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CHAPTER 2: POLAR COVALENT BONDS; ACIDS AND BASES
2.1 Polar Covalent Bonds: Electronegativity ● Up to this point, we’ve treated chemical bonds as either ionic or covalent. REVIEW ON WHAT IS COVALENT AND IONIC BOND: ● COVALENT BONDS - is between 2 nonmetals. ● IONIC BONDS - is between a metal and a nonmetal. ● The bond in sodium chloride , for instance, is ionic. Sodium transfers an electron to chlorine to produce Na+^ and C_^ ions, which are held together in the solid by electrostatic attractions between unlike charges. ● The C—C bond in ethane, however, is covalent. The two bonding electrons are shared equally by the two equivalent carbon atoms, resulting in a symmetrical electron distribution in the bond. ● Most bonds, however , are neither fully ionic nor fully covalent but are somewhere between the two extremes. POLAR COVALENT BONDS ● Bonds that are somewhere in between the two extremes. ● meaning that the bonding electrons are attracted more strongly by one atom than the other so that the electron distribution between atoms is not symmetrical. ● The continuum in bonding from covalent to ionic is a result of an unequal distribution of bonding electrons between atoms. The symbol δ (lowercase Greek delta) means partial charge, either partial positive (δ+) for the electron-poor atom or partial negative (δ-) for the electron-rich atom. BOND POLARITY ● is due to differences in electronegativity (EN) , the intrinsic ability of an atom to attract the shared electrons in a covalent bond. NOTE: ● fluorine has the most electronegative (EN = 4.0) ● and cesium the least (EN = 0.7) ● Metals on the left side of the periodic table attract electrons weakly and have lower electronegativities, ● while oxygen, nitrogen, and halogens on the right side of the periodic table attract electrons strongly and have higher electronegativities. ● Carbon , the most important element in organic compounds, has an electronegativity value of 2.5. ELECTRONEGATIVITY VALUES AND TRENDS ● increases from left to right across the periodic table ● decreases from top to bottom. ● Elements in red are the most electronegative , ● those in yellow are medium ● those in green are the least electronegative. Bonds between atoms whose electronegativities differ by Type of Bonds Less than 0.5 Nonpolar covalent 0.5 to 2.0 Polar covalent More than 2.0 Largely ionic FOR EXAMPLE:
- Carbon–hydrogen bonds are relatively nonpolar because carbon (EN=2.5) and hydrogen (EN=2.1) have similar electronegativities.
- Bonds between carbon and more electronegative elements such as oxygen (EN= 3.5) and nitrogen (EN=3.0), by contrast, are polarized so that the bonding electrons are drawn away from carbon toward the electronegative atom. OxygenEN ➖ CarbonEN = 3.5 - 2.5 = 1. NitrogenEN ➖ CarbonEN = 3.0 - 2.5 = 0. This leaves carbon with a partial positive charge, denoted by δ+ , and the electronegative atom with a partial negative charge , δ-. In methanol, oxygen carries a partial negative charge and is colored red, while the carbon and hydrogen atoms carry partial positive charges and are colored blue-green.
