Me cchapter 3

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Chapter 3
Structure and Properties of Ionic and
Covalent Compounds

Denniston
Topping
Caret
7th Edition

3.1 Chemical Bonding
• Chemical bond - the force of attraction
between any two atoms in a compound
• This attractive force overcomes the
repulsion of the positively charged nuclei of
the two atoms participating in the bond
• Interactions involving valence electrons are
responsible for the chemical bond

3.1 Chemical Bonding Lewis Symbols
• Lewis symbol (Lewis structure) - a way to
represent atoms using the element symbol
and valence electrons as dots
• As only valence electrons participate in
bonding, this makes it much easier to work
with the octet rule
• The number of dots used corresponds
directly to the number of valence electrons
located in the outermost shell of the atoms
of the element

3.1 Chemical Bonding Lewis Symbols
• Each “side” of the symbol represents an atomic
orbital, which may hold up to two electrons
• Using Lewis symbols
– Place one dot on each side until there are four dots
around the symbol
– Now add a second dot to each side in turn
– The number of valence electrons limits the number of
dots placed
– Each unpaired dot (unpaired electron of the valence
shell) is available to form a chemical bond

3.1 Chemical Bonding Lewis Dot Symbols for
Representative Elements

Principal Types of Chemical Bonds:
Ionic and Covalent
• Ionic bond - a transfer of one or more
electrons from one atom to another
• Forms attractions due to the opposite charges of
the atoms
• Covalent bond - attractive force due to the
sharing of electrons between atoms
• Some bonds have characteristics of both
types and not easily identified as one or the
other

3.1 Chemical Bonding Ionic Bonding
• Representative elements form ions that
obey the octet rule
• Ions of opposite charge attract each other
creating the ionic bond
• When electrons are lost by a metal and
electrons are gained by a nonmetal
– Each atom achieves a “Noble Gas”
configuration
– 2 ions are formed; a cation and anion, which
are attracted to each other

3.1 Chemical Bonding Ionic Bonding
Consider the formation of
NaCl
Chlorine has a high
Na + Cl  NaCl electron affinity
When chlorine gains
Sodium has a low an electron, it gains
ionization energy it the Ar configuration
readily loses this
electron .. −  ..  −
: Cl⋅ + e → : Cl :
Na  Na+ + e- ..  .. 
When sodium loses the
electron, it gains the
Ne configuration

Essential Features of Ionic Bonding
• Atoms with low I.E. and low E.A. tend to form
positive ions
• Atoms with high I.E. and high E.A. tend to form
negative ions
• Ion formation takes place by electron transfer
• The ions are held together by the electrostatic
force of the opposite charges
• Reactions between metals and nonmetals
(representative elements) tend to be ionic

3.1 Chemical Bonding Ion Arrangement in a Crystal
• As a sodium atom loses one electron, it becomes a
smaller sodium ion
• When a chlorine atom gains that electron, it
becomes a larger chloride ion
• Attraction of the Na cation with the Cl anion
forms NaCl ion pairs that aggregate into a crystal

3.1 Chemical Bonding Covalent Bonding
Let’s look at the formation of H2:
H + H  H2
• Each hydrogen has one electron in its valance
shell
• If it were an ionic bond it would look like this:
H ⋅ + H ⋅ → H + [ H :]
+ −

• However, both hydrogen atoms have an equal
tendency to gain or lose electrons
• Electron transfer from one H to another usually
will not occur under normal conditions

• Instead, each atom attains a noble gas
configuration by sharing electrons

H⋅ + ⋅H →H : H
Each hydrogen
atom now has two The shared
electrons around it electron
and attained a He pair is called a
configuration Covalent Bond

