CHEM 1120 EXAM 3

Ch. 16 & 17

General Properties of Acids

• Sour taste

• ability to dissolve many metals

• ability to neutralize bases

• change blue litmus paper to red

General properties of bases

• Taste bitter (alkaloid is a plant product that is alaline in nature and often poisonous)

• feels slippery to the touch

• ability to neutralize acids

• change red litmus paper to blue

Hydronium Ion: H₃O⁺

• The H⁺ ions (protons) produced by the acid are very reactive and cant exist in water.

• They react with water molecules to produce complex ions with the formula H₃O⁺ (interchangeable with H⁺

Arrhenius defintion

• simplest and most restrictive

• Acids are substances that when dissolved in water produce a hydronium (H⁺)

• Bases are substances that when dissolved in water produce a hydroxide ion (OH^_)

Bronsted-Lowry Definition

• based on reactions in water

• Acids are substances that when dissolved in water donate protons (H⁺)

• Bases are substances that accept protons

Lewis Defintion

• Most expansive and based on electron donors and acceptors

• Acids are substances that accept for need an electron pair

• Bases are substances that donate an electron pair to another substance

Arrhenius Acid-Base Reactions

• H⁺ ions from acid combine with OH^- ions from the base to make H₂O

• H—OH

• In a neutralization reaction, the cation from the base (Na+ from NaOh) combines with the anion (Cl- from HCl) from the acid to make a salt compound

• acid + base —> salt + water

Problems with Arrhenius Theory

• Doesnt expalin why molecular substances (NH3) dissolve in water to form basic solutions

• Doesnt explain how some ionic compounds (Na2CO3) dissolve in water to form basic solutions

• Doesnt explain why molecular susbatnces (CO2) dissolve in water to form acidic solutions

• What if we arent in water?

Bronsted-Lowry Acid-Base Theory

• Defines acids and bases based on what happens in the chemical reaction

• Any reaction that transfers H+ (proton) from one molecule to another is classified as an acid-base reaction, regardless of whether is occurs in aqueous solution or if there are OH- ions present

• All acid-base reactions that fit the Arrhenius defintion also fit bronsted-lowry

Bronsted-Lowry Theory: Acid Base reactions

• Acid is a H+ ion donor

• Base is H+ ion accepter (must have lone pair of electrons)

• H—A + B: —> :A^- + H—B⁺

(acid) (base) (conjugate base) (conjugate acid)

Conjugate Acid-Base Pairs

• The original base becomes a conjugate acid in the reverse reaction

• The original acid becomes a conjugate base in the reverse reaction

• Each reactant and the product it becomes is called a conjugate pair

Acid strength and molecular structure of acids

• Binary acids (H—Y) have acidic hydrogens attached to a nonmetal atom (HCl and HF)

• The more polarized the bond the more acidic the bond

• The stronger the H—X bond the weaker the acid

• Binary acid strength increase to the right across a period

• Binary acid strength increases down a column

Oxyacids

• Oxyacids have acidic hydrogen atoms attached to an oxygen atom

• The more electronegative the Y atom, the stronger the oxyacid (HClO>HBrO>HIO)

• Acidity of oxyacids decreases down a group

• Acidity of oxyacids inceases to the right (larger oxidation number = stronger acid)

• The more oxygen atoms attached to Y (nonmetal) the stronger the oxyacid as it further weakens and polarizes the H—O bond (HClO₃>HClO₂)

Strong Acid/Base vs Weak

• A strong acid is a strong electrolyte (complete ionization in water)

• A weak acid is a weak electrolyte (small percentage or partial ionization of the acid molecule in water) and equilibrium is established

• A strong base is a strong electrolyte (complete ionization in water)

• A weak base is a weak electrolyte (small percentage or partial ionization of the base molecules in water) and equilibrium is established

General trends in acidity

• The stronger an acid is at donating a proton, the weaker the conjugate base is at accepting a proton

• Higher oxidation number = stronger oxyacid

• Cation stronger acid than neutral molecule, neutral molecule stronger than anion (H₃O+ > H₂O > OH-)

Strong acids

• Strong acids completely ionizes in solution (strong electrolyte)

• [H₃O⁺] = [H+]

• 6 strong acids: HCl, HBr, HI, HNO₃, NClO₄, H₂SO₄

Weak Acids

• Only partially ionize and acid molecules stay largely intact

• [H₃O⁺] « [weak acid]

• Ex: acetic acid, carbonic acid, formic acid

Acid Ionization Constant K_a

• Acid strