Chemical
Bonds
6-1 Ionic Bonding

When is an atom unlikely to
react?

What is one way in which
elements can achieve stable
electron configurations?

How does the structure of an
ionic compound affect its
properties?
Stable Electron Configurations

When the highest occupied
energy level of an atom is
filled with electrons, the
atom is stable and not likely
to react.


The highest occupied energy
level of a noble gas atom is
filled.
The noble gases have stable
electron configurations with
eight valence electrons (or 2
in the case of helium)
The Electron Dot Diagram

Since the chemical
properties of an element
depend on the number
of valence electrons, it is
useful to have a model
that focuses on valence
electrons.


A model of an atom in
which dot represents
valence electrons
Symbol in the center
represents the nucleus
and all the other
electrons
Ionic Bonds



Elements that do not have complete sets of valence
electrons tend to react.
By reacting, they achieve electron configurations similar to
those of noble gases.
Some elements achieve stable electron configurations
through the transfer of electrons between atoms.

Transfer of Electrons
Chlorine to Argon
Sodium to Neon
Ionic Bonds Continued

Formation of Ions



When atoms gain or lose an electron, the number of protons is no
longer equal to the number of electrons.
The charge is not balanced and atom is not neutral.
An atom that has a net positive or negative electric charge is an
ion. This is represented by a + or – sign next to the symbol.


Chlorine gains an electron =17 protons, 18 electrons = 1- charge
Sodium has 11 protons and 10 electrons (one extra proton) = 1+ charge
Negative charge - anion…suffix = ide
Positive charge – cation… just use element name
Ionic Bonds Continued

Formation of Ionic Bonds


Particle with negative charge will
attract a particle with a positive
charge
When anion and cation are close,
a chemical bond forms.



Chemical bond: force that holds
atoms or ions together as a unit
(attraction between them)
Ionic bond: force that holds cations
and anions together (when
electrons are transferred from one
atom to another)
Ionization Energy:




The amount of energy used to
remove an electron
An electron can move to a higher
energy level when an atom
absorbs energy because the
energy allows the electron to
overcome the attraction of the
protons in the nucleus..
Varies from element to element
The lower the energy, the easier it is
to remove an electron from an
atom
Ionic Compounds

Compounds that contain ionic bonds are ionic
compounds.



can be represented by chemical formulas (a notation
that shows what elements a compound contains and
the ratio of the atoms or ions of those elements in the
compound)
Example: NaCl = one sodium ion for each chloride ion
in sodium chloride
Magnesium chloride = ?
Ionic Compounds Continued
 To
A chemical formula for an ionic compound
tells you the ratio of the ions in the
compound
 It does not tell you how the ions are
arranged in the compound
Solids arranged in a lattice structure are
called crystals


review
Ionic Compounds Continued

A: sodium chloride: ions arranged in an
orderly three-dimensional structure (cubic
shape)
Crystals: solids arranged in a lattice structure

B: sodium chloride: crystals look like cubes
The properties of an ionic compound can be
explained by the strong attractions among ions
within a crystal lattice.
Class Participation Opportunity: Research
report on how to make rubies
6-2 Covalent Bonding
 How
are atoms held together in a
covalent bond?
 What happens when atoms don’t share
electrons equally?
 What factors determine whether a
molecule is polar?
 How do attractions between polar
molecules compare to attractions
between nonpolar molecules?
Covalent Bonds

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
Plants absorb water through roots
Carbon dioxide from the air enters
through stomata in leaves
Plants use energy from sun to convert
water and carbon dioxide into sugar
Energy is stored in the chemical bonds of
the sugar
The elements of sugar are carbon,
oxygen and hydrogen(nonmetals with
high ionization energies so a transfer of
electrons does not tend to occur).
When nonmetals join, they share
valence electrons forming a covalent
bond.
Covalent Bonds Continued
 The
attractions between the shared electrons
and the protons in each nucleus hold the
atoms together in a covalent bond.
 Sharing Electrons:



A hydrogen atom has one electron
Two hydrogen atoms can achieve a stable
electron configuration by sharing their electrons
When two atoms share one pair of electrons the
bond is called a single bond
Covalent Bonds Continued
Four different ways to represent a covalent bond.
• Electron dot model = pair of dots
• Structural formula = line
• Electron cloud model and space-filling model = orbitals of atoms
overlap
Hydrogen atoms bonded together form a molecule (a neutral group of
atoms joined together by one or more covalent bonds).
The attraction between the share electrons and the protons in each
nucleus hold the atoms together in a covalent bond.
Covalent Bonds Continued
Many nonmetal elements exist as
diatomic molecules (two atoms).
 Why are the atoms in the models
of diatomic molecules not
complete spheres?
The space-filling models show that
orbitals of atoms overlap when they form
covalent bonds.

