Chapters 6 and 16 Covalent Bonding Objectives – Section 16.1 Use electron dot structures to show...
If you can't read please download the document
Chapters 6 and 16 Covalent Bonding Objectives – Section 16.1 Use electron dot structures to show the formation of single, double and triple covalent bonds
Chapters 6 and 16 Covalent Bonding Objectives Section 16.1 Use
electron dot structures to show the formation of single, double and
triple covalent bonds Describe and give examples of coordinate
covalent bonding, resonance structures, and exceptions to the octet
rule.
Slide 2
CA Standards Students know atoms combine to form molecules by
sharing electrons to form covalent or metallic bonds or by
exchanging electrons to form ionic bonds. Students know chemical
bonds between atoms in molecules such as H 2, CH 4, NH 3, H 2 CCH
2, N 2, Cl 2, and many large biological molecules are covalent.
Students know how to draw Lewis dot structures. Students know atoms
combine to form molecules by sharing electrons to form covalent or
metallic bonds or by exchanging electrons to form ionic bonds.
Students know chemical bonds between atoms in molecules such as H
2, CH 4, NH 3, H 2 CCH 2, N 2, Cl 2, and many large biological
molecules are covalent. Students know how to draw Lewis dot
structures.
Slide 3
Single Covalent Bonds Lets look at Hydrogen as the simplest
model of a covalent bond H + H H:H Hydrogen Hydrogen Hydrogen atom
atom molecule H : H Each Hydrogen has one electron and they share
them to form a single covalent bond. The single covalent bond can
be represented by the pair of electrons or as a dash as shown below
H:H or H-H Each dash represents a pair of shared electrons.
Slide 4
Conventions for naming The chemical formulas of ionic compounds
describe formula units (Example: NaCl is a formula unit) The
chemical formulas of covalent compounds describe molecules.
(Example H 2 O is a molecule) Ionic compounds do not have molecular
formulas because they are not composed of molecules. What does that
mean? Example: Ionic copper(II) oxide is composed of equal numbers
of Cu 2+ and O 2- ions in a crystal lattice. The formula unit shows
the lowest whole-number ratio of Cu 2+ to O 2-, which is 1:1, CuO.
In contrast, individual H 2 atoms do exist, and their subscripts
show actual number of atoms, not a lowest-number ratio.
Slide 5
Ionic versus Molecular Compounds
Slide 6
Covalent Molecules Combinations of atoms of the nonmetallic
elements in groups 4A, 5A, 6A and 7A of the periodic table are
likely to form covalent bonds. octet rule for covalent bonding:
Chemist Gilbert Lewis summarized this tendency in his formulation
of the octet rule for covalent bonding: Sharing of electrons occurs
if the atoms involved acquire the electron configuration of noble
gases. The configurations often contain an octet (eight) valence
electrons. [H 2 is of course an exception to this rule.]
Slide 7
Covalent Bonding Diatomic Gas Fluorine
Slide 8
In a water molecule, two hydrogen atoms form one single
covalent bond each with one oxygen atom. Note how the O atom ends
up with eight electrons around it. Covalent molecules will form if
each atom will end up with 8 electrons around it (except H). Each
dot is one electron. Each line is two electrons.
Slide 9
In an ammonia molecule, NH 3, three H atoms form single
covalent bonds with one N atom. Note how the N has eight electrons
around it. When you consider the N and the three Hs, you can see
the 2s and 2p orbitals are now full with eight electrons.
Slide 10
A methane molecule has four carbon-hydrogen bonds. In each
bond, C and H share the 1s e- from the hydrogen and a 2s or 2p e-
from the carbon. Normally C would start with 1s 2 2s 2 2p 2
configuration, but by promoting one 2s e- to 2p, resulting in 1s 2
2s 1 2p 3 it can create a stable octet with the four H atoms.
Slide 11
Carbon bonding 4 covalent bonds This is what Carbon normally
looks like. This is what it looks like when it promotes an e- to 2p
so it can bond with four other atoms. CH 4 is much more stable than
CH 2 so having four covalent bonds is better for Carbon.
