91 Bonding Models and AIDS Drugs X RAY BEAM CRYSTALLIZED DNA MOLECULE X RAY BEAM Film CRYSTALLIZED DNA MOLECULE X RAY BEAM DIFFRACTED RAYS Film CRYSTALLIZED DNA MOLECULE ID: 812516
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Slide1
LEWIS STRUCTURES
BONDS
IONIC BONDING
Slide29.1 Bonding Models and AIDS Drugs
Slide3Slide4X RAY BEAM
▽
Slide5△
CRYSTALLIZED DNA MOLECULE
X RAY BEAM
▽
Slide6Film
▽
△
CRYSTALLIZED DNA MOLECULE
X RAY BEAM
▽
DIFFRACTED RAYS
▽
Slide7Film
▽
△
CRYSTALLIZED DNA MOLECULE
X RAY BEAM
▽
FILM
▽
DIFFRACTED RAYS
▽
Slide8ACTIVE SITE ＞
PHARMACISTS PLANNED TO STOP HIV PROTEASE FROM SPREADING
Slide9ACTIVE SITE ＞
＜MOLECULE THAT WOULD PLUG ACTIVE SITE
PHARMACISTS PLANNED TO STOP HIV PROTEASE FROM SPREADING
Slide10ACTIVE SITE ＞
＜MOLECULE THAT WOULD PLUG ACTIVE SITE
PHARMACISTS PLANNED TO STOP HIV PROTEASE FROM SPREADING
No HIV Protease,
HIV Doesn’t
Spread to AIDS
BOOM! PROBLEM SOLVED!
Slide11Bonding theories are essential to chemistry because…
Explain how atoms bond together to form molecules
They explain why some combinations of atoms are stable and why others are not
Help predict shapes of molecules which determine many physical and chemical properties of compounds
Slide129.2 Types of Bonds
Slide13Why Do Chemical Bonds Form?
Slide14Chemical bonds form because they lower the potential energy between charged particles that compose atoms.
Slide15Chemical bonds form because they lower the potential energy between charged particles that compose atoms.
?
Slide16AS YOU ALREADY KNOW:
Atoms are composed of protons (+) and electrons (-)
(ignoring neutrons for now)
When two atoms get close to each other, electrons of one atom are attracted to the nucleus of another (Coulomb's Law)
However, at the same time, the electrons of each atom repel the electrons of the other and the protons of each atom repel the protons of the other
The result is a complex set of interactions among a potentially large number of charged particles
If these interactions lead to an overall net reduction of energy between the charged particles, a chemical bond forms
(energy is released)
Slide17AS YOU ALREADY KNOW:
Atoms are composed of protons (+) and electrons (-)
(ignoring neutrons for now)
When two atoms get close to each other, electrons of one atom are attracted to the nucleus of another (Coulomb's Law)
However, at the same time, the electrons of each atom repel the electrons of the other and the protons of each atom repel the protons of the other
The result is a complex set of interactions among a potentially large number of charged particles
Bonding theories help us predict the circumstances under which bonds form and also the properties of the resultant molecules
Slide18Classification of Bonds
Types of atoms
Types of Bond
Characteristics of Bond
Metal + Metal
Metallic
Electrons pooled
Metal + Nonmetal
Ionic
Electrons transferred
Nonmetal + Nonmetal
Covalent
Electrons shared
Slide19△
H
2
O Molecule
△
Na
(s)
NaCl
▽
Slide20△
H
2
O Molecule
△
Na
(s)
NaCl
▽
COVALENT
Slide21△
H
2
O Molecule
△
Na
(s)
NaCl
▽
COVALENT
IONIC
Slide22△
H
2
O Molecule
△
Na
(s)
NaCl
▽
COVALENT
IONIC
METALLIC
Slide23Electronegativity = measure of tendency of an atom to attract a bonding pair of electrons
Ionization Energy = how much energy it takes to remove the electron from its atom
Metals have low electronegativity and ionization energies
Nonmetals have high electronegativity and high ionization energies
(ionization energy is the reason in a covalent bond, no atom gives up any electrons, instead they share)
IONIZATION ENERGY AND ELECTRONEGATIVITY DETERMINE WHY METALS BECOME CATIONS (TRANSFER ELECTRONS) AND NONMETALS BECOME ANIONS (ACCEPT ELECTRONS)
Thus, one element accepting an electron and one element transferring an electron creates an ionic bond
Slide24QUICK REVIEW
(OF VALENCE ELECTRON)
Slide25What are valence electrons?
