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3.1 Ionic Bonding
18:57
Learning Objectives
By the end of this section, you will be able to:
Explain the formation of cations, anions, and ionic compounds.
Predict the charge of common metallic and nonmetallic elements and write their electron
configurations.
Ionic Bonding
As you have learned, ions are atoms or molecules bearing an electrical charge. A cation (a positive
ion) forms when a neutral atom loses one or more electrons from its valence shell, and an anion (a
negative ion) forms when a neutral atom gains one or more electrons in its valence shell.
Compounds composed of ions are called ionic compounds (or salts), and their constituent ions are
held together by ionic bonds: electrostatic forces of attraction between oppositely charged cations
and anions. The properties of ionic compounds shed some light on the nature of ionic bonds. Ionic
solids exhibit a crystalline structure and tend to be rigid and brittle; they also tend to have high
melting and boiling points, which suggests that ionic bonds are very strong. Ionic solids are also poor
conductors of electricity for the same reason—the strength of ionic bonds prevents ions from moving
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freely in the solid state. Most ionic solids, however, dissolve readily in water. Once dissolved or
melted, ionic compounds are excellent conductors of electricity and heat because the ions can move
about freely.
Neutral atoms and their associated ions have very different physical and chemical properties. Sodium
atoms form sodium metal, a soft, silvery-white metal that burns vigorously in air and reacts
explosively with water. Chlorine atoms form chlorine gas, Cl2, a yellow-green gas that is extremely
corrosive to most metals and very poisonous to animals and plants. The vigorous reaction between
the elements sodium and chlorine forms the white, crystalline compound sodium chloride, common
table salt, which contains sodium cations and chloride anions. The compound composed of these
ions exhibits properties entirely different from the properties of the elements sodium and chlorine.
Chlorine is poisonous, but sodium chloride is essential to life; sodium atoms react vigorously with
water, but sodium chloride simply dissolves in water.
Figure 3.1 (a) Sodium is a soft metal that must be stored in mineral oil to prevent reaction with air or
water. (b) Chlorine is a pale yellow-green gas. (c) When combined, they form white crystals of sodium
chloride (table salt). (credit a: modification of work by “Jurii”/Wikimedia Commons)
The Formation of Ionic Compounds
Binary ionic compounds are composed of just two elements (thus, the prefix “bi,” meaning two): a
metal (which forms the cations) and a nonmetal (which forms the anions). For example, NaCl is a
binary ionic compound. We can think about the formation of such compounds in terms of the periodic
properties of the elements. Many metallic elements have relatively low ionization potentials and lose
electrons easily. These elements lie to the left in a period or near the bottom of a group on the
periodic table. Nonmetal atoms have relatively high electron affinities and, thus, readily gain electrons
lost by metal atoms, thereby filling their valence shells. Nonmetallic elements are found in the upper-
right corner of the periodic table.
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As all substances must be electrically neutral, the total number of positive charges on the cations of
an ionic compound must equal the total number of negative charges on its anions. The formula of an
ionic compound represents the simplest ratio of the numbers of ions necessary to give identical
numbers of positive and negative charges. For example, the formula for aluminum oxide, Al2O3,
indicates that this ionic compound contains two aluminum cations, Al3+, for every three oxide anions,
O2− ,
[thus, (2 × +3) + (3 × –2) = 0].
It is important to note, however, that the formula for an ionic compound does not represent the
physical arrangement of its ions. It is incorrect to refer to a sodium chloride (NaCl) “molecule”
because there is not a single ionic bond, per se, between any specific pair of sodium and chloride
ions. The attractive forces between ions are isotropic—the same in all directions—meaning that any
particular ion is equally attracted to all of the nearby ions of opposite charge. This results in the ions
arranging themselves into a tightly bound, three-dimensional lattice structure. Sodium chloride, for
example, consists of a regular arrangement of equal numbers of Na+ cations and Cl– anions.
Figure 3.2 The atoms in sodium chloride (common table salt) are arranged to (a)
maximize opposite charges interacting. The smaller spheres represent sodium ions,
the larger ones represent chloride ions. In the expanded view (b), the geometry can
be seen more clearly. Note that each ion is “bonded” to all of the surrounding ions—
six in this case.
The strong electrostatic attraction between Na+ and Cl– ions holds them tightly together in solid NaCl.
It requires 769 kJ of energy to dissociate one mole of solid NaCl into separate gaseous Na+ and Cl–
ions:
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NaCl(s) + 769 kJ Na+(g) + Cl–(g)
Electronic Structures of Cations
When forming a cation, an atom of a main group element tends to lose all of its valence electrons,
thus assuming the electronic structure of the noble gas that precedes it in the periodic table. For
groups 1 (the alkali metals) and 2 (the alkaline earth metals), the group numbers are equal to the
numbers of valence shell electrons and, consequently, to the charges of the cations formed from
atoms of these elements when all valence shell electrons are removed. For example, calcium is a
group 2 element whose neutral atoms have 20 electrons and a ground state electron configuration of
1s22s22p63s23p64s2. When a Ca atom loses both of its valence electrons, the result is a cation with
18 electrons, a 2+ charge, and an electron configuration of 1s22s22p63s23p6. The Ca2+ ion is,
therefore, isoelectronic with the noble gas Ar.
