solids, but many are liquids or gases. Since all pure ionic compounds are solids at room temperature
(25°C), you can classify any pure liquid or gas substance at room temperature as molecular.
Melting point : Generally, molecular substances like sugar melt at temperatures below
300°C, whereas ionic substances tend to have higher melting points. However, melting point data alone is usually insufficient evidence to classify a substance as ionic or molecular.
Molecular solids high melting points high boiling points conduct electricity when low melting points low boiling points do not conduct electricity molten or dissolved in water many are soluble in water few are soluble in non-polar few are soluble in water many are soluble in non-polar solvents solvents
Solubility of solids in water : Both ionic and molecular compounds may or may not dissolve in water, so this evidence alone cannot be used to classify a solid compound as ionic or molecular. However, combined with an electrical conductivity test , solubility tests are an excellent way of classifying solids.
Drawing Lewis Structures of Molecules and Polyatomic Ions
The electronic structure of atoms, ions, and molecules, which is closely allied to the properties of these substances, can conveniently be represented using Lewis structures, or electron-dot diagrams based on the octet rule.

Covalent bond - a bond between two atoms, in which the two atoms share a pair of electrons. Covalent bonds can join atoms together to form molecules.
Lewis-electron-dot or Lewis structure of a molecule or ion
Lewis electron dot structures are representations of the distribution of electrons in molecules and polyatomic ions. They are useful in determining the three-dimensional shape of a molecule or ion. A Lewis structure can be drawn for a molecule or ion by following three steps:
1. Calculate the number of valence electrons (including charges, if any)
2. Write skeleton structure
• Lowest EN (electronegative) atom, largest atom, and/or atom forming most bonds is usually central atom
• Hydrogen cannot be central
• Connect all atoms with a single bond. Single bond is represented by line or two dots.
• In oxoacids the oxygen bonds to central, the H to oxygen. Like in sulfuric acid,
H
2
SO
4
.
• Compounds are usually compact and symmetrical structures
3 . Count up used electrons
Distribute remaining electrons to terminal atoms filling their eight electrons(octet rule)
4. Central atom octet is filled last
Remaining electrons become lone pairs on central.
If central atoms do not have an octet, move terminal a pair to form multiple bonds on at a time forming double and triple bonds.
Problem : Draw the Lewis structures for the following molecules: H
2
O a) H
2
O sum of valence electrons: 2 + 6 =8 ; 4 pairs
Pick the central atom: usually the biggest atom and the atom forming most bonds.
O is the central atom
Connect central atom to other terminal atoms with a line representing a covalent bond
(two electrons). Fill octet to central atom:
A dot represents one electron and
line represents two electrons.
Check to see whether duet on H
Count electrons and check valence electrons on C.
2 bond pairs = 2 x 2 =
2 lone pairs = 2 x 2 = pairs (an Octet)
4
4
Total 8 = 4 electron
Bond pairs : an electron pair shared by two atom in a bond. E.g. two pairs between O--H in water.
Lone pair : an electron pair found solely on a single atom.
H
2
S has the same Lewis Structure as H
2
O since S is in the same group as O.
As a rule if an elements from the period 2 is replaced by another element from the period
3 and same group in the periodic table they will have same Lewis structure.
E.g. two pairs found on the
O atom at the top and the bottom.