What are valence electrons?
Slide26What are valence electrons?
The electrons in the outermost principal energy level
What are valence electrons?
Slide27What are valence electrons?
The electrons in the outermost principal energy level
What does this mean?
What are valence electrons?
Why do we care?
Slide28What are valence electrons?
The electrons in the outermost principal energy level
What does this mean?
They are the least attracted to the nucleus
What are valence electrons?
Why do we care?
Slide29What are valence electrons?
The electrons in the outermost principal energy level
What does this mean?
They are the least attracted to the nucleus
So they are important for bonding because...
What are valence electrons?
Why do we care?
So they are important for bonding because...
Slide30What are valence electrons?
The electrons in the outermost principal energy level
What does this mean?
They are the least attracted to the nucleus
So they are important for bonding because...
The valence are the ones that will be shared or transferred during a bond
What are valence electrons?
Why do we care?
So they are important for bonding because...
Slide31Groups 1-3 usually...
Groups 5-8 usually...
Group 4...
Give up their electrons
(cation)
Accept other elements electrons
(anion)
Act as anions and cations
Slide32Back to Lesson
Slide33Steps to making Lewis Dot Diagrams
Identify the number of valence electrons of given element
Write symbol of element
Surround element symbol with number of valence electrons
Slide34ELEMENT
# OF VALENCE E-
E- CONFIG.
SUBLEVEL NOTATION
ORBITAL DIAGRAM
Oxygen
6
2-6
1
s2
2
s2
2
p4
1s 2s 2p
What would the Lewis Structure look like?
Slide35ELEMENT
# OF VALENCE E-
E- CONFIG.
SUBLEVEL NOTATION
ORBITAL DIAGRAM
Oxygen
6
What would the Lewis Structure look like?
What would the Lewis Structure look like?
Slide36ELEMENT
# OF VALENCE E-
E- CONFIG.
SUBLEVEL NOTATION
ORBITAL DIAGRAM
Oxygen
6
2-6
1
s2
2
s2
2
p4
1s 2s 2p
↑↓
↑↓
↑↓
↑↓
↑↓
What would the Lewis Structure look like?
What would the Lewis Structure look like?
(if you don’t know number of valence e-, you can use sublevel notation to figure it out)
Slide37ELEMENT
# OF VALENCE E-
E- CONFIG.
SUBLEVEL NOTATION
ORBITAL DIAGRAM
Oxygen
6
2-6
1
s2
2
s2
2
p4
1s 2s 2p
↑↓
↑↓
↑↓
↑↓
↑↓
O
What would the Lewis Structure look like?
What would the Lewis Structure look like?
(if you don’t know number of valence e-, you can use sublevel notation to figure it out)
Slide38Now try...
Na
N
Slide39Now try...
Na
N
Slide40Now try...
Na
N
Slide41Goal of Bonding is to Gain a Stable Electron Configuration
Atoms create a stable electron configuration by sharing or transferring electrons
Stable electrons are usually 8 valence e- (octet rule)
Helium has 2 e- for a stable configuration (called a duet)
Slide42Why do we use Lewis Structures?
Lewis Structures accurately predict how atoms will likely bond to one another
The fancy math of calculating repulsions and attractions is way too complicated
Slide439.4 Ionic Bonding: Lewis Symbols and Lattice Energies
Slide44To represent an ionic bond, we move electron dots from the metal to the nonmetal and then allow resultant ions to form a crystalline lattice compound
Slide45To represent an ionic bond, we move electron dots from the metal to the nonmetal and then allow resultant ions to form a crystalline lattice compound
(so how do we do that?)
Slide46Consider KCl
Symbol
Number of Valence E-
Potassium
K
1
Chlorine
Cl
7
SO...
Slide47Consider KCl
Symbol
Number of Valence E-
Potassium
K
1
Chlorine
Cl
7
SO...
K +
Cl
K
Cl
+
[ ]
-
Slide48K +
Cl
K
Cl
+
[ ]
-
LET’S ANALYZE THIS
So Chlorine is clearly stable….
why does this work for potassium?
