© 2006 Thomson Higher Education Chapter 2 Polar Covalent Bonds; Acids and Bases

2.1 Polar Covalent Bonds: Electronegativity Chemical bonds Ionic bonds Ions held together by electrostatic attractions between unlike charges Bond in sodium chloride Sodium transfers an electron to chlorine to give Na + and Cl - Nonpolar Covalent bonds Two electrons are shared equally by the two bonding atoms Carbon-carbon bond in ethane Symmetrical electron distribution in the bond

Polar Covalent Bonds: Electronegativity Most bonds neither fully ionic or covalent Polar covalent bonds A covalent bond in which the electron distribution between atoms is unsymmetrical Bond polarity due to difference in electronegativity (EN)

Polar Covalent Bonds: Electronegativity Electronegativity (EN) The ability of an atom to attract shared electrons in a covalent bond Generally increases across the periodic table from left to right and from bottom to top

Polar Covalent Bonds: Electronegativity Bonds between atoms whose electronegativities differ by less than 0.5 are nonpolar covalent Bonds between atoms whose electronegativities differ by 0.5 to 2.0 are polar covalent Bonds between atoms whose electronegativities differ by more than 2.0 are largely ionic Carbon hydrogen bonds are nonpolar. Bonds between carbon (EN = 2.5) and more electronegative elements, such as oxygen (EN = 3.5) and nitrogen (EN = 3.0) are polar covalent bonds with the bonding electrons drawn towards the more electronegative atoms

Polar Covalent Bonds: Electronegativity Electrostatic potential maps Show calculated charge distributions Colors indicate electron-rich (red;  - ) and electron-poor (blue;  + ) regions Methanol, CH 3 OH, has a polar covalent C-O bond, and methyllithium has a polar covalent C-Li bond A crossed arrow is used to indicate direction of bond polarity Electrons are displaced in the direction of the arrow

Polar Covalent Bonds: Electronegativity Atoms ability to polarize a bond is known as the inductive effect Inductive effect The electron-attracting or electron-withdrawing effect transmitted through  bonds. Electronegative elements have an electron-withdrawing inductive effect Metals inductively donate electrons Reactive nonmetals inductively withdraw electrons Inductive effects play a major role in understanding chemical reactivity

2.2 Polar Covalent Bonds Molecules as a whole are often polar Molecular polarity results from the vector summation of all individual bond polarities and lone-pair contributions in the molecule Strongly polar substances soluble in polar solvents like water Dipole moment (  ) Magnitude of charge Q at either end of molecular dipole times distance r between charges  = Q  r, in debyes ( D ) 1 D =  10  30 coulomb meter (C m) A measure of the net polarity of a molecule Arises when the centers of mass of positive and negative charges within a molecule do not coincide If one positive and one negative charge were separated by just less than the length of an average covalent bond (100 pm), the dipole moment would be  = (1.60  10  19 C)(100  m)( 1 D /  C m) = 4.80 D

Polar Covalent Bonds

Lone-pair electrons on oxygen and nitrogen stick out into space away from positively charged nuclei giving rise to a considerable charge separation and contributing to the dipole moment Symmetrical structures of molecules cause the individual bond polarities and lone-pair contributions to exactly cancel

Worked Example 2.1 Predicting the Direction of a Dipole Moment Make a three-dimensional drawing of methylamine, CH 3 NH 2, and show the direction of its dipole moment (  = 1.31)

Worked Example 2.1 Predicting the Direction of a Dipole Moment Strategy Look for any lone-pair electrons Identify any atom with an electronegativity substantially different from that of carbon (usually O, N, F, Cl, or Br) Electron density will be displaced in the general direction of the electronegative atoms and the lone pairs

Worked Example 2.1 Predicting the Direction of a Dipole Moment Solution Methylamine has an electronegative nitrogen atom and a lone pair of electrons. The dipole moment thus points generally from –CH 3 toward the nitrogen