For groups 13–17, the group numbers exceed the number of valence electrons by 10 (accounting for
the possibility of full d subshells in atoms of elements in the fourth and greater periods). Thus, the
charge of a cation formed by the loss of all valence electrons is equal to the group number minus 10.
For example, aluminum (in group 13) forms 3+ ions (Al3+).
Exceptions to the expected behavior involve elements toward the bottom of the groups. In addition to
the expected ions Tl3+, Sn4+, Pb4+, and Bi5+, a partial loss of these atoms’ valence shell electrons can
also lead to the formation of Tl+, Sn2+, Pb2+, and Bi3+ ions. The formation of these 1+, 2+, and 3+
cations is ascribed to the inert pair effect, which reflects the relatively low energy of the valence s-
electron pair for atoms of the heavy elements of groups 13, 14, and 15. Mercury (group 12) also
exhibits an unexpected behavior: it forms a diatomic ion, Hg22+ (an ion formed from two mercury
atoms, with an Hg-Hg bond), in addition to the expected monatomic ion Hg2+ (formed from only one
mercury atom).
Transition and inner transition metal elements behave differently than main group elements. Most
transition metal cations have 2+ or 3+ charges that result from the loss of their outermost s
electron(s) first, sometimes followed by the loss of one or two d electrons from the next-to-outermost
shell. For example, iron (1s22s22p63s23p64s23d6) forms the ion Fe2+ (1s22s22p63s23p63d6) by the
loss of the 4s electrons and the ion Fe3+ (1s22s22p63s23p63d5) by the loss of the 4s electrons and
one of the 3d electrons. Although the d orbitals of the transition elements are—according to the
Aufbau principle—the last to fill when building up electron configurations, the outermost s electrons
are the first to be lost when these atoms ionize. When the inner transition metals form ions, they
usually have a 3+ charge, resulting from the loss of their outermost s electrons and a d or f electron.
Example: Determining the Electronic Structures of Cations
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3.1 Ionic Bonding
18:57
Learning Objectives
By the end of this section, you will be able to:
Explain the formation of cations, anions, and ionic compounds.
Predict the charge of common metallic and nonmetallic elements and write their electron
configurations.
Ionic Bonding
As you have learned, ions are atoms or molecules bearing an electrical charge. A cation (a positive
ion) forms when a neutral atom loses one or more electrons from its valence shell, and an anion (a
negative ion) forms when a neutral atom gains one or more electrons in its valence shell.
Compounds composed of ions are called ionic compounds (or salts), and their constituent ions are
held together by ionic bonds: electrostatic forces of attraction between oppositely charged cations
and anions. The properties of ionic compounds shed some light on the nature of ionic bonds. Ionic
solids exhibit a crystalline structure and tend to be rigid and brittle; they also tend to have high
melting and boiling points, which suggests that ionic bonds are very strong. Ionic solids are also poor
conductors of electricity for the same reason—the strength of ionic bonds prevents ions from moving
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freely in the solid state. Most ionic solids, however, dissolve readily in water. Once dissolved or
melted, ionic compounds are excellent conductors of electricity and heat because the ions can move
about freely.
Neutral atoms and their associated ions have very different physical and chemical properties. Sodium
atoms form sodium metal, a soft, silvery-white metal that burns vigorously in air and reacts
explosively with water. Chlorine atoms form chlorine gas, Cl2, a yellow-green gas that is extremely
corrosive to most metals and very poisonous to animals and plants. The vigorous reaction between
the elements sodium and chlorine forms the white, crystalline compound sodium chloride, common
table salt, which contains sodium cations and chloride anions. The compound composed of these
ions exhibits properties entirely different from the properties of the elements sodium and chlorine.
Chlorine is poisonous, but sodium chloride is essential to life; sodium atoms react vigorously with
water, but sodium chloride simply dissolves in water.
Figure 3.1 (a) Sodium is a soft metal that must be stored in mineral oil to prevent reaction with air or
water. (b) Chlorine is a pale yellow-green gas. (c) When combined, they form white crystals of sodium
chloride (table salt). (credit a: modification of work by “Jurii”/Wikimedia Commons)
The Formation of Ionic Compounds
Binary ionic compounds are composed of just two elements (thus, the prefix “bi,” meaning two): a
metal (which forms the cations) and a nonmetal (which forms the anions). For example, NaCl is a
binary ionic compound. We can think about the formation of such compounds in terms of the periodic
properties of the elements. Many metallic elements have relatively low ionization potentials and lose
electrons easily. These elements lie to the left in a period or near the bottom of a group on the
periodic table. Nonmetal atoms have relatively high electron affinities and, thus, readily gain electrons
lost by metal atoms, thereby filling their valence shells. Nonmetallic elements are found in the upper-
right corner of the periodic table.