Slide 12
Covalent Bonding Different bonding models for methane, CH 4.
Models are NOT reality. Each has its own strengths and
limitations.
Slide 13
Review: The Octet Rule and Covalent Compounds Covalent
compounds tend to form so that each atom, by sharing electrons, has
an octet of electrons in its highest occupied energy level.
Covalent compounds involve atoms of nonmetals only. The term
molecule is used exclusively for covalent bonding
Slide 14
A single bond The Diatomic Fluorine Molecule F F 1s 2s 2p seven
Each has seven valence electrons FF
Slide 15
Some double bonds: Carbon Dioxide A carbon dioxide molecule has
two C=O bonds Note how C and the Os each have 8 electrons now
Slide 16
An exception to the rule: The Diatomic Oxygen Molecule O O 1s
2s 2p six Oxygen has six valence electrons. You would think O 2
would form a double bond by the looks of it, but experiments show
its nonstandard and has two unpaired e- O O
Slide 17
A triple bond: The Diatomic Nitrogen Molecule N N 1s 2s 2p five
Each has five valence electrons N N
Slide 18
I Bring Clay For Our New House
Slide 19
Lewis structures show how valence electrons are arranged among
atoms in a molecule. Lewis structures reflect the central idea that
stability of a compound relates to noble gas electron configuration
(atoms will react if they can arrange themselves to have 8
electrons around them). Shared electron pairs are covalent bonds
and can be represented by two dots (:) or by a single line ( - )
Lewis Dot Structures
Slide 20
The HONC Rule HH Hydrogen (and Halogens) form one covalent bond
O Oxygen (and sulfur) form two covalent bonds One double bond, or
two single bonds N Nitrogen (and phosphorus) form three covalent
bonds One triple bond, or three single bonds, or one double bond
and a single bond C Carbon (and silicon) form four covalent bonds.
Two double bonds, or four single bonds, or a triple and a single,
or a double and two singles
Slide 21
C H H H Cl........ Completing a Lewis Structure: CH 3 Cl Add up
available valence electrons: C = 4, H = (3)(1), Cl = 7 Total = 14
Join peripheral atoms to the central atom with electron pairs.
Complete octets on atoms other than hydrogen with the remaining
electrons Draw carbon as the central atom (it wants the most bonds,
4)...... Check: Final structure should have 14 e-
Slide 22
Coordinate Covalent Bonds A covalent bond in which one atom
contributes both bonding electrons is called a coordinate covalent
bond. This is signified by showing coordinate covalent bonds as
arrows that point from the atom donating the pair of electrons to
the atom receiving the bond. Many polyatomic cations and anions
contain both covalent and coordinate bonds. NH 4 + is an
example.
Slide 23
Carbon Monoxide coordinate covalent bonding In a coordinate
covalent compound, one atom contributes both electrons of a bonding
pair. In carbon monoxide, which atom contributes two electrons in
one of the carbon-oxygen bonds? :C O: Triple bond one of them is a
coordinate bond
Slide 24
Slide 25
exceptions to octet rule
Slide 26
How to represent ions (covalent bonds within the ion)
Slide 27
Electron dot structure of the sulfite ion (SO 3 2- ) each O has
6 each S has 6 plus 2 extra = 26 e- Sulfur has to create one
coordinate covalent bond to make this work.
Slide 28
Slide 29
Bond Dissociation Energies Large amounts of heat are given off
when hydrogen atoms combine to make H 2, which implies that the
product is more stable than the reactants. If you try to break H 2
apart, it will require a large amount of energy to do it. The same
thing is true for Carbon. A typical C-C single covalent bond has a
bond dissociation energy of 347 kJ. Since Carbon forms such strong
C-C bonds, that explains why its compounds are so stable. See table
of bond dissociation energies on the next page. Note which one is
weakest.
Slide 30
Bond Length and Bond Energy What trend do you notice?
