LEWIS STRUCTURES BONDS IONIC BONDING

LEWIS STRUCTURES BONDS IONIC BONDING

9.1 Bonding Models and AIDS Drugs

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Bonding 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

9.2 Types of Bonds

Why Do Chemical Bonds Form?

Chemical bonds form because they lower the potential energy between charged particles that compose atoms.

Chemical bonds form because they lower the potential energy between charged particles that compose atoms. ?

AS 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)

AS 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

Classification 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

NaCl ▽ △ H2O Molecule △ Na(s)

COVALENT NaCl ▽ △ H2O Molecule △ Na(s)

COVALENT NaCl ▽ △ H2O Molecule △ Na(s) IONIC

METALLIC COVALENT NaCl ▽ △ H2O Molecule △ Na(s) IONIC

Metals have low electronegativity and ionization energies
Electronegativity = 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

QUICK REVIEW (OF VALENCE ELECTRON)

What are valence electrons? What are valence electrons?

The electrons in the outermost principal energy level

The electrons in the outermost principal energy level Why do we care? What does this mean?

The electrons in the outermost principal energy level Why do we care? What does this mean? They are the least attracted to the nucleus

The electrons in the outermost principal energy level Why do we care? What does this mean? They are the least attracted to the nucleus So they are important for bonding because... So they are important for bonding because...

The electrons in the outermost principal energy level Why do we care? What does this mean? They are the least attracted to the nucleus So they are important for bonding because... So they are important for bonding because... The valence are the ones that will be shared or transferred during a bond

Act as anions and cations Give up their electrons (cation)
Groups 1-3 usually... Group 4... Groups 5-8 usually... Act as anions and cations Give up their electrons (cation) Accept other elements electrons (anion)

Back to Lesson

Steps 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

What would the Lewis Structure look like?
ELEMENT # OF VALENCE E- E- CONFIG. SUBLEVEL NOTATION ORBITAL DIAGRAM Oxygen 6 2-6 1s2 2s2 2p4 1s s p What would the Lewis Structure look like?

ELEMENT # 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?

ELEMENT # OF VALENCE E- E- CONFIG. SUBLEVEL NOTATION ORBITAL DIAGRAM Oxygen 6 2-6 1s2 2s2 2p4 1s s p ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ (if you don’t know number of valence e-, you can use sublevel notation to figure it out) What would the Lewis Structure look like? What would the Lewis Structure look like?

ELEMENT # OF VALENCE E- E- CONFIG. SUBLEVEL NOTATION ORBITAL DIAGRAM Oxygen 6 2-6 1s2 2s2 2p4 1s s p ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ (if you don’t know number of valence e-, you can use sublevel notation to figure it out) What would the Lewis Structure look like? What would the Lewis Structure look like? O

Now try... Na N

Goal 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)

Why 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

9.4 Ionic Bonding: Lewis Symbols and Lattice Energies

To 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

To 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?)

Consider KCl Symbol Number of Valence E- Potassium K 1 Chlorine Cl 7
SO...

[ ] K + Cl K Cl Consider KCl Symbol Number of Valence E- Potassium K 1
Chlorine Cl 7 SO... [ ] + - K + Cl K Cl

[ ] + - K + Cl K Cl LET’S ANALYZE THIS So Chlorine is clearly stable….why does this work for potassium? K’s original e- configuration: 1s2 2s2 2p6 3s2 3p6 4s1 ( ) 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+: 1s2 2s2 2p6 3s2 3p6 (2-8-8) △ Regents Configuration △ Regents Configuration △ 8 e- in 3rd principal energy level

[ ] + - K + Cl K Cl Identify the Cation: Why? Identify the Anion: Why?

[ ] K + Cl K Cl + - Identify the Cation: K Why? Identify the Anion:
[ ] + - K + Cl K Cl Identify the Cation: K Why? Identify the Anion: Why?

[ ] K + Cl K Cl + - Identify the Cation: K Why?
[ ] + - K + Cl K Cl Identify the Cation: K Why? K becomes a cation because it gave up its electron and became positive Identify the Anion: Why?

[ ] + - K + Cl K Cl Identify the Cation: K Why? K becomes a cation because it gave up its electron and became positive Identify the Anion: Cl Why?