K’s original e- configuration:
1
s2
2
s2
2
p6
3
s2
3
p6
4
s1
(2-8-8-1)
Since potassium gave up the 1 electron in the 4th principal energy level, it has 8 electrons in the 3rd principal energy level giving potassium stable electron configuration as well:
New electron configuration for K+:
1
s2
2
s2
2
p6
3
s2
3
p6
(2-8-8)
△
8 e- in 3rd principal energy level
△
Regents Configuration
△
Regents Configuration
Slide49K +
Cl
K
Cl
+
[ ]
-
Identify the Cation:
Why?
Identify the Anion:
Why?
Slide50K +
Cl
K
Cl
+
[ ]
-
Identify the Cation: K
Why?
Identify the Anion:
Why?
Slide51K +
Cl
K
Cl
+
[ ]
-
Identify the Cation: K
Why?
Identify the Anion:
Why?
K becomes a cation because it gave up its electron and became positive
Slide52K +
Cl
K
Cl
+
[ ]
-
Identify the Cation: K
Why?
Identify the Anion: Cl
Why?
K becomes a cation because it gave up its electron and became positive
Slide53K +
Cl
K
Cl
+
[ ]
-
Identify the Cation: K
Why?
Identify the Anion: Cl
Why?
K becomes a cation because it gave up its electron and became positive
Cl becomes an anion because it accepted K’s electron and became negative
Slide54The positive charge of K+ attracts to the negative charge of Cl-, forming an
ionic bond
This equation also shows us that it only takes
one
potassium and one chlorine to form KCl
Slide55Practice: Na and S
Draw Lewis Structure for each element:
(HINT: USE CRISS CROSS METHOD FOR BONDING TO FIGURE OUT RATIO)
Practice: Na and S
Slide56Practice: Na and S
Draw Lewis Structure for each element:
Draw Combined Lewis Structure:
(HINT: USE CRISS CROSS METHOD FOR BONDING TO FIGURE OUT RATIO)
Slide57What’s the ratio of sodium to sulfur?
What’s the cation?
Why?
What’s the anion?
Why?
Why is sulfur stable?
Why is sodium stable?
Slide58What’s the ratio of sodium to sulfur? 2:1
What’s the cation?
Why?
What’s the anion?
Why?
Why is sulfur stable?
Why is sodium stable?
Slide59What’s the ratio of sodium to sulfur? 2:1
What’s the cation? Na
Why?
What’s the anion?
Why?
Why is sulfur stable?
Why is sodium stable?
Slide60What’s the ratio of sodium to sulfur? 2:1
What’s the cation? Na
Why? Gives up electron to sulfur (positive charge)
What’s the anion?
Why?
Why is sulfur stable?
Why is sodium stable?
Slide61What’s the ratio of sodium to sulfur? 2:1
What’s the cation? Na
Why? Gives up electron to sulfur (positive charge)
What’s the anion? S
Why?
Why is sulfur stable?
Why is sodium stable?
Slide62What’s the ratio of sodium to sulfur? 2:1
What’s the cation? Na
Why? Gives up electron to sulfur (positive charge)
What’s the anion? S
Why? Accepts electron from sodium (negative charge)
Why is sulfur stable?
Why is sodium stable?
Slide63What’s the ratio of sodium to sulfur? 2:1
What’s the cation? Na
Why? Gives up electron to sulfur (positive charge)
What’s the anion? S
Why? Accepts electron from sodium (negative charge)
Why is sulfur stable? Has an octet in outer energy level
Why is sodium stable?
Slide64What’s the ratio of sodium to sulfur? 2:1
What’s the cation? Na
Why? Gives up electron to sulfur (positive charge)
What’s the anion? S
Why? Accepts electron from sodium (negative charge)
Why is sulfur stable? Has an octet in outer energy level
Why is sodium stable? Has an octet in new highest principal energy level after giving up the 1 valence e-
Slide65Lattice Energy
Slide66For example: Energy is released when NaCl is formed
Na
(s)
+ ½ Cl
2(g) ⟶ NaCl
(s) Δ
H = -411 kJ/mol
Slide67For example: Energy is released when NaCl is formed
Na
(s)
+ ½ Cl
2(g) ⟶ NaCl
(s) Δ
H = -411 kJ/mol
Where does this energy comes from?
Slide68For example: Energy is released when NaCl is formed
Na
(s)
+ ½ Cl
2(g) ⟶ NaCl
(s) Δ
H = -411 kJ/mol
Where does this energy comes from?
