bond energy- energy required to break a chemical bond -we...
TRANSCRIPT
bond energy- energy required to
break a chemical bond
-We can measure bond energy to
determine strength of interaction
ionic compound- a metal reacts
with a nonmetal
• Ionic bonds form when an atom that loses
electrons easily reacts with an atom that has
a high affinity for electrons. The charged
ions are held together by their mutual
attraction (Coulombic attraction).
• Ionic bonds form because the ion pair has
lower energy than the separated ions. All
bonds form in order to reach a lower energy
level.
Bond length- the distance where the
energy is at a minimum. We have a
balance among proton-proton repulsion,
electron-electron repulsion, and proton-
electron attraction.
In H2, the two e− will usually be found
between the two H atoms because they are
spontaneously attracted to both protons.
Therefore, electrons are shared by both
nuclei. This is called covalent bonding.
Polar covalent bonds occur when
electrons are not shared equally. One
end of the molecule may have a partial
charge. This is called a dipole.
+
H F H H
+ -
O
-
+ −
H—F polar H—H nonpolar
has dipole moment
O
+ +
H H O S
O bent, polar O
− has dipole moment planar
no dipole moment
CH4 tetrahedral NH3 trigonal pyramidal
no dipole moment has dipole moment
Electron Configurations:
Stable compounds usually have atoms with noble
gas electron configurations.
Two nonmetals react to form a covalent bond by
sharing electrons to gain valence electron
configurations.
When a nonmetal and a group A metal react to
form a binary ionic compound, the ions form
so that the valence electron configuration of
the nonmetal is completed and the valence
orbitals of the metal are emptied to give both
noble gas configurations.
Ions form to get noble gas configurations.
-exceptions in Group A metals:
Sn2+ & Sn4+
Pb2+ &Pb4+
Bi3+ & Bi5+
Tl+ & Tl 3+
Metals with d electrons will lose their highest
numerical energy level electrons before losing
their inner d electrons.
Size of Ions Positive ions (cations) are smaller than their
parent atoms since they are losing electrons. (More protons than electrons=greater nuclear pull)
Negative ions (anions) are larger than their
parent atoms since they are gaining electrons.
(Fewer protons than electrons= lower nuclear pull)
Think: Monster Ants & miniature cats
Isoelectronic ions
–ions containing the same number of
electrons
O2−, F−, Na+, Mg2+, Al3+ all have the
Ne configuration. They are
isoelectronic.
*** For an isoelectronic series, size
decreases as Z increases.
Lattice energy- the change in energy that
takes place when separated gaseous ions
are packed together to form an ionic solid.
Na+(g) + Cl−(g) NaCl(s)
If exothermic, the sign will be negative and the
ionic solid will be the stable form.
We can use a variety of steps to determine the
heat of formation of an ionic solid from its
elements. This is called the Born-Haber cycle.
See examples on pages 366 & 368.
Lattice energy can be calculated using the following:
where k is a proportionality constant that depends on the
structure of the solid and the electron configuration of
the ions. Q1 & Q2 are the charges on the ions. r is
the distance between the center of the cation and the
anion.
Since the ions will have opposite charges, lattice energy
will be negative (exothermic).
The attractive force between a pair of oppositely
charged ions increases with increased charge on the
ions or with decreased ionic sizes.
r
QQk
21energy Lattice
The Relationship Between the Ionic Character of a Covalent Bond
and the Electronegativity Difference of the Bonded Atoms
Polyatomic ions are held together by covalent
bonds. We call Na2SO4 ionic even though it
has 4 covalent bonds and 2 ionic bonds.
Ionic compound- any solid that conducts an
electrical current when melted or dissolved in
water
Salt- an ionic compound
A chemical bond is a model “invented” by
scientists to explain stability of compounds. A
bond really represents an amount of energy.
The bonding model helps us understand and
describe molecular structure. It is supported
by much research data. However, some data
suggests that electrons are delocalized. That
is, they are not associated with a particular
atom in a molecule.
• Single bond- one pair of shared electrons
• Double bond- two pair of shared electrons
• Triple bond- three pair of shared electrons
These values may be slightly different from those in your text. Use the
textbook values for your homework.
Looking at the chart on the
previous slide, what is the
relationship between bond
length and bond energy?
