Chemistry 110
Chapter Four Notes
Chemical Bonds
Valence Electrons
„ The electrons occupying the outermost energy level of an atom are called the valence electrons; all other
electrons are called the core electrons.
„ The valence electrons, as we will see, are responsible for chemical bonding.
{ Knowing the number of valence electrons an atom has is the single most important information
you can have in predicting how an atom will react chemically.
„ For elements in group IA through VIIIA, the number of valence electrons an atom has is simply its
group number.
{ So phosphorus, in group VA, has five valence electrons.
„ Helium is the only significant exception; it has 2 valence electrons, even though it is in group VIIIA
Lewis Dot Diagrams
„ Lewis dot diagrams are simple symbols which show the symbol of an element surrounded by as many
dots as that atom has valence electrons.
„ The first four electrons are drawn (in any order) on each of the four “sides” of the symbol.
„ The next four electrons are “paired up” with the other four atoms (again, in any order).
„ These symbols will be used to model chemical bonding.
Examples
Draw the Lewis dot diagrams for calcium, arsenic, and iodine atoms.
The Octet Rule
„ The Octet Rule states that atoms react in such a way as to give them eight electrons in their outermost
energy level (their valence shell).
„ The Noble Gases, those elements in the period on the far right of the Periodic Table, generally do not
react with other atoms, as their valence shell is filled with eight electrons (or two in the case of helium).
„ By gaining or losing enough electrons to give an atom eight electrons in its valence shell, the atom gains
the stable features of a noble gas.
Getting an Octet
„ Atoms usually achieve an octet in the following way:
{ Metals lose their valence electrons, giving them an empty outer shell.
„
Note that the ion now has the electron configuration of a noble gas.
{ Nonmetals gain or share enough electrons to give themselves eight electrons in their outer shell.
„
Nonmetals generally share with other nonmetal atoms, or take electrons away from metal
atoms.
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Chemistry 110
Chapter Four Notes
Example
How many electrons will usually be gained or lost by each of the following atoms?
Aluminum
Bromine
Oxygen
Barium
Potassium
Neon
The Duet Rule
„ Hydrogen atoms attempt to pick up another electron, giving them the same electron configuration as the
noble gas helium; this is called The Duet Rule.
„ Likewise, when lithium loses an electron to become Li+, it has the same electron configuration as
helium.
A Cautionary Note
„ Gaining and losing electrons does not occur unless there are other atoms around to encourage this.
{ A sodium atom cannot just toss its valence electron away; it must give it to another atom which
wants to take it, like chlorine.
„ Considering this, be sure that you do not confuse an atom of an element (which has its usual number of
electrons) with an ion (which has gained or lost electrons).
What is Bonding?
„ Bonding describes how atoms interact with each other in an attractive sense.
„ There are two types of bonding which will be discussed in this class:
{ Ionic bonding
{ Covalent bonding
Ionic Bonding
„ Ionic bonding occurs between ions of different charges.
„ Often, but not always, ionic bonding occurs between the ion of metal with the ion of a nonmetal.
2{ Sometimes, one of these ions is a polyatomic ion, such as sulfate (SO4 ).
„ Ionic bonding is generally the strongest of the bonding types.
{ It requires a great amount of energy to separate the ions from each other, resulting in high melting
points (500 oC to 2,000 oC).
Examples of Ionic Compounds
NaCl
CaO
CuBr2
LiF
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Chemistry 110
Chapter Four Notes
Formulas of Ionic Compounds
„ An ionic compound should have no total charge in its formula.
„ Ions of opposite charge are paired in such a way as to cancel out charges.
{
Positively charged ions are called _____________________________
{
Negatively charged ions are called _____________________________
„ Example:
One Na+ combines with one Cl- to give the neutral ionic compound NaCl.
Two Na+ combine with one O2- to give the neutral ionic compound Na2O.
„ Notice that the cation is always listed first in the formula.
„ Further examples
{ Note that the formula of an ionic compound shows the smallest whole number ratio between the
ions.
{ So, when Ca2+ and O2- combine, the formula of the product is CaO, not Ca2O2.
{ Sometimes, the ratio between ions is more complex, like with Al3+ and O2-.