2.3 Formal Charges Formal charge The difference in the number of electrons owned by an atom in a molecule and by the same atom in its elemental state Assigned to specific atoms within a molecule Dimethyl sulfoxide CH 3 SOCH 3 Sulfur atom has three bonds rather than the usual two and has a formal positive charge Oxygen atom has one bond rather than the usual two and has a formal negative charge

Formal Charges

2.4 Resonance Two different ways to draw the acetate ion Double bond placement Neither structure correct by itself True structure is intermediate between the two Two structures are known as resonance forms

Resonance Resonance forms Individual line-bond structures of a molecule or ion that differ only in the placement of  and nonbonding valence electrons Indicated by “ ” Resonance forms contribute to a single, unchanging structure that is the resonance hybrid of the individual forms and exhibits the characteristics of all contributors

Resonance Benzene has two equivalent resonance forms The true structure of benzene is a hybrid of the two individual forms, and all six carbon-carbon bonds are equivalent Symmetrical distribution of electrons is evident in an electrostatic potential map of benzene

2.5 Rules for Resonance Forms Rules for drawing and interpreting resonance forms: Rule 1 – Individual resonance forms are imaginary, not real Real structure is a composite Rule 2 – Resonance forms differ only in the placement of their  or nonbonding electrons

Rules for Resonance Forms Rule 3 – Different resonance forms of a substrate do not have to be equivalent

Rules for Resonance Forms Rule 4 – Resonance forms obey normal rules of valency (follow the octet rule) Rule 5 – The resonance hybrid is more stable than any individual resonance form Resonance leads to stability

In general any three-atom grouping with a p orbital on each atom has two resonance forms: The atoms X,Y, and Z in the general structure might be C,N,O,P,or S The asterisk (*) on atom Z for the resonance form on the left might mean that the p orbital is: Vacant Contains a single electron Contains a lone pair of electrons 2.6 Drawing Resonance Forms

Drawing Resonance Forms Reaction of pentane-2,4-dione with a base H + is removed An anion is formed Resonance of the anion product:

Worked Example 2.2 Drawing Resonance Forms of an Anion Draw three resonance forms for the carbonate ion, CO 3 2-.

Worked Example 2.2 Drawing Resonance Forms of an Anion Strategy Look for one or more three-atom groupings that contain a multiple bond next to an atom with a p orbital. Then exchange the positions of the multiple bond and the electrons in the p orbital In the carbonate ion, each of the singly bonded oxygen atoms with its lone pairs and negative charge is next to the C=O double bond, giving the grouping O=C-O: -

Worked Example 2.2 Drawing Resonance Forms of an Anion Solution Exchanging the position of the double bond and an electron lone pair in each grouping generates three resonance structures:

Worked Example 2.3 Drawing Resonance Forms for a Radical Draw three resonance forms for the pentadienyl radical. A radical is a substance that contains a single, unpaired electron in one of its orbitals, denoted by a dot ( ).

Worked Example 2.3 Drawing Resonance Forms for a Radical Strategy Find the three-atom groupings that contain a multiple bond next to a p orbital.

Worked Example 2.3 Drawing Resonance Forms for a Radical Solution The unpaired electron is on a carbon next to a C=C bond, giving a typical three-atom grouping that has two resonance forms:

2.7 Acids and Bases: The Brønsted-Lowry Definition Two frequently used definitions of acidity The Brønsted-Lowry definition Lewis definition Brønsted-Lowry acid A substance that donates a hydrogen ion (proton; H + ) to a base Brønsted-Lowry base A substance that accepts a hydrogen ion (proton; H + ) from an acid

Acids and Bases: The Brønsted-Lowry Definition Conjugate acid The product that results from protonation of a Brønsted- Lowry base Conjugate base The anion that results from deprotonation of a Brønsted- Lowry acid In a general sense

Acids and Bases: The Brønsted-Lowry Definition Water can act either as an acid or as a base

2.8 Acid and Base Strength Acids differ in their ability to donate H + The exact strength of a given acid, HA, in water solution is described using the equilibrium constant K eq for the acid- dissociation equilibrium