[ ] + - K + Cl K Cl Identify the Cation: K Why? K becomes a cation because it gave up its electron and became positive Identify the Anion: Cl Why? Cl becomes an anion because it accepted K’s electron and became negative

The 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

Practice: Na and S Practice: Na and S
Draw Lewis Structure for each element: (HINT: USE CRISS CROSS METHOD FOR BONDING TO FIGURE OUT RATIO)

(HINT: USE CRISS CROSS METHOD FOR BONDING TO FIGURE OUT RATIO)
Practice: Na and S Draw Lewis Structure for each element: Draw Combined Lewis Structure: Do on white board (HINT: USE CRISS CROSS METHOD FOR BONDING TO FIGURE OUT RATIO)

What’s the ratio of sodium to sulfur?
What’s the cation? Why? What’s the anion? Why is sulfur stable? Why is sodium stable?

What’s the ratio of sodium to sulfur? 2:1
What’s the cation? Why? What’s the anion? Why is sulfur stable? Why is sodium stable?

What’s the cation? Na Why? What’s the anion? Why is sulfur stable? Why is sodium stable?

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?

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?

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?

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?

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-

Lattice Energy

For example: Energy is released when NaCl is formed
Na(s) + ½ Cl2(g) ⟶ NaCl(s) ΔH = -411 kJ/mol

Where does this energy comes from? Na(s) + ½ Cl2(g) ⟶ NaCl(s) ΔH = -411 kJ/mol

Where does this energy comes from? Na(s) + ½ Cl2(g) ⟶ NaCl(s) ΔH = -411 kJ/mol 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)

Where does this energy comes from? Na(s) + ½ Cl2(g) ⟶ NaCl(s) ΔH = -411 kJ/mol 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

Where does this energy comes from? Na(s) + ½ Cl2(g) ⟶ NaCl(s) ΔH = -411 kJ/mol 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 So how does this work?

BECAUSE OF LATTICE ENERGY
For example: Energy is released when NaCl is formed Where does this energy comes from? Na(s) + ½ Cl2(g) ⟶ NaCl(s) ΔH = -411 kJ/mol 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 So how does this work? BECAUSE OF LATTICE ENERGY (the energy associated with the formation of a crystalline lattice of alternating cations and anions from the gaseous ions)

As the lattice forms, heat is released
The ΔH = lattice energy

(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 △ Lattice Structure (Has alternating cations and anions forming a lattice structure)

The Born-Haber Cycle LET'S GO BACK TO OUR EQUATION:
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) ΔH = -411 kJ/mol

STEP 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

ΔHf = ΔH(step 1) + ΔH(step 2) + ΔH(step 3) + ΔH(step 4)
ΔHlattice= ΔHf - (ΔH(step 1) + ΔH(step 2) + ΔH(step 3) + ΔH(step 4))

ΔHlattice= ΔHf - (ΔH(step 1) + ΔH(step 2) + ΔH(step 3) + ΔH(step 4)) Na(s) + ½ Cl2(g) ⟶ NaCl(s) ΔH = -411 kJ/mol

ΔHlattice= ΔHf - (ΔH(step 1) + ΔH(step 2) + ΔH(step 3) + ΔH(step 4)) Na(s) + ½ Cl2(g) ⟶ NaCl(s) ΔH = -411 kJ/mol = (+108 kJ kJ +476 kJ kJ) = -788 kJ < the value of ΔHlattice 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

Ionic radius increases as you move down a column
Metal 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

What does this mean? WE KNOW THAT….
Metal 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

As 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***

Whys is there much more lattice energy in CaO than NaF?
Compound Lattice Energy kJ/mol NaF -910 CaO -3414 Whys is there much more lattice energy in CaO than NaF?

Why is there much more lattice energy in CaO than NaF?
Compound 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

SUMMARIZING LATTICE TRENDS:
LATTICE ENERGIES BECOME LESS EXOTHERMIC (LESS NEGATIVE) WITH INCREASING IONIC RADIUS Cl- K+ Cl- Na+ 231 pm 239 pm Since the distance is less more energy is released Since the distance is more, less energy is released LATTICE ENERGIES BECOME MORE EXOTHERMIC (MORE NEGATIVE) WITH INCREASING MAGNITUDE OF IONIC CHARGE Ca2+ O 2- Na+ F- Overall Charge: -4 You can look at your charts for exact data Overall Charge: -1 CaO has a higher charge, so more energy is released NaF has a lower charge, so less energy is released

THE CLOSER THE IONS CAN GET, THE MORE ENERGY IS RELEASED
THE HIGHER THE COMBINED CHARGE, THE MORE ENERGY IS RELEASED

IONIC BONDING: MODELS AND REALITY

Does 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?

Boiling 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

Conducting 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

Conducting 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

LEWIS STRUCTURES BONDS IONIC BONDING

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