Chemistry
chemical bond: force that holds two atoms together
-creates stability in the atom
Bonds may form in two ways:
1. Attraction between a positive nucleus and negative
electrons (covalent bonding)
2. Attraction between a positive ion and a negative ion
(ionic bonding)
Remember: It is the valence electrons that are involved in
this bonding.
ionic bond: electrostatic force that holds oppositely
charged particles together
-called ionic compounds
-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
Example:
Sodium reacts with chlorine to form sodium chloride.
Electron Configuration Notation:
Orbital Notation:
Lewis Dot Notation:
Try this # 1:
Magnesium reacts with oxygen to form magnesium oxide.
Electron Configuration Notation:
Orbital Notation:
Lewis Dot Notation:
Try this # 2:
Lithium reacts with nitrogen to form lithium nitride.
Electron Configuration Notation:
Orbital Notation:
Lewis Dot Notation:
It is the chemical bonds between atoms that determines
many of the physical properties of the compound.
-alternating positive and negative ions form an ionic
crystal
-the ratio of positive to negative ions is determined
by the number of electrons transferred
-strong attraction
results in a crystal
lattice, a 3-D
arrangement of
atoms.
Other characteristics include:
-high melting and boiling points
-very hard and rigid
-brittle
-electrolyte when dissolved in water
During chemical reactions, energy is either absorbed
(endergonic) or released (exergonic)
-the formation of ionic bonds is always exothermic
(exergonic)
lattice energy: energy required to separate one mole
of ions of an ionic compound
-the more negative the lattice energy, the stronger the
bond
Lattice Energyies of Some Ionic Compounds
Compound
Name
Lattice Energy
(kJ/mol)
Compound
Name
Lattice Energy
(kJ/mol)
KI
-632
KF
-808
KBr
-671
AgCl
-910
RbF
-774
NaF
-910
NaI
-682
LiF
-1030
NaBr
-732
SrCl2
-2142
NaCl
-769
MgO
-3795
Compound
Name
KI
KBr
RbF
NaI
NaBr
NaCl
Lattice Energyies of Some Ionic Compounds
Lattice Energy
Compound
Lattice Energy
(kJ/mol)
Name
(kJ/mol)
-632
KF
-808
-671
AgCl
-910
-774
NaF
-910
-682
LiF
-1030
-732
SrCl2
-2142
-769
MgO
-3795
Depends on:
1. smaller ions -more negative value because the attraction is
stronger between the nucleus and valence electrons
2. larger the positive/negative charge, the more
negative the lattice energy because the attraction is
stronger when more electrons are lost/gained
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.
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.
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
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.
*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
Metallic bonds are similar to ionic bonds because they
often form lattices in the solid state.
-eight to twelve metal atoms surround another, central
metal atom
Instead of sharing electrons or losing electrons, the outer
orbitals overlap.
-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
metallic bond: attraction of a metallic cation for the
delocalized electrons that surround it
This bonding contributes to the unique properties of
metals:
1. 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
2. malleable (hammered into sheets) and 3. ductile
(drawn into wire)
-mobile electrons can easily be pulled and pushed
past each other
4. durable
-though electrons move freely, they are strongly
attracted to the metal cations and are not easily
removed from the metal
5. good conductors
-free movement of the delocalized electrons,
allowing heat and electricity to move from one
place to another very quickly
6. luster
-interaction between light and delocalized
electrons
As the number of delocalized electrons increases, as in
transition metals (d electrons), the hardness and
strength also increases.
-alkali and alkaline earth metals are soft (s valence
electrons only)
It is easy to combine 2 or more different metals to make a
metallic crystal
-alloy: mixture of elements with metallic properties
-the properties of alloys differ from those of the
individual elements that make it up
TEST
Remember that atoms bond to increase stability, which
occurs when an atom gets a full outer shell of electrons.
-in ionic bonding, one atom loses electrons (metal)
and another gains electrons (nonmetal) to form
oppositely charged ions with a full outer shell
However, sometimes there is not a transfer of electrons,
but a sharing of electrons.