Acid and Base Strength The concentration of water, [H 2 O], remains nearly constant at 55.5 M at 25 °C Can rewrite equilibrium expression using new quantity called the acidity constant K a Acidity constant K a A measure of acid strength in water For any weak acid HA, the acidity constant is given by the expression K a

Acid and Base Strength Equilibria for stronger acids favor the products (to the right) and thus have larger acidity constants Equilibria for weaker acids favor the reactants (to the left) and thus have smaller acidity constants Acid strengths are normally expressed using p K a values pK a The negative common logarithm of the K a pK a = -log K a Stronger acids (larger K a ) have smaller pK a Weaker acids (smaller K a ) have larger pK a

Acid and Base Strength

Self ionization of water [H 3 O + ][OH - ] = K w = ion product constant for water K w = 1.00 x [H 2 O] = 55.4 M at 25.0 °C

Acid and Base Strength Strong acid (Br Ø nsted-Lowry) One that loses H + easily Conjugate base holds on to the H + weakly (weak base) Strong acid has weak conjugate base Weak acid (Br Ø nsted-Lowry) One that loses H + with difficulty Conjugate base holds on to the H + strongly (strong base) Weak acid has strong conjugate base

2.9 Predicting Acid-Base Reactions from pK a Values An acid will donate a proton to the conjugate base of a weaker acid The conjugate base of a weaker acid will remove the proton from a stronger acid

Predicting Acid-Base Reactions from pK a Values Product of conjugate base must be weaker and less reactive than the starting acid Product of conjugate base must be weaker and less reactive than the starting base

Worked Example 2.4 Predicting Acid Strengths from pK a Values Water has pK a = 15.74, and acetylene has pK a = 25. Which is the stronger acid? Does hydroxide ion react with acetylene?

Worked Example 2.4 Predicting Acid Strengths from pK a Values Strategy In comparing two acids, the one with the lower pK a is stronger Thus water is a stronger acid than acetylene and gives up H + more easily

Worked Example 2.4 Predicting Acid Strengths from pK a Values Solution Because water is a stronger acid and gives up H + more easily than acetylene does: The HO - ion must have less affinity for H + than the HC≡C: - ion has The anion of acetylene is a stronger base than the hydroxide ion The reaction will not proceed as written

Worked Example 2.5 Calculating K a from pK a According to the data in Table 2.3, acetic acid has pK a = What is its K a ?

Worked Example 2.5 Calculating K a from pK a Strategy pK a is the negative logarithm of K a Use a calculator with an ANTILOG or INV LOG function Enter value of the pK a (4.76) Change the sign (-4.76) Find the antilog

Worked Example 2.5 Calculating K a from pK a Solution K a = 1.74 x 10 -5

2.10 Organic Acids and Organic Bases Most biological reactions involve organic acids and organic bases Organic acid Positively polarized hydrogen atom Two main kinds of organic acids 1. Contains a hydrogen atom bonded to an oxygen atom (O-H) 2. Contains a hydrogen atom bonded to a carbon atom next to a C=O double bond (O=C-C-H)

Organic Acids and Organic Bases Conjugate base Anion stabilized by having its negative charge on a highly electronegative atom Anion stabilized by resonance Methanol Acetic Acid Acetone

Organic Acids and Organic Bases Conjugate bases from methanol, acetic acid, and acetone The electronegative oxygen atoms stabilize the negative charge in all three

Organic Acids and Organic Bases Carboxylic acids Contain the –CO 2 H grouping Occur abundantly in all living organisms Involved in almost all metabolic pathways At cellular pH of 7.3 carboxylic acids are usually dissociated and exist as their carboxylate anions, –CO 2 -

Organic Acids and Organic Bases Organic bases Characterized by the presence of an atom with a lone pair of electrons that can bond to H + Nitrogen-containing compounds are common organic bases and are involved in almost all metabolic pathways Oxygen-containing compounds can act both as acids and as bases

