Chemical Bonding
• Chemical Bond-force that holds two atoms
together.
Atoms either transfer electrons and then form
ionic compounds or they share electrons to
form covalent compounds.
• In both cases, the bond forms because of an
increase in stability.
Chemical Bonding
• By forming bonds, atoms acquire an octet of
electrons and the stable electron
configuration of a noble gas.
Electron-dot configuration
Formation of Positive Ions
• Atoms can lose one or more valence electrons
to obtain noble gas electron configuration and
form a positive ion called a Cation.
• Sodium has one 3s valence electron which it
loses to form a Na+ cation.
Transition Elements
Transition elements have full s orbitals and as you move from left to
right on the periodic table electrons are being added to inner d orbitals.
Transition metals generally lose the two s orbital electrons to form
2+ cations but in some cases d orbital electrons can be lost to form
cations with charges greater than 2+
Pseudo-Noble Gas Configuration
• Elements from periods 4-6 lose electrons to form
outer orbitals with full s, p and d orbitals.
Zinc ion formation
Zinc ([Ar]4s23d10)
Zinc2+ ([Ar]3d10) + 2e-
pseudo-noble gas configuration
Formation of Negative Ions
• Nonmetals have large attraction for electrons and
attain stability by gaining electrons to form negative
ion called an Anion.
ex. Chlorine (1s22s22p63s23p5)
Chloride anion (1s22s22p63s23p6)
[Ar] valence configuration
• Nomenclature for anion is the addition of –ide to
name.
i.e. Cl-1 is named chloride
The charge on the ion is known as the
oxidation number of the atom.
Formation of Ionic Bond
• Ionic Bond- electrostatic force that holds
opposite charged particles together in an
ionic compound.
Formation of Ionic Compound NaCl
• Sodium is in Group 1, so it has one valence electron. Chlorine
is in Group 17 and has seven valence electrons.
• The one valence electron of sodium is transferred to the
chlorine atom, both become stable with noble gas electron
configurations.
• The chlorine atom now has an extra electron and has a
negative charge.
• Sodium lost an electron and has a positive charge.
Ionic Compounds Nomenclature
• Most ionic compounds are called salts.
• Metal + Oxygen form ionic compounds called
oxides.
Binary Ionic Compounds
• Ionic compounds with only two different
elements are termed Binary.
Binary ionic compounds can contain more
than one ion of each element, as in CaF2, but
they are not composed of three or more
different elements.
Binary Ionic Compounds
To name a binary ionic compound, first write the name
of the positively charged ion, usually a metal, and
then add the name of the nonmetal or negatively
charged ion, whose name has been modified to end
in -ide.
Potassium combines with chlorine to form potassium
chloride salt.
Magnesium combines with oxygen to form
magnesium oxide.
Formation of Aluminum Oxide
• Aluminum is a Group 13 metal, so it loses its three
outer electrons to become an Al3+ ion; oxygen is in
Group 16 and has six valence electrons, so it gains
two electrons to become an O2- ion.
• All the electrons must be accounted for, therefore
more than one oxygen atom must be involved in the
reaction.
Formation of Aluminum Oxide
• In all, two Al3+ ions must combine with three
O2- ions to form Al2O3.
•Remember that the charges in the formula
for aluminum oxide must add up to zero.
Properties of Ionic Compounds
• Chemical bonds between atoms in a
compound determine many physical
properties of the compound.
• In an ionic compound the positive and
negative ions are packed into a regular
repeating pattern that balances the forces of
attraction and repulsion between ions forming
Ionic Crystals.
Ionic Crystal
Properties of Ionic Compounds
• Ionic compounds are composed of wellorganized, tightly bound ions forming a
strong, three-dimensional crystal structure.
• Ionic compounds are crystalline solids at room
temperature with relatively high melting and
boiling points.
• In the solid-state ionic compounds are
nonconductive due to fixed positions of the
ions.
Properties of Ionic Compounds
• Another physical property of ionic compounds is
their tendency to dissolve in water and conduct
electricity while in solution.
• Any compound that conducts electricity when
dissolved in water is an electrolyte.
In order to conduct electricity, ions must be free to
move because they must take on or give up
electrons.
Energy and Ionic Bonds
• Endothermic- energy is absorbed during a
reaction.
• Exothermic- energy is released during a reaction.
All ionic reactions of cations and anions are
exothermic.
Resultant compounds are more stable
configuration (i.e. lower energy level) so excess
energy is released. The amount of energy released is
equal to amount needed to break the resultant
bond.
Energy and Ionic Bonds
• The energy required to separate one mole of the ions
of an ionic compound is called lattice energy,
which is expressed as a negative quantity.
i.e. The greater (more negative) the lattice
energy is, the stronger the force of attraction
between the ions.
Lattice energy tends to be greater for more-highlycharged ions and for small ions than for ions of lower
charge or large size.
Ionic Compound Nomenclature
• No single particle of an ionic compound
exists so they are represented by a formula
that provides the simplest ratio of the ions in
an ionic compound and is called a formula
unit.
The overall charge of any formula unit is
zero.