In methyllithium, lithium carries a partial positive charge (blue), while carbon and the hydrogen atoms carry partial negative charges (red). NOTE:
- The computer generated representations, called electrostatic potential maps, use color to show calculated charge distributions, ranging from ● red (electron-rich; δ-) to ● blue (electron-poor; δ+)
- a crossed arrow is used to indicate the direction of bond polarity. By convention, electrons are displaced in the direction of the arrow. The tail of the arrow (which looks like a plus sign) is electron-poor (δ+) , and the head of the arrow is electron-rich (δ-). INDUCTIVE EFFECT ● When speaking of an atom’s ability to polarize a bond
● is simply the shifting of electrons in a σ bond in
response to the electronegativity of nearby atoms. ○ Metals, such as lithium and magnesium , inductively donate electrons ○ reactive nonmetals, such as oxygen and nitrogen , inductively withdraw electrons. ● Inductive effects play a major role in understanding chemical reactivity. 2.2 Polar Covalent Bonds: Dipole Moments Just as individual bonds are often polar, molecules as a whole are often polar as well. Molecular polarity results from the vector summation of all individual bond polarities and lone-pair contributions in the molecule. As a practical matter,
- strongly polar substances are often soluble in polar solvents like water ,
- whereas less polar substances are insoluble in water. Net molecular polarity ● is measured by a quantity called the dipole moment and can be thought of in the following way: assume that there is a center of mass of all positive charges (nuclei) in a molecule and a center of mass of all negative charges (electrons). If these two centers don’t coincide, then the molecule has a net polarity. The dipole moment, 𝜇 (Greek mu) ● is defined as the magnitude of the charge Q at either end of the molecular dipole times the distance r between the charges, 𝜇 = Q x r ● Dipole moments are expressed in debyes (D), where 1D = 3.336 x 10-30^ coulomb meters (C · m) in SI units. ● For example, the unit charge on an electron is 1. x10-19^ C. Thus, if one positive charge and one negative charge are separated by 100 pm (a bit less than the length of a typical covalent bond), the dipole moment is 1.60 x 10-29^ C · m, or 4.80 D. Dipole moments for some common substances are given in Table 2-1. Of the compounds shown in the table, sodium chloride has the largest dipole moment (9.00 D) because it is ionic. Even small molecules like water (H 2 O), methanol (CH 3 OH), and ammonia (NH 3 ), have substantial dipole moments , however, both because they contain strongly electronegative atoms (oxygen and nitrogen) and because all three molecules have lone-pair electrons. The lone-pair electrons on oxygen and nitrogen stick out into space away from the positively charged nuclei, giving rise to a considerable charge separation and making a large contribution to the dipole moment.
2.3 Formal Charges A formal charge (FC) is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. ➔ Closely related to bond polarity and dipole moment ➔ assigning formal charges to specific atoms within a molecule, particularly atoms that have an “abnormal” number of bonds. ➔ A typical covalent bond is formed when each atom donates one electron. ➔ Although the bonding electrons are shared by both atoms, each atom can still be considered to “own” one electron for bookkeeping purposes. Example: Methane The carbon atom owns one electron in each of the four C-H bonds. Because a neutral, isolated carbon atom has four valence electrons, and because the carbon atom in methane still owns four, the methane carbon atom is neutral and has no formal charge Example: Ammonia It has three covalent N-H bonds and two nonbonding electrons (a lone pair). Atomic nitrogen has five valence electrons, and the ammonia nitrogen also has five—one in each of three shared N]H bonds plus two in the lone pair. Thus, the nitrogen atom in ammonia has no formal charge Example: dimethyl sulfoxide (CH3SOCH3) The sulfur atom in dimethyl sulfoxide has three bonds rather than the usual two and has a formal positive charge. The oxygen atom, by contrast, has one bond rather than the usual two and has a formal negative charge. Note that an electrostatic potential map of dimethyl sulfoxide shows the oxygen as negative (red) and the sulfur as relatively positive (blue), in accordance with the formal charges Explanation : Atomic sulfur has six valence electrons, but the dimethyl sulfoxide sulfur owns only five—one in each of the two S]C single bonds, one in the S]O single bond, and two in a lone pair. Thus, the sulfur atom has formally lost an electron and therefore has a positive charge. A similar calculation for the oxygen atom shows that it has formally gained an electron and has a negative charge. Atomic oxygen has six valence electrons, but the oxygen in dimethyl sulfoxide has seven—one in the O]S bond and two in each of three lone pairs. formal charge is calculated: A summary of commonly encountered formal charges and the bonding situations in which they occur is given in Table 2-2.
Problem 2- Calculate formal charges for the non hydrogen atoms in the following molecules: Formula: Solution: Problem 2- Organic phosphate groups occur commonly in biological molecules. Calculate formal charges on the four O atoms in the methyl phosphate dianion Solution:
● Resonance forms obey normal rules of valency.
- octet rule still applies to second-row, main-group atoms ● The resonance hybrid is more stable than any individual resonance form.
- the larger the number of resonance forms, the more stable a substance is, because its electrons are spread out over a larger part of the molecule and are closer to more nuclei. 2.6 Drawing Resonance Forms Useful technique for drawing resonance forms
- any three-atom
- grouping with a p orbital on each atom has two resonance forms: The two resonance forms differ simply by an exchange in position of both the multiple bond and the asterisk from one end of the three-atom grouping to the other.
