Ionic and Covalent Bonding
CHAPTER 12
I. Electron Configuration in Ionic Bonding.
A. Valence Electrons.
1. The knowledge about the electron structure of atoms and the
organization of the periodic table will help in understanding the
the chemical bonding that occurs between atoms.
* Electron dot structures are diagrams that show valence
electrons as dots.
* Examples:
2. All of the elements within each group of the periodic table behave
similarly because they have the same number of valence electrons.
a. Valence electrons are the electrons in the highest occupied
energy levels of an element’s atoms.
b. The number of valence electrons determines the chemical
properties of an element.
c. One way to determine the number of valence electrons is to
look at the element’s electron configuration.
d. The number of valence electrons is also related to the group
numbers in the periodic table.
* Examples:
B. Electron Configurations for Cations.
1. In 1916, chemist Gilbert Lewis explained why atoms form certain
kinds of ions and molecules.
a. The octet rule states that in forming compounds, atoms tend to
achieve the electron configuration of a noble gas.
b. An octet is a set of eight.
c. Noble gases, except helium, have eight electrons in its highest
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energy level and a general electron configuration of ns np .
d. Atoms of metallic elements tend to lose their valence electrons,
leaving a complete octet in the next-lowest energy level.
e. Atoms of nonmetallic elements tend to gain electrons of share
electrons with another nonmetallic element to achieve a
complete octet.
f. Exceptions to the octet rule do exist (more later).
e. Valence electrons are usually the only electrons used in chemical
bonding (those electrons in the outermost s and p sublevels).
f. As a general rule only the valence electrons are shown in
electron dot structures.
2. An atom’s loss of valence electrons produces a cation, or a positively
charged ion.
a. The most common cations are those produced by the loss of
valence electrons from metal atoms.
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b. Most of these atoms have one to three valence electrons.
c. Examples:
b. Ions with charges of three or greater are uncommon and are
extremely unlikely, although some do exist.
c. 18 electrons in the outer energy with all of its orbitals filled is a
favorable orientation called a pseudo noble-gas configuration.
d. Examples:
3. For metals, the atoms lose enough valence electrons to attain the
the electron configuration of a noble gas, and the charge reflects
this.
* Examples:
C. Electron Configurations for Anions.
1. An anion is an atom or a group of atoms with a negative charge.
a. The gain of negatively charged electron by a neutral atom
produces an anion.
b. Atoms of nonmetallic elements attain noble-gas electron
configurations more easily by gaining electrons than by losing
them.
c. Examples:
4. For transition metals, the charge of cations vary.
a. Some ions formed by transition metals do not have a noble-gas
electron configuration and are therefore exceptions to the
octet rule.
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2. Ions that are produced when the halogens gain electrons are called
halides.
a. All halogens ions have seven valence electrons and need to gain
only one electron to achieve the electron configuration of a
noble gas.
d. Examples:
b. Halogen ions have a -1 charge.
3. Other groups.
* Examples:
II. Ionic Bonds.
A. Formation of Ionic Compounds.
1. Anions and cations have opposite charges and attract each other by
electrostatic forces.
a. The forces of attraction that bind oppositely charged ions are
called ionic bonds.
b. Compounds that consist of electrically neutral groups of ions
joined by electrostatic forces are called ionic compounds.
c. The total positive charge must equal the total negative charge in
all ionic compounds.
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2. Problems.
a. Use electron dot structures to determine chemical formulas of
the ionic compound formed when the following elements combine.
* Magnesium and chlorine
* Aluminum and sulfur
* Example:
d. The coordination number of an ion determines the number of
ions of opposite charges that surround it in a crystal.
* Potassium and iodine
* Strontium and oxygen
2. When melted, ionic compounds can conduct an electric current.
a. When ionic compounds are melted, the orderly crystal structure
breaks down.
b. If a voltage is applied across a molten mass, cations migrate
freely to one electrode and anions migrate to the other.
c. The ion movement produces a flow of electricity between the
electrodes through an external wire.
3. Ionic compounds also conduct electricity when dissolved in water.
a. When ions dissolve, they are free to move about in aqueous
solution and electricity is conducted.
b. Unless melted or dissolved in water, an ionic compound will not
conduct electricity as no movement of ions is possible.
B. Properties of Ionic Compounds.
1. At room temperature, most ionic compounds are crystalline solids.
a. The component ions are arranged in repeating three-dimensional
patterns.
b. The simplest repeating unit of the crystal is called a unit cell.
c. The crystalline form is favored in ionic compounds because it
is approximately 1.76 times stronger than the single formula
unit.
III. Bonding in Metals.
A. Metallic Bonds and Metallic Properties.
1. Metals are made up of closely packed cations rather than neutral
atoms.
a. The cations are surrounded by mobile valence electrons which
can drift freely from one part of the metal to another.
b. This theory of metallic bonding is called “The Electron-Sea
Model.”
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c. A metallic bond consists of the attraction of the free-floating
valence electrons for the positively charged metal ions.
d. These bonds are the forces of attraction that hold metals
together.
e. Diagram:
B. Alloys.
1. Very few “metallic” objects in life are pure metals.
a. Most metals are alloys.
b. Alloys are mixtures composed of two or more elements, at least
one is a metal.
2. Alloys tend to be harder and more durable than their component
elements.
3. Common examples:
2. Properties of metals as explained by the electron-sea model.
a. Metals are good conductors of electrical current because
electrons can flow freely in them
b. Metals are ductile (drawn into a thin wire) and malleable
(hammered or forced into shapes) because of the mobility of
valence electrons.
