CHEMICAL BONDING
CHAPTERS 8-9
(IONIC, COVALENT)
Chemistry
WHAT IS A CHEMICAL BOND?
chemical bond: force that holds two atoms together
-determines the properties of compounds
-creates stability in the atom
►nature tends to favor lower energy systems
►bonded atoms are lower energy
Bond breaking is endergonic and bond formation is
exergonic!!!
FORMING CHEMICAL BONDS
Bonds may form in three ways:
1. ionic bond: electrostatic force that holds
oppositely charged particles together
-called ionic compunds
2. covalent bond: attractive force between atoms
due to the sharing of valence electrons
-called molecules
3. metallic bond: attraction of a metallic cation for
the delocalized electrons that surround it
IONIC BONDS
-forms between metals and nonmetals
◊metals lose electrons, forms a cation
~cation: positive ion from loss of electrons
◊nonmetals gain electrons, forms an anion
~anion: negative ion formed from gain of
electrons
-most are binary, which means they contain 2
different elements, such as MgO, Al2O3
PROPERTIES OF IONIC COMPOUNDS
-alternating positive and negative ions form an ionic
crystal
-the ratio of positive to negative ions is determined
by the number of electrons transferred
◊due to high difference in electronegativity
-strong attraction
results in a crystal
lattice, a 3-D
arrangement of
atoms.
-high melting and boiling points
-hard, rigid,brittle solids at room temperature
-electrolyte when dissolved in water or in molten state
-formulas are in smallest whole number ratio of
elements
-creates very strong bonds
METALLIC BONDS
-similar to ionic bonds because they often form lattices
in the solid state.
◊ outer orbitals overlap
~no sharing/transfer of electrons
-electron sea model: all metal atoms in a metallic
solid contribute their valence electrons to form a
‘sea’ of electrons around the metal atoms.
-valence electrons are free to move from atom to
atom (delocalized electrons), forming metallic
cations
PROPERTIES OF METALLIC BONDS
-formula written as an atom
-generally have high melting and boiling points, with
especially high boiling points
~due to the amount of energy needed to separate
the electrons from the group of cations
~varies due to # valence electrons
-malleable & ductile
~mobile electrons can easily be pulled and pushed
past each other
-durable
~though electrons move freely, they are strongly attracted to
the metal cations and are not easily removed from the
metal
-good conductors
~free movement of the delocalized electrons, allowing heat
and electricity to move from one place to another very
quickly
-luster
~interaction between light and delocalized electrons
-forms alloys, a mixture of elements with metallic properties
-properties differ from those of the individual elements
COVALENT BONDS & THEIR PROPERTIES
-form between:
-atoms with small difference in electronegativity
~2 or more nonmetal atoms
~metalloids and nonmetals
-formulas give true ratio of atoms (molecular formula)
-low melting and boiling points.
-many vaporize readily at room temperature
MORE PROPERTIES OF COVALENT BONDS
-may exist as liquids, gases or relatively soft solids
-some can form weak crystal lattices (sugar)
-nonelectrolytes when dissolved in water
-weakest of the three types
~low bond strength
STRENGTH OF COVALENT BONDS
What affects bond strength?
bond length: distance that separates the bonded
nuclei
-determined by the size of the atoms and how many
electron pairs are shared
♦larger the atom, the longer the bond length, the
weaker the bond
♦more shared electrons gives a shorter, stronger
bond
TYPES OF COVALENT BONDS
Single Covalent
-2 electrons shared between atoms
-represented by a single line
C C
-sigma bond (s): single covalent bond formed when
an electron pair is shared by the direct overlap of
orbitals
♦can occur between s & s, s & p , or p & p orbitals
MULTIPLE BONDS
-two atoms share more than 2 electrons.
~double bond: 4 electrons shared ( 2 pairs)
O=O
~triple bond: 6 electrons shared (3 pairs)
N
N
-commonly formed by C, N, O, P, S
pi bond (p): parallel orbitals overlap
-only occurs with multiple bonds
SINGLE VS MULTIPLE BONDS
-the more electrons shared, the stronger the bond
~triple bond, shortest, strongest
~single bond, longest, weakest
-due to increase in electron density between the 2
nuclei, which increases the attraction between the
nuclei
N N
O
O
C
C
MOLECULAR STRUCTURES (LEWIS STRUCTURES)
structural formula: uses letter symbols and bonds
to show relative positions of atoms
-can be predicted for many molecules by drawing
Lewis structures (covalent only)
-H is always an end (terminal) atom, never a
central atom
-less electronegative atom is the central atom
-nature favors symmetry
RULES FOR DRAWING STRUCTURAL FORMULAS
Once you have the central atom:
1. Find the total number of valence electrons
-for negative ions, add electrons
-for positive ions, subtract electrons
2. Determine the number of bonding pairs by dividing
the total number by 2
3. Place one bonding pair (single bond) between the
central atom and each terminal atom.