2.11 Acids and Bases: The Lewis Definition The Lewis definition is broader than the Brønsted-Lowry definition Lewis acid A substance with a vacant low energy orbital that can accept an electron pair from a base All electrophiles are Lewis acids Lewis base A substance that donates an electron lone pair to an acid All nucleophiles are Lewis bases

Lewis Acids and the Curved Arrow Formalism To accept an electron pair a Lewis acid must have either: A vacant, low-energy orbital A polar bond to hydrogen so that it can donate H + Various metal cations, such as Mg 2+, are Lewis acids because they accept a pair of electrons when they form a bond to a base Acids and Bases: The Lewis Definition

Compounds of group 3A elements, such as BF 3 and AlCl 3 are Lewis acids Have unfilled valence orbitals and can accept electron pairs from Lewis bases Many transition metals, such as TiCl 4, FeCl 3, ZnCl 2, and SnCl 4 are Lewis acids

Acids and Bases: The Lewis Definition Curved arrow formalism Indicates the direction of electron flow from the base to the acid Always means that a pair of electrons moves from the atom at the tail of the arrow to the atom at the head of the arrow For the reaction of boron trifluoride, a Lewis acid, with dimethyl ether, a Lewis base. All movement of electrons from the Lewis base to the Lewis acid is indicated by a curved arrow

Acids and Bases: The Lewis Definition Further examples of Lewis acids

Acids and Bases: The Lewis Definition Lewis bases A compound with a pair of nonbonding electrons that it can use in bonding to a Lewis acid Definition of Lewis base similar to Brønsted-Lowry definition H 2 O acts as a Lewis base Has two nonbonding electrons on oxygen

Acids and Bases: The Lewis Definition Most oxygen- and nitrogen- containing organic compounds are Lewis bases They have lone pair electrons

Acids and Bases: The Lewis Definition Some compounds can act as both acids and bases Some compounds have more than one atom with a lone pair of electrons Reaction normally occurs only once in such instances The more stable of the two possible protonation products is formed Occurs with carboxylic acids, esters, and amides

Worked Example 2.6 Using Curves Arrows to Show Electron Flow Using curved arrows, show how acetaldehyde, CH 3 CHO, can act as a Lewis base.

Worked Example 2.6 Using Curves Arrows to Show Electron Flow Strategy A Lewis base donates an electron pair to a Lewis acid Locate the electron lone pairs on acetaldehyde Use a curved arrow to show the movement of a pair toward the H atom of the acid

Worked Example 2.6 Using Curves Arrows to Show Electron Flow Solution

2.12 Noncovalent Interactions between Molecules Noncovalent interactions Also called intermolecular forces or van der Waals forces Dipole-dipole forces, dispersion forces, and hydrogen bonds Dipole-Dipole forces Occur between polar molecules as a result of electrostatic interactions among dipoles Forces are either attractive or repulsive Attractive Repulsive Attractive geometry is lower in energy and therefore predominates

Noncovalent Interactions between Molecules Dispersion forces Attractive dispersion forces in nonpolar molecules are caused by temporary dipoles One side of the molecule may have a slight excess of electrons relative to the opposite side, giving the molecule a temporary dipole Temporary dipole in one molecule causes a nearby molecule to adopt a temporarily opposite dipole resulting in a small attraction between the two molecules Arise because the electron distribution within molecules is constantly changing

Noncovalent Interactions between Molecules Hydrogen Bond A weak attraction between a hydrogen atom bonded to an electronegative O or N and an electron lone pair on another O or N atom Strong dipole-dipole interaction involving polarized O-H and N-H bonds Important noncovalent interaction in biological molecules

Noncovalent Interactions between Molecules Hydrogen bond Causes water to be a liquid rather than a gas at room temperature Holds enzymes in the shapes necessary for catalyzing biological reactions Causes strands of deoxyribonucleic acid (DNA) to pair up and coil into a double helix Hydrophilic (water-loving) Dissolves in water Table sugar Has ionic charges, polar –OH groups, in its structure Hydrophobic (water-fearing) Does not dissolve in water Vegetable oil Does not have groups that form hydrogen bonds