Is there a relationship between
number of bonds and bond
energy?
Bond energies and bond lengths are given in
tables on page 374.
We can use bond energies to calculate heats of
reaction.
H = D(bonds broken)- D(bonds formed)
2H2 + O2 2H2O
Ex. H = [2(432) + 495] –[4(467)] = −509 kJ
2 H−H O=O 4H−O
exothermic
Bonding Models:
Molecular Orbital Model-
Electrons occupy orbitals in a molecule in
much the same way as they occupy
orbitals in atoms.
Electrons do not belong to any one atom.
-very complex model
Localized electron model-
• molecules are composed if atoms that are
bound together by sharing pairs of electrons
using the atomic orbitals of the bound atoms
• traditional model
lone pair- pair of electrons
localized on an atom
(nonbonding)
shared pair or bonding pair-
electrons found in the space
between atoms
The most important requirement for the
formation of a stable compound is that the atoms
achieve noble gas configurations
ionic [ Na ]+ [Cl]−
only valence electrons are included
molecular H2O
H – O - H
duet rule- hydrogen forms stable molecules
when it shares two electrons
H:H
-filled valence shell
Why does He not form bonds?
Its valence orbitals are already filled.
octet rule – most elements need 8 electrons to
complete their valence shell
Cl-Cl
Rules for writing Lewis structures
1. Add up the number of valence electrons from
all atoms.
2. Use 2 electrons to form a bond between each
pair of bound atoms. A dash represents a pair of
shared electrons.
3. Arrange the remaining electrons to satisfy the
duet rule for H and the octet rule for most others.
Exceptions:
Boron and beryllium tend to form compounds
where the B or Be atom have fewer than 8
electrons around them.
BF3 = 24 valence electrons
F
B F
F Common AP equation:
NH3 + BF3 H3NBF3
Some elements in Period 3 and beyond exceed
the octet rule.
Ex. SF6 S has 12 electrons around it
48 valence electrons
F
F F
F S F
F
d orbitals are used to accommodate the
extra electrons.
Elements in the 1st or 2nd period of the
table can’t exceed the octet rule
because there is no d sublevel.
If the octet rule can be exceeded, the
extra electrons are placed on the
central atom.
Resonance-
-occurs when more than one valid Lewis
structure can be written for a particular molecule
actual structure is an average of all resonance
structures
-this concept is needed to fit the localized
electron model (electrons are really delocalized)
Formal Charge
-used to determine the most accurate
Lewis structure
-is the difference between the # of
valence electrons on the free atom and
the # of valence electrons assigned to
the atom in the molecule
-atoms try to achieve formal
charges as close to zero as possible
-any negative formal charges are
expected to reside on the most
electronegative atoms
-Sum of the formal charges must
equal the overall charge on the
molecule (zero) or ion.
Ex. SO42−
O 2− O 2−
O S O O S O
O O
Formal charge only needs to be considered on the AP test if it is specifically asked for.
VSEPR-Valence Shell Electron Pair
Repulsion
-allows us to use electron dot structures to
determine molecular shapes
-the structure around a given atom is
determined primarily by minimizing
electron repulsions
-bonding and nonbonding pairs of electrons
around an atom position themselves as far
apart as possible
Steps:
1. Draw Lewis structure
2. Count effective electron pairs on central atom
(double and triple bonds count as one)
3. Arrange the electron pairs as far apart as
possible
Shapes
AX2 (A represents central atom, X represents
attached atom, E represents unshared electron
pair)
X – A – X linear 180o bond angle
O=C=O Cl – Be – Cl
AX3 Shape is trigonal planar
X X
A 120o bond angle
F F
X BF3 B
F
Any resonance SO3
structure can
be used to O− S = O
determine shape. O
AX2E2 Shape is bent Bond angle is < 109.5o
Unshared electron pairs repel more than shared pair.
Lone pairs require more space than share pairs.
E Ex. H2O
X A X
E
H – O − H
AX5 Shape is trigonal bipyramidal
Bond angles are 120o(equatorial) and 90o(axial)
X
X A X
X
X
Ex. PCl5
Cl
Cl P Cl
Cl Cl
AX4E2 Shape is square planar.
Bond angle is 90o.
E
X X
A
X X
E
Youtube VSEPR annimation
VSEPR OKState