4. Subtract the number of pairs you used in step 3
from the number of bonding pairs determined in
step 2.
5. Take the remaining electron pairs and place them
around the terminal atoms so each satisfies the
octet rule.
-place any remaining pairs on the central atom
6. If the central atom is not surrounded by 4 electron
pairs, it does not have an octet
-convert one or two of the lone pairs on a terminal
atom to a double or triple bond between that
terminal atom and the central atom
(remember which can form multiple bonds)
7. Exceptions:
-reduced octet (H & B can have less than 8)
-expanded octet (period 3-7 central atoms)
SHAPE & HYBRIDIZATION
1. Count areas of electron density around the central
atom
-multiple bonds count as 1 area
2. Count the number of lone pairs on the central a
3. Identify the shape & hybridization
4. Identify the polarity:
-polar molecules have uneven electron forces,
caused by the presence of lone pairs on the
central atom or different terminal atoms.
RESONANCE STRUCTURES (& AN EXAMPLE)
-when one or more valid Lewis structure can be
written for a molecule, resonance occurs
~let’s look at NO3-1
-each molecule/ion that undergoes resonance
behaves as if it only has one Lewis structure
MOLECULAR SHAPE & HYBRIDIZATION
The shape of molecules determines if two or more
molecules can get close enough for a reaction to
occur.
VSEPR (Valence Shell Electron Pair Repulsion)
model: atoms in a molecule are arranged so that
the pairs of electrons (bonded and lone) minimize
repulsion.
-unshared electron pairs have greater repulsive
force than shared electron pairs
VSEPR MODEL
The repulsion between electron pairs result in fixed
angles between atoms
-bond angle: angle formed by any two terminal
atoms and the central atom
♦lone pairs take up slightly more space than bonded
pairs
♦multiple bonds have no affect on the geometry
because they exist in the same region as single
bonds
-example: H2O
ELECTRONEGATIVITY AND POLARITY
Remember that atoms have different attractions for
electrons (electronegativity).
-electronegativity increases left to right and
decreases
down a period
The character and type of bond can be predicted using
the difference in electronegativities between bonded
atoms.
-pure covalent bond: electronegativity difference
=0
(usually occurs between identical atoms, H2)
Most atoms do not have equal sharing of electrons,
producing a purely covalent bond.
-polar covalent bond: unequal sharing of electrons
♦the larger the electronegativity difference, the more
ionic the bond character
-ionic bonds form when the electronegativity
difference is > 1.7 and nonpolar covalent bonds form
when the difference is < 0.5
-the cutoff between polar covalent, nonpolar, and
ionic is sometimes inconsistent with experimental
data
PROPERTIES OF MOLECULES
These properties are due as a result of differences in
attractive forces
-attraction between atoms within a molecules is
strong
-attraction between different molecules is weak
~called intermolecular forces or van der Walls
forces
Types of Intermolecular Forces (van der Walls forces)
1.
dispersion force (induced dipole)
2.
dipole-dipole force
3.
hydrogen bonding
PROPERTIES OF MOLECULES
dispersion force (induced dipole)
-occurs between nonpolar molecules
-very weak
dipole-dipole force
-occurs between polar molecules
-the more polar the molecule, the stronger the
force
hydrogen bonding
-strong intermolecular force between the hydrogen
end of one dipole and a fluorine, oxygen or
nitrogen atom on another molecule’s dipole
IONIC BONDING REVIEW 1
1.
2.
3.
Define chemical bond.
What is an ionic bond? How does it form?
What are two ways bonding can occur?
Describe each.
4. Draw the orbital notation and Lewis dot
notation showing the bonding between sodium
and sulfur. (you may use noble gas notation).
IONIC BONDING REVIEW 2 (FINISH FOR HW)
1. How do positive ions form? How do negative ions
form? What are each called?
2. Why do atoms bond? .
3. What determines the properties of an element?
4. What is a crystal lattice?
5. List 5 characteristics of ionic compounds.
6. What is the difference between endothermic and
exothermic? Which occurs in ionic reactions?
7. What is lattice energy?
8. What does lattice energy depend on?
9. Which substance has a stronger bond: NaCl or
MgO? Why?
METALLIC BONDING REVIEW
1.
2.
3.
4.
5.
6.
7.
8.
9.
What is a metallic bond?
What is an alloy?
