chapter 9 covalent bonding. 9.1 the covalent bond covalent bond – a chemical bond that results...
TRANSCRIPT
Chapter 9
Covalent Bonding
9.1 The Covalent Bond Covalent Bond – A chemical bond that
results from the sharing of valence electrons.– Between nonmetallic elements of similar
electronegativity. A group of atoms united by a covalent bond is
called a molecule. A substance made up of molecules is called a
molecular substance. Molecules may have as few as two atoms.
(CO), (O2)
Polar Covalent Bonds: Unevenly matched, but willing to share.
To describe the composition of a molecular compound we use a molecular formula.
The molecular formula tells you how many atoms are in a single molecule of the compound.– The molecular formula for oxygen is
O2, which tells you that this molecule contains 2 atoms of oxygen.
The molecular formula for sucrose (table sugar) is C12H22O11. This indicates that one sucrose molecule contains 12 carbon atoms, 22 hydrogen atoms and 11 oxygen atoms.
The empirical formula shows the ratio of atoms in the molecule. You get it by dividing all subscripts by the smallest subscript and rounding. C6H12O6 becomes CH2O.
Molecules are often shown using a structural formula.
A structural formula specifies which atoms are bonded to each other in a molecule.
– One type is the Lewis structures.– The Lewis structure uses dots as
before and has a circle drawn around the dots that represent valence electrons that are joined in a covalent bond.
The principle that describes covalent bonding is the Octet rule.
– Atoms within a molecule will share electrons so as to acquire a full set (8) of valence electrons.
If each atom has 7 valence electrons, than they will share one so that they both will have 8. The sharing of that electron is a Single covalent bond.
Single Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals
F F
Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals
F F8 Valence electrons
Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals
F F8 Valence electrons
If more than two atoms join in a covalent bond the atoms with the fewest valence electrons will pair up with the atom that has the most up to a total of 8.
If a pair of electrons is not shared they are called an unshared pair.
If two atoms share two pairs of electrons it is called a double covalent bond.
If two atoms share three pairs of electrons it is called a triple covalent bond.
The Lewis structures that you have seen so far use a pair of dots to represent covalent bonds.
Another kind of Lewis structure uses dashes to represent covalent bonds.
– A single dash represents a single covalent bond.
– A double dash represents a double covalent bond
– A triple dash represents a triple covalent bond.
Covalent compounds (molecules) tend to have low melting and boiling points
When an electron pair is shared by the direct overlap of bonding orbitals, a sigma bond results. Single bonds are sigma bonds.
The overlap of parallel orbitals forms a pi bond. Multiple bonds involve both sigma and pi bonds.
Bond length depends on the sizes of the bonded atoms and the number of electron pairs they share.
Bond dissociation energy is the energy needed to break a covalent bond.– Double bonds have larger bond
dissociation energies than single. Triple even larger
Bond length and bond dissociation energy are directly related.
9.2 Naming Molecules– Covalent compounds (Molecules) are named by adding
prefixes to the element names.• Covalent’ (in this context) means both elements are nonmetals.
– The names of molecular compounds will include prefixes that indicate the number of atoms in the molecule. (CO2 would be called Carbon Dioxide, CCl4 is called Carbon Tetrachloride)
– The following prefixes are used:mono-1 di-2tri-3 tetra-4penta-5 hexa-6hepta-7 octa-8nona-9 deca-10
Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.
– Remember to use the “ide” suffix. (Cl is chloride not chlorine)
– Exceptions to the rule:• The prefix “mono” is usually not written with the
first word of the compounds name. – CO2 is not called monocarbon dioxide, just carbon
dioxide.
• Some compounds have common names. (H2O is water or dihydrogen monoxide)
Naming Binary Covalent Compounds: Examples
N2S4 dinitrogen tetrasulfide
NI3 nitrogen triiodide
XeF6 xenon hexafluoride
CCl4 carbon tetrachloride
P2O5 diphosphorus pentoxide
SO3 sulfur trioxide
1 mono
2 di
3 tri
4 tetra
5 penta
6 hexa
7 heptaa
8 octa
9 nona
10 deca
* Second element in ‘ide’ from
* Drop –a & -o before ‘oxide’
Naming Acids An acid is a molecular substance that
dissolves in water to produce hydrogen ions (H+).– Acids behave like an ionic compound
because when they dissolve in water, they create cations and anions.
– The cation is always H+ and the anion depends on the particular acid.
• Example: When the acid HCL dissolves in water it forms H+ cations and CL- anions.
• HNO3 acid will form H+ and NO3- ions.
Naming Acids The name of an acid usually results from the name of the
anion(-).– For acids, where the anion ends with the suffix
“ide”(Chloride) the name of the acid begins with the prefix hydro.
