Chapter 10 - Chemical Bonding: The Lewis Model

HIV-Protease

• HIV-protease is a protein synthesized by the human immunodeficiency virus (HIV).

• This protein is crucial for the virus’s ability to multiply and cause AIDS.

• Pharmaceutical companies designed molecules to disable HIV-protease by binding to its active site, known as protease inhibitors.

• Researchers utilized bonding theories to simulate the shape of potential drug molecules and their interaction with the protease molecule.

Bonding Theories

• Explain how and why atoms attach together to form molecules.

• Explain why some combinations of atoms are stable while others are not (e.g., why water is $H<em>2O$ and not $HO$ or $H</em>3O$).

• Can be used to predict the shapes of molecules.

• Can be used to predict the chemical and physical properties of compounds.

Lewis Model

• One of the simplest bonding theories is the Lewis model.

• The Lewis model emphasizes valence electrons to explain bonding.

• Valence electrons are represented as dots, and Lewis structures are drawn to depict molecules (also known as electron dot structures).

• Lewis structures allow the prediction of many properties of molecules, such as molecular stability.

Why Do Atoms Bond?

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

• A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms.

• To calculate this potential energy, you need to consider the following interactions:

  • Nucleus-to-nucleus repulsions

  • Electron-to-electron repulsions

  • Nucleus-to-electron attractions

Types of Bonds

• Bonds can be classified based on the kinds of atoms that are bonded together:

  • Metal and nonmetal: Ionic bonds (electrons transferred)

  • Nonmetal and nonmetal: Covalent bonds (electrons shared)

  • Metal and metal: Metallic bonds (electrons pooled)

Ionic Bonds

• When a metal atom loses electrons, it becomes a cation.

  • Metals have low ionization energy, making it relatively easy to remove electrons from them.

• When a nonmetal atom gains electrons, it becomes an anion.

  • Nonmetals have high electron affinities, making it advantageous to add electrons to these atoms.

• The oppositely charged ions are then attracted to each other, resulting in an ionic bond.

Covalent Bonds

• Nonmetal atoms have relatively high ionization energies, so it is difficult to remove electrons from them.

• When nonmetals bond together, it is more favorable in terms of potential energy for the atoms to share valence electrons.

  • Potential energy is lowest when the electrons are between the nuclei.

• Shared electrons hold the atoms together by attracting the nuclei of both atoms.

Metallic Bonds

• The relatively low ionization energy of metals allows them to lose electrons easily.

• The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared as a pool by all the atoms/ions in the metal.

  • This creates an organization of metal cation "islands" in a "sea" of electrons.

  • Electrons are delocalized throughout the metal structure.

• Bonding results from the attraction of the cations for the delocalized electrons.

Valence Electrons and Bonding

• Valence electrons are held most loosely.

• Chemical bonding involves the transfer or sharing of electrons between two or more atoms.

• Because of the above facts, valence electrons are most important in bonding.

• The Lewis model focuses on the valence electrons.

Determining the Number of Valence Electrons in an Atom

• The main-group column number on the periodic table indicates the number of valence electrons a main group atom has.

• Transition elements all have two valence electrons.

Lewis Structures of Atoms

• In a Lewis structure, valence electrons of main-group elements are represented as dots surrounding the symbol for the element (also known as electron dot structures).

• The symbol of the element represents the nucleus and inner electrons.

• Dots around the symbol represent valence electrons:

  • Pair the first two dots for the s orbital electrons.

  • Put one dot on each open side for the first three p electrons.

  • Then, pair the rest of the dots for the remaining p electrons.

Lewis Bonding Theory

• Atoms bond because bonding results in a more stable electron configuration (lower potential energy).

• A chemical bond forms by atoms either transferring or sharing electrons.

• Usually, this results in all atoms obtaining an outer shell with eight electrons (octet rule).

  • There are some exceptions to this rule. The key is to try to get an electron configuration like a noble gas.

Stable Electron Arrangements and Ion Charge

• Metals form cations by losing valence shell electrons.

• Nonmetals form anions by gaining valence electrons.

Lewis Theory and Ionic Bonding

• Lewis symbols can be used to represent the transfer of electrons from a metal atom to a nonmetal atom, resulting in ions that are attracted to each other and, therefore, bond.

Lewis Model Predictions for Ionic Bonding

• The Lewis model predicts the number of electrons a metal atom should lose or a nonmetal atom should gain in order to attain a stable electron arrangement (octet rule).