Describe the electron sea model.
What occurs with orbitals in metals?
How is metallic bonding similar to ionic bonding?
What are delocalized electrons?
What contributes to a metal’s high boiling point,
malleability, ductility and conductivity?
List the other 2 properties of metals.
What happens to strength and hardness as you
decrease the number of delocalized electrons?
COVALENT BONDING REVIEW 1
1.
2.
3.
4.
5.
Describe a covalent bond.
What types of atoms do covalent bonds form between?
Describe single and double bonds.
What do we mean by sigma bonds?
What do we call covalent compounds?
SINGLE BOND PRACTICE
1. PH3
2. H2S
3. HCl
4. SCl2
5. SiH4
MULTIPLE BONDS PRACTICE
1. CO2
2. CH2O
3. C2H2
NAMES AND FORMULAS-IONIC COMPOUNDS
A universal set of rules must be used so chemists
around the world can communicate.
formula unit: simplest ratio of ions represented in
an ionic compound
-remember that ionic compounds form a crystal
lattice, consisting of many cations and anions.
-the overall charge for the compound is 0
Most ionic compounds are binary, consisting of two
monatomic ions.
-monatomic ion: one atom ion, either positively or
negatively charged
Remember that we determine the charge of each
ion by its oxidation number.
Formula Rules for Ionic Compounds
1. write the cation first, followed by the anion
2. state the charges of both ions
3. cross the number for the charge of one ion to
become
the subscript for the other ion.
-subscripts are used to state the number of each
atom
in the compound
Example: Determine the formula for the ionic
compound formed when potassium reacts with
oxygen.
1. Cation = potassium = K
Anion = oxygen = O
2. K+1 O-2
3. K+1 O-2
K2O1
K2O
You try: Determine the formula for the ionic
compound formed when aluminum reacts with
chlorine.
IONIC BONDING PRACTICE 2
Write the correct formula for the following pairs of
atoms:
1. potassium and iodine
2. magnesium and chloride
3. aluminum and bromide
4. cesium and nitride
5. barium and sulfide
IONIC BONDING REVIEW 3
1. Why do we need a universal set of rules for naming and
writing formulas?
2. Define monatomic and binary.
3. What is meant by a formula unit?
4. Briefly describe the steps to writing ionic formulas.
5. Explain how we determine the charge of the cation and
anion.
6. What is the purpose of subscripts.
7. Determine the formula for the ionic compound formed when
lithium reacts with nitrogen.
IONIC COMPOUNDS WITH POLYATOMIC IONS
We write formulas for ionic compounds containing
polyatomic ions the same way as in binary
compounds.
-the cation comes first, followed by the anion
-state the charges
-cross over the number for the charges
However:
-if you have more than one polyatomic ion, place
parenthesis around the polyatomic ion, with the
subscript outside the parenthesis.
Example: Determine the formula for the ionic
compound formed when beryllium reacts with
cyanide.
1. Cation = beryllium = Be
Anion = cyanide = CN2. Be+2 CN-1
3. Be+2 CN-1
Be1(CN)2
Be(CN)2
You try: Determine the formula for the ionic
compound formed when ammonium reacts with
iodine.
IONIC BONDING PRACTICE 3
Write the correct formula for the following pairs of
atoms:
1. ammonium and oxygen
2. lithium and nitrate
3. aluminum and hydroxide
4. ammonium and phosphate
5. strontium and acetate
IONIC BONDING PRACTICE 4
Write the correct formula for the following pairs of
atoms:
1. aluminum and carbon
2. ammonium and carbonate
3. calcium and oxygen
4. aluminum and chromate
5. sodium and phosphate
6. potassium and hydrogen sulfate
7. magnesium and phosphorus
IONIC BONDING REVIEW 4
1. Describe what a polyatomic ion is?
2. When do we use parenthesis for writing ionic
compounds with polyatomic ions?
3. Determine the formula for the ionic compound
formed when lead reacts with sulfur.
4. Determine the formula for the ionic compound
formed when magnesium reacts with phosphate.
NAMING IONIC COMPOUNDS
The names of ionic compounds include the ions of
which they are composed.
1. The element whose symbol appears first in the
formula also appears first in the name.
-this is always the positively charged ion, or metal
2. The name of the second ion follows, with its ending
changed to –ide for single atom ions.
Ex: What is the name of MgCl2?
magnesium chloride
IONIC COMPOUNDS PRACTICE 5
Write the formula and the name.
1. Na2S
2. Ga2S3
3. CaSe
4. LiF
NAMING WITH POLYATOMIC IONS
You follow the same rules when naming polyatomic
ions as when you have binary ionic compounds,
however:
-you do not change the ending of the polyatomic
ions, even when they are the second atom.