Example
What is the formula of the ionic compound which contains ions of each of the two elements below?
sodium and sulfur
barium and nitrogen
selenium and calcium
Structure of Ionic Compounds
„ At room temperature, ionic compounds are always found as solids
„ Ions come together in well-organized patterns at the submicroscopic level
„ This organization is seen in the near-perfect cubic shape of salt (NaCl) crystals
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Chemistry 110
Chapter Four Notes
Naming Binary Ionic Compounds
„ Binary ionic compounds contain ions from exactly two different elements
„ As we have seen, one ion is derived from a metal, the other a nonmetal
„ Before naming the compound, determine whether the metal produces ions of only one type of charge
„ Elements in Group 1A always form ions with a 1+ charge
„ Elements in Group 2A always form ions with a 2+ charge
„ Both aluminum and gallium in Group 3A always form ions with a 3+ charge
„ Zinc only forms a 2+ ion: Zn2+
„ Cadmium only forms a 2+ ion: Cd2+
„ Silver only forms a 1+ ion: Ag1+
„ Any metal not mentioned so far may make ions with different charges
„ Consider these examples:
{ Copper produces Cu+ and Cu2+
{ Iron produces Fe2+ and Fe3+
{ Lead produces Pb2+ and Pb4+
„ We name binary ionic compounds by stating the name of the first element (the cation) and then the
name of the second element (the anion)
„ We change the ending of the name of the anion to “-ide”.
„ If the first element can form more than one charge, we specify the charge in the name with Roman
Numerals in parenthesis right after the name of the element
How do we name MgCl2?
How do we name CuCl2?
Name each of the following compounds
KBr
AlCl3
CoF2
Ag2O
Ba3N2
SnCl4
Cr2O3
ZnS
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Chemistry 110
Chapter Four Notes
Covalent Bonding
„ Covalent bonds result from the sharing of electrons between two atoms.
„ Recall that covalent bonding is usually found between nonmetal atoms.
{
This is a bit oversimplified, but will work well for our purposes in this class.
„ In following the octet rule, atoms share enough electrons with each other to give each eight valence
electrons (or two in the case of hydrogen).
„ Covalent bonds are generally not as strong as ionic bonds
{
Example: Compare heating salt (ionic) and sugar (covalent).
Molecules
„ Compounds which are held together by covalent bonds are called molecules
„ This term is not appropriate for ionic compounds, which are made up of repeating unit cells
„ For example an H2O molecule contains exactly three elements, but there is no “NaCl molecule”
consisting of only two atoms
Examples of Covalent Compounds and Elements
H2O
C6H12O6
C6H6
H2S
The elements O2, N2, and H2
SO3
CO2
Electronegativity
„ The electronegativity of an atom indicates how strong its attraction for electrons is.
{ The greater an atom’s electronegativity, the more strongly it pulls electrons toward it.
„ We can determine how electronegative an atom is from its position on the periodic table
{ Electronegativity increases up a group and from left to right across a period.
„ Fluorine is the most electronegative atom; in general, the closer a main group element is to
fluorine on the periodic table, the more electronegative it is.
„ The noble gases have virtually no electronegativity.
Lewis Structures
„ Lewis Structures are simple drawings that show the covalent bonds between atoms.
„ As in Lewis dot diagrams, electrons are represented by dots.
{ Since two electrons are required to make a bond, a pair of electrons between two atoms indicates a
bond.
Multiple Bonds
„ Remember, atoms bond in order to gain an octet (or duet for hydrogen).
„ In order to do this, atoms can also share more than one pair of electrons, up to a maximum of three
pairs.
{ Sharing one pair of electrons forms a single bond.
{ Sharing two pairs forms a double bond.
{ Sharing three pairs forms a triple bond.
{ There are no quadruple bonds!
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Chemistry 110
Chapter Four Notes
„ It is much simpler when drawing Lewis structures to represent electron pairs as lines rather than dots.
{ However, always keep in mind that the lines stand for electron pairs!
Drawing Lewis Structures
There are several steps involved in taking a formula and converting it to a Lewis structure; just keep in
mind the goals set out by the octet and duet rules.