Draw Lewis Structure of NH
3 i) Valence electrons: 5 + 3 x 1 = 8 = 4 electron pairs ii) N is the central atom: iii) Add terminal Hs vi) Give octet to N v)Check duet on Hs iv) Count electrons:
3 bond pairs = 3 pairs
1 lone pairs = 1 pairs
4 electron pairs
Draw Lewis structure of Nitrogen (N
2
)
Lewis Structures of NH
3
, PH
3
has the same
Lewis Structure
3 bond pairs and 1 lone pair are found on central N.
: N ::: N : triple bond between N and N
Draw Lewis structure of CO
2
i) Valence electrons: 4 + 2 x 6 = 16 ( 8 pairs) ii) Central atom C; O -- C -- O iii) Give octet to carbon
--
O -- C -- O
--
Draw Lewis structure of CH
4
i) Sum of valence electrons:
4 ( from one C) + 4 (from four H) = 8 electrons = 4 electron pairs ii) Central atom is C iii) Octet on C atom and duet on H are already complete. iv) Count valence electrons on the Lewis structure:
Draw Lewis structure of Ethane( CH
Count valence electrons:
2C = 8
3
CH
3
)
Draw two carbons and attach six hydrogen.
6H = 6
14 (7 electron pairs)
HNO3
H
N
= 1 valence electron
= 5 valence electrons
O x 3 = 6 x 3 = 18 valence electrons
Total of 24 valence electrons
Try to fill octet to O
iv) Count electrons:
4 bond pairs = 4 pairs
4 lone pairs = 4 pairs
8 electron pairs bond pairs = 2 x 4 = 8 = 4 electron pairs
lone pairs = 0
Octet on each carbon and duet on each hydrogen
H H
H
H H
Bond pairs : An electron pair shared by two atoms in a bond.
Lone pair : An electron pair found solely on a single atom.
Single covalent bond - Bond between two atoms when they shared 1 pair
Double covalent bond - Bond between two atoms when they shared 2 pairs.
Triple covalent bond - Bond between two atoms when they shared 3 pairs.

Lewis Structure, Stability, Multiple Bonds, and Bond Energies
Bond order
The stability of a covalent compound is related to the bond energy. The magnitude of the bond energy increases and the bond length decreases in the order: single bond > double bond > triple bond.
Bond Energy order : single < double < triple
Bond length order : single (1) < double (2) < triple (3)
Lewis Structures of Polyatomic Ions
Draw Lewis structure of cyanide ion (CN ) [ : C ::: N : ]
Draw Lewis structure of carbonate ion
(CO
3
2) [Note: 4 valence electrons from carbon, 6 from each of three oxygen atoms, and two for the -2 charge equal 24
O
O
C valence electrons total .]
O
2-
Resonance is a phenomenon invoked when more than one Lewis formula is needed to adequately describe the electronic distribution in a molecule. Resonance is the use of contributing structures to represent the electronic structure of a molecule.
Benzene consists of 6 carbon atoms in a hexagon. Each C atom is attached to two other C atoms and one hydrogen atom. There are alternating double and single bonds between the C atoms.
Experimentally, the C-C bonds in benzene are all the same length. Experimentally, benzene is planar. The six electrons in three π -bonds are delocalized over the entire ring and sometime shown inside the hexagon as a circle.
Show that it is necessary to draw several Lewis structures (resonance structures) to adequately describe the structures of the following molecules: CO
3
2, NO
3
, NO
2
-

NO -
3
Lewis Structure of CO
3
2-
The double bond could also be between C and any other oxygen atom.
Therefore, there are three resonance structures for
CO
3
2ion.
Lewis Structures of
NO
3
-
Lewis Structures and Exceptions to the Octet Rule
There are three classes of exceptions to the octet rule.
1) Molecules with an odd number of electrons;
2) Molecules in which one atom has less than an octet;
3) Molecules in which one atom has more than an octet.
Odd Number of Electrons Odd Number of Electrons
Few examples. Generally molecules such as ClO
2
, NO, and NO electrons .
2
have an odd number of
Less than an Octet Less than an Octet
Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A.
Most typical example is BF
3
.
More electrons than an Octet More than an Octet
This is the largest class of exceptions. Atoms from the 3rd period onwards can accommodate more than an octet. Beyond the third period, the d -orbitals are low enough in energy to participate in bonding and accept the extra electron density.
Lewis Structures and Molecular Geometry; VSEPR Theory
The valence shell electron pair repulsion (VSEPR) theory helps to explain (or predict) molecular geometry, including linear, trigonal planar, and tetrahedral arrangements of attached atoms. Exceptions to the octet rule exist.
Valence-Shell Electron-Pair
Repulsion Theory (VSEPR Theory)
In the theory, valence shell electron pairs are assumed to repel each other, assuming orientations to minimize repulsions and establish certain groups of molecular shapes.
Basic Geometries created by molecules with 2,3,4,5 and 6 electron pairs .