- The atoms X, Y, and Z in the general structure might be C, N, O, P, S, or others, and the asterisk (*) might mean that the p orbital on atom Z is vacant, that it contains a single electron, or that it contains a lone pair of electrons. The anion produced when H⁺ is removed from 2,4-pentanedione by reaction with a base. How many resonance structures does the resultant anion have? ● The 2,4-pentanedione anion has a lone pair of electrons and a formal negative charge on the central carbon atom, next to a C=O bond on the left. The O=C-C:⁻ grouping is a typical one for which two resonance structures can be drawn. ● Just as there is a C=O bond to the left of the lone pair, there is a second C=O bond to the right. ● Thus, we can draw a total of three resonance structures for the 2,4-pentanedione anion. Drawing Resonance Forms for an Anion
STRATEGY
● Look for 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 singly bonded oxygen atom with three lone pairs and a negative charge is adjacent to the C=O double bond, giving the grouping SOLUTION ● Exchanging the position of the double bond and an electron lone pair in each grouping generates three resonance structures. Drawing Resonance Forms for a Radical STRATEGY ● Find the three-atom groupings that contain a multiple bond next to an atom with a p orbital. SOLUTION ● The unpaired electron is on a carbon atom next to a C=C bond, giving a typical three-atom grouping that has two resonance forms.
In Second Resonance Form
- the unpaired electron is next to another double bond, giving another three-atom grouping and leading to another resonance form Thus, the three resonance forms for the pentadienyl radical are: 2.7 ACIDS AND BASES: THE BRONSTED-LOWRY DEFINITION Brønsted-Lowry Acid ● a substance that donates a hydrogen ion, H+ Brønsted-Lowry Base ● a substance that accepts a hydrogen ion Gaseous Hydrogen Chloride Dissolves in Water ● HCl molecule acts as an acid and donates a proton ● water molecule acts as a base and accepts the proton, yielding chloride ion (Cl-) and hydronium ion (H 3 O+) ● acid-base reactions are reversible so a double, forward-and-backward arrows should be used ● chloride ion (the product after the acid, HCl, loses a proton) is called the conjugate base of the acid ● hydronium ion (the product after the base, H 2 O, gains a proton) is called the conjugate acid of the base General Reaction Examples: ● water can act either as an acid or a base, depending on the circumstances ● water becoms a base in its reaction with HCl since it accepts a proton to give to the hydronium ion, H 3 O+ ● water is an acid in its reaction with ammonia, NH 3 , since it donates a proton to give ammonium ion (NH 4 +) and hydroxide ion (HO-) 2.8 ACID AND BASE STRENGTH Strong Acids ● react almost completely with water Weak Acids ● react only slightly with water, do not dissociate completely Acidity Constant (Ka) ● is the exact strength of a given acid HA in water solution ● the concentration of solvent is ignored in the equilibrium expression and that brackets [ ] around a substance refer to the concentration of the enclosed species in moles per liter ● stronger acids have their equilibria toward the right and have larger acidity constants ● weaker acids have their equilibria toward the left and have smaller acidity constants ● the range of Ka values run from 10^15 for the strongest acids to 10-60^ for the weakest ● common inorganic acids (i.e. H 2 SO 4 , HNO 3 , HCl) have Ka’s ranging from 10^2 to 10^9 ● organic acids generally have Ka in the range of 10-5^ to 10 - pKa ● acid strengths are expressed using pKa ● is the negative common logarith of Ka: pKa = – log Ka ● a stronger acid (larger Ka) has a smaller pKa ● a weaker acid (smaller Ka) has a larger pKa
Methanol contains an O–H bond and is a weak acid , while acetic acid also contains an O–H bond and is a somewhat stronger acid. ACIDITY ● due to the fact that the conjugate base resulting from loss of H+^ atoms is stabilized by having its negative charge on a strongly electronegative oxygen atom; ● also because the conjugate base of acetic acid is stabilized by resonance. Anion is stabilized by having negative charge on a highly electronegative atom. Anion is stabilized both by having negative charge on a highly electronegative atom and by resonance. ● The acidity of acetone and other compounds with C=O bonds is due to the fact that the conjugate base resulting from loss of H+^ is stabilized by resonance. ● In addition, one of the resonance forms stabilizes the negative charge by