Example:
Al2(SO4)3
aluminum (III) sulfate
Rule: You must state the charge of all metals not
included in groups 1 and 2 because many have
multiple charges.
RULES FOR TRANSITION METALS
*According to the previous rules, FeO and Fe2O3
would both be named iron oxide,even though they
are not the same compound*
Since many transition metals can have more than one
charge, the name must show this. This is done
using roman numerals.
-FeO
is named iron (II) oxide because Fe has a
+2
charge
-Fe2O3 is named iron (III) oxide because Fe has a
+3 charge
*The roman numeral states the charge of the metal*
Q: How do I know the iron in FeO has a +2 charge?
A: The oxide ion has a –2 charge, so the Fe must have
a +2 charge so the compound is overall neutral.
Q: How do I know the iron in Fe2O3 has a +3 charge?
A: There are three oxide ions with a –2 charge:
(3 ions)(-2 charge/ion) = a total of –6 charge
Since the overall charge must be neutral, the iron
must have a total charge of +6. Therefore:
(2 ions)(x charge/ion) = +6
x = +3
IONIC COMPOUNDS PRACTICE 6
Write the formula given & the name of each compound.
1. FeCl3
2. Zn3P2
3. CuS
4. AuF
5. CuC2H3O2
6. AgHCO3
7. ZnSO4
8. Pb(CO3)2
IONIC COMPOUNDS PRACTICE 7
Name the following compounds:
1. NaBr
2. CaCl2
3. KOH
4. Cu(NO3)2
5. Ag2CrO4
6. PbNO2
7. AlCl3
IONIC BONDING REVIEW 5
1. Describe what a polyatomic ion is?
2. What is the relationship between lattice energy and
the strength of ionic bonds?
3. What is the ending of the second atom changed to
when naming ionic compounds?
4. Write the name for (NH4)3P
5. Write the name for AlS.
6. Determine the formula for the ionic compound
formed when magnesium reacts with phosphate.
TEST
COVALENT BONDING REVIEW 2
1. Describe single, double, and triple bonds.
2. How is a pi bond different from a sigma bond?
3. Can a molecule with single bonds have a pi bond?
Why or why not?
4. What affects bond strength?
5. What two things determines bond length?
Describe them.
6. What is bond dissociation energy and what does it
indicate?
7. What occurs when a bond forms or breaks?
8. List the properties of molecules.
NAMING MOLECULES (9.2)
Molecules are represented by both names and
formulas.
Rules for Naming Binary Molecular Compounds
1. The first element in the formula is named first,
using
the entire element name.
2. The second element in the formula is named using
the root of the element and adding the suffix –ide.
3. Prefixes are used to indicate the number of atoms of
each type that are present in the compound.
-exception: 1st element never uses the prefix mono-drop the final letter of the prefix if element name
begins with a vowel.
Prefixes you need to know:
# atoms
prefix
1
mono2
di3
tri4
tetra5
penta6
hexa7
hepta8
octa9
nona10
deca-
NAMING BINARY MOLECULES-EXAMPLE
Name the compound P2O5, which is used as a drying
and dehydrating agent.
1st atom: P = phosphorus
2nd atom: O = oxygen = oxide
There are 2 phosphorus = diphosphorus
There are 5 oxygens = pentoxide (drop the –a of penta)
Put it together: diphosphorus pentoxide
NAMING BINARY MOLECULES PRACTICE
Name the following molecules:
1. CCl4
2. As2O3
3. CO
4. SO2
5. NF3
NAMING ACIDS
(We will talk more about acids in Ch 19)
There are two types of acids:
1. binary acid: contains hydrogen and one other
element
-when naming use the prefix hydro- plus the root of
the second element with the suffix –ic, followed by
the word acid.
-ex: HCl
H = hydroCl = chloride = chloric
hydrochloric acid
Some acids are not binary, but are named according
to the binary acid rules when oxygen is not present,
as in HCN.
H = hydro
CN = cyanide = cyanic
hydrocyanic acid
2. oxyacid: an acid that contains an oxyanion
(oxygen
containing polyatomic ion)
-the name depends on the oxyanion present
-the name consists of the root of the anion, a suffix,
and the word acid
♦if the anion suffix is –ate, it is replaced with -ic
♦if the anion suffix is –ite, it is replaced with -ous
-examples:
~HNO3
NO3 = nitrate
= nitric
nitric acid
~HNO2
NO2 = nitrite
= nitrous
nitrous acid
NAMING ACIDS PRACTICE
Name the following acids:
1. HBr
2. H3PO4
3. H2SO4
4. H2SO3
5. H2CO3
WRITING FORMULAS
Use the prefixes in the molecule’s name to determine
the subscript for each atom in the compound.