Ionic Compound Nomenclature
• The charges of monatomic ions, or ions
containing only one atom, can often be
determined by referring to the periodic table
or table of
common
ions
based on
group
number.
Ionic Compound Nomenclature
• The charge of a monatomic ion is equal to
its oxidation number.
• The oxidation number, or oxidation state, of
an ion in an ionic compound is numerically
equal to the number of electrons that were
transferred to or from an atom of the
element in forming the compound.
Ionic Compound Nomenclature
• In the formula for an ionic compound, the
symbol of the cation is written before that
of the anion.
• Subscripts, or small numbers written to the
lower right of the chemical symbols, show
the numbers of ions of each type present in
a formula unit.
Ionic Compound Nomenclature
• If the ions in the ionic compound have the
same charge, the formula unit contains one of
each ion.
– Na+ and Cl- combine to form NaCl.
– Mg2+ and S2- combine to form MgS.
• If the charges are not equal, we must balance
the positive and negative charges.
– Ca2+ and Cl- combine to form CaCl2.
– Na+ and O2- combine to form Na2O.
Ionic Compound Nomenclature
• In naming ionic compounds, name the cation
first, then the anion.
• Monatomic cations use the element name.
• Monatomic anions use the root of the
element name plus the suffix -ide.
• If an element can have more than one
oxidation number, use a Roman numeral in
parentheses after the element name, for
example, iron(II) to indicate the Fe2+ ion.
Crossover Rule
• You can quickly verify that the chemical formula is
written correctly by crossing over the charge on each
ion.
• The charge on the aluminum ion becomes the
subscript for the oxygen, and the charge on the oxide
ion becomes the subscript for the aluminum ion.
Compounds with Polyatomic Ions
• Some ions contain more than one element.
• An ion that has two or more different
elements is called a polyatomic ion.
• Although the individual atoms have no charge,
the group as a whole has an overall charge.
Nick the Camel
Polyatomic Ions
• Ionic compounds may contain:
• Positive metal ions bonded to negative polyatomic
ions, such as in NaOH sodium hydroxide
• Negative nonmetal ions bonded to positive
polyatomic ions, such as in NH4I ammonium iodide
• Positive polyatomic ions bonded to negative
polyatomic ions, such as in NH4NO3 ammonium
nitrate.
Nomenclature with polyatomics
• Follow the same rules as binary ionic compounds; if the
charges are equal, the formula has one of each ion.
– Mg2+ and SO42- combine to form MgSO4
Magnesium sulfate
– K+ and ClO3- combine to form KClO3
Potassium Chlorate
• If the charges are not equal, the total charge must equal
zero. If you have more than one polyatomic ion, it is
placed in parentheses.
– Al3+ and CO32- combine to form Al2(CO3)2.
Aluminum carbonate
Polyatomic Ions
• To name a compound containing a polyatomic
ion, follow the same rules as used in naming
binary compounds.
However, do not change the ending of the
negative polyatomic ion name.
i.e. To form a neutral compound, one calcium
ion (Ca2+) must combine with one carbonate
ion (CO3 2–) to give calcium carbonate with the
formula CaCO3.
Ions of Transition Elements
• Transition elements form positive ions just as other
metals do, but most transition elements can form
more than one type of positive ion and have multiple
oxidation states.
• Zinc and silver are two exceptions each forms one
type of ion. Zinc ion is Zn2+ and the silver ion is Ag+.
• A Roman numeral is used to indicate the oxidation
number of a transition element ion.
Naming Ionic Compounds
• Certain polyatomic ions, called
oxyanions, contain oxygen and another
element.
• If two different oxyanions can be formed by
an element, the suffix -ate is used for the
oxyanion containing more oxygen atoms,
and the suffix -ite for the oxyanion
containing fewer oxygens.
Oxyanions
• In the case of the oxyanions of the halogens, the
following special rules are used.
– four oxygens, per+root+ate (ex: perchlorate, ClO4–)
– three oxygens, root + -ate (ex: chlorate, ClO3-)
– two oxygens, root + -ite (ex: chlorite, ClO2-)
– one oxygen, hypo+root+ite (ex: hypochlorite, ClO–)
Naming Ionic Compounds
1. NaBrO3
(sodium bromate)
2. Mg(NO3)2
(magnesium nitrate)
3. NH4ClO4
(ammonium perchlorate)
4. Al(ClO)3
(aluminum hypochlorite)
Metallic Bonds and Properties of Metals
• The bonding in metals is explained by the
electron sea model, which proposes
that the atoms in a metallic solid contribute
their valence electrons to form a “sea” of
electrons that surrounds metallic cations.
• These delocalized electrons are not
held by any specific atom and can move
easily throughout the solid.
• A metallic bond is the attraction
between these electrons and a metallic
cation.
Properties of Metals
• Metals generally have extremely high boiling points
because it is difficult to pull metal atoms completely
away from the group of cations and attracting
electrons.
Metals are also malleable (able to be hammered into
sheets) and ductile (able to be drawn into wire)
because of the mobility of the particles.
The delocalized electrons make metals good
conductors of electricity.