3 Chapter Lecture
General, Organic, and Biological
Chemistry: An Integrated Approach
Laura Frost, Todd Deal and Karen Timberlake
by Richard Triplett
Chapter 3
Compounds:
PuttingParticlesTogether
2011 Pearson Education, Inc.
Chapter Outline
3.1 The Octet Rule
3.2 In Search of an Octet Part 1: Ion Formation
3.3 Ionic CompoundsElectron Give and Take
3.4 In Search of an Octet Part 2: Covalent Bonding
3.5 Getting Covalent Compounds into Shape
3.6 Electronegativity and Molecular Polarity
3.1 The Octet Rule
What drives most elements to form
compounds?
Consider Group 8A elementsthe noble gases.
Noble gases are chemically unreactive and
highly stable.
If an element is more stable, it is less likely to
react with another element, and if an element is
less stable, it is likely to react.
3.1 The Octet Rule, Continued
The stability of an element and its reactivity is
inversely related.
What makes Group 8A elements more stable
and less reactive than other elements?
Recall that the group number indicates the
number of valence electrons for any element in
that group. Each noble gas has eight valence
electrons (except helium, which has two).
3.1 The Octet Rule, Continued
An electron arrangement of eight valence
electrons is unusually stable. Chemists do not
completely understand why this arrangement is
so stable.
The octet rule: Most atoms react to achieve a
total of eight valence electrons in their valence
shell.
3.2 In Search of an Octet Part 1: Ion Formation
Atoms gain or lose electrons to achieve a
valence octet.
An atom contains an even number of electrons
and protons, which means it is electrically
neutral.
If an atom gains or loses electrons to form a
valence octet, the new particle would have an
unequal number of protons and electrons and
would have a net charge.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
An ion is an atom that has gained or lost electrons
and has a charge.
Group 7A elements, like chlorine, have seven
valence electrons, gain an electron to have an
octet, and form an ion with a 1- charge (ions have
17 protons and 18 electrons).
NOTE: When chlorine forms an ion it has the
same number of valence electrons and total
electrons as argon. Therefore, the chlorine ion is
said to be isoelectronic (iso, a prefix meaning
same) to argon.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
An anion is any atom that contains a net negative
charge given that it has an unequal number of
protons and electrons.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
Group 1A elements, like sodium, have one
valence electron. In order to have an octet, the
elements in this group must lose one electron to
form ions with a 1+ charge (ions have 11 protons
and 10 electrons).
NOTE: When sodium forms an ion it has the
same number of valence electrons and total
electrons as argon. Therefore, the sodium ion is
isoelectronic to neon.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
A cation is any atom that contains a net positive
charge, given that it has an unequal number of
protons and electrons.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
Trends in Ion Formation
How do we determine which atoms will gain
electrons to form anions and which atoms will
lose electrons to form cations?
Metals lose electrons to form cations and
nonmetals gain electrons to form anions.
The main group metals in Groups 1A, 2A, and
3A form cations. Main group nonmetals in
Groups 5A, 6A, and 7A form anions.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
The main group metals form cations with charges
of 1+ to 3+, whereas the main group nonmetals
form anions with charges of 1- to 3-.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
Group 1A elements lose one electron to form
cations with a 1+ charge. Group 2A elements
lose two electrons to form cations with a 2+
charge. Group 3A elements give up three
electrons to form cations with a 3+ charge.
With main group metals, the group number
indicates the charge of the cation that is formed.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
The charge on the main group elements can be
determined by subtracting eight from the group
number. For example, for Group 6A elements,
the ions formed have a 2- charge (6 8 = 2-).
The rules for predicting charges on ions formed
by the main group elements are as follows:
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
Predicting the charge of metals (other than the
main group metals) is not as simple.
Transition metals form more than one ion. For
example, iron will form the Fe2+ and Fe3+ cation.
Copper will form the Cu+ and Cu2+ cation.
Another group of common ions whose charge
cannot be determined by the periodic table are
those formed when two or more elements
combine to form an ion.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
Polyatomic ions are a special group of ions
formed when two or more nonmetal elements
interact with one another. The ions formed have a
single charge.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
Naming Ions
To distinguish between an element and its ion,
a different name is given to the ion.
To name metal ions, add the word ion to the
name of the metal, so Na+ becomes the sodium
ion.
Transition metals form more than one ion. A
roman numeral in parentheses following the
name of the metal indicates the charge. F2+ is
iron(II) and Fe3+ is iron(III).
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
To name nonmetal ions, the suffix ide replaces
the last few letters of the name of the element.
The anion formed from chlorine becomes the
chloride ion. The anion formed from oxygen is
the oxide ion.