• This allows us to predict the formulas of ionic compounds that result.

• It also allows us to predict the relative strengths of the resulting ionic bonds from Coulomb’s law.

Ionic Bonding and the Crystal Lattice

• A crystal lattice is the alternating cations and anions from the gaseous ions coming together in a crystalline structure.

• The crystal lattice is held together by the electrostatic attraction of the cations for all the surrounding anions.

  • Attractions between cations and anions are maximized, leading to the most stable arrangement.

Lattice Energy

• The extra stability that accompanies the formation of the crystal lattice is measured as the lattice energy.

• The lattice energy is the energy released when the solid crystal forms from separate ions in the gas state.

  • Always exothermic

  • Hard to measure directly but can be calculated from knowledge of other processes

• Lattice energy depends directly on the size of charges and inversely on the distance between ions.

Determining Lattice Energy: The Born–Haber Cycle

• The Born–Haber cycle is a hypothetical series of reactions that represents the formation of an ionic compound from its constituent elements.

• The reactions are chosen so that the change in enthalpy of each reaction is known except for the last one, which is the lattice energy.

• Use Hess’s law to add up enthalpy changes of other reactions to determine the lattice energy.

  • Don’t forget to add together all the ionization energies to get to the desired cation.

• $\Delta H<em>f^\circ\text{(salt)} = \Delta H</em>f^\circ\text{(metal atoms, g)} + \Delta H<em>f^\circ\text{(nonmetal atoms, g)} + \Delta H</em>f^\circ\text{(cations, g)} + \Delta H_f^\circ\text{(anions, g)} + \Delta H^\circ\text{(crystal lattice)}$

• $\Delta H^\circ\text{(crystal lattice)} = \text{lattice energy}$

  • For metal atom (g) → cation (g), $\Delta H_f^\circ = \text{first ionization energy}$

  • For nonmetal atoms (g) → anions (g), $\Delta H_f^\circ = \text{electron affinity}$

Trends in Lattice Energy

• The force of attraction between charged particles is inversely proportional to the distance between them.

• Larger ions mean the center of positive charge (nucleus of the cation) is farther away from the negative charge (electrons of the anion).

  • Less exothermic lattice energy with larger ionic radius

  • More exothermic lattice energy with increasing magnitude of ionic charge

Lattice Energy Versus Ion Size

• Lithium Chloride (LiCl): -834 kJ/mol

• Sodium Chloride (NaCl): -788 kJ/mol

• Potassium Chloride (KCl): -701 kJ/mol

• Cesium Chloride (CsCl): -657 kJ/mol

Trends in Lattice Energy: Ion Charge

• The force of attraction between oppositely charged particles is directly proportional to the product of the charges.

• Larger charge means the ions are more strongly attracted.

  • Larger charge = stronger attraction

  • Stronger attraction = larger lattice energy

• Of the two factors (ion size and ion charge), ion charge is generally more important.

Ionic Bonding Model Versus Reality

• Lewis theory implies that the attractions between ions are strong.

• Lewis theory predicts ionic compounds should have high melting points and boiling points because breaking down the crystal should require a lot of energy.

  • The stronger the attraction (larger the lattice energy), the higher the melting point.

• Ionic compounds have high melting points and boiling points.

Ionic Solids

• Electrostatic attraction is nondirectional, with no direct anion–cation pair.

• Therefore, there is no ionic molecule.

  • The chemical formula is an empirical formula, simply giving the ratio of ions based on charge balance.

Properties of Ionic Compounds

• Hard and brittle crystalline solids that are solids at room temperature.

• Melting points are generally > 300°C.

• The liquid state conducts electricity, but the solid state does not.

• Many are soluble in water, and the solution conducts electricity well.

Conductivity of NaCl

• In solid NaCl ($NaCl(s)$), ions are fixed in place and not allowed to move to the charged rods.

• In aqueous NaCl ($NaCl(aq)$), ions are separated and allowed to move to the charged rods.

Covalent Bonding: Lewis Structures

• The Lewis model provides us with a simple and useful model for covalent bonding.

• Lewis structures depict neighboring atoms as sharing some (or all) of their valence electrons in order to attain octets (or duets for hydrogen).

• The shared electrons would then count toward each atom’s octet.

• A shared pair of valence electrons is called a covalent bond.

Covalent Bonding: Bonding and Lone Pair Electrons

• Electrons that are shared by atoms are called bonding pairs.