* A sea of drifting electrons insulates the metal cations from
one another.
* When a metal is subjected to pressure, the metal cations
easily slide past one another like ball bearings immersed in oil.
* Diagrams:
4. The most important alloys today are steels.
a. They contain a mixture of iron, carbon, boron, chromium,
manganese, molybdenum, nickel, tungsten, and vanadium.
b. Properties include corrosion resistance, ductility, hardness, and
toughness.
5. Types of alloys.
a. Substitutional alloys form if atoms of the components in an alloy
are about the same size and can replace each other in the
crystal.
b. Interstitial alloys form if the atomic sizes are different and
the smaller atoms can fit in to the interstices (spaces) between
the larger atoms.
* Steels are examples.
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IV. The Nature of Covalent Bonding.
A. Single Covalent Bonds.
1. A single covalent bond is a bond in which two atoms share a pair of
electrons.
2. One difference between covalent compounds and ionic compounds is
that covalent compounds exist as molecules in which the subscripts
are not necessarily lowest whole-number ratios; ionic compounds
exist as formula units in which the subscripts are always the lowest
whole-number ratio.
3. The atoms of the nonmetallic elements are most likely to form
covalent bonds.
4. Gilbert Lewis summarized covalent bonding in his octet rule: Sharing
of electrons occurs if the atoms involved acquire the electron
configurations of noble gases.
* Often the configurations contain eight valence electrons (an
octet).
5. Notations of covalent molecules or compounds are shown using
structural formulas, chemical formulas that show the arrangement
of atoms in molecules or polyatomic ions.
a. Electrons are shown as dots.
* Write the symbol for the atoms in the molecule. In simple
molecules, one atom will be the central atom surrounded by
the other atoms. The central atom is often the first atom in
the formula. Hydrogen can form only one covalent bond, so it
cannot be the central atom.
* Draw a dash between each pair of atoms covalently bonded.
* For each dash drawn, subtract two electrons for the total
electrons as dots around the atoms. Arrange the dots so that
most atoms have eight valence electrons, and hydrogen has
two.
* If there are not enough electrons to give the atoms eight
electrons, shift unbonded electrons as necessary or change
single bonds to double or triple bonds. Verify that each atom
has a noble gas structure (two electrons for hydrogen and
eight for the others).
f. Examples:
b. Covalent bonds are shown as dashes.
c. Electrons between atoms in a molecule are called shared
electron pairs.
d. Electrons that are not shared between atoms are called
unshared electrons pairs, also known as lone pairs or nonbonding
pairs.
e. Rules for Lewis Electron-Dot Structures for Molecules.
* Add up the valence electrons for each atom in the molecule.
use the periodic table to determine the number of valence
electrons for each atom.
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6. Problems.
a. The following covalent molecules have only single covalent bonds.
Draw a Lewis dot structure for each.
NF3
SBr2
HCl
PCl3
c. Examples:
C. Coordinate Covalent Bonds.
1. A coordinate covalent bond is a covalent bond in which one atom
contributes both bonding electrons.
a. In structural formulas, coordinate covalent bonds can be shown
as arrows that point from the atom donating the pair of
electrons to the atom receiving it.
b. Examples:
CH4
B. Double and Triple Covalent Bonds.
1. Atoms sometimes share more than one pair of electrons to attain a
stable noble-gas electron configuration.
a. Double covalent bonds are bonds that involve two shared pairs
of electrons.
b. Triple covalent bonds are bonds that involve three shared pairs
of electrons.
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2. Most polyatomic ions are examples of covalent and coordinate
covalent bonds.
a. When drawing Lewis dot structures for polyatomic ions, the
structure is placed in brackets.
b. An electron is added to the total valence electrons for each
negative charge of the ion.
c. An electron is subtracted from the total valence electrons for
each positive charge of the ion.
e. Examples:
D. Resonance.
1. Resonance structures are structures that occur when it is possible to
write two or more valid electron dot formulas that have the same
number of electron pairs for a molecule or ion.
2. Examples:
3. Problem.
* Draw Lewis dot structures for hydroxide, boron tetrafluoride
-
(BF4 ), chlorate, sulfate, and hydrogen carbonate.
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E. Exceptions to the Octet Rule.
1. For some molecules or ions, it is impossible to write structures that
satisfy the octet rule.
2. This may occur whenever the total number of valence electrons is an
odd number.
a. Commonly, exceptions to the rule exists when the compound
contains a Group VA central atom (usually N).
a. This is common among the nonmetal in the third energy level and
beyond.
b. Examples:
b. Example:
5. The concept of electron pairing is important in understanding the
bonding and properties of molecules.
a. Substances in which all of the electrons are paired are
diamagnetic, weakly repelled by an external magnetic field.
3. Some molecules are electron deficient.
a. This commonly occurs when Be or B are the central atoms.
b. Examples:
b. Paramagnetic substances are those that contain one or more
unpaired electrons and show a strong attraction for a magnetic
field.
c. Paramagnetism should not be confused with ferromagnetism in
2+
2+
which permanent magnets are created when Fe , Co , and Ni
ions line up in their metallic structure.
2+
V. Bonding Theories.
A. VSEPR Theory.
1. VSEPR (valence shell electron pair repulsion) theory states that
because electron pairs repel, molecular shape adjusts so the valence
electron pairs are as far apart as possible.
4. Other molecules have expanded valence shell and can hold 10 or 12
valence electrons.
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2. Examples
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