Basic Concepts of Chemical Bonding Ionic, Covalent, Metallic Bonding Polarity &...

Basic Concepts of Chemical BondingBasic Concepts of Chemical Bonding

Ionic, Covalent, Metallic Bonding

Polarity & ElectronegativityLewis Structures

Ionic, Covalent, Metallic Bonding

Polarity & ElectronegativityLewis Structures

Three Types of Chemical BondsThree Types of Chemical Bonds

Ionic Bond: electrostatic forces that exist between ions of opposite charge. Metals + nonmetals. Ex: NaCl

Covalent Bond: sharing of electrons between two atoms. Nonmetals. Ex: CO2

Metallic Bond: metal ions in a “sea of valence electrons.” Explains high electrical and thermal conductivity of pure metals. Ex: Cu

Ionic Bond: electrostatic forces that exist between ions of opposite charge. Metals + nonmetals. Ex: NaCl

Covalent Bond: sharing of electrons between two atoms. Nonmetals. Ex: CO2

Metallic Bond: metal ions in a “sea of valence electrons.” Explains high electrical and thermal conductivity of pure metals. Ex: Cu

Lewis SymbolsLewis Symbols Only valence electrons are involved in chemical bonding.

We use Lewis electron-dot symbols (Lewis symbols)to show valence electrons.

Lewis symbol = element symbol + dots representing electrons. Can use all four sides around element symbol and no more than two electrons per side. Usually balance one dot on one side with dot on opposite side.

Only valence electrons are involved in chemical bonding.

We use Lewis electron-dot symbols (Lewis symbols)to show valence electrons.

Lewis symbol = element symbol + dots representing electrons. Can use all four sides around element symbol and no more than two electrons per side. Usually balance one dot on one side with dot on opposite side.

Examples of Lewis Symbols

Examples of Lewis Symbols

The Octet RuleThe Octet Rule

Octet Rule: Atoms tend to lose, gain, or share electrons until they are surrounded by eight valence electrons.

Octet Rule: Atoms tend to lose, gain, or share electrons until they are surrounded by eight valence electrons.

Atoms often gain, lose, or share electrons to achieve noble gas configuration of eight valence electrons. Full s and p orbitals.Very stable electron arrangement.

Sodium Chloride Formation

Sodium Chloride Formation

Na(s) + 1/2 Cl2(g) NaCl(s) There is a transfer of an electron from the Na atom to the Cl atom.

Na(s) + 1/2 Cl2(g) NaCl(s) There is a transfer of an electron from the Na atom to the Cl atom.

Ionic Bonding: Formation of NaCl

Ionic Bonding: Formation of NaCl

NaCl is composed of Na+ and Cl- ions that are arranged in a regular three-dimensional array.

Stability of Ionic Compouds

Stability of Ionic Compouds

The heats of formation of ionic compounds are very negative.

Ions are drawn together, releasing energy as a solid lattice forms.

Ionic compounds are very stable.

Ionic compounds are hard & brittle with high melting points.

The heats of formation of ionic compounds are very negative.

Ions are drawn together, releasing energy as a solid lattice forms.

Ionic compounds are very stable.

Ionic compounds are hard & brittle with high melting points.

Lattice EnergyLattice Energy Lattice energy: the energy required to completely separate a mole of solid ionic compound into its gaseous ions.

Note: This is an endothermic process.

***The higher the lattice energy, the stronger the ionic bond.***

NaCl(s) Na+(g) + Cl-

(g)

∆H = +788 kJ/mol

Lattice energy: the energy required to completely separate a mole of solid ionic compound into its gaseous ions.

Note: This is an endothermic process.

***The higher the lattice energy, the stronger the ionic bond.***

NaCl(s) Na+(g) + Cl-

(g)

∆H = +788 kJ/mol

Lattice Energy, Cont.

Lattice Energy, Cont.

Lattice energy depends on: Eelectric = (Q1Q2)/d Where: Q1 and Q2 = ion charges; = constant; d = distance between center of ions

Lattice energy increases as the charges on the ions increase and as their radii decrease.

Lattice energy depends on: Eelectric = (Q1Q2)/d Where: Q1 and Q2 = ion charges; = constant; d = distance between center of ions

Lattice energy increases as the charges on the ions increase and as their radii decrease.

Magnitude of lattice energy depends mainly on ionic charge because radii do not vary greatly.

Magnitude of lattice energy depends mainly on ionic charge because radii do not vary greatly.