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As all substances must be electrically neutral, the total number of positive charges on the cations of
an ionic compound must equal the total number of negative charges on its anions. The formula of an
ionic compound represents the simplest ratio of the numbers of ions necessary to give identical
numbers of positive and negative charges. For example, the formula for aluminum oxide, Al2O3,
indicates that this ionic compound contains two aluminum cations, Al3+, for every three oxide anions,
O2− ,
[thus, (2 × +3) + (3 × –2) = 0].
It is important to note, however, that the formula for an ionic compound does not represent the
physical arrangement of its ions. It is incorrect to refer to a sodium chloride (NaCl) “molecule”
because there is not a single ionic bond, per se, between any specific pair of sodium and chloride
ions. The attractive forces between ions are isotropic—the same in all directions—meaning that any
particular ion is equally attracted to all of the nearby ions of opposite charge. This results in the ions
arranging themselves into a tightly bound, three-dimensional lattice structure. Sodium chloride, for
example, consists of a regular arrangement of equal numbers of Na+ cations and Cl– anions.
Figure 3.2 The atoms in sodium chloride (common table salt) are arranged to (a)
maximize opposite charges interacting. The smaller spheres represent sodium ions,
the larger ones represent chloride ions. In the expanded view (b), the geometry can
be seen more clearly. Note that each ion is “bonded” to all of the surrounding ions—
six in this case.
The strong electrostatic attraction between Na+ and Cl– ions holds them tightly together in solid NaCl.
It requires 769 kJ of energy to dissociate one mole of solid NaCl into separate gaseous Na+ and Cl–
ions:
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NaCl(s) + 769 kJ Na+(g) + Cl–(g)
Electronic Structures of Cations
When forming a cation, an atom of a main group element tends to lose all of its valence electrons,
thus assuming the electronic structure of the noble gas that precedes it in the periodic table. For
groups 1 (the alkali metals) and 2 (the alkaline earth metals), the group numbers are equal to the
numbers of valence shell electrons and, consequently, to the charges of the cations formed from
atoms of these elements when all valence shell electrons are removed. For example, calcium is a
group 2 element whose neutral atoms have 20 electrons and a ground state electron configuration of
1s22s22p63s23p64s2. When a Ca atom loses both of its valence electrons, the result is a cation with
18 electrons, a 2+ charge, and an electron configuration of 1s22s22p63s23p6. The Ca2+ ion is,
therefore, isoelectronic with the noble gas Ar.
For groups 13–17, the group numbers exceed the number of valence electrons by 10 (accounting for
the possibility of full d subshells in atoms of elements in the fourth and greater periods). Thus, the
charge of a cation formed by the loss of all valence electrons is equal to the group number minus 10.
For example, aluminum (in group 13) forms 3+ ions (Al3+).
Exceptions to the expected behavior involve elements toward the bottom of the groups. In addition to
the expected ions Tl3+, Sn4+, Pb4+, and Bi5+, a partial loss of these atoms’ valence shell electrons can
also lead to the formation of Tl+, Sn2+, Pb2+, and Bi3+ ions. The formation of these 1+, 2+, and 3+
cations is ascribed to the inert pair effect, which reflects the relatively low energy of the valence s-
electron pair for atoms of the heavy elements of groups 13, 14, and 15. Mercury (group 12) also
exhibits an unexpected behavior: it forms a diatomic ion, Hg22+ (an ion formed from two mercury
atoms, with an Hg-Hg bond), in addition to the expected monatomic ion Hg2+ (formed from only one
mercury atom).
Transition and inner transition metal elements behave differently than main group elements. Most
transition metal cations have 2+ or 3+ charges that result from the loss of their outermost s
electron(s) first, sometimes followed by the loss of one or two d electrons from the next-to-outermost
shell. For example, iron (1s22s22p63s23p64s23d6) forms the ion Fe2+ (1s22s22p63s23p63d6) by the
loss of the 4s electrons and the ion Fe3+ (1s22s22p63s23p63d5) by the loss of the 4s electrons and
one of the 3d electrons. Although the d orbitals of the transition elements are—according to the
Aufbau principle—the last to fill when building up electron configurations, the outermost s electrons
are the first to be lost when these atoms ionize. When the inner transition metals form ions, they
usually have a 3+ charge, resulting from the loss of their outermost s electrons and a d or f electron.
Example: Determining the Electronic Structures of Cations
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