- phosphorus tribromide
P
Br
1 (no prefix)
3 (tri)
PBr3
- the formula for an acid can be derived from the
name as well
♦charge of the oxyanion or anion gives the number
of hydrogens
hydrofluoric acid = HF
(1 H because fluorine has a -1 charge)
WRITING FORMULAS PRACTICE
1. oxygen difluoride
2. dinitrogen tetrasulfide
3. phosphorus pentachloride
4. iodic acid
5. phosphoric acid
MOLECULAR STRUCTURE OBJECTIVES
1. Draw structural formulas.
2. Explain why resonance occurs and identify
resonance structures.
3. Discuss exceptions to the octet rule.
STRUCTURAL FORMULAS-PRACTICE 1
1.
2.
3.
4.
5.
6.
7.
8.
SO3
N 2O
SF6
ClF3
SiF4
PO4-3
BF3
SO3-2
MOLECULAR STRUCTURE REVIEW 1
1. What is a structural formula?
2. Describe resonance.
3. List three reasons for exceptions to the octet rule.
4. Name the following:
a. BH3
b. SO2
c. PO4-3
5. Write formulas for the following:
a. sulfur trioxide
c. chlorous acid
b. hydrosulfuric acid
6. Draw structural formulas
a. SO2
b. H2O
c. BrCl5
TEST/QUIZ
MOLECULAR SHAPE & VSEPR OBJECTIVES
1. Discuss the VSEPR bonding theory.
2. Predict the shape of and the bond angles in a
molecule.
3. Define hybridization.
4. Describe how electronegativity is used to
determine
bond type.
5. Compare and contrast polar and nonpolar bonds.
SHAPES AND POLARITY REVIEW 1
1. What determines many of the physical and chemical
properties of molecules?
2. Describe the VSEPR model.
3. What does the repulsion between electron pairs
result in?
4. Why do multiple bonds have no affect on geometry
of a molecule?
5. Why do molecules with lone pairs have shorter bond
angles?
6. What is electronegativity and what does it predict?
7. What is the difference between a nonpolar covalent
bond and a polar covalent bond?
ELECTRONEGATIVITY PRACTICE
Remember: bonding is not clearly ionic or covalent, with ionic
character increasing as the difference in electronegativity
increases.
Decide if the following pairs of atoms are polar covalent,
nonpolar covalent or ionic.
1.
N-H
3.04-2.20 = 0.84
polar covalent
2.
C-Cl
2.55-3.16 = 0.61
polar covalent
3.
S-Se
2.58-2.55 = 0.03
nonpolar covalent
When a polar bond forms the shared electrons are
pulled more strongly toward one atom.
-this creates partial charges at opposite ends of the
molecule, which is called a dipole
♦ d- indicates a partial negative
d+ indicates a partial positive
Polar molecule or not?
A molecule can have individual polar bonds, but make
a nonpolar molecule. How?
We look at the shape of the molecule.
Let’s look at H2O and CCl4.
O—H
C—Cl
dd+
d+
d1.24
0.61
both O-H and C-Cl have polar covalent bonds
One molecule is polar and the other is nonpolar? How
do we know?
We look at the shape of the molecule and the terminal
atoms.
-symmetric molecules like CCl4 are nonpolar because
the polar bonds cancel each other out.
CCl4
-asymmetric molecules like H2O are polar because the
polar bonds do not cancel each other out.
H2O
If water is polar, why will oil not dissolve in it?
Oil must be nonpolar because
A substance is only soluble (dissolvable) when
combined with a like molecule.
“Like Dissolves Like”
hydrophobic- “fear of water”
hydrophilic- “likes water”
SHAPE AND POLARITY REVIEW 2
1. What is a dipole and what indicates them?
2. How do we know if a molecule is polar or nonpolar?
3. Describe the electronegativity trend both across a
period and down a group.
4. Are the following bonds polar or nonpolar covalent?
a. H-Br
b. C-O
c. S-C
5. Describe the relationship between polarity and
solubility.
6. What do we mean by symmetric and asymmetric?
7. True or False? Explain your answer if false.
a. “Orbital hybridization theory can describe both
the
shape and bonding of molecules.”
b. “Covalent bonds differ in the way electrons are
shared by the bonded atoms, depending on the
kind and number of atoms.”
8. Draw Lewis structures for the following and
determine if they are polar or nonpolar? Why?
a. CO2
b. NH3
c. HCl
TEST