2+; 1-

2+; 2-

3+; 3-

1+; 1-

Practice With Lattice Energy

Practice With Lattice Energy

Arrange the following in order of increasing lattice energy: NaF, CsI, CaO

CsI < NaF < CaO

Which substance has the greatest lattice energy? Why? AgCl, CuO, CrN

CrN, because of larger 3+/3- charges

Arrange the following in order of increasing lattice energy: NaF, CsI, CaO

CsI < NaF < CaO

Which substance has the greatest lattice energy? Why? AgCl, CuO, CrN

CrN, because of larger 3+/3- charges

Born-Haber CycleBorn-Haber Cycle Lattice energies cannot be measured experimentally.

We can use Hess’s Law to construct a Born-Haber cycle.

A Born-Haber cycle shows the energetic relationship in the formation of ionic solids from the elements. Enthalpy of formation of a compound is equal to the sum of the energies of several individual steps.

Lattice energies cannot be measured experimentally.

We can use Hess’s Law to construct a Born-Haber cycle.

A Born-Haber cycle shows the energetic relationship in the formation of ionic solids from the elements. Enthalpy of formation of a compound is equal to the sum of the energies of several individual steps.

The principles of Hess’s Law can be applied in the Born-Haber cycle to determine the lattice energy of NaCl.

The principles of Hess’s Law can be applied in the Born-Haber cycle to determine the lattice energy of NaCl.

Covalent BondingCovalent Bonding

Covalent Bond: chemical bond formed by sharing of electrons to form noble gas configurations. Ex: H2 molecule

Covalent Bond: chemical bond formed by sharing of electrons to form noble gas configurations. Ex: H2 molecule

Multiple BondsMultiple Bonds Single bond: sharing of 2 electrons. Shown as H-H or H:H

Double bond : sharing of 4 electrons. Shown as O::C::O or O=C=O

Triple bond: sharing of 6 electrons. Shown as :N:::N: or :NN:

Single bond: sharing of 2 electrons. Shown as H-H or H:H

Double bond : sharing of 4 electrons. Shown as O::C::O or O=C=O

Triple bond: sharing of 6 electrons. Shown as :N:::N: or :NN:

Examples of Multiple Bonds

Examples of Multiple Bonds

Bond LengthsBond Lengths Average bond distance varies with number of bonds. Generally, as number of shared electron pairs increases, bond length decreases.

N-N (1.47Å) N=N (1.24Å) NN (1.10Å) Note: A triple bond is NOT 3x stronger than a single bond. The first bond in a triple bond is stronger than a “normal” single bond. The “second” and “third” bonds in a triple bond are weaker than the “first” bond.

Average bond distance varies with number of bonds. Generally, as number of shared electron pairs increases, bond length decreases.

N-N (1.47Å) N=N (1.24Å) NN (1.10Å) Note: A triple bond is NOT 3x stronger than a single bond. The first bond in a triple bond is stronger than a “normal” single bond. The “second” and “third” bonds in a triple bond are weaker than the “first” bond.

Bond PolarityBond Polarity Bond polarity helps to describe the sharing of electrons between atoms.

Nonpolar covalent bond: electrons shared equally between atoms.

Polar covalent bond: one of the atoms in a compound exerts a greater attraction for the bonding electrons than the other. If the difference is big enough, we get an ionic bond.

Bond polarity helps to describe the sharing of electrons between atoms.

Nonpolar covalent bond: electrons shared equally between atoms.

Polar covalent bond: one of the atoms in a compound exerts a greater attraction for the bonding electrons than the other. If the difference is big enough, we get an ionic bond.

ElectronegativityElectronegativity We use electronegativity to estimate whether a bond will be nonpolar covalent, polar covalent, or ionic.

Electronegativity: the ability of atoms in a molecule to attract electrons to itself.

The greater the electronegativity, the greater its ability to attract electrons.

High ionization energy & very negative electron affinity = high electronegativity.

We use electronegativity to estimate whether a bond will be nonpolar covalent, polar covalent, or ionic.

Electronegativity: the ability of atoms in a molecule to attract electrons to itself.

The greater the electronegativity, the greater its ability to attract electrons.

High ionization energy & very negative electron affinity = high electronegativity.

Electronegativity Scale

Electronegativity Scale

Linus Pauling makes electronegativity scale.

Values are unitless Scale: 0.7 (cesium) - 4.0 (fluorine)

Metals are less electronegative; nonmetals are more electronegative.

Linus Pauling makes electronegativity scale.

Values are unitless Scale: 0.7 (cesium) - 4.0 (fluorine)

Metals are less electronegative; nonmetals are more electronegative.

Electronegativity & Bond Polarity

Electronegativity & Bond Polarity

The greater the difference in electronegativity between two bonded atoms, the more polar the bond.