• Electrons that are not shared by atoms but belong to a particular atom are called lone pairs (also known as nonbonding pairs).

Single Covalent Bonds

• When two atoms share one pair of electrons, it is called a single covalent bond (two electrons).

• One atom may use more than one single bond to fulfill its octet, to different atoms (H only duet).

Double Covalent Bond

• When two atoms share two pairs of electrons, the result is called a double covalent bond (four electrons).

Triple Covalent Bond

• When two atoms share three pairs of electrons, the result is called a triple covalent bond (six electrons).

Covalent Bonding: Model Versus Reality

• The Lewis model implies that some combinations should be stable, whereas others should not, as stable combinations result in “octets.”

• Using these ideas from the Lewis theory allows us to predict the formulas of molecules of covalently bonded substances.

• Hydrogen and the halogens are all diatomic molecular elements, as predicted by Lewis theory.

• Oxygen generally forms either two single bonds or a double bond in its molecular compounds, as predicted by Lewis theory.

Covalent Bonding: Model Versus Reality

• Lewis model accounts for or predicts:

  • Why particular combinations of atoms form molecules and others do not.

  • That $H<em>2O$ should be stable, while $H</em>2O_3^+$ should not be, and that is in fact the case.

  • That $H<em>2O</em>3$ should be stable, and in fact it is.

  • Why covalent bonds are highly directional, in contrast to ionic bonds, which are nondirectional and hold together an entire array of ions.

  • Therefore, Lewis model predicts covalently bonded compounds will be found as individual molecules rather than an array like ionic compounds.

  • Compounds of nonmetals are made of individual molecule units.

Covalent Bonding: Model Versus Reality

Lewis model accounts for or predicts that:

• Molecular compounds have low melting points and boiling points (melting points generally < 300°C).

• Molecular compounds are found in all three states at room temperature.

Covalent Bonding: Model Versus Reality

Lewis model accounts for or predicts that:

• The hardness and brittleness of molecular compounds should vary depending on the strength of intermolecular attractive forces.

  • The kind and strength of the intermolecular attractions vary based on many factors.

  • Some molecular solids are brittle and hard, but many are soft and waxy.

• Neither molecular solids nor liquids should conduct electricity.

  • There are no charged particles around to allow the material to conduct.

  • Molecular compounds do not conduct electricity.

Electronegativity and Bond Polarity

• Covalent bonding between unlike atoms results in unequal sharing of the electrons.

  • One atom pulls the electrons in the bond closer to its side.

  • One end of the bond has a larger electron density than the other.

• The result is a polar covalent bond (bond polarity).

  • The end with the larger electron density gets a partial negative charge.

  • The end that is electron deficient gets a partial positive charge.

Bond Polarity

• Most bonds have some degree of sharing and some degree of ion formation to them.

• Bonds are classified as covalent if the amount of electron transfer is insufficient for the material to display the classic properties of ionic compounds.

• If the sharing is unequal enough to produce a dipole in the bond, the bond is classified as polar covalent.

Electronegativity

• The ability of an atom to attract bonding electrons to itself is called electronegativity.

• Electronegativity increases across the period (left to right) and decreases down the group (top to bottom).

  • Fluorine is the most electronegative element.

  • Francium is the least electronegative element.

  • Noble gas atoms are not assigned values.

  • This trend is opposite of the atomic size trend.

• The larger the difference in electronegativity, the more polar the bond, with the negative end toward the more electronegative atom.

Electronegativity Difference and Bond Type

• If the difference in electronegativity between bonded atoms is 0, the bond is pure covalent (equal sharing, also nonpolar).

• If the difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent.

• If the difference in electronegativity between bonded atoms is 0.4 to 1.9, the bond is polar covalent.

• If the difference in electronegativity between bonded atoms is larger than or equal to 2.0, the bond is ionic.

Bond Dipole Moments

• Dipole moment, $\mu$, is a measure of bond polarity.

  • A dipole is a bond in a molecule with a $\delta +$ and $\delta -$ end.

  • It is directly proportional to the size of the partial charges and directly proportional to the distance between them: $\mu = q \times r$

• Measured in Debyes (D).

• Generally, the more electrons two atoms share and the larger the atoms are, the larger the dipole moment.

Percent Ionic Character

• The percent ionic character is the percentage of a bond’s measured dipole moment compared to what it would be if the electrons were completely transferred.

• The percent ionic character indicates the degree to which the electron is transferred.