Why are the spheres in the models
of fluorine, chlorine, and bromine
different sizes?
The different sizes of spheres model the
different atomic radii of the atoms.
Covalent Bonds Continued
 Multiple



Covalent Bonds
Nitrogen = five valence electrons
Two nitrogen atoms = shared a pair of electrons,
each would have seven
Nitrogen molecule = each molecule shared
three pairs of electrons, each atoms has eight
valence electrons

triple bond
Unequal Sharing of Electrons
Generally, elements on the right and on the top of the
periodic table have a greater attraction for electrons
that elements on the left (except noble gases).

In Polar Covalent Bonds, electrons are not shared
equally.

When atoms form a polar covalent bond, the atom with
the greater attraction for electrons has a partial
negative charge. The other atom has a partial positive
charge.

Polar bonds occur in molecules have only two atoms.
Unequal Sharing of Electrons
Continued

Polar and Nonpolar Molecules

The type of atoms in a molecule and its shape
are factors that determine whether a molecule is
polar or nonpolar.

In carbon dioxide, there are double bonds
between each oxygen atom and the central
carbon atom (because oxygen has a greater
attraction for electrons than carbon does, each
double bond is polar). The molecule is linear.
Oxygen atoms are opposite each other. There is
an equal pull on the electrons from opposite
directions. These pulls cancel out and the
molecule is nonpolar.

In a water molecule, there are two single polar
bonds. Oxygen has a greater attraction for
electrons than hydrogen does. The molecule is
bent; the polar bonds do not cancel out.
Oxygen has partial negative charge; hydrogen
has a partial positive charge.
Attraction Between Molecules
There are forces of attraction
between molecules but they
are not as strong as the
attractions between ionic and
covalent bonds. They are
strong enough to hold the
molecules together in a liquid or
solid.

Attractions between polar
molecules are stronger than
attractions between
nonpolar molecules.
6-3 Naming Compounds and
Writing Formulas
 What
information do the name and
formula of an ionic compound provide?
 What
information do the name and
formula of a molecular compound
provide?
Chemists name compounds based on
their compositions to avoid confusion.

Thomas Drummond discovered the white solid
lime(calcium oxide) that emits a bright light when
heated to a high temperature.


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
Lime wash
Quicklime
Unslaked lime
Can be mixed
with paint for
whitewashing
To avoid confusion, chemists call lime calcium oxide.
Describing Ionic Compounds
The vase and the plate are both coated in oxides
of copper (copper and oxygen), one red and one
black.
Calling both oxides by the same name won’t
work, so there must be two names.


The name of an ionic compound must distinguish
the compound from other ionic compounds
containing the same elements.
The formula of an ionic compound describes the
ratio of the ions in the compound.
Describing Ionic Compounds
Continued
Binary Ionic Compounds
A compound made from two elements is a binary compound.
The name of the cation is followed by the name of the anion.
cation = metal without any change
sodium atom and sodium ion
anion = uses part of the name of the nonmetal with the
suffix ide
iodine atom and iodide ion
Describing Ionic Compounds Continued

Metals with Multiple Ions

Many transition metals form more than on type of ion


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Copper (I) 1+
Copper (II) 2+
Red copper (I)oxide – “copper one oxide”
Cu2O →Cu1+O2-

Black copper (II) oxide – “copper two oxide”
CuO→ Cu2+ + O2-
Describing Ionic Compounds Continued
 Polyatomic



Ions – ammonium ion in picture (below right)
A covalently bonded group of atoms
that has a positive or negative charge on
and acts as a unit is a polyatomic ion.
Most are anions (below left)
In iron (III) hydroxide, the subscript 3
indicates that there are three hydroxide ions
for each iron (III) ion.
Writing Formulas for Ionic Compounds
 Write
the cation first, followed by
the symbol of the anion
 Use subscripts to show the ratio of
the ions in the compound
 Since all compounds are neutral,
the total charges on the cations
and anions must add up to zero
Two Na (Na+) for each S(S2-)
Na2S
Writing a Formula for Ionic Compounds
Describing Molecular Compounds

The name and formula of a molecular
compound describe the type and number of
atoms in a molecule of the
compound(unlike ionic compounds).