You would think that because the first ionization energy of Na is +496 kJ/mol and Cl’s is -349 kJ/mol the heat of reaction would be +147 kJ/mol (endothermic)
Slide69For example: Energy is released when NaCl is formed
Na
(s)
+ ½ Cl
2(g) ⟶ NaCl
(s) Δ
H = -411 kJ/mol
Where does this energy comes from?
You would think that because the first ionization energy of Na is +496 kJ/mol and Cl’s is -349 kJ/mol the heat of reaction would be +147 kJ/mol (endothermic)
On top of that, we know that as bonds form, energy is released
Slide70For example: Energy is released when NaCl is formed
Na
(s)
+ ½ Cl
2(g) ⟶ NaCl
(s) Δ
H = -411 kJ/mol
Where does this energy comes from?
So how does this work?
You would think that because the first ionization energy of Na is +496 kJ/mol and Cl’s is -349 kJ/mol the heat of reaction would be +147 kJ/mol (endothermic)
On top of that, we know that as bonds form, energy is released
Slide71For example: Energy is released when NaCl is formed
Na
(s)
+ ½ Cl
2(g) ⟶ NaCl
(s) Δ
H = -411 kJ/mol
Where does this energy comes from?
BECAUSE OF LATTICE ENERGY
(the energy associated with the formation of a crystalline lattice of alternating cations and anions from the gaseous ions)
So how does this work?
You would think that because the first ionization energy of Na is +496 kJ/mol and Cl’s is -349 kJ/mol the heat of reaction would be +147 kJ/mol (endothermic)
On top of that, we know that as bonds form, energy is released
Slide72As the lattice forms, heat is released
The Δ
H
= lattice energy
Slide73△
Lattice Structure
(Has alternating cations and anions forming a lattice structure)
When a lattice forms, potential energy decreases (according to Coulomb's Law)
The potential is released in the form of heat when the lattice forms
The easiest way to calculate heat is with the Born-Haber Cycle
Slide74The Born-Haber Cycle
Born-Haber Cycle - a hypothetical series of steps that represents the formation of an ionic compound from its constituent elements
The steps are chosen so that the change is enthalpy is known except for the last change in enthalpy
Using Hess’s Law, we can determine the enthalpy change for the unknown step (the lattice energy)
(Reminder: what is Hess’s Law?)
LET'S GO BACK TO OUR EQUATION:
Na(s) + ½ Cl2(g) ⟶ NaCl(s) Δ
H
= -411 kJ/mol
Slide75STEP 1: Formation of solid sodium to gaseous sodium
Na(s) ⟶ Na(g) ΔH = 108 kJ
STEP 2: Formation of a chlorine atom from a chlorine molecule
1/2Cl2(g) ⟶ Cl(g) ΔH = 122 kJ
STEP 3: Ionization of gaseous sodium. The enthalpy change for this step is the ionization energy of sodium.
Na+(g) ⟶ Na+(g) + e- ΔH = 496 kJ
STEP 4: Addition of an electron to gaseous chlorine. The enthalpy change for this step is the electron affinity of chlorine.
Cl(g) + e- ⟶Cl-(g) ΔH = -349 kJ
STEP 5: Formation of crystalline solid from the gaseous ion. The enthalpy change for this step is the lattice energy, the unknown quantity.
Na+(g) + Cl-(g) ⟶ NaCl(s) ΔH = ? kJ
Slide76ΔH
f
= ΔH
(step 1) + ΔH(step 2)
+ ΔH(step 3) + ΔH(step 4)
ΔH
lattice
=
ΔH
f
- (
ΔH(step 1) + ΔH
(step 2) + ΔH(step 3) +
ΔH(step 4))
Slide77ΔH
f
= ΔH
(step 1) + ΔH(step 2)
+ ΔH(step 3) + ΔH(step 4)
ΔH
lattice
=
ΔH
f
- (
ΔH(step 1) + ΔH
(step 2) + ΔH(step 3) +
ΔH(step 4))
Na(s) + ½ Cl2
(g) ⟶ NaCl(s) ΔH = -411 kJ/mol
Slide78ΔH
f
= ΔH
(step 1) + ΔH(step 2)
+ ΔH(step 3) + ΔH(step 4)
ΔH
lattice
=
ΔH
f
- (
ΔH(step 1) + ΔH
(step 2) + ΔH(step 3) +
ΔH(step 4))
Na(s) + ½ Cl2
(g) ⟶ NaCl(s) ΔH = -411 kJ/mol
= -411 - (+108 kJ + 122 kJ +476 kJ - 349 kJ)
= -788 kJ
< the value of ΔH
lattice is negative
It’s exothermic because of the large amounts of heat released when sodium and chlorine ions come together to form a crystalline lattice
Slide79Metal Chloride
Lattice Energy kJ/mol
LiCl
-834
NaCl
-788
KCl
-701
CsCl
-657
WE KNOW THAT….