3.1 Chemical Bonding Covalent Bonding in Hydrogen

3.1 Chemical Bonding Features of Covalent Bonds
• Covalent bonds form between atoms with
similar tendencies to gain or lose electrons
• Compounds containing covalent bonds are
called covalent compounds or molecules
• The diatomic elements have completely
covalent bonds (totally equal sharing)
– H2, N2, O2, F2, Cl2, Br2, I2
Each fluorine is
.. .. .. .. surrounded by 8
: F⋅ + ⋅ F : → : F : F : electrons – Ne
.. .. .. ..
configuration

3.1 Chemical Bonding Examples of Covalent Bonding
.. ..
2H ⋅ + ⋅ O⋅ → H : O : H
.. ..
2e– from 2H 2e– for H
6e– from O 8e– for O

H
. ..
4H ⋅ + ⋅ C⋅ → H : C : H
⋅ ⋅⋅
H
4e– from 4H 2e– for H
4e– from C 8e– for C

Polar Covalent Bonding and
Electronegativity

• The Polar Covalent Bond
– Ionic bonding involves the transfer of
electrons
– Covalent bonding involves the sharing of
electrons
– Polar covalent bonding - bonds made up
of unequally shared electron pairs

somewhat positively charged somewhat negatively charged
3.1 Chemical Bonding .. ..
H⋅ + ⋅F:→ H : F:
⋅⋅ ⋅⋅
These two
electrons are
not shared equally
• The electrons spend more time with fluorine
• This sets up a polar covalent bond
• A truly covalent bond can only occur when
both atoms are identical

3.1 Chemical Bonding Polar Covalent Bonding in Water
• Oxygen is electron rich = δ-
• Hydrogen is electron
deficient = δ+
• This results in unequal
sharing of electrons in the
pairs = polar covalent bonds
• Water has 2 covalent bonds

3.1 Chemical Bonding Electronegativity
• Electronegativity - a measure of the
ability of an atom to attract electrons in a
chemical bond
• Elements with high electronegativity have
a greater ability to attract electrons than do
elements with low electronegativity
• Consider the covalent bond as competition
for electrons between 2 positive centers
– The difference in electronegativity determines
the extent of bond polarity

Electronegativities of Selected
Elements
• The most electronegative elements are found in the upper
right corner of the periodic table
• The least electronegative elements are found in the lower
left corner of the periodic table

electronegativity increases
electronegativity increases

3.1 Chemical Bonding Electronegativity Calculations
• The greater the difference in electronegativity
between two atoms, the greater the polarity of
their bond
• Which would be more polar, a H-F bond or H-Cl
bond?
• H-F … 4.0 - 2.1 = 1.9
• H-Cl … 3.0 - 2.1 = 0.9
• The HF bond is more polar than the HCl bond

3.2 Naming Compounds and
Writing Formulas of Compounds

• Nomenclature - the assignment of a correct
and unambiguous name to each and every
chemical compound
• Two naming systems:
– ionic compounds
– covalent compounds

3.2 Naming Compounds and Formulas of Compounds
• A formula is the representation of the
fundamental compound using chemical
symbols and numerical subscripts
– The formula identifies the number and type
of the various atoms that make up the
compound unit
– The number of like atoms in the unit is
shown by the use of a subscript
– Presence of only one atom is understood
when no subscript is present

3.2 Naming Compounds and Ionic Compounds
• Metals and nonmetals usually react to form
ionic compounds
• The metals are cations and the nonmetals
are anions
• The cations and anions arrange themselves
in a regular three-dimensional repeating
array called a crystal lattice
• Formula of an ionic compound is the
smallest whole-number ratio of ions in the
substance

3.2 Naming Compounds and Writing Formulas of Ionic Compounds
Writing Formulas of Compounds from the Identities of the Component Ions
• Determine the charge of each ion
– Metals have a charge equal to group number
– Nonmetals have a charge equal to the group
number minus eight
• Cations and anions must combine to give a
formula with a net charge of zero
• It must have the same number of positive
charges as negative charges

3.2 Naming Compounds and Predict Formulas
Predict the formula of the ionic compounds
formed from combining ions of the
following pairs of elements:

1. sodium and oxygen

2. lithium and bromine

3. aluminum and oxygen

4. barium and fluorine

3.2 Naming Compounds and Writing Names of Ionic Compounds
Writing Formulas of Compounds from the Formula of the Compound
• Name the cation followed by the name2
of the anion
• A positive ion retains the name of the
element; change the anion suffix to
-ide

3.2 Naming Compounds and Writing Names of Ionic Compounds
Writing Formulas of Compounds from the Formula of the Compound
• If the cation of an element has several ions of
different charges (as with transition metals) use a
Roman numeral following the metal name
• Roman numerals give the charge of the metal
• Examples:
• FeCl3 is iron(III) chloride
• FeCl2 is iron(II) chloride
• CuO is copper(II) oxide

Writing Formulas of Compounds Common Nomenclature System
• Use -ic to indicate the higher of the
charges that ion might have
• Use -ous to indicate the lower of the
charges that ion might have
• Examples:
• FeCl2 is ferrous chloride
• FeCl3 is ferric chloride

Iron and Copper Ions
Stock and Common Names for

Writing Formulas of Compounds Common Monatomic Cations
and Anions

• Monatomic ions - ions consisting of a
single charged atom

3.2 Naming Compounds and Polyatomic Ions
• Polyatomic ions - ions composed of 2 or
more atoms bonded together with an
overall positive or negative charge
– Within the ion itself, the atoms are bonded
using covalent bonds
– The positive and negative ions will be
bonded to each other with ionic bonds
• Examples:
• NH4+ ammonium ion
• SO42- sulfate ion

Writing Formulas of Compounds Anions
Common Polyatomic Cations and

Writing Formulas of Compounds Name These Compounds
1. NH4Cl

2. BaSO4

3. Fe(NO3)3

4. CuHCO3

5. Ca(OH)2

3.2 Naming Compounds and Writing Formulas of Ionic Compounds
From the Name of the Compound
• Determine the charge of each ion
• Write the formula so that the resulting
compound is neutral
• Example:
Barium chloride:
Barium is +2, Chloride is -1
Formula is BaCl2

3.2 Naming Compounds and Determine the Formulas From
Writing Formulas of Compounds Names
Write the formula for the following ionic
compounds:

1. sodium sulfate

2. ammonium sulfide

3. magnesium phosphate

4. chromium(II) sulfate

3.2 Naming Compounds and Covalent Compounds
• Covalent compounds are typically formed
from nonmetals
• Molecules - compounds characterized by
covalent bonding
• Not a part of a massive three-dimensional
crystal structure
• Exist as discrete molecules in the solid, liquid,
and gas states

Writing Formulas of Compounds Naming Covalent Compounds
1. The names of the elements are written
in the order in which they appear in
the formula
2. A prefix indicates the number of each
kind of atom

3.2 Naming Compounds and Naming Covalent Compounds
3. If only one atom of a particular element is
present in the molecule, the prefix mono- is
usually omitted from the first element
Example: CO is carbon monoxide
4. The stem of the name of the last element is
used with the suffix –ide
5. The final vowel in a prefix is often dropped
before a vowel in the stem name

3.2 Naming Compounds and Name These Covalent Compounds
1. SiO2

2. N2O5

3. CCl4

4. IF7

3.2 Naming Compounds and Writing Formulas of Covalent
Writing Formulas of Compounds Compounds
• Use the prefixes in the names to determine the
2
subscripts for the elements
• Examples:
• nitrogen trichloride NCl3
• diphosphorus pentoxide P2O5
• Some common names that are used:
– H2O water
– NH3 ammonia
– C2H5OH ethanol
– C6H12O6 glucose

3.2 Naming Compounds and Provide Formulas for These
Writing Formulas of Compounds Covalent Compounds
1. nitrogen monoxide