To name polyatomic ions, keep in mind that
most polyatomic ions end in -ate, but the ending
ite is used for names of related ions that have
one less oxygen atom. The hydroxide (OH-),
hydronium (H3O+), and cyanide (CN-) ions are
exceptions to this rule.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
Important Ions in the Body
Ions in the body perform important functions.
Main cations in the body are Na+, K+, Ca2+, and
Mg2+. They are important in maintaining solution
concentrations inside and outside cells.
Main anions in the body are Cl-, HCO3-, and
HPO42-. These help maintain charge neutrality
between blood and fluids in the cell.
3.2 In Search of an Octet Part 1: Ion Formation,
Continued
3.3 Ionic CompoundsElectron Give and Take
Cations do not form unless anions form and vice
versa. When metals and nonmetals combine,
electrons are transferred between the atoms. Ion
formation is shown as follows:
3.3 Ionic CompoundsElectron Give and Take,
Continued
Opposite charges attract, so any cation and
anion formed are strongly attracted to one
another.
An ionic bond is formed when a cation is
attracted to an anion. The resulting combination
of cation and anion is called an ionic
compound.
3.3 Ionic CompoundsElectron Give and Take,
Continued
Formation of table salt, sodium chloride, is shown
as follows:
3.3 Ionic CompoundsElectron Give and Take,
Continued
Formulas of Ionic Compounds
When ions are attracted to one another to form
ionic compounds, the overall charge is zero.
3.3 Ionic CompoundsElectron Give and Take,
Continued
Na+ and Cl- interact to form a neutral ionic
compound, NaCl. The formula is written with no
charges shown.
When a formula is written, the number of each
ion in the compound is represented as a
subscript.
Just as a charge of 1+ on an ion is represented
as +, the 1 is understood in the formula.
Sodium chloride is written as NaCl, not as
Na1Cl1.
3.3 Ionic CompoundsElectron Give and Take,
Continued
How do you write the formula if one ion has a
charge greater than the other?
Cations and anions combine so the total charge
on the ionic compound has no net charge.
Consider calcium ions with a charge of 2+ and
fluorine ions with a charge of 1-.
A calcium atom reacts with two fluorine atoms
giving up two electrons, one to each of the
fluorine atoms.
One Ca2+ ion is formed and two F- ions are
formed.
3.3 Ionic CompoundsElectron Give and Take,
Continued
The overall charge resulting from the ionic
compound formed when one calcium atom reacts
with two fluorine atoms is shown as:
3.3 Ionic CompoundsElectron Give and Take,
Continued
How do we determine the formula for the ionic
compound formed with Al3+ reacts with O2-?
Use the method known as the crossover approach.
Using the crossover approach, the charge
number of the cation becomes the subscript of the
anion, and the charge number of the anion
becomes the subscript of the cation.
3.3 Ionic CompoundsElectron Give and Take,
Continued
Naming Ionic Compounds
When naming ionic compounds, combine the
names of the ions involved after dropping the word
ion from the name of the individual ions.
CATION
ANION
COMPOUND NAME
Na+(sodium ion)
Cl-(chloride ion)
NaCl, sodium chloride
Ca2+(calcium ion)
F-(fluoride ion)
CaF2, calcium fluoride
Al3+(aluminum ion)
O2-(oxide ion)
Al2O3, aluminum oxide
3.4 In Search of an Octet Part 2:
Covalent Bonding
To form an octet in the valence shell, some
atoms transfer or accept electrons to form ions,
and some atoms chemically combine with other
atoms to share valence electrons.
A covalent bond is formed by the sharing of
electrons between two nonmetals.
A covalent compound is formed when
nonmetals share electrons. The smallest unit of
a covalent compound is called a molecule.
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
Covalent Bond Formation
The rule of thumb in covalent bond formation is
that the number of covalent bonds that an atom
will form equals the number of electrons that it
needs to complete its octet.
The electron dot symbol (elemental symbol
plus a dot for each valence electron) for an atom
is used to determine covalent bond formation.
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
Below is the list of steps for the construction of an
electron dot symbol:
1. Write the symbol of the element.
2. Place one valence electron on the top, bottom, right,
and left of the element symbol.
3. Continue placing the remaining valence electrons
around the element symbol by forming ion pairs with
the first four electrons.
4. Chlorine has seven valence electrons and its
electron dot symbol is:
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
One chlorine atom can form a covalent bond
with another chlorine atom.
Chlorine needs only one electron to complete its
octet, so it shares an electron from another
chlorine atom.
A bonding pair is a shared pair of electrons.
The remaining valence electrons known are as
the lone pair of electrons.
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
The bonding and lone pair of electrons are
shown for the newly formed molecule Cl2.