As you can see, the magnitude of energy decreases as you move down the column
Ionic radius increases as you move down a column
Coulomb’s Law says that potential energy of oppositely charged ions become less negative as the distance between ions increase
LATTICE TRENDS: ION SIZE
Slide80Metal Chloride
Lattice Energy kJ/mol
LiCl
-834
NaCl
-788
KCl
-701
CsCl
-657
WE KNOW THAT….
As you can see, the magnitude of energy decreases as you move down the column
Ionic radius increases as you move down a column
Coulomb’s Law says that potential energy of oppositely charged ions become less negative as the distance between ions increase
What does this mean?
LATTICE TRENDS: ION SIZE
LATTICE TRENDS: ION SIZE
Slide81As the size of alkali metals increase down a column, so does the distance between the metal and chloride ions
As a result, the magnitude of the lattice energy of chlorides decreases accordingly, making the formation of chlorides less exothermic
***In other words, as the ionic radii increase as we move down a column, the ions cannot get as close to each other and therefore do not release as much energy when the lattice forms***
Slide82Compound
Lattice Energy kJ/mol
NaF
-910
CaO
-3414
Whys is there much more lattice energy in CaO than NaF?
Slide83Compound
Lattice Energy kJ/mol
NaF
-910
CaO
-3414
Why is there much more lattice energy in CaO than NaF?
SO…
Na+ has a radius of 95 pm and F- has a radius of 136 pm, resulting in a distance between the ions of 231 pm.
Ca 2+ has a radius of 99pm and O 2- has a radius of 140 pm, resulting in a distance between ions 239 pm.
< has a greater distance between ions
Even though the distance is only slightly larger, the lattice energy is about 4x greater
This is because of Coulomb’s Law E = q1q2
r
NaF is portional to (1-)(1+) =1-, while CaO = (2+)(2-) = 4- , so the relative stabilization for CaO is 4x
greater
Slide84SUMMARIZING LATTICE TRENDS:
LATTICE ENERGIES BECOME LESS EXOTHERMIC (LESS NEGATIVE) WITH INCREASING IONIC RADIUS
LATTICE ENERGIES BECOME MORE EXOTHERMIC (MORE NEGATIVE) WITH INCREASING MAGNITUDE OF IONIC CHARGE
Na+
Cl-
K+
231 pm
239 pm
Since the distance is less more energy is released
Since the distance is more, less energy is released
Cl-
Ca2+
O 2-
Na+
F-
1+ 1-
2+ 2-
Overall Charge: -4
Overall Charge: -1
CaO has a higher charge, so more energy is released
NaF has a lower charge, so less energy is released
Slide85THE CLOSER THE IONS CAN GET, THE MORE ENERGY IS RELEASED
THE HIGHER THE COMBINED CHARGE, THE MORE ENERGY IS RELEASED
Slide86IONIC BONDING: MODELS AND REALITY
Slide87Does our ionic bonding model explain the properties of ionic compounds, including their high melting and boiling points, their
tendency not to conduct electricity as solids
, and their tendency to conduct electricity when dissolved in water
?
Slide88Boiling and Melting Points
We modeled an ionic solid as a lattice held together by coulombic forces that are
non-directional
= meaning as you move away from the ion, all forces are equally strong in all directions
To melt a solid, these forces must be overcome, which requires a lot of heat
Therefore our model accounts for high melting and high boiling points in ionic solids
Slide89Conducting Electricity
In the model, electrons are transferred from one element from a metal to a nonmetal, but the transfer of an electron is localized on one atom
SO…..
Our model does not include any free electrons that might conduct
electricity
(the movement or flow of electrons or other charged particles in response to an electric potential or voltage is electrical current)
In addition the ions themselves are fixed in place
Our model accounts for the non-conductivity of ionic solids
Slide90Conducting Electricity
However, when the ions are freed by introducing water to the ionic solid, the cations and anions are free to move in the solution
SO…..
Our model predicts that solutions of ionic compounds can conduct electricity