2. dinitrogen tetroxide

3. diphosphorus pentoxide

4. nitrogen trifluoride

3.3 Properties of Ionic and
Covalent Compounds
• Physical State
– Ionic compounds are usually solids at room
temperature
– Covalent compounds can be solids, liquids, and
gases
• Melting and Boiling Points
– Melting point - the temperature at which a
solid is converted to a liquid
– Boiling point - the temperature at which a
liquid is converted to a gas

Physical Properties
Covalent Compounds
• Melting and Boiling Points
– Ionic compounds have much higher melting points
and boiling points than covalent compounds
– A large amount of energy is required to break the
electrostatic attractions between ions
– Ionic compounds typically melt at several hundred
degrees Celsius
• Structure of Compounds in the Solid State
– Ionic compounds are crystalline
– Covalent compounds are crystalline or amorphous –
having no regular structure

Physical Properties
Covalent Compounds
• Solutions of Ionic and Covalent
Compounds
– Ionic compounds often dissolve in water,
where they dissociate - form positive and
negative ions in solution
– Electrolytes - ions present in solution
allowing the solution to conduct electricity
– Covalent solids usually do not dissociate and
do not conduct electricity - nonelectrolytes

Comparison of Ionic vs. Covalent
Compounds
Ionic Covalent
Composed of Metal + nonmetal 2 nonmetals
Electrons Transferred Shared
Physical state Solid / crystal Any / crystal
OR amorphous
Dissociation Yes, electrolytes No,
nonelectrolytes
Boiling/Melting High Low

3.4 Drawing Lewis Structures on
Molecules and Polyatomic Ions
Lewis Structure Guidelines
1. Use chemical symbols for the various
elements to write the skeletal structure of
the compound
– The least electronegative atom will be placed in
the central position
– Hydrogen and halogens occupy terminal
positions
– Carbon often forms chains of carbon-carbon
covalent bonds

Structures of Molecules Lewis Structure Guidelines
2. Determine the number of valence
3.4 Drawing Lewis

electrons associated with each atom in
the compound
– Combine these valence electrons to
determine the total number of valence
electrons in the compound
– Polyatomic cations, subtract one electron for
every positive charge
– Polyatomic anions, add one electron for
every negative charge

Structures of Molecules Lewis Structure Guidelines
3. Connect the central atom to each of the
3.4 Drawing Lewis

surrounding atoms using electron pairs
• Next, complete octets of all the atoms
bonded to the central atom
• Hydrogen needs only two electrons
• Electrons not involved in bonding are
represented as lone pairs
• Total number of electrons in the structure
must equal the number of valence electrons
in step 2

Lewis Structure Guidelines
Structures of Molecules
4. Count the number of electrons you have
3.4 Drawing Lewis

and compare to the number you used
• If they are the same, you are finished
• If you used more electrons than you have,
add a bond for every two too many you used
• Then, give every atom an octet
• If you used less electrons than you
have….see later exceptions to the octet rule
5. Recheck that all atoms have the octet rule
satisfied and that the total number of
valance electrons are used

Structures of Molecules Drawing Lewis Structures of
Covalent Compounds
3.4 Drawing Lewis

Draw the Lewis structure of carbon dioxide, CO2
Draw a skeletal structure of the molecule
1. Arrange the atoms in their most probable order
C-O-O and/or O-C-O
2. Find the electronegativity of O=3.5 & C=2.5
3. Place the least electronegative atom as the central
atom, here carbon is the central atom
4. Result is the O-C-O structure from above

Structures of Molecules Drawing Lewis Structures
3.4 Drawing Lewis

5. Find the number of valence electrons for each
atom and the total for the compound
1 C atom x 4 valence electrons = 4 e-
2 O atoms x 6 valence electrons = 12 e-
16 e- total
6. Use electron pairs to connect the C to each O
with a single bond

O:C:O
7. Place electron pairs around the atoms
: :

: :