Difference < 0.5 = nonpolar covalent Ex: F2

4.0 - 4.0 = 0 0.5 ≤ difference < 2.0 = polar covalent Ex: HF 4.0 - 2.1 = 1.9

Difference ≥ 2.00 = ionic bond Ex: LiF 4.0 - 1.0 = 3.0

Use + and - to show partial positive and negative charges on atoms.

The greater the difference in electronegativity between two bonded atoms, the more polar the bond.

Difference < 0.5 = nonpolar covalent Ex: F2

4.0 - 4.0 = 0 0.5 ≤ difference < 2.0 = polar covalent Ex: HF 4.0 - 2.1 = 1.9

Difference ≥ 2.00 = ionic bond Ex: LiF 4.0 - 1.0 = 3.0

Use + and - to show partial positive and negative charges on atoms.

Polar covalent bondPolar covalent bondNonpolar covalent bond Ionic bond

+ -

+ -

Polarity PracticePolarity Practice Tell which bond is more polar and indicate in each case which atom has the partial negative charge.

B-Cl or C-Cl? B-Cl, with - on Cl P-F or P-Cl? P-F, with - on F

Tell which bond is more polar and indicate in each case which atom has the partial negative charge.

B-Cl or C-Cl? B-Cl, with - on Cl P-F or P-Cl? P-F, with - on F

Polar MoleculesPolar Molecules

Entire molecules can be polar, not just bonds within the molecules. Ex: HF

Polarity is important--it helps to determine many properties of a cmpd!

Entire molecules can be polar, not just bonds within the molecules. Ex: HF

Polarity is important--it helps to determine many properties of a cmpd!

Dipole MomentsDipole Moments

Dipole: two electrical charges of equal magnitude but opposite sign separated by a distance.

Dipole moment ():quantitative measure of dipole magnitude in unit of debyes (D).

= Qr Q= charge; r= distance

Dipole: two electrical charges of equal magnitude but opposite sign separated by a distance.

Dipole moment ():quantitative measure of dipole magnitude in unit of debyes (D).

= Qr Q= charge; r= distance

Drawing Lewis Structures

Drawing Lewis Structures

1. Sum valence e-’s of all atoms. Add e-’s for anions.

2. Sum number of e-’s each atom needs for an octet. All atoms need 8 e-’s except for H that needs only 2 e-’s.

3. (e-’s needed-valence e-’s)/2 = minimum number of bonds in the molecule

Note: bonds may be single or multiple. One double bond=two single bonds

1. Sum valence e-’s of all atoms. Add e-’s for anions.

2. Sum number of e-’s each atom needs for an octet. All atoms need 8 e-’s except for H that needs only 2 e-’s.

3. (e-’s needed-valence e-’s)/2 = minimum number of bonds in the molecule

Note: bonds may be single or multiple. One double bond=two single bonds

Lewis Structures, Cont.

Lewis Structures, Cont.

4. Draw symbols, attaching bonds as needed. Note: H only takes a single bond. C takes 4 bonds, is often central.

5. Complete the octets of atoms bonded to central atom. Use dots for lone pairs of e-’s.

6. Place any leftover e-’s on central atom even if it has more than an octet.

4. Draw symbols, attaching bonds as needed. Note: H only takes a single bond. C takes 4 bonds, is often central.

5. Complete the octets of atoms bonded to central atom. Use dots for lone pairs of e-’s.

6. Place any leftover e-’s on central atom even if it has more than an octet.

Draw Structure for PBr3

Draw Structure for PBr3

Which is your central atom? Determine number of valence e-’s. P has 5, Br has (3 x 7). 5 + 21 = 26 e-’s

Determine number of e-’s needed for octets. P has (1 x 8), Br has (3 x 8). 4 x 8 = 32 e-’s

(Needs - valence)/2 = minimum bonds needed (32-26)/2 = 3 bonds needed at minimum

Which is your central atom? Determine number of valence e-’s. P has 5, Br has (3 x 7). 5 + 21 = 26 e-’s

Determine number of e-’s needed for octets. P has (1 x 8), Br has (3 x 8). 4 x 8 = 32 e-’s

(Needs - valence)/2 = minimum bonds needed (32-26)/2 = 3 bonds needed at minimum

PBr3 Lewis StructurePBr3 Lewis Structure

Note: bonds + lone pairs = number of valence e-’s.

Note: bonds + lone pairs = number of valence e-’s.

Formal ChargeFormal Charge Sometimes multiple Lewis structures can be drawn for a molecule. Formal charge can help determine correct one.