A = central atom;
{ SHAPE * MERGEFORMAT }
X= terminal atoms
.
Molecular Structures with Lone Pairs
E= lone pair of electrons
The molecular configuration is not exactly the same as the electron configuration. The reason for this is that lone pairs of electrons do not show up in the molecular configuration, although their effects are still seen (such as pushing bonding pair electrons closer together). This gives rise to 6 more configurations (in addition to the 5 basic electron configurations).
Angular or Bent: AX
2
E
2
H
2
O, OF
2
Triangular Pyramidal or
Trigonal Bipyamid
AX
3
E
1
NH
3
, NCl
3
Seasaw AX
4
E
1
SF
4
T-shape AX
4
E
2
CF
3
Square pyramidal AX
5
E
1
BrF
5

Square Planar AX
4
E2
XeF
4
First case of angular structures has three electron pairs surrounding the central atom.
The possible molecular configurations are either trigonal planar (if all of the electron pairs are bonding pairs) or angular Angular configuration could be based on trigonal planar or triangular planar electron geometry or bent (If one is a lone pair).
The second case of angular structure has four electron pairs surrounding the central atom. The water molecule (H:O:H) has two lone pairs and two bond pairs (4 total pairs on O). The shape is angular or bent with a bond angle which is less than that in NH
3
.
This is due to the greater repulsions with two lone pairs.
Lewis Structures and Polarity
Concepts of polarity are better understood when viewed from the perspective of a Lewis structure. A polar covalent molecule has at least one polar covalent bond. An understanding of the concept of electronegativity helps to assess the polarity of a bond.
A molecule containing only non-polar bonds must be non-polar. a molecule containing polar bonds may be polar or non-polar, depending on the relative position of the bonds.
Molecular Polarity
Molecules composed of covalently bonded atoms may also be polar or nonpolar. For the molecule to be polar, it must, of course, have polar bonds. But the key factor for determining the polarity of a molecule is its shape. If the polar bonds (dipoles) are symmetrical around the central atom, they offset each other and the resulting molecule is non-polar. However, if the dipoles are not symmetrical around the central atom, the electrons will be pulled to one end of the molecule. The resulting molecule is polar. Ball and stick models are often used to demonstrate molecular shape. In this exercise you will build several covalent molecules and predict each molecule's polarity on the basis of its molecular shape.
The geometry of the molecule affects its dipole moment or polarity - the asymmetric distribution of positive and negative charges.
Polar Molecules will orient in a magnetic field
When two elements with different electronegativities are bonded together, one obtains a polar bond.
Example: HF
Depending on the molecular shape, polar bonds can give rise to a polar molecule, one with an overall uneven distribution of electron charge. We can see this charge imbalance when the molecules are placed in an electric field. Polarity is a key feature of a molecule because it can influence physical, chemical, and even biological properties.
For a diatomic molecule, a polar bond must lead to a polar molecule. Consider hydrogen fluoride, shown here as the Lewis structure changes to a ball-and-stick model enclosed within the space-filling shape. Note the polar arrows and the colors. If red indicates high electron density and blue indicates low, you can see that the F end of the