:O:C:O:
This satisfies the rule for the O atoms, but not for C

Drawing Lewis Structures of
Structures of Molecules Covalent Compounds
3.4 Drawing Lewis

8. Redistribute the electrons moving 2 e- from each O,
placing them between C:O

: :

: :
C::O::C
9. In this structure, the octet rule is satisfied
• This is the most probable structure
• Four electrons are between C and O
• These electrons are share in covalent bonds
• Four electrons in this arrangement signify a double
bond
10. Recheck the electron distribution
• 8 electron pairs = 16 valence electrons, number
counted at start
• 8 electrons around each atom, octet rule satisfied

Lewis Structures Practice
Using the guidelines presented, write Lewis
3.4 Drawing Lewis

structures for the following:

1. H2O

2. NH3

3. CO2

4. NH4+

5. CO32-

6. N2

Lewis Structures of
Polyatomic Ions 6
3.4 Drawing Lewis

• Prepare Lewis structures of polyatomic
ions as for neutral compounds, except:
• The charge on the ion must be
accounted for when computing the
total number of valence electrons

Lewis Structure of
Structures of Molecules Polyatomic Cations
3.4 Drawing Lewis

Draw the Lewis structure of ammonium ion, NH4+
1. Ammonium has this structure and charge:
2. The total number of valence electrons is determined by
subtracting one electron for each unit of positive
charge
1 N atom x 5 valence electrons = 5 e-
4 H atoms x 1 valence electron = 4 e-
- 1 electron for +1 charge = -1 e-
8 e- total
3. Distribute these 8 e- around the skeletal structure

Lewis Structure of Polyatomic
Structures of Molecules Anions
3.4 Drawing Lewis

Draw the Lewis structure of carbonate ion, CO32-
1. Carbon is less electronegative than oxygen
• This makes carbon the central atom
• Skeletal structure and charge:
2. The total number of valence electrons is determined by
adding one electron for each unit of negative charge
1 C atom x 4 valence electrons = 4 e-
3 O atoms x 6 valence electron = 18 e-
+ 2 negative charges = 2 e-
24 e- total
3. Distribute these e- around the skeletal structure

Lewis Structure of Polyatomic
Structures of Molecules Anions
3.4 Drawing Lewis

Draw the Lewis structure of carbonate ion, CO32-
4. Distributing the electrons around the central carbon
atom (4 bonds) and around the surrounding O atoms
attempting to satisfy the octet rule results in:

5. This satisfies the octet rule for the 3 oxygen, but not
for the carbon
6. Move a lone pair from one of the O atoms to form
another bond with C

Lewis Structure, Stability, Multiple
Bonds, and Bond Energies
3.4 Drawing Lewis

• Single bond - one pair of electrons are
shared between two atoms
• Double bond - two pairs of electrons are
• Triple bond - three pairs of electrons are
• Very stable

.. ..
H : H or H - H O :: O or O = O
Structures of Molecules ⋅⋅ ⋅⋅
.. ..
3.4 Drawing Lewis

N  N or N ≡ N
Bond energy - the amount of energy
required to break a bond holding two
atoms together
triple bond > double bond > single bond
Bond length - the distance separating the
nuclei of two adjacent atoms
single bond > double bond > triple bond

Lewis Structures and Resonance
• Write the Lewis structure of CO32-
3.4 Drawing Lewis

• If you look around you, you will probably
see the double bond put in different places
• Who is right? All of you!
• In some cases it is possible to write more
than one Lewis structure that satisfies the
octet rule for a particular compound

.. ..
:O: :O: :O:
Structures of Molecules : :: :
.. .. .. .. ..
: O : C :: O ↔ : O : C : O : ↔ : O :: C : O :
3.4 Drawing Lewis

⋅⋅ ⋅⋅ ⋅⋅ ⋅⋅
• Experimental evidence shows all bonds are
the same length, meaning there is not really
any double bond in this ion
• None of theses three Lewis structures exist,
but the actual structure is an average or
hybrid of these three Lewis structures
• Resonance - two or more Lewis structures
that contribute to the real structure