Formal charge: charge that an atom would have in a molecule if all atoms had the same electronegativity (bonding e-’s are shared equally).

Sometimes multiple Lewis structures can be drawn for a molecule. Formal charge can help determine correct one.

Formal charge: charge that an atom would have in a molecule if all atoms had the same electronegativity (bonding e-’s are shared equally).

Calculating Formal Charge

Calculating Formal Charge

Electrons are assigned as follows: 1) All unshared (nonbonding) e-’s are assigned to the atom on which they are found.

2) Half of the bonding e-’s are assigned to each atom in the bond.

Formal charge = # of valence e-’s - # of e-’s assigned in Lewis structure.

Electrons are assigned as follows: 1) All unshared (nonbonding) e-’s are assigned to the atom on which they are found.

2) Half of the bonding e-’s are assigned to each atom in the bond.

Formal charge = # of valence e-’s - # of e-’s assigned in Lewis structure.

Formal Charge, Cont.Formal Charge, Cont.

It is important to note that oxidation number overstates the importance of electronegativity and formal charge ignores the role of electronegativity completely.

It is important to note that oxidation number overstates the importance of electronegativity and formal charge ignores the role of electronegativity completely.

Oxidation NumberOxidation Number

Formal Charge

Polarity

Formal Charge as the “Tie Breaker”

Formal Charge as the “Tie Breaker”

As a general rule, when several Lewis structures are possible, the most stable (and favored) one will be the one in which…

1) the atoms bear formal charges closest to zero.

2) any negative charges reside on the more electronegative atoms.

As a general rule, when several Lewis structures are possible, the most stable (and favored) one will be the one in which…

1) the atoms bear formal charges closest to zero.

2) any negative charges reside on the more electronegative atoms.

Formal Charge & NO2Formal Charge & NO2

Resonance StructuresResonance Structures Sometimes molecules and ions cannot be described by a single Lewis structure.

Resonance structure: atoms keep same arrangement but placement of e-’s changes. Look for changes in placement of double bonds.

Sometimes molecules and ions cannot be described by a single Lewis structure.

Resonance structure: atoms keep same arrangement but placement of e-’s changes. Look for changes in placement of double bonds.

ResonanceResonance All of the resonance structures are taken together as a description of a molecule or ion.

All of the resonance structures are taken together as a description of a molecule or ion.

A molecule or ion does not switch among resonance structures, it is more of a blending.Not all resonance structures are equivalent. Formal charges can help to predict favored structures.

Resonance StructuresResonance Structures

Resonance In BenzeneResonance In BenzeneBenzene (C6H6) is an important example of an aromatic organic molecule. It is the “poster child” for resonance.

Benzene (C6H6) is an important example of an aromatic organic molecule. It is the “poster child” for resonance.

Exceptions to Octet Rule

Exceptions to Octet Rule

Odd number of electrons: Odd number of electrons:

Less than an octet:

Exceptions, Cont.Exceptions, Cont.

More than an octet: More than an octet:

Bond EnthalpyBond Enthalpy Stability of a molecule is related to the strengths of its covalent bonds.

Bond Enthalpy: ∆H for the breaking of a particular bond in a mole of gaseous substance.

Bond enthalpy is always endothermic. Greater bond enthalpy = stronger bond

Stronger bonds = less reactivity

Stability of a molecule is related to the strengths of its covalent bonds.

Bond Enthalpy: ∆H for the breaking of a particular bond in a mole of gaseous substance.

Bond enthalpy is always endothermic. Greater bond enthalpy = stronger bond

Stronger bonds = less reactivity

Molecular Orbital Model

Molecular Orbital Model

Some problems with localized electron model:

1. Incorrectly assumes e-’s are localized.

2. Not good with unpaired e-’s. 3. No direct info on bond energies. Another model is molecular orbital model.

Some problems with localized electron model:

1. Incorrectly assumes e-’s are localized.

2. Not good with unpaired e-’s. 3. No direct info on bond energies. Another model is molecular orbital model.

MO TheoryMO Theory MO’s are solutions to molecular problems such as why does O2 interact with a magnetic field and why are some molecules colored?)

• Just as e-’s in atoms are found in atomic orbitals, e-’s in molecules are found in MO’s.

Square of molecular orbital wave function indicates e- probability.

MO’s are solutions to molecular problems such as why does O2 interact with a magnetic field and why are some molecules colored?)

• Just as e-’s in atoms are found in atomic orbitals, e-’s in molecules are found in MO’s.

Square of molecular orbital wave function indicates e- probability.

Molecular OrbitalsMolecular Orbitals

•each contain a maximum of two electrons with opposite spin

•have definite energies•are associated with an entire molecule.