In other words, the names and formulas
identify (match) specific compounds.
Most metallic elements appears first in the
name (farther to the left in periodic table
or if in same group, closer one to the
bottom)
 Second element is changed t end in the
suffix –ide
 Examples:
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carbon dioxide - CO2
Dinitrogen tetraoxide - N2O4
Mononitrogen dioxide, commonly written
as nitrogen dioxide - NO2
Diphosphorus tetraflouride - P2F4
6-4 The Structure of Metals
 What
are the forces that give a metal its
structure as a solid?
 How do metallic bonds produce some of
the typical properties of metals?
 How are the properties of alloys
controlled?
The Light Bulb: A New Technology?
Just before the year1900, the light bulb was
cutting edge technology.
Researchers were busy working on
finding the best material for the
filaments
*had to be ductile enough
*couldn’t melt
*had to have low vapor
pressure
Found Tungsten (W)
*highest melting point
*lowest vapor pressure
Metallic Bonds

Metal atoms achieve stability by losing electrons.

If there are no nonmetal atoms available to accept the
electrons, valence electrons move among the atoms. (The
metal atoms become cations surrounded by a pool of
shared electrons.)

This attraction between metal cation and the shared
electrons is a metallic bond.

The cations in a metal form a lattice that is held in place
by strong metallic bonds between the cations and the
surrounding valence electrons.
Metallic Bonds Continued
Even though the electrons are moving among
the atoms, the total number of electrons does
not change so the metal is neutral.
The more valence electrons an atom can
contribute to the pool, the stronger the bond
will be.
alkali metal = weak bond (1 valence
electron), soft with low melting point
transition metal = stronger bond )more
valence electrons), harder with higher
melting points
Explaining Properties of Metals


The structure within a metal affects its properties.
The mobility of electrons within a metal lattice explains some of the
properties of metals.

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Conducting electricity:

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Examples: ability to conduct an electric current and malleability
Metals have a pool of electrons, a built-in supply of charged particles that can
flow from one location to another
Electric currents are carried by these free flowing shared electrons.
Malleability

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
The lattice in metals is flexible (unlike in rigid ionic compounds).
This is why tungsten and copper can be made into thin wires without breaking.
Picture below illustrates how ions shift in position and shape changes but the
metal does not shatter because of metallic bonds between the ions and
electrons.
Alloys

Scientists can design alloys with specific properties
by varying the types and amounts of elements in
an alloy.
 Gold
100% pure gold is expressed in karats…24 karat
gold
 Soft metal that can be easily worn away


Gold alloys
Gold mixed with silver, copper, nickel or zinc
gold is more resistant to wear
 12 karat gold is only 50% gold
 18 karat gold is 75% gold

Copper and Steel Alloys

Copper Alloys

First most important alloy of copper = bronze

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Brass is another alloy of copper (contains copper and zinc)

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Associated with the Bronze Age
In its simplest form, bronze contains only copper and tin,
which when mixed together, are stronger than they are
alone
Can be used for ships, statues and bells
Bronze and brass have distinctly different properties
Brass shiner, softer, weathers more quickly
Steel Alloys



1900s = Steel Age = skyscrapers, automobiles and ships
Alloy of iron and some carbon
Properties depend on an element other than iron and
carbon

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Stainless steel = more than 10% chromium = durable and
resists rust
Steel cables = sulfur, manganese phosphorus and silicon =
stronger
Another Example
 Airplane


parts
Aluminum is lightweight but too soft and
dents too easily
With a small amount of copper or
manganese, the aluminum is stronger but
still lighter than steel
A Historical Perspective