Slide 31
Resonance Occurs when more than one valid Lewis structure can
be written for a particular molecule, such as ozone, below. These
are resonance structures. The actual structure is an average or a
blend of the resonance structures.
Slide 32
Resonance in Benzene, C 6 H 6 Each of these junctions
represents where a Carbon is
Slide 33
Exceptions to the Octet Rule NO 2 (also known as smog in LA)
has one unpaired electron, so it is an exception to the octet
rule.
Slide 34
Oxygen: an exception to the octet rule O O O The measured
distance between oxygen atoms indicates that O 2 does have some
double bond character. This suggests that oxygen is a hybrid of the
two structures shown on this page.
Slide 35
Other exceptions to the octet rule F B F F BF 3 is deficient by
2 e-. Sometimes Phosphorus or Sulfur expand the octet to include 10
or 12 electrons. Examples are PCl 5 and SF 6 P S
Slide 36
Diamagnetic vs. Paramagnetic Molecules When e- spins are paired
one up/one down as shown in these box diagrams, their moving
electric charges create a magnetic fields that oppose each other,
cancelling out the magnetic effects. Normal molecules with all e-
in pairs are weakly repelled by a magnetic field. These are called
diamagnetic substances. But when there is a lone unpaired e-, then
the substance shows a strong attraction to an external magnetic
field. These substances are paramagnetic. (But they are not
permanent magnets like we see with Fe).
Slide 37
Section 16.3: Polar bonds and molecules Covalent bonds involve
sharing electrons between two atoms. Sometimes the sharing is equal
and the electron resides halfway in between the atoms, as in a
diatomic gas like N 2, Cl 2, etc. This is called a nonpolar
covalent bond. This sharing isnt always equal, because one atom may
pull harder than the other atom, and then the electron will not be
in the middle. If the bonding electron is shared unequally, this is
called a polar covalent bond (or just a polar bond).
Slide 38
Polar Bonds The greater the electronegativity value, the
greater the ability of an atom to attract electrons to itself. A
high electronegativity atom is not stealing electrons as in the
ionic case, but it is moving them in its direction. Consider HCl.
Hydrogen has an electronegativity of 2.1 and Chlorine has an
electronegativity of 3.0. These values are quite different, so the
covalent bond in HCl is polar. The shared electron is pulled in the
direction of Cl, because it is more electronegative. This can be
represented as follows: H Cl H - Cl
Slide 39
Bond Polarity Electronegativity Differences and Bond Types
Electronegativity Difference Range Most probable type of bond
Example 0.0 0.4Nonpolar covalentH H (0.0) 0.4 1.0Moderately polar
covalent H Cl (0.9) 1.0 2.0Very polar covalent H F (1.9) 2.0IonicNa
+ Cl - (2.1)
Slide 40
Sample Problem 16-4 Find which type of bond (nonpolar covalent,
moderately polar covalent, very polar covalent or ionic) will form
between each of the following? A)N (3.0) and H (2.1) = 0.9,
moderately polar covalent B) F (4.0) and F (4.0) = 0.0, nonpolar
covalent C) Ca (1.0) and O (3.5) = 2.5, ionic D) Al (1.5) and Cl
(3.0) = 1.5, very polar covalent
Slide 41
Slide 42
Polar Molecules The presence of a polar bond in a molecule
makes the entire molecule polar. That means one end of the molecule
is slightly negative and the other end slightly positive. A
molecule that has two poles is called a dipolar molecule or a
dipole. If this kind of molecule is placed in an electric field,
they orient themselves with respect to the positive and negative
plates creating the field.
Slide 43
Polar Molecules Some molecules are polar, but their polarities
line up in such a way that they cancel. Carbon dioxide is one such
example O = C = O
Slide 44
Slide 45
Intermolecular Attractions- van der Waals forces The weakest
intermolecular attraction is the van der Waals forces. These
consist of two possible types, London dispersion forces and dipole
interactions. London dispersion forces, (weakest of all
intermolecular interactions) are caused by the motion of electrons.