•When two AOs overlap, two MOs form.

•each contain a maximum of two electrons with opposite spin

•have definite energies•are associated with an entire molecule.

•When two AOs overlap, two MOs form.

Hydrogen ExampleHydrogen Example• 1s (H) + 1s (H) must result in two MOs for H2:

• one has electron density between nuclei (bonding MO);

• one has little electron density between nuclei (antibonding MO).

• MOs resulting from s orbitals are MOs.

(bonding) MO is lower energy than * (antibonding) MO.

• 1s (H) + 1s (H) must result in two MOs for H2:

• one has electron density between nuclei (bonding MO);

• one has little electron density between nuclei (antibonding MO).

• MOs resulting from s orbitals are MOs.

(bonding) MO is lower energy than * (antibonding) MO.

Hydrogen Molecular Orbitals

MO DiagramMO Diagram• Energy level diagram or MO diagram show energies and electrons in an orbital.

• Total number of electrons in all atoms are placed in the MOs starting from lowest energy (1s) and ending when you run out of electrons.• Note that electrons in MOs have opposite spins.

• H2 has two bonding electrons.

• Energy level diagram or MO diagram show energies and electrons in an orbital.

• Total number of electrons in all atoms are placed in the MOs starting from lowest energy (1s) and ending when you run out of electrons.• Note that electrons in MOs have opposite spins.

• H2 has two bonding electrons.

Hydrogen and Helium MO’s

Two molecular orbitals form from the 2 separated H atoms: MO1 = 1sA + 1sB (Bonding orbital) MO2 = 1sA - 1sB (Antibonding orbital)

Hydrogen’s Molecular Orbitals

Hydrogen’s Molecular Orbitals

MO1 has greatest e- density between nuclei.

MO2 has greatest e- density on either side of nuclei.

Called sigma () molecular orbitals.

Note: sigma MO’s replace 1s atomic orbitals that no longer exist.

MO1 has greatest e- density between nuclei.

MO2 has greatest e- density on either side of nuclei.

Called sigma () molecular orbitals.

Note: sigma MO’s replace 1s atomic orbitals that no longer exist.

Hydrogen’s Molecular Orbitals

Hydrogen’s Molecular Orbitals

MO1 lower in energy than 1s atomic orbitals; MO2 higher in energy.

If 2 e-’s occupy MO1, energy will be lower than in 2 separate H atoms.

Lower energy = driving force for molecular formation = probonding.

MO1 lower in energy than 1s atomic orbitals; MO2 higher in energy.

If 2 e-’s occupy MO1, energy will be lower than in 2 separate H atoms.

Lower energy = driving force for molecular formation = probonding.

Bonding & Antibonding MO’s

Bonding & Antibonding MO’s

Bonding MO: lower in energy than atomic orbitals of which it is composed.

Antibonding MO: higher in energy than the atomic orbitals of which it is composed.

Bonding MO: lower in energy than atomic orbitals of which it is composed.

Antibonding MO: higher in energy than the atomic orbitals of which it is composed.

MO LabelingMO Labeling

Labels on MO’s indicate symmetry, parent atomic orbitals, and if they are bonding or antibonding.

MO’s for H2 are:

MO1 = bonding = 1s

MO2 = antibonding = 1s*

Since H2 has 2 e-’s, we write the electron configuration as 1s

2

Labels on MO’s indicate symmetry, parent atomic orbitals, and if they are bonding or antibonding.

MO’s for H2 are:

MO1 = bonding = 1s

MO2 = antibonding = 1s*

Since H2 has 2 e-’s, we write the electron configuration as 1s

2

Bond OrderBond Order

• Bond order indicates bond strength. • Larger bond order = greater strength.

• Bond order = 1 for single bond.• Bond order = 2 for double bond.• Bond order = 3 for triple bond.• Fractional bond orders are possible.

•For H2

• Therefore, H2 has a single bond.

• Bond order indicates bond strength. • Larger bond order = greater strength.

• Bond order = 1 for single bond.• Bond order = 2 for double bond.• Bond order = 3 for triple bond.• Fractional bond orders are possible.

•For H2

• Therefore, H2 has a single bond.

electrons gantibondin-electrons bondingorder Bond21

102order Bond21

Bond Order & StabilityBond Order & Stability

•For H2:

• Therefore, H2 has a single bond.

•For He2:

• Therefore He2 is not a stable molecule.

•For H2:

• Therefore, H2 has a single bond.

•For He2:

• Therefore He2 is not a stable molecule.

022order Bond21

102order Bond21