Below are the electron dot symbols of the main
elements found in living systems:
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
Carbon, the element of life on Earth, has four
unpaired valence electrons.
Carbon can form:
Four single bonds (a sharing of two electrons)
One double and two single bonds, or two
double bonds (a sharing of four electrons)
One single bond and one triple bond (a
sharing of six electrons)
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
Formulas and Structures of Covalent Compounds
A molecular formula identifies all the components
in a molecule of a covalent compound.
Table sugar, sucrose (molecular formula C12H22O11)
has 12 carbon atoms, 22 hydrogen atoms, and 11
oxygen atoms.
Molecular formula only tells us the number of atoms
in a molecule not how they are joined or what the
structure is.
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
Consider methane, which has a molecular
formula of CH4. It contains one carbon atom and
four hydrogen atoms.
To determine how atoms are connected, start by
drawing the electron dot symbols for all the
atoms involved.
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
Carbon has four unpaired valence electrons and
needs four to complete its octet.
Each hydrogen atom has one valence electron
and needs two electrons for completion (an
exception to the octet rule), therefore hydrogen
can form one single bond.
Carbon can form one single bond with each of
the hydrogen atoms in its molecular formula.
This representation is called a Lewis structure.
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
Naming Covalent Compounds
A binary compound is a covalent compound
composed of only two elements.
The following are rules for naming binary
compounds:
Step 1. Name the first element in the formula.
Step 2. Name the second element and change
the ending to -ide.
Step 3. Designate the number of each element
present using a Greek prefix.
3.4 In Search of an Octet Part 2:
Covalent Bonding, Continued
Indicating the number of each element present is
important. For example, carbon and oxygen can
form both CO (carbon monoxide) and CO2 (carbon
dioxide). The prefix mono- for carbon is
understood.
3.5 Getting Covalent Compounds into
Shape
To understand molecules it is important to
identify the three-dimensional shapes of each
molecule.
The Lewis structure of methane appears flat on
paper, which implies the four hydrogen bond
angles are 90o.
Electrons are negatively charged and repel each
other, so the clouds of negative charge want to
be as far away from each other as possible.
3.5 Getting Covalent Compounds into
Shape, Continued
When the electrons obtain maximum distance from
each other, the three-dimensional structure of
methane shows a molecule whose atoms achieve a
tetrahedral shape, with bond angles of 109.5o
between the clouds of electrons.
3.5 Getting Covalent Compounds into
Shape, Continued
The valence-shell electron-pair repulsion
(VSEPR) model is used to predict the shape of
a molecule based on the number of electroncharged clouds on a given atom.
VSEPR states that in the valence shell of an
atom, the charge clouds formed by groups of
electrons will arrange themselves to be as far
away from each other as possible in order to
reduce repulsions.
3.5 Getting Covalent Compounds into
Shape, Continued
Determining the Shape of a Molecule
The following are steps used to determine the
shape of a molecule using VSEPR:
Step 1. Determine the central atom of a molecule.
The central atom is the one the shape will be
centered around. It is represented as A in the
VSEPR form.
Step 2. Determine the number of charge clouds
around the central atom. Single and multiple
bonds are counted as one charge cloud. Charge
clouds in bonds are called bonding clouds.
They are represented by B in the VSEPR form.
3.5 Getting Covalent Compounds into
Shape, Continued
Lone pair of electrons are counted as one
charge cloud and affect the shape of the
molecule.
Nonbonding clouds are charge clouds from
lone pairs of electrons and are represented as N
in the VSEPR form.
3.5 Getting Covalent Compounds into
Shape, Continued
One example is formaldehyde, which has a VSEPR
form of AB3a central atom with three bonding
charge clouds. Its three-dimensional structure is
shown as:
3.5 Getting Covalent Compounds into
Shape, Continued
Another example is ammonia, which has a VSEPR
form of AB3Na central atom with three bonding
charge clouds and one nonbonding charge cloud.
Its three-dimensional structure is shown as:
3.5 Getting Covalent Compounds into
Shape, Continued
The following lists molecular shapes associated
with the VSEPR forms.
3.5 Getting Covalent Compounds into
Shape, Continued
Nonbonding Pairs and Their Effect on Molecular
Shape
Nonbonding pairs affect the shape of a molecule.
A molecule like methane has a central atom with four
atoms around it and assumes the tetrahedral shape.
Ammonia is a molecule with three atoms and a
nonbonding pair around central atoms. Its shape
changes to pyramidal.
3.5 Getting Covalent Compounds into
Shape, Continued
The lone pair of ammonia forces the bonded pair
together because they repel the bonded pairs more
than the bonded pairs repel each other, causing a
small angle between the bonded pairs.