molecule is much more negative than the H end, and thus HF is highly polar. Between two electric plates with the field off, the molecules lie every which way. With the field on, however, they become oriented with their negative ends facing the positive plate and their positive ends facing the negative plate.
If a molecule has more than two atoms, its shape can affect the polarity in a crucial way. For example, in carbon dioxide, since oxygen is more electronegative than carbon, each bond is highly polar. But the linear molecular shape makes the bond polarities cancel each other, so the CO
2
molecule is nonpolar. Notice that the orientation of the molecules is random, whether the field is off or on.
The situation is very different for water. As in CO
2
, the highly electronegative oxygen pulls electron density toward itself, so each bond is polar. But the V-shape of the molecule allows the bond polarities to reinforce each other, so water is highly polar, as the large polar arrow shows. With the field off, the molecules are oriented randomly, but with it on, their poles become oriented toward the oppositely charged plates.
The case of boron trifluoride is similar to that of CO
2
. Because the three highly polar bonds point to the corners of an equilateral triangle, the bond polarities cancel each other, and BF
3 or on.
is nonpolar. The molecules are oriented randomly with the field off
Like BF
3
, ammonia has four atoms and three polar bonds, but the trigonal pyramidal shape means that the bond polarities reinforce each other. Thus, ammonia is highly polar for the same reason that water is. Sometimes two molecules have similar overall shapes, but their slightly different compositions lead to very different polarities.
Carbon tetrachloride has four polar carbon-chlorine bonds that point to the corners of a tetrahedron, so they cancel each other and give a nonpolar molecule. Substituting a hydrogen for one of the chlorines gives chloroform, another tetrahedral molecule, but now the polar bonds reinforce each other. As a result, chloroform is highly polar.
Using VSEPR molecular structures predict the polarity (net dipole moment or zero dipole moment) of the following molecules: H
XeF
4
2
O, NH
3
, CO
2
, SO
3
, SF
6
, PCl
5
,
Molecule Shape Electron pair distribution
Polarity
H
2
O
NH
3
CO
2
SO
3
SF
6
Trigonal pyramid
Trigonal planer asymmetric symmetric
PCl
5
XeF
4
Trigonal bipyramid symmetric
Square planar symmetric
Noncovalent Interactions and Forces between Molecules polar
Non-polar octahedral symmetric Non-polar
Non-polar
Non-polar
Intermolecular forces: dipole-dipole, hydrogen bond and London dispersion.

Intermolecular forces are the attractions that exist among particles or molecules of matter in any of the three distinct phases of matter(gases, liquids and solids) found on earth. The strength of these forces increases as the matter changes phase from a gas to a liquid and then to a solid. Thus, gases have the weakest intermolecular forces compare to liquids and solids and solids have the strongest forces. The intermolecular forces in ionic solids are so strong that they exist only as solids even at very higher temperatures.
Ionic solids have ionic interactions; the strong electrostatic attractions between oppositely charged ions form an ionic lattice as in sodium chloride, NaCl(s). Other forms of weaker intermolecular forces exist resulting from polar covalent bonds in molecules and the polarization of electrons on non-polar covalent molecules. These weaker intermolecular forces are classified according to their strength as shown below: i) Dipole-Dipole Intermolecular forces: This includes the attraction between all polar molecules through their dipoles (except F-H, O-H and N-H dipoles). Dipolar covalent bonds are formed by unequal sharing of electrons in bonds in a molecule. Because of symmetry of a molecular structure, canceling the dipoles in the molecule can make it non-polar even though it may have polar covalent bonds. Therefore, there are no dipoledipole interactions in non-polar molecules. ii) Hydrogen bonding Three special cases of stronger dipole-dipole interactions resulting from F-H, O-H and N-H dipolar covalent bond are called hydrogen bonds.
This is because F, O and N have the largest electronegativities among all other elements. Hydrogen bonding (between O-H and O-H dipoles) in water allows water to exist as a liquid at room temperature. N-H dipole in proteins allows the formation of double helix structures in DNA and other complex structures found in living cells. iii) London Dispersion Forces: These forces are the weakest of all intermolecular attractions and occur in non-polar molecules without dipoles or dipolar covalent bonds.
The London dispersion forces result from instantaneous shifts of electron clouds of nonpolar molecules. These shifts in electron cloud create instantaneous dipoles with a very short lifespan. A weaker attractive force results because of the short lived dipole attractions between two molecules. Non-polar molecules such as H
2
and N
2
can be cooled to liquids at very low temperature due to the existence of London Dispersion forces. As the number of electrons increases or the molecular weight of a substance increases, the London Dispersion forces tend to increase which makes these substances exist as solids. Even though London dispersion forces exist in ionic solids and polar covalent compounds, their effect is masked by stronger ionic and dipole-dipole interactions.
Melting points and boiling points of molecular compounds
The melting point of a compound is the temperature at which a compound turns from a solid to a liquid or a liquid to a solid.
The boiling point of a compound is the temperature at which a compound turns from a liquid to a gas or a gas to a liquid. This temperature is a true measure of the forces of attractions between molecules as molecules separate from one another when they turn from a liquid to a gas.

