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Chapter 9:
Chemical Bonds
Compounds and
Chemical Change
• Chemistry is all about bonds being made or broken.
• Bonds are simply the sharing or exchanging of electrons.
• A chemical reaction occurs when substances are created by
the breaking or formation of chemical bonds.
• Example: H2 and O2 can react (bonds formed) to make H2O
or H2O can decompose (bonds broken) to form H2 and O2.
• Left: The nobel gases, example Neon above, are
monatomic- meaning they naturally occur as stable, single
atoms. Right: Many gases, such as Hydrogen, are diatomic,
which means they naturally occur as a molecule with two
atoms. Not shown: Some gasses, such as Ozone- a form of
oxygen, occur in a three atom arrangement called triatomic.
Decomposition of Water
• The figure shows the decomposition of water into H2
and O2. The reaction also works in reverse to make
water from H2 and O2.
• Note: both H2 and O2. are diatomic gases..
Atoms and Molecules
• An atom is the smallest unit of an element that
can exist alone or in combination with other
elements.
• A molecule is the smallest particle of a
compound, or a gaseous element, that can exist
and still retain the characteristic chemical
properties of a substance.
Chemical Reaction
• A chemical reaction is a change in matter in which
different chemical substances are created by forming
or breaking chemical bonds.
– Chemical bonds are formed when atoms of elements bind
together to form compounds.
– Chemical bonds are broken when a compound is
decomposed into simpler substances.
Chemical Equations
• A chemical equation is a symbolic representation of
a chemical reaction.
– The reactants are the substances that are changed in a
reaction . They are always on the left side of a written
equation.
– The products are the substances that are formed in the
reaction. They are always written on the right side of the
equation.
Valence Electrons
• Valence electrons are those electrons that reside in
the outer shell of an atom.
– These are the electrons that are lost/gained or shared in a
chemical reaction.
– Inner electrons are in filled orbitals in lower energy
levels.
• These electrons are therefore not available for interactions with
other electrons.
The Octet Rule
– Atoms attempt to acquire an outer orbital with eight
electrons through chemical reactions.
– This gives them an outer shell configuration like their
nearest noble gas and therefore they become stable.
– From the family number of the representative elements,
you can determine the number of valence electrons, and
therefore the number of electrons necessary to gain the
stable configuration.
Octet Rule Example
Sodium is in Group IA and therefore has 1
electron in its outer shell -1 valence electron.
• If sodium loses one electron if becomes Na+
and has 8 electrons in its outer shell.
• It then has the electron configuration of Ne
(neon) and has a filled outer shell
configuration.
Chemical Bonds
• Introduction
– A chemical bond that is an attractive force that holds
the atoms in a compound together.
– An Ionic bond forms when an atom transfers an
electron to another atom during a chemical reaction.
– A covalent bond forms when atoms share electrons in
a chemical bond.
– Metallic bonds form in metals.
Bond Theories
– Bonds that form in compounds can be described in
several ways.
– Molecular orbital theory describes the electrons as
belonging to the whole molecule which gives the orbital
its own shape, orientation, and energy levels.
– Isolated atom description considers the electrons
around the atoms as being isolated from the rest of the
molecule.
Ionic Bonds
– Ionic bonding occurs when one atom transfers an
electron to another atom.
• The difference in electrical charge results in an
electrostatic attraction between unlike electrical charges.
• This occurs when a metal reacts with a nonmetal.
• A crystal of sodium chloride is composed of sodium and
chlorine ions held together by electrical attraction. Each
positive sodium ion is surrounded by six negative
chlorine ions, and each chlorine ion is surrounded by six
sodium ions. This atomic arrangement gives the sodium
chloride crystal a cubic structure.
Salt Crystals
• Left: The cubic structure of table salt crystals is easily
seen when magnified about ten times.
• Right: Model of a salt crystal.
Heat of Formation
– Energy and Electrons in Ionic Bonding
– energy + Na+  Na+ + e– Cl + e-  Cl- + energy
– Na+ + Cl-  Na+Cl- + energy
• The energy that is released in steps 2 and 3 is greater
that that absorbed in step one and an ionic bond is
formed.
– This energy is called the heat of formation.
Ionic Bonding Rules
• Two rules for keeping track of electrons in ionic
bonding reactions.
– Ions are formed when atoms gain or lose electrons
to achieve a noble gas configuration.
– The number of electrons that are lost must equal
the number of electrons that are gained.
Compounds/Formulas
– Ionic Compounds and Formulas
• The formula of a compound describes what elements
are in the compound and in what proportions.
• Compounds that are held together by ionic bonds are
called ionic compounds.
• The elements in Group IA and IIA tend to lose
electrons for form positive ions
• The elements in Group VIA and VIIA tend to gain
electrons to form negative ions.
Covalent Bonds
– A covalent bond is a chemical bond that is formed when
two atoms share a pair of electrons.
– H. + H.  H:H
– Covalent Compounds and Formulas
• Since a pair of electrons is shared in a covalent bond, the
electrons move throughout the entire molecular orbital.
• In the above example, both hydrogen atoms gain the electron
configuration of helium.
• Covalent compounds are compounds with covalent bonds.
• Covalent compounds form from atoms on the right side of the
periodic table (nonmetals).
Multiple Bonds
• Bonding pairs can also be represented by lines
connecting atoms. H:H = H—H
• When one pair of electrons is shared, it is called
a single bond.
• When two pairs of electrons are shared it is
called a double bond.
• When three pairs of electrons are shared it is
called a triple bond.
Electronegativity
– Electronegativity is a measure of the relative ability of
an atom to attract bonding electrons.
– Elements with higher values have the greatest affinity for
bonding electrons.
– Differences in electronegativity can be used to predict
whether a bond will be ionic or covalent.
• If the absolute difference is 0.5 or less, the bond will be
covalent.
• If the absolute difference is 1.7 or more the bond will be ionic.
Ionic Compound Names
– Ionic compounds that are formed from metal ions are
named by naming the metal ion (electropositive ion) first,
followed by the nonmetal (electronegative ion).
– The ending of the nonmetal is changed to end in -ide
– When a metal can have various oxidation states the
oxidation state is give by roman numerals in parenthesis
after the name of the metal.
Ionic Compound Formulas
– Rules
• The symbol for the positive element is written first,
followed by the symbol of the negative element
• Subscripts are used to indicate the numbers of ions needed
to produce an electrically neutral compound.
• Example calcium chloride
– calcium is Ca2+ and chlorine is Cl– there needs to be two negative charges to balance the
2+ on the calcium
– the formula is therefore CaCl2
Covalent Compound Names
– Names for covalent compounds uses Greek prefixes to indicate
numbers of atoms of each element
• The first element in the formula is named first with a prefix
indicating the number of atoms if the number is greater than
one.
• The stem name of the second element in the formula is named
next, with a prefix used with the stem name if two elements can
form more than one compound
– The suffix –ide is again used.
• Example
– CO = carbon monoxide
– CO2 = carbon dioxide
Covalent Compound Formulas
– The systematic name tells you the formula.
– Formulas indicate how many atoms of one element
combine with atoms of another element.
– The number of covalent bonds that an atom can form is
called its valence.
– Lone pairs create the possibility of creating coordinate
covalent compounds.