– The acids name also includes the root name of the anion(-) and the word acid.
– In addition you need to change the suffix of the anion from “ide” to “ic”.
• By these rules HCL is named hydrochloric acid.
Other acids are named without the prefix “hydro”. The acid HNO3
- derives its name from NO3- which is nitrate. HNO3 is
named Nitric Acid. The following are names and formulas of common acids:
Naming Acids Anion Corresponding Acid
F-, Fluoride HF, hydrofluoric Acid CL-, Chloride HCL, hydrochloric acid Br-, Bromine HBr, hydrobromic acid I-, iodide HI, hyroiodic acid S2
-, sulfide H2S, hydrosulfuric acid NO3
-, nitrate HNO3, nitric acid CO3
-2, carbonate H2CO3, carbonic acid SO4
-2, sulfate H2SO4, sulfuric acid PO4
-3, phoshate H3PO4, Phosphoric acid C2H3O2
- , acetate HC2H3O2, acetic acid
9.3 Molecular Structures Structural formulas (Lewis Structures)
use letter symbols and bonds to show the position of atoms.
Drawing Lewis Structures:
Water
H
O
Each hydrogen has 1 valence electron
and wants 1 more
The oxygen has 6 valence electrons
and wants 2 more
They share to make each other “happy”
Water Put the pieces together The first hydrogen is happy The oxygen still wants one more
H O
Water The second hydrogen attaches Every atom has full energy levels
H OH
Multiple Bonds Sometimes atoms share more than one
pair of valence electrons. A double bond is when atoms share two
pair (4) of electrons. A triple bond is when atoms share three
pair (6) of electrons.
Carbon dioxide CO2 - Carbon is central
atom ( I have to tell you)
Carbon has 4 valence electrons
Wants 4 more Oxygen has 6 valence
electrons Wants 2 more
O
C
Carbon dioxide Attaching 1 oxygen leaves the oxygen 1
short and the carbon 3 short
OC
Carbon dioxide Attaching the second oxygen leaves
both oxygen 1 short and the carbon 2 short
OCO
Carbon dioxide The only solution is to share more
OCO
Carbon dioxide The only solution is to share more
OCO
Carbon dioxide The only solution is to share more
OCO
Carbon dioxide The only solution is to share more
OCO
Carbon dioxide The only solution is to share more
OCO
Carbon dioxide The only solution is to share more
OCO
Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in
the bond
OCO
Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in
the bond
OCO8 valence electrons
Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in
the bond
OCO8 valence electrons
Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in
the bond
OCO
8 valence electrons
How to draw them To figure out if you need multiple bonds Add up all the valence electrons. Count up the total number of electrons to
make all atoms happy. Subtract. Divide by 2 Tells you how many bonds - draw them. Fill in the rest of the valence electrons to
fill atoms up.
Examples NH3
N - has 5 valence electrons wants 8
H - has 1 valence electrons wants 2
NH3 has 5+3(1) = 8
NH3 wants 8+3(2) = 14
(14-8)/2= 3 bonds 4 atoms with 3 bonds
N
H
N HHH
Examples Draw in the bonds All 8 electrons are accounted for Everything is full
Examples HCN C is central atom N - has 5 valence electrons wants 8 C - has 4 valence electrons wants 8 H - has 1 valence electrons wants 2
HCN has 5+4+1 = 10
HCN wants 8+8+2 = 18
(18-10)/2= 4 bonds 3 atoms with 4 bonds -will require multiple
bonds - not to H
HCN Put in single bonds Need 2 more bonds Must go between C and N
NH C
HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add
NH C
HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add Must go on N to fill octet
NH C
Another way of indicating bonds
Often use a line to indicate a bond Called a structural formula Each line is 2 valence electrons
H HO =H HO
Structural Examples
H C N
C OH
H
C has 8 electrons because each line is 2 electrons
Ditto for N
Ditto for C here Ditto for O
Coordinate Covalent Bond A Coordinate Covalent Bond is when one
atom donates both electrons in a covalent bond.
Carbon monoxide CO
OC
Coordinate Covalent Bond When one atom donates both electrons
in a covalent bond. Carbon monoxide CO
OC
Coordinate Covalent Bond When one atom donates both electrons
in a covalent bond. Carbon monoxide CO
OCOC
Resonance Resonance is a condition when more
than one dot diagram with the same connections is possible.
Resonance structures differ only in the position of the electron pairs, never the atom positions.
9.4 Molecular Shape (VSEPR) Valence Shell Electron Pair Repulsion. Predicts three dimensional geometry of
molecules. The name tells you the theory. Valence shell - outside electrons. Electron Pair repulsion - electron pairs try
to get as far away as possible. Can determine the angles of bonds. And the shape of molecules
VSEPR Based on the number of pairs of
valence electrons both bonded and unbonded.