The stronger the attractions between particles (molecules or ions), the more difficult it will be to separate the particles. When substances melt, the particles are still close to one another but the forces of attraction that held the particles rigidly together in the solid state have been sufficiently overcome to allow the particles to move. When substances boil, the particles are completely separated from one another and the attractions between molecules are completely overcome. The energy required to cause substances to melt and to boil, and thus disrupt the forces of attraction, comes from the environment surrounding the material. If you place a piece of ice in your hand, the ice will melt more quickly than if it is placed on a cold counter top. The energy required to melt the ice comes from your hand, your hand gets colder and the ice gets warmer.
Nitrogen Monoxide and Methanol
Nitrogen monoxide (NO) and methanol (CH
3
OH) are similar in size and thus have similar London forces. Nitrogen monoxide and methanol are polar covalent molecules and thus have dipole-dipole forces. Since methanol has the higher melting point and boiling point, it must have the stronger intermolecular forces. The difference in these molecules is the presence of a certain extremely polar bond present in methanol that is not present in nitrogen monoxide. This is the oxygen - hydrogen bond.
Oxygen is more electronegative than hydrogen and pulls the electron density in the oxygen - hydrogen bond towards it. This leaves very little electron density around the hydrogen since hydrogen has no core electrons. The part of hydrogen directed away from the oxygen - hydrogen bond has very little electron density shielding the nucleus.
Thus that part of the hydrogen nucleus which is exposed can interact with the nonbonding electrons on another methanol molecule. This interaction of a non-bonding pair with a hydrogen attached to an electronegative element such as oxygen is called a hydrogen bond .
For each of the following pairs of compounds, circle the compound that you would predict to have the higher boiling point. State a reason for your choice. a. CH b. P c. HCl d.
4
NaBr or or or
Ne CH
4
is more polarizable or AsH
3 present more e - a larger London forces
Hydrogen bonding dominates when
HF
ClBr NaBr is ionic and ClBr is molecular
Biomolecules: DNA and the Importance of Molecular Structure
Hydrogen-bonding pairs found in DNA. Note that the A-T pair has only two H-bonds because of the adenine molecule. (The lone pair of electrons on the oxygen atom in thymine molecule could act to form a third H-bond, but adenine does not have the required donor). The arrangement of donors and acceptors in thymine and guanine precludes formation of H-bonding pairs between these two compounds.
Covalent bonds versus Noncovalent Associations. covalent bond requires a chemical reaction for its formation or for it to be broken.
Noncovalent associations (electrostatic interactions, hydrogen bonds, dipole-dipole associations, Van der Waals associatons) do NOT require a chemical reaction to be either formed or broken.

What is the relative strength of the four kinds of noncovalent associations?
kcal Relative Scale
C-C covalent 83.0
Electrostatic 20.6
100
25
Hydrogen Bond 4.8
Dipole-Dipole 2.2
5
3
London Forces 0.1 less than 1

1. Which of the following is not a correct Lewis symbol?
F a.
Be b. c. d.
2. The Sulfite ion is ,
a. SO
3
2 b. NO
3
c. SO
4
2 d. CrO
4
-
3.
The acetate ion is ,
a. CO
3
2 b. CN c. NH
4
+ d. C
2
H
3
O
2
-
S
4.
The correct name for the compound CoCl
3 a) Cobalt chlorate
is, b) cobalt (III) chloride c) cobalt (I) chloride d) cobalt (III) chloride
5.
Which name/formula combination is wrong? a) phosphate/PO
4
3 b)
4 c) sulfurous acid/H
2
SO
3
d) sodium nitrate/NaNO
2 e) dinitrogen tetroxide/ N
2
O
4
.
6.
Which name/formula combination is wrong? a) phosphorous acid/H
3
PO
3 c) chloric acid/HClO
4 b) sodium cyanide/NaCN d) calcium hydrogensulfite/Ca(HSO
3
)
2 e) ammonium fluoride/NH
4
F
7.
How many valence electrons does a phosphorus atom have?
a. 3 b. 5 c. 15 d. 18 e. 31
8. The total number of valence electrons in the Lewis formula for PBr3 is a. 8 b. 12 c. 18 d. 26 e. 30
9.
The total number of non-bonding-pairs of electrons in the Lewis formula for
NH
3
is a. 0 b. 1 c. 2 d. 3 e. 4
10. Resonance Lewis structures can be drawn for all the ions and molecules of the following list except, a. CO
3
2- b. NO
3
- c. NO
2
- d. CO
2
11.
Which species of the following list possesses a central atom with more than octet of electrons?
a. SO
2 b. PF
5 c. BF
3 d. H
2
O e. NH
3
12.
The (H-O) bond in water is best characterized as a. polar covalent. c. coordinate covalent. b. ionic. d.
pure covalent.
13.
Which molecule will not contain a multiple bond?
a. N
2 b. CO c. SO
2 d. F
2 e. HCN
14.
molecule? 5 a. Square pyramidal c. tetrahedral e. see-saw b. octahedral d. trigonal bipyramidal
15.
According to VSEPR theory, what is the molecular geometry of the SF
4 molecule? a. trigonal planar b. T-shaped c. tetrahedral d. trigonal pyramidal e. see-saw