3.5 Getting Covalent Compounds into
Shape, Continued
Molecular Shape of Larger Molecules
The molecular shape of larger molecules with
several charge clouds can be determined as
long as a central atom is identified.
The shape around a single atom that is bonded
to at least two other atoms can be determined.
C2H5OH, ethanol, has the molecular shape
shown below.
3.6 Electronegativity and Molecular Polarity
Electronegativity
Electronegativity is the ability of an atom to
attract bonding electrons of a covalent bond to
itself.
Electronegativity of the elements in the periodic
table increases as you move to the right and up,
excluding the noble gases. Fluorine is the most
electronegative element.
3.6 Electronegativity and Molecular Polarity,
Continued
3.6 Electronegativity and Molecular Polarity,
Continued
Unequal sharing exists when one element in a
molecular formula has a higher electronegativity
than another element in the formula.
The unequal sharing of electrons results in a
covalent bond called a polar covalent bond.
When electrons in a covalent bond are shared
equally, the bond is called a nonpolar covalent
bond.
3.6 Electronegativity and Molecular Polarity,
Continued
An uneven sharing of electrons in a polar covalent
bond is denoted using the Greek symbol delta, ,
meaning partial.
3.6 Electronegativity and Molecular Polarity,
Continued
The type of bond formed between two elements
can be predicted by the difference between the
electronegativities of the two elements.
An electronegativity difference of 1.8 or more is
an ionic bond resulting from the transfer of
electrons.
If the difference is less than 1.8, the resulting
bond is a covalent bond. The greater the
difference in electronegativity of the two atoms,
the more polar the bond.
3.6 Electronegativity and Molecular Polarity,
Continued
The farther apart two elements are on the periodic
table, the greater the polarity of the bond between
them.
3.6 Electronegativity and Molecular Polarity,
Continued
Molecular Polarity
The interaction of electronegativity and
molecular shape is the basis for determining if a
molecule is polar or nonpolar.
Carbon dioxide is linear and water is bent.
Because of the bent shape of water, the net
polarity of water is around the oxygen atom.
3.6 Electronegativity and Molecular Polarity,
Continued
The rules for determining the polarity of a
molecule are as follows:
If the central atom has one or more pairs of
nonbonding electrons, the molecule is polar.
If the central atom has no nonbonding pairs of
electrons and all other atoms connected to the
central atom are identical, the molecule is
nonpolar.
If the central atom has no nonbonding pairs of
electrons and at least one of the atoms
connected to the central atom is different from
the others, the molecule is polar.
Chapter Summary
3.1 The Octet Rule
The noble gases are stable because they have
eight valence electrons.
Atoms react with other atoms to obtain a stable
valence octet.
Hydrogen and helium are exceptions. They have
two electrons in the full shell.
Chapter Summary, Continued
3.2 In Search of an Octet Part 1: Ion Formation
Gaining or losing electrons is one way atoms
can attain a valence octet.
Metals form cations, which are positively
charged, by losing electrons. Nonmetals form
anions, which are negatively charged, by
gaining electrons.
Polyatomic ions are groups of nonmetals
combined together to form ions.
Chapter Summary, Continued
3.3 Ionic CompoundsElectron Give and Take
Ionic compounds are formed when one atom
loses an electron(s) and another atom gains an
electron(s). Ionic compounds are neutral.
To name an ionic compound, drop the word ion
on the cation and then name the anion.
If a transition metal is present, a Roman numeral
is placed in parentheses following the name of
the metal.
Chapter Summary, Continued
3.4 In Search of an Octet Part 2: Covalent Bonding
Atoms share electrons with one another to achieve a
valence octet. A bond formed by sharing electrons is
called a covalent bond.
The number of covalent bonds that an atom can
form is determined by the number of electrons it
needs to complete its octet.
A molecular formula gives the number of each type
of atom in a covalent compound.
In naming molecular formulas, name the first
element with its name and name the second
element by changing the ending to -ide.
Chapter Summary, Continued
3.5 Getting Covalent Compounds into Shape
Molecules have three-dimensional shape.
Valence-shell electron-pair repulsion
(VSEPR) theory explains that electrons around
any given atom arrange themselves to get as far
away from each other as possible.
Presence of nonbonding electron pairs on an
atom reduces the bond angle around that atom
and changes the shape of the molecule.
Chapter Summary, Continued
3.6 Electronegativity and Molecular Polarity
The unequal sharing of electrons in covalent
bonds is determined by comparing the
electronegativities of the atoms involved.
Nonpolar covalent bonds have equal sharing
of electrons.
Polar covalent bonds have unequal sharing of
electrons.