The strength of dispersion forces increases as the number of
electrons increases. For halogens, which have more e- in their
outer shell, the major attraction between them is dispersion
forces. These forces are weaker for F and Cl (gases at STP). They
are stronger for Bromine, a liquid at STP, and even stronger for
Iodine, a solid at STP.
Slide 46
Dipole interaction forces The second type of van der Waals
force is the dipole interaction, when polar molecules are attracted
to one another. The positive region of one molecule is attracted to
the negative region of another. HCl molecules
Slide 47
Hydrogen Bonds A hydrogen bond is an attractive force where a
hydrogen which is covalently bonded to a very electronegative atom
(meaning the H has a slight + charge on it) is also weakly bonded
to an unshared electron pair of another atom (pair has charge to
it). This happens because when H bonds to O, F or N, the very polar
bond leaves the H very electron deficient, with essentially an
exposed nucleus with no electrons. The H nucleus is then attracted
to a negatively charged unshared electron pair on another atom. The
resulting hydrogen bond is only about 5% of the strength of a
regular covalent bond, but it is still the strongest of the
intermolecular forces. This is what causes water to be a liquid at
room temperature
Slide 48
Hydrogen bonds in water Hydrogen has valence e- that are not
shielded from the nucleus by another layer of electrons. Water has
this type of interaction because the hydrogens have a slightly +
charge and the oxygen has a slightly charge. This relatively strong
interaction is called a hydrogen bond.
Intermolecular Forces Summary Weakest London dispersion forces
Middle dipole-dipole interactions Strongest hydrogen bonds But all
three are still much weaker than a covalent bond (max 5% of the
strength of an average covalent bond)
Slide 53
Intermolecular Attractions and Molecular Properties The
physical properties of a compound depend on the type of bonding it
has ionic or covalent. Here are some comparisons of physical
properties Characteristics of Ionic and Covalent Compounds
CharacteristicIonic CompoundCovalent Compound Representative
unitFormula unitMolecule Bond formationTransfer of one or more
electrons between atoms Sharing of electron pairs between atoms
Type of elementsMetallic & nonmetallicNonmetallic Physical
stateSolidSolid, liquid and gas Melting pointHigh (usually > 300
C)Low (usually < 300 C) Solubility in waterUsually highHigh to
low Electrical conductivity of aqueous solution Good conductorPoor
conductor or nonconducting
Slide 54
Section 6.5 Molecular Compounds Binary molecular compounds are
composed of two nonmetallic atoms. Because atoms can combine in
different ratios (for example CO and CO 2 ) we use prefixes to help
distinguish between compounds. CO is carbon monoxide CO 2 is carbon
dioxide CCl 4 is carbon tetrachloride Note the ide ending (similar
to how an anion works, but these arent ionic compounds)
PrefixNumber mono-1 di-2 tri-3 tetra-4 penta-5 hexa-6 hepta-7
octa-8 nona-9 deca-10
Slide 55
Section 6.5 Molecular Compound Naming To convert a name to a
formula, write the correct symbols for the two elements, then add
appropriate subscripts. If there is just one of the first atom, you
dont need to write the mono-, it is assumed. But for the second
atom, if there is one, use mono- Ex: tetraiodine nonoxide is I 4 O
9 sulfur trioxide SO 3 phosphorus pentafluoridePF 5 Ex: N 2 O
dinitrogen monoxide PCl 3 phosphorus trichloride SF 6 sulfur
hexafluoride H 2 O dihydrogen monoxide
Slide 56
Molecular Bonding - Acids Here is a list of some of the most
common acids which have covalent bonds and their names, which dont
always follow the standard naming convention. HCl Hydrochloric acid
H 2 SO 4 Sulfuric acid HNO 3 Nitric acid CH 3 COOH Acetic acid
(also written HC 2 H 3 O 2 ) H 3 PO 4 Phosphoric acid H 2 CO 3
Carbonic acid These are the most common ones and good to
memorize