16.
Which molecular geometry is matched to CORRECT bond angles?
a. octahedral - 90º and 180º b. trigonal planar - 90º and 120º c. tetrahedral - 120º d. trigonal bipyramidal- 109.5º e. linear - 120º
17.
Which of the following is a non-polar molecule?
a. CCl
4 b. NH
3 c. H
2
S d. NF
3 e. OF
2
18. Which of the listed molecules possesses a square planar geometry?
a. SiCl
4 b. SF
4 c. XeF
4 d. CCl
4 e. CH
4
19.
Which of the following molecules is no polar?
a. NH
3 b. H
2
S c. SF
6 d.HCl e. CO
20.
What is the major type of force that must be overcome to allow the processes below?
I. the evaporation of propanol (CH
3
CH
2
CH
2
OH)
II. the melting of solid Br
2
III. the boiling of liquid BF
3
IV. the boiling of liquid CH
2
Cl
2 a. London forces, covalent bonding, dipole- dipole, dipole-dipole b. London forces, London forces, London forces, London forces c. hydrogen bonding, dipole-dipole, dipole-dipole, London forces d. hydrogen bonding, London forces, London forces, dipole-dipole e. dipole-dipole, dipole-dipole, London forces, London forces

Sample Test Chapter 4.
Structure and Properties of Ionic and Covalent
Compounds
1. In a Lewis structure, what do the dots represent?
Ans. valence electrons
2. Draw the Lewis structure of the bromine atom.
Ans.
3. How many dots are shown in the Lewis structure for the sulfur atom?
Ans. six
4. What are the two principal types of bonding called?
Ans. ionic bonding and covalent bonding
5. Name the two classes of element which are most likely to form an ionic compound if they are allowed to react with each other.
Ans. metal, nonmetal
6. Draw the Lewis structure of the Pb 2+ ion.
Ans.
7. What constitutes a covalent bond between two atoms?
Ans. a shared pair of electrons
8. In what way is a polar covalent bond similar to a nonpolar covalent bond? In what way are they different?
Ans. In each case, the bond consists of an electron pair shared between the bonded atoms.
The difference is that the sharing is unequal in the case of the polar covalent bond, equal in a nonpolar covalent bond.
9. What does it mean if an atom is said to have a high electronegativity?
Ans. The atom has a strong attraction for shared electron pairs
(electrons in covalent bonds).
10. The elements with the lowest electronegativities are found in the
_______ _______ region of the periodic table.
Ans. bottom left
11. Who first assigned electronegativity values to many of the elements?
Ans. Pauling
12. What do we call the three-dimensional arrangement of positive and negative ions in an ionic solid?
Ans. crystal lattice
13. Predict the formula of the compound formed when ions of sodium and sulfur combine.
Ans Na
2
S
14. Predict the formula of the compound formed when ions of barium and nitrogen combine.