1.
Add up the total number of valence electrons in the formula.
Ex. NH3 has 5 + (3×1)=8 v.e.
Note that if the molecule is an ion, each negative charge adds one electron to this count, and each
positive charge takes one away.
2. Draw a simple “skeleton” of the molecule, showing which atoms are connected to which. If there is
one central atom (that is, an atom bonded to two or more different atoms) it is usually the least
electronegative.
Note that hydrogen cannot be a central atom!
3.
Figure out how many electrons you have left. Remember that each bond stands for two electrons,
so subtract two from the total number of valence electrons for every bond you drew in the
previous step.
4.
Give enough electron pairs to each atom in the structure to give each atom (except hydrogen) an
octet, using the following guidelines:
-Electron pairs go to the most electronegative atoms first
-Electron pairs go on the central atom last.
-You must stop adding electron pairs when you run out of electrons (from the original number you
started with in step 1).
5.
This step is not always necessary:
If you have run out of electrons, and some atoms still need an octet, change a lone pair on the least
electronegative adjacent atom (except hydrogen) to a double bond between the two atoms. If
necessary, you may need to make a triple bond by taking away another lone pair.
Note that fluorine will not share its lone pairs.
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Chemistry 110
Chapter Four Notes
Examples
Draw Lewis Structures for each of the following molecules:
CH4
H2O
HCN
SO2
PF3
The Shapes of Molecules
„ The shape of a molecule is predicted by Valence Shell Electron Pair Repulsion (VSEPR) theory
{ We pronounce this as “ves-per”
„ According to this theory, electron pairs will orient themselves as far away from each other as possible
{ Since electrons have like charge, we expect them to repel each other
„ There are two types of VSEPR electron pairs which we use to determine the shape of a molecule
„ There is one bonding VSEPR electron pair for each atom a central atom is bonded to
{ It does not matter whether the bond joining them is a single bond, double bond, or triple bond
„ There is one nonbonding VSEPR electron pair for each lone pair of electrons on an atom
VSEPR Pairs
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Chemistry 110
Chapter Four Notes
VSEPR Electron Pair Arrangements
„ Molecules tend to take on three different arrangements about the central atom, depending on the number
of VSEPR electron pairs
„ The linear arrangement is found about central atoms with two VSEPR electron pairs. The electron pairs
are 180° apart in space.
„ The trigonal planar arrangement is found about central atoms with three VSEPR electron pairs. The
electron pairs are 120° apart in space.
„ The tetrahedral arrangement is found about central atoms with four VSEPR electron pairs. The electron
pairs are 109° apart in space.
The Shapes of Molecules
„ We will consider six different molecular geometries, based on the VSEPR pair arrangements
„ We determine the molecular geometry by counting the number of bonded and nonbonded VSEPR
electron pairs on an atom
„ In some cases, the name of the molecular geometry is the same as its VSEPR pair arrangement
„ The molecular geometry around a central atom with 2 bonding VSEPR electron pairs is called linear
„ The molecular geometry around a central atom with 3 bonding VSEPR electron pairs is called trigonal
planar
Page 8
Chemistry 110
Chapter Four Notes
„ The molecular geometry around a central atom with 2 bonding VSEPR electron pairs and 1 nonbonding
pair is called angular
„ The molecular geometry around a central atom with 4 bonding VSEPR electron pairs is called
tetrahedral
„ The molecular geometry around a central atom with 3 bonding VSEPR electron pairs and 1 nonbonding
pair is called trigonal pyramid
„ The molecular geometry around a central atom with 2 bonding VSEPR electron pairs and 2 nonbonding
pairs is called angular
{ This has the same name as the geometry with 2 bonding and 1 nonbonding pairs, but the angles are
different!
Summary of Shapes
Electron Pair
Arrangement
Angles
Molecular
Geometries
linear
180o
linear
2
0
trigonal
planar
3
0
angular
2
1
tetrahedral
4
0
trigonal
pyramid
3
1
angular
2
2
trigonal
planar
tetrahedral
120o
109o
number of
number of
bonding pairs nonbonding pairs
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Chemistry 110
Chapter Four Notes
Polarity
„ If two atoms which are covalently bonded have substantially different electronegativity values, then the
bond is said to be polar covalent.