Lewis Structures and Exceptions
Structures of Molecules to the Octet Rule
3.4 Drawing Lewis

1. Incomplete octet - less then eight
electrons around an atom other than H
– Let’s look at BeH2
1 Be atom x 2 valence electrons = 2 e-
2 H atoms x 1 valence electrons = 2 e-
total 4 e-

– Resulting Lewis structure:
H : Be : H or H – Be – H

Structures of Molecules Odd Electron
2. Odd electron - if there is an odd number of
3.4 Drawing Lewis

valence electrons, it is not possible to give
every atom eight electrons
• Let’s look at NO, nitric oxide
• It is impossible to pair all electrons as the
compound contains an ODD number of valence
electrons
N - O

Structures of Molecules Expanded Octet
3.4 Drawing Lewis

3. Expanded octet - an element in the 3rd period or
below may have 10 and 12 electrons around it
• Expanded octet is the most common exception
• Consider the Lewis structure of PF5
• Phosphorus is a third period element

1 P atom x 5 valence electrons = 5 e-
5 F atoms x 7 valence electrons = 35 e-
40 e- total
• Distributing the electrons results in this Lewis structure

Structures of Molecules Lewis Structures and Molecular
Geometry: VSEPR Theory
3.4 Drawing Lewis

• Molecular shape plays a large part in
determining properties and shape
• VSEPR theory - Valance Shell Electron Pair
Repulsion theory
• Used to predict the shape of the molecules
• All electrons around the central atom arrange
themselves so they can be as far away from
each other as possible – to minimize electronic
repulsion

Structures of Molecules VSEPR Theory
3.4 Drawing Lewis

• In the covalent bond, bonding electrons
are localized around the nucleus
• The covalent bond is directional, having
a specific orientation in space between
the bonded atoms
• Ionic bonds have electrostatic forces
which have no specific orientation in
space

Structures of Molecules Molecular Bonding
3.4 Drawing Lewis

• Bonding pair = two electrons shared
by 2 atoms
– H:O
• Nonbonding pair = two electrons
belonging to 1 atom, pair not shared
– N:
• Maximal separation of bonding pairs
= 4 corners of a TETRAHEDRON

3.4 Molecular Geometry A Stable Exception to the Octet
Rule
• Consider BeH2
– Only 4 electrons surround the beryllium atom
– These 2 electron pairs have minimal repulsion
when located on opposite sides of the structure
– Linear structure having bond angles of 180°

3.4 Molecular Geometry Another Stable Exception to the
Octet Rule
• Consider BF3
– There are 3 shared electron pairs around the central
atom
– These electron pairs have minimal repulsion when
placed in a plane, forming a triangle
– Trigonal planar structure with bond angles of 120°

3.4 Molecular Geometry Basic Electron Pair Repulsion of
a Full Octet
• Consider CH4
– There are 4 shared electron pairs around the central
Carbon
– Minimal electron repulsion when electrons are placed at
the four corners of a tetrahedron
– Each H-C-H bond angle is 109.5°
• Tetrahedron is the primary structure of a full octet

a Full Octet with One Lone Pair
Consider NH3
• There are 4 electron pairs around the central Nitrogen
• 3 pairs are shared electron pairs
• 1 pair is a lone pair
– A lone pair is more electronegative with a greater electron repulsion
– The lone pair takes one of the corners of the tetrahedron without being
visible, distorting the arrangement of electron pairs
• Ammonia has a trigonal pyramidal structure with 107° angles

a Full Octet with Two Lone Pairs
Consider H2O
• There are 4 electron pairs around the central Oxygen
• 2 pairs are shared electron pairs
• 2 pairs are lone pairs
– All 4 electron pairs are approximately tetrahedral to each other
– The lone pairs take two of the corners of the tetrahedron without being
visible, distorting the arrangement of electron pairs
• Water has a bent or angular structure with 104.5° bond angles