Ans. Ba
3
N
2
15. What is the name of Fe 2+ in the Stock system?
Ans. iron(II) ion
16. What does the suffix "-ous" on the common names of ions mean?
Ans. lower positive charge
17. What is the term used for ions that are composed of two or more atoms bonded together?
Ans. polyatomic
18. What is the formula of the sulfate ion?
Ans. SO
4
2-
19. What is the name of the ion HCO
3
?
Ans. hydrogen carbonate or bicarbonate
20. What is the name of the ion NH
4
+ ?
Ans. ammonium
21. Provide the name of Na
3
PO
4
.
Ans. sodium phosphate
22. What is the name of Cu
2
O in the Stock system?
Ans. copper(I) oxide
23. Write the formula of sodium carbonate.
Ans. Na
2
CO
3
24. What kind of compound results when two or more different nonmetals share electrons?
Ans. covalent
25. What kind of bonding exists in substances which consist of discrete molecules?
Ans. Covalent
26. Provide the formula of sulfur trioxide.
Ans. SO
3
27. Write the formula of ammonia.
Ans NH
3
28. Provide the name of CCl
4
.
Ans. carbon tetrachloride
29. Provide the name of the compound whose formula is N
2
O
Ans. dinitrogen pentoxide
5
.
30. At what temperature is a liquid converted into a gas?

Ans. boiling point
31. What is the term that describes a solid with no regular structure?
Ans. amorphous
32. What is the term that describes a compound that, when dissolved in water conducts an electric current?
Ans. electrolyte
33. What is the term that describes a compound that, when dissolved in water does not conduct an electric current?
Ans. nonelectrolyte
34. What kind of bonding is present in substances which are nonelectrolytes?
Ans. covalent
35. How many bonding electrons are shown in the Lewis structure for the bicarbonate ion, HCO
3
?
Ans. ten
36. Draw the Lewis structure of methylamine, CH
3
NH
2
.
37. Draw the Lewis structure of hydrogen sulfide, H
2
S.
38. What is wrong with the Lewis structure shown below for sulfur trioxide, SO
3
?
The structure shows 26 valence electrons, but there should only be 24.
39. Ozone, O3, has two resonance forms. Draw them, given the skeletal arrangement O-O-O
40. What is defined as the amount of energy needed to break a bond holding two atoms together?
Ans. energy
41. What is defined as the distance of separation of two nuclei in a

covalent bond?
Ans. bond length
42. What do the letters VSEPR stand for?
Ans. Valence Shell Electron Pair Repulsion
43. If the shape of a molecule is trigonal planar, what are the values of the bond angles?
Ans. 120°
44. In the molecule AX2, the central atom A has two lone pairs of electrons in addition to the two bond pairs in the A-X bonds. What is the shape of this molecule?
Ans. bent or angular
45. The ammonia molecule, NH
3
, is polar. Why does this fact suggest that its shape is trigonal pyramidal, rather than trigonal planar?
Ans. If the molecule were trigonal planar, the symmetry would result in a nonpolar molecule. The centers of positive and negative charge would coincide.
46. Which of the following Lewis structures of neutral atoms is correct?
Ans. C
47. Which of the following Lewis structures of ions is incorrect?
Ans. C
48. Which of the following has the greatest electronegativity?
A. H B. Cl C. O D. F E. Na
Ans. D
49. Which of the following has the greatest electronegativity?
A. Si
Ans. C
B. P C. Cl D. Ar E. Br
50. In the compound CH
3
Cl the bond between carbon and chlorine is
A. intermolecular B. ionic
Ans. D
C. nonpolar covalent D. polar covalent
51. Which one of the following is NOT true about elements that form cations?
A. The atoms lose electrons in forming ions.

Ans. E
B. The elements are metals.
C. They are located to the left of the periodic table.
D. They have low ionization energies.
E. They have high electron affinities.
1.
Which of the following pairs of atoms are least likely to form an ionic compound?
A. Ni, O B. Na, F C. Cu, Cl D. Li, Mg E. Li, F
Ans. D. Li, Mg
2.
The bond in dinitrogen (N
2
) is a: a.
double bond. B. single bond. C. triple bond. D. lone pair. E. none of the above
Ans. C
3.
Which of the following formulas are incorrect on the basis of simple Lewis dot structures? a.
LiCl B. MgO C. Na
2
O D. CO
2
E. none of these
Ans. E
4.
Which of the following bonds is most polar? a.
H-F B. H-Cl C. H-H D. F-F E. H-I
Ans. H-F
56. Assuming reactions between the following pairs of elements, which pair is most likely to form an ionic compound?
A. copper and tin
B. chlorine and oxygen
C. cesium and iodine
Ans. C
D. carbon and chlorine
E. fluorine and iodine
57. What kind of bond results when electron transfer occurs between atoms of two different elements?
A. ionic
B. covalent
C. nonpolar
D. single
Ans. A
E. double
58. What is the old name of Cu+?
A. cupric ion
B. cuprous ion
C. copper(I) ion