Writing Lewis Structures of Molecules

1.  Write the correct skeletal structure for the molecule.
    *   Hydrogen atoms are always terminal.
    *   The more electronegative atoms are placed in terminal positions.
2.  Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule.
3.  Distribute the electrons among the atoms, giving octets (or duets in the case of hydrogen) to as many atoms as possible.
4.  If any atoms lack an octet, form double or triple bonds as necessary to give them octets.

Resonance

• Lewis model localizes the electrons between the atoms that are bonding together.

• Extensions of Lewis model suggest that there is some degree of delocalization of the electrons; we call this concept resonance.

• Delocalization of charge helps to stabilize the molecule.

Resonance Structures

• When there is more than one Lewis structure for a molecule that differs only in the position of the electrons, they are called resonance structures.

• The actual molecule is a combination of the resonance forms—a resonance hybrid.

  • The molecule does not resonate between the different forms, though we often draw it that way.

• Look for multiple bonds or lone pairs.

Formal Charge

• Formal charge is a fictitious charge assigned to each atom in a Lewis structure that helps us to distinguish among competing Lewis structures.

• In a Lewis structure, calculate an atom’s formal charge, which indicates the charge it would have if all bonding electrons were shared equally between the bonded atom.

• The sum of all the formal charges in a molecule = 0. In an ion, the sum equals the charge.

• \text{FC} = # \text{valence } e^- - [# \text{nonbonding } e^- + \frac{1}{2} # \text{bonding } e^-]

Evaluating Resonance Structures

In general, when we calculate formal charges, these four rules apply:

1. The sum of all formal charges in a neutral molecule must be zero.

2. The sum of all formal charges in an ion must equal the charge of the ion.

3. Small (or zero) formal charges on individual atoms are better than large ones.

4. When formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom.

Exceptions to the Octet Rule

• Odd number electron species (e.g., NO):

  • Will have one unpaired electron

  • Free-radical

  • Very reactive

• Incomplete octets (B, Al).

• Expanded octets (elements with empty d orbitals can have more than eight electrons).

Bond Energies

• Bond energy, or bond enthalpy: the amount of energy it takes to break one mole of a bond in the gas phase.

  • Always positive, endothermic

  • Average for that bond in a large number of molecules

• Chemical reactions involve breaking bonds in reactant molecules and making new bonds to create the products.

• The overall enthalpy change for the reaction, $\Delta H_{rxn}^\circ$, can be calculated as the sum of the enthalpy changes associated with breaking the required bonds in the reactants and forming the required bonds in the products.

Trends in Bond Energies

• In general, the more electrons two atoms share, the stronger (higher bond energy) the covalent bond (must be comparing bonds between like atoms).

  • $\text{C} \equiv \text{C} > \text{C} = \text{C} > \text{C} - \text{C}$

  • $\text{C} \equiv \text{N} > \text{C} = \text{N} > \text{C} - \text{N}$

  • $\text{Br} - \text{F} > \text{Br} - \text{Cl} > \text{Br} - \text{Br}$

• In general, the shorter the covalent bond, the stronger the bond (must be comparing similar types of bonds).

  • Bonds get weaker down the column.

  • Bonds get stronger across the period.

Estimating the Enthalpy Change of a Reaction From Bond Energies

The actual bond energy depends on the surrounding atoms and other factors.

• We often use average bond energies to estimate $\Delta_{rxn}H$. (works best when all reactants and products are in the gas state).

• Bond breaking is endothermic, $\Delta H_\text{(breaking)} = +$.

• Bond forming is exothermic, $\Delta H_\text{(forming)} = -$.

Bond Lengths

• The distance between the nuclei of bonded atoms is called the bond length.

• Because the actual bond length depends on the other atoms around the bond, we often use the average bond length, averaged for similar bonds from many compounds.

Trends in Bond Lengths

• In general, the more electrons two atoms share, the shorter the covalent bond (must be comparing bonds between like atoms).

  • $\text{C} \equiv \text{C} < \text{C} = \text{C} < \text{C} - \text{C}$

  • $\text{C} \equiv \text{N} < \text{C} = \text{N} < \text{C} - \text{N}$

  • $\text{C} - \text{C} > \text{C} - \text{N} > \text{C} - \text{O}$

  • $\text{F} - \text{F} < \text{Cl} - \text{Cl} < \text{Br} - \text{Br}$

• Generally, bond length decreases from left to right across the period.

• Generally, bond length increases down the column.

• In general, as bonds get longer, they also get weaker.