Predicting Geometric Shape Using
Structures of Molecules Electron Pairs
3.4 Drawing Lewis

Basic Procedure to Determine
Structures of Molecules Molecular Shape
3.4 Drawing Lewis

1. Write the Lewis structure
2. Count the number of shared electron pairs and
lone pairs around the central atom
3. If no lone pairs are present, shape is:
• 2 shared pairs - linear
• 3 shared pairs - trigonal planar
• 4 shared pairs - tetrahedral
1. Look at the arrangement and name the shape
• Linear
• Trigonal planar
• Bent
• Trigonal pyramid
• Tetrahedral

Determine the Molecular
Geometry
3.4 Drawing Lewis

• PCl3

• SO2

• PH3

• SiH4

Lewis Structures and Polarity
• A molecule is polar if its centers of positive and
3.4 Drawing Lewis

negative charges do not coincide
• Polar molecules when placed in an electric field
will align themselves in the field
• Molecules that are polar behave as a dipole (having
two “poles” or ends)
• One end is positively charged the other is negatively
charged
• Nonpolar molecules will not align themselves in
an electric field

Structures of Molecules Determining Polarity
To determine if a molecule is polar:
3.4 Drawing Lewis

• Write the Lewis structure
• Draw the geometry
• Use the following symbol to denote the polarity
of each bond

Positive end of the Negative end of the bond,
bond, the less more electronegative atom
electronegative atom attracts the electrons more
strongly towards it

Practice Determining Polarity
Determine whether the following bonds and
3.4 Drawing Lewis

molecules are polar:
1. Si – Cl 1. O2
2. H–C 2. HF
3. C–C 3. CH4
4. S – Cl 4. H2O

3.5 Properties Based on Electronic
Structure and Molecular Geometry
• Intramolecular forces – attractive forces
within molecules – Chemical bonds
• Intermolecular forces – attractive forces
between molecules
• Intermolecular forces determine many
physical properties
– Intermolecular forces are a direct consequence
of the intramolecular forces in the molecules

Solubility and Intermolecular
Forces
Solubility - the maximum amount of solute
that dissolves in a given amount of
solvent at a specific temperature
• “Like dissolves like”
– Polar molecules are most soluble in polar
solvents
– Nonpolar molecules are most soluble in
nonpolar solvents
• Does ammonia, NH3, dissolve in water?
• Yes, both molecules are polar

Interaction of Water and
Ammonia

• The δ- end of ammonia, N, is attracted to the δ+ end of
the water molecule, H
• The δ+ end of ammonia, H, is attracted to the δ- end of
the water molecule, O
• The attractive forces, called hydrogen bonds, pull
ammonia into water, distributing the ammonia
molecules throughout the water, forming a
homogeneous solution

Interaction of Water and Oil
• What do you know about oil
and water?
– “They don’t mix”
• Why?
– Because water is polar and
oil is nonpolar
• Water molecules exert their
attractive forces on other water
molecules
• Oil remains insoluble and
floats on the surface of the
water as it is less dense

Boiling Points of Liquids
and Melting Points of Solids
• Energy is used to overcome the
intermolecular attractive forces in a
substance, driving the molecules into a
less associated phase
• The greater the intermolecular force, the
more energy is required leading to
– Higher melting point of a solid
– Higher boiling point of a liquid

Factors Influencing Boiling and
Melting Points
• Strength of the attractive force holding the
substance in its current physical state
• Molecular mass
• Larger molecules have higher m.p. and b.p. than
smaller molecules as it is more difficult to convert a
larger mass to another phase

• Polarity
• Polar molecules have higher m.p. and b.p. than
nonpolar molecules of similar molecular mass due to
their stronger attractive force

Melting and Boiling Points –
Selected Compounds by Bonding Type

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