placing it on an electronegative oxygen atom. Anion is stabilized both by resonance and by having negative charge on a highly electronegative atom. Electrostatic Potential Maps of the Conjugate Bases of (a) Methanol, (b) Acetic Acid, and (c) Acetone As you might expect, all three show a substantial amount of negative charge (red) on oxygen. The electronegative oxygen atoms stabilize the negative charge in all three. CARBOXYLIC ACIDS ● compounds which contain the –CO 2 H grouping ● occur abundantly in all living organisms and are involved in almost all metabolic pathways ● ex. acetic acid, pyruvic acid, and citric acid ● usually dissociated and exist as their carboxylate anions, –CO 2 –^ at the typical pH of 7.3 found within cells ORGANIC BASES ● characterized by the presence of an atom (reddish in electrostatic potential maps) with a lone pair of electrons that can bond to H+ METHYLAMINE ● nitrogen-containing compounds ● most common organic bases ● involved in almost all metabolic pathways OXYGEN-CONTAINING COMPOUNDS ● can also act as bases when reacting with a sufficiently strong acid ● some can act both as acids and as bases depending on the circumstances, just as water can ● Methanol and Acetone – act as acids when they donate a proton but as bases when their oxygen atom accepts a proton AMINO ACIDS ● so-named because they are both amines (–NH 2 ) and carboxylic acids (–CO 2 H–) ● building blocks from which proteins in all living organisms are made – twenty different amino acids (ex. alanine , etc.) – exist primarily in a doubly charged form called a zwitterion rather than in the uncharged form ● Zwitterion – arises because amino acids have both acidic and basic sites within the same molecule and therefore undergo an internal acid–base reaction
2.11 (^) Acids and Bases: The Lewis Definition A Lewis acid is a substance that accepts an electron pair , and a Lewis base is a substance that donates an electron pair. The donated electron pair is shared between the acid and the base in a covalent bond.
Lewis Acids and the Curved Arrow Formalism
● Lewis acid must have either a vacant, low-energy orbital or a polar bond to hydrogen so that it can donate H 1 (which has an empty 1s orbital) ● the Lewis definition of acidity includes many species in addition to H 1 ● various metal cations, such as Mg2+, are Lewis acids because they accept a pair of electrons when they form a bond to a base ● compounds of group 3A elements, such as BF 3 and AlCl 3 , are Lewis acids because they have unfilled valence orbitals and can accept electron pairs from Lewis bases ● many transition-metal compounds, such as TiCl 4 , FeCl 3 , ZnCl 2 , and SnCl 4 , are Lewis acids Dimethyl ether, the Lewis base, donates an electron pair to a vacant valence orbital of the boron atom in BF3, a Lewis acid. The direction of electron-pair flow from base to acid is shown using curved arrows. A curved arrow 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.
Lewis Bases
● a compound with a pair of nonbonding electrons that it can use to bond to a Lewis acid ● H 2 O, with its two pairs of nonbonding electrons on oxygen, acts as a Lewis base by donating an electron pair to an H+^ in forming the hydronium ion, H 3 O^1. ● most oxygen- and nitrogen-containing organic compounds can act as Lewis bases because they have lone pairs of electrons ● A divalent oxygen compound has two lone pairs of electrons, and a trivalent nitrogen compound has one lone pair. ● some compounds can act as both acids and bases , just as water can. Alcohols and carboxylic acids, for instance, act as acids when they donate an H+^ but as bases when their oxygen atom accepts an H+.
● Hydrogen bonding has enormous consequences for living organisms. ● Hydrogen bonds cause water to be a liquid rather than a gas at ordinary temperatures, they hold enzymes in the shapes necessary for catalyzing biological reactions, and they cause strands of deoxyribonucleic acid (DNA) to pair up and coil into the double helix that stores genetic information. Hydrophilic
- meaning “ water-loving ,” to describe a substance that is strongly attracted to water
- hydrophilic substances, such as table sugar, usually have a number of ionic charges or polar -OH groups in their structure so they can form hydrogen bonds, Hydrophobic
- meaning “ water-fearing ,” to describe a substance that is not strongly attracted to water
- hydrophobic substances, such as vegetable oil, do not have groups that form hydrogen bonds, so their attraction to water is limited to weak dispersion forces.