-covalent bond: attractive force between atoms due
to the sharing of valence electrons
Covalent bonds can form between:
-2 or more nonmetal atoms
-metalloids (especially the ones to the right of the
metalloid line) and nonmetals
molecule: when two or more atoms bond covalently
Covalent bonds can have either single bonds or multiple
bonds.
-single bonds: 2 shared electrons (1 pair)
-multiple bonds: 4 or 6 electrons shared (2 pair=
double or 3 pair = triple)
When we show bonding, shared electron pairs can be
shown by either a pair of dots or a single line.
-Lewis Structures are used to show how bonding
electrons are arranged in molecules
-example: NH3
-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
A multiple bond forms when two atoms share more than 2
electrons.
-double bond: 4 electrons shared ( 2 pairs)
♦ O2
-triple bond: 6 electrons shared (3 pairs)
♦ N2
Some molecules have both single and multiple bonds.
♦HCN
pi bond (p): forms when parallel orbitals overlap to share
electrons
-only occurs with multiple bonds because the first
overlap is always a sigma bond
All bonds can be broken, though some more easily than
others.
-due to the strength of the bond
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
When a bond forms or breaks, an energy change occurs.
-bond formation: energy released (exergonic)
-bond breaking: energy absorbed (endergonic)
bond dissociation energy: amount of energy required to
break a specific covalent bond
-always a positive number
-indicates the strength of a covalent bond
larger the bond dissociation energy, stronger the bond
(see p 246 for examples)
1. low melting and boiling points.
2. many vaporize readily at room temperature
3. relatively soft solids (but not all, some are gases/liq.)
4. can form weak crystal lattices
5. do not conduct electricity when dissolved in water
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
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
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-
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
(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
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)
structural formula: uses letter symbols and bonds to
show relative positions of atoms
-one of the most useful
-can be predicted for many molecules by drawing
Lewis structures
-H is always an end (terminal) atom, never a central
atom
-less electronegative atom is the central atom
(nm or metalloid closest to the left of the PT-usually)
CH2O
1. Predict the location of the atoms
C is least electronegative & farthest to left on PT,
therefore it is the central atom
2. Find the total number of electrons available for
bonding.
1 C-4, 2 H-2, 1 O-6 for a total of 12 valence e3. Determine the number of bonding pairs
12 valence e- / 2 = 6 electron pairs
4. Place one bonding pair (single bond) between the
central atom and each terminal atom.
H
C
O
H
5. Subtract the number of pairs you used in step 4 from
the number of bonding pairs determined in step 3.
6 – 3 used = 3 e- pairs left
5. Subtract the number of pairs you used in step 4 from
the number of bonding pairs determined in step 3.
-take the remaining electron pairs and place electron
pairs around the terminal atoms to satisfy the octet
rule
H
C O
H
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
H C O
H
Practice:
1. CH3Cl
2. NBr5
Writing structural formulas for polyatomic ions is the
same with one exception:
-the total number of electrons may differ due to the
negative and positive charge.
♦negative charge, more electrons are present
SO4-2
add two electrons
♦positive charge, less electrons are present
NH4+1 subtract one electron
Let’s look at CO3-2.
-when one or more valid Lewis structure can be written for
a molecule, resonance occurs
-let’s look at another resonance molecule/ion: NO3-1
-each molecule/ion that undergoes resonance behaves as
if it only has one Lewis structure
Some molecules do not obey the octet rule.
Three reasons exist:
1. when a small group of molecules have an odd number of
valence electrons:
-NO2 for a total of 17 valance electrons-one unpaired
electron on N
2. Some form with fewer than eight, though this is
relatively rare:
-B in BH3 is stable with six because it only has 3 valence
electrons.
3. When the central atom has more than 8 electrons,
which is referred to as an expanded octet.
-can occur in elements that are found in period three
or higher elements (because of the d orbitals).
-P in PCl5
(1 s orbital, 3 p orbitals, and 1 d orbital)
TEST/QUIZ
Many of the physical and chemical properties of molecules
is determined by the shape of the molecule.
-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.
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
See page 260 for the Molecular Geometries (Shapes)
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
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”
TEST