D. copper(II) ion
E. ferrous ion
Ans. B
59. Give the name of FeSO4 in the Stock system.
A. iron monosulfuric acid
B. iron(II) sulfate
C. iron(III) sulfate
D. ferrous sulfate
E. ferric sulfate
Ans. B
60. Assuming reactions between the following pairs of elements, which pair is most likely to form a covalent compound?
A. lithium and iodine
B. sodium and oxygen
C. calcium and chlorine
D. copper and tin
E. carbon and oxygen
Ans. E
61. A double bond between two atoms, A and B
A. is longer than a single bond between the same two atoms
B. has a lower bond energy than a single bond between the same two atoms
C. arises when two electrons are transferred from A to B
D. consists of two electrons shared between A and B
E. consists of four electrons shared between A and B
Ans. E
62. What is the correct formula of phosphorus pentachloride?
A. PCl B. PCl
3
C. PCl
5
D. P
2
Cl
5
E. P
5
Cl
Ans. C
63. What term describes the temperature at which a solid is converted into a liquid?
A. critical point
B. flash point
C. sublimation point
D. melting point
E. boiling point
Ans. D
64. What term describes a solution of a compound in water that conducts an electric current?
A. amorphous solution
B. an electrolyte solution
C. a nonelectrolyte solution

D. superconducting solution
E. isoelectric solution
Ans. B
65. How many bonding electrons are in CO
2
?
A. 1 B. 2 C. 3 D. 4 E. 8
Ans. E
66. How many nonbonding electrons are in CH
4
?
A. 0 B. 1 C. 2 D. 3 E. 8
4 62 62 A
67. How many valence electrons are in SO
4
2?
A. 2 B. 64 C. 32 D. 12 E. 16
Ans. C
68. According to VSEPR theory, if the valence electrons on a central atom are 3 bond pairs and one nonbonding (lone) pair, the geometry (shape) at this atom will be
A. linear B. bent (angular)
C. trigonal planar D. trigonal pyramidal
E. tetrahedral
Ans. D
69. T F In Lewis structures, the chemical symbol of an element represents both the nucleus and the lower energy (nonvalence) electrons.
Ans. T
70. T F The name of SnO
2
is tin(I) oxide.
Ans. F
71. T F The old name of iron(III) chloride is ferrous chloride.
Ans. F
72. T F The are three atoms of iodine represented in the formula NaIO
3
.
Ans. F
73. T F In solid NaCl, no molecules of NaCl exist.
Ans. T
74. T F Ionic solids are amorphous.
Ans. F
75. T F Molecular compounds usually involve ionic bonding.
Ans. F
76. T F As a rule, ionic compounds tend to have lower melting and boiling points than covalent compounds consisting of small molecules.
Ans. F
77. T F In the water molecule, the oxygen atom is an exception to the

octet rule.
Ans. F
78. T F The NO2 molecule can never satisfy the octet rule.
4 74 74 T
79. T F Six electrons shared between two atoms corresponds to a bond order of three.
4 75 75 T
80. T F Resonance occurs when two or more different, valid Lewis structures can be drawn for a molecule+.
Ans. T
81. T F The existence of resonance makes a molecule less stable than would otherwise be the case.
Ans. F
82. T F Because the C-H bond in methane is polar, the CH
4
molecule will also be polar.
Ans. F
83. T F Chemical bonds are intramolecular forces.
Ans. T
84. T F In determining properties such as solubility, melting point and boiling point, intramolecular forces are more important than intermolecular forces.
Ans. F
85. T F As a rule, a polar substance will be a good solvent for nonpolar solutes, and vice versa.
Ans F