Unbonded pair are called lone pair.
CH4 - draw the structural formula
VSEPR Single bonds fill
all atoms. There are 4 pairs
of electrons pushing away.
The furthest they can get away is 109.5º.
C HH
H
H
4 atoms bonded Basic shape is
tetrahedral. A pyramid with a
triangular base. Same basic shape
for everything with 4 pairs. CH H
H
H109.5º
3 bonded - 1 lone pair
N HH
H
NH HH
<109.5º
Still basic tetrahedral but you can’t see the electron pair.
Shape is calledtrigonal pyramidal.
2 bonded - 2 lone pair
OH
H
O HH
<109.5º
Still basic tetrahedral but you can’t see the 2 lone pair.
Shape is calledbent.
3 atoms no lone pair
CH
HO
The farthest you can the electron pair apart is 120º.
Shape is flat and called trigonal planar.
Will require 1 double bond
C
H
H O
120º
2 atoms no lone pair With three atoms the farthest they can
get apart is 180º. Shape called linear. Will require 2 double bonds or one triple
bond
C OO180º
Molecular Orbitals The overlap of atomic orbitals from
separate atoms makes molecular orbitals
Each molecular orbital has room for two electrons
Two types of MO
– Sigma ( σ ) between atoms
– Pi ( π ) above and below atoms
Sigma bonding orbitals From s orbitals on separate atoms
+ +
s orbital s orbital
+ ++ +
Sigma bondingmolecular orbital
Sigma bonding orbitals From p orbitals on separate atoms
p orbital p orbital
Sigma bondingmolecular orbital
Pi bonding orbitals P orbitals on separate atoms
Pi bondingmolecular orbital
Sigma and pi bonds All single bonds are sigma bonds A double bond is one sigma and one pi
bond A triple bond is one sigma and two pi
bonds.
Hybridization The mixing of several atomic orbitals to
form the same number of hybrid orbitals.
When an atom approaches another atom to form a bond, the orbitals of its electrons may be changed.
The changing of the orbitals forms what is called hybrid orbitals. This is a cross between all of the orbitals involved (s,p,d,f).
Hybridization When an “s” and “p” orbital come together they
form a linear shaped molecule.
When an “s” and “p2” orbital come together they form a trigonal planer molecule.
When an “s” and “p3” orbital come together they form tetrahedral, pyramidal or bent molecules.
Atoms with higher orbital numbers form longer bonds. Multiple bonds are shorter than single bonds because bonds with more electrons attract the nuclei of the bonding atoms more strongly.
9.5 Electronegativity and Bond Type
Electro negativity
Difference
Bond Type
<.4 Non-Polar Covalent
0.5 – 1.9 Polar Covalent
>2.0 Ionic
Polar Bonds When the atoms in a bond are the
same, the electrons are shared equally. This is a nonpolar covalent bond. When two different atoms are
connected, the electrons may not be shared equally.
This is a polar covalent bond.
Electronegativity A measure of how strongly the atoms
attract electrons in a bond. The bigger the electronegativity
difference the more polar the bond.
0.0 - 0.4 Covalent nonpolar
0.5 - 1.0 Covalent moderately polar
1.0 -2.0 Covalent polar
>2.0 Ionic
How to show a bond is polar Isn’t a whole charge just a partial charge means a partially positive means a partially negative
The Cl pulls harder on the electrons The electrons spend more time near the Cl
H Cl
Polar Molecules Polar molecules have a partially
positive end and a partially negative end Requires two things to be true:
– The molecule must contain polar bonds.This can be determined from differences in electronegativity.
– Symmetry can not cancel out the effects of the polar bonds.
Must determine geometry first.
Polar Molecules Symmetrical shapes are those without
lone pair on central atom – Tetrahedral– Trigonal planar– Linear
Will be nonpolar if all the atoms are the same
Shapes with lone pair on central atom are not symmetrical
Can be polar even with the same atom
Intermolecular Forces They are what make solid and liquid
molecular compounds possible. The weakest are called van der Waal’s
forces - there are two kinds
– Dispersion forces
– Dipole Interactions
Dispersion Force
Depends only on the number of electrons in the molecule
Bigger molecules more electrons More electrons stronger forces
• F2 is a gas
• Br2 is a liquid
• I2 is a solid
Dipole interactions Occur when polar molecules are
attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely
hooked like in ionic solids.
Dipole interactions Occur when polar molecules are
attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely
hooked like in ionic solids.
H F
H F
Properties of Molecular Compounds
Made of nonmetals Poor or nonconducting as solid, liquid or
aqueous solution Low melting point Two kinds of crystals
– Molecular solids held together by IMF– Network solids- held together by bonds– One big molecule (diamond, graphite)