{ The electrons are not shared equally in the bond.
{ The electrons will tend to stay closer to the more electronegative atom.
„ If the electronegativities are the same or very close, the bond is described as nonpolar covalent
„ The polarity may be shown by drawing a special arrow pointing towards the negative side of the
molecule.
„ The polarity will cancel out if all electron pair groups around the atom are the same.
„ Molecules that are polar are attracted to one another
„ Energy is required to pull apart polar molecules
„ Molecules without a dipole moment do not have as much attraction for each other and require less
energy to separate
Naming Binary Molecular Compounds
„ Binary Molecular Compounds are compounds containing atoms of exactly two different non-metals
„ Nonmetals can form covalent bonds with one another in many possible ratios.
{
For example, nitrogen and oxygen make compounds such as NO, NO2, and N2O.
„ Although these compounds contain the same elements (and may even contain them in the same ratio),
these chemicals have considerably different chemical and physical properties.
„ Therefore, each must be given its own distinct name to distinguish it.
„ The compound is named by putting prefixes before the name of the element to indicate how many atoms
of each type are in the molecule.
„ The ending of the name of the second element is changed to “-ide”.
prefix[1st element] prefix[2nd element]
„ The first six prefixes and the tenth must be memorized:
Number of
Atoms
in Molecule
Prefix
1
mono-
2
di-
3
tri-
4
tetra-
5
penta-
6
hexa-
10
deca-
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Chemistry 110
Chapter Four Notes
Naming Binary Molecular Compounds
„ For example, a compound with formula S2O3 would be named disulfur trioxide.
„ Exceptions:
{ Do not use the prefix “mono-” with the first element in the formula; just stating the name of the
element implies that there is only one of it in the formula.
„ For example, CO2 is carbon dioxide, not monocarbon dioxide.
„ What is the name of SO3?
{
If the last letter of the prefix is an “a” or an “o”, and the first letter of the element after the prefix is
an “o”, drop the last letter of the prefix; it looks and sounds awkward.
„ Example: CO is carbon monoxide, not carbon monooxide.
„ P2O5 should be named diphosphorus pentoxide, which is preferable to diphosphorus
pentaoxide (which is technically acceptable).
„ In a few cases, the naming rules are totally ignored, and the common name is used.
„ These three are
{ H2O, which is always called water.
{ NH3, which is always called ammonia.
„ Other common names do exist, but you are not required to know them for this class
„ When hydrogen appears first in the formula of a binary nonmetal, do not use prefixes at all. Simply
name the compound without them.
{ There will be as many hydrogens bonded to the other nonmetal as are necessary to give it a
complete octet.
{ So, H2S is named hydrogen sulfide, and HCl is named hydrogen chloride.
Polyatomic Ions
„ Polyatomic ions can be described as molecules which have collectively gained or lost electrons to
become an ion.
„ Being molecules, the ions themselves are held together by covalent bonds.
„ From this, you would expect the atoms in a polyatomic ion to all be nonmetals.
+
{ Examples: NO3 , ClO4 , NH4
„ However, some polyatomic ions contain metals which are able to covalently bond.
2{ Examples: Cr2O7 , MnO4
„ All but one of the more common polyatomic ions are anions.
„ The ammonium ion, NH4+, is the only common polyatomic cation.
„ The common polyatomic ions must be memorized, and you must learn to recognize them in a formula
on sight. In addition to ammonium, you must memorize the following ions:
1- anions
OHCNC2H3O2-
hydroxide
cyanide
acetate
CO32-
carbonate
NO3NO2ClO3-
nitrate
nitrite
chlorate
SO42-
sulfate
2- anions
3- anion
PO43-
phosphate
Page 11
Chemistry 110
Chapter Four Notes
Naming Compounds Containing Polyatomic Ions
„ The naming rules for these compounds are the same as those for binary ionic compounds
„ When writing their formulas, we must put parenthesis around polyatomic ions when two or more occur
in the formula
„ Examples:
{ Na2SO4
{
Ca(NO3)2
{
Mg(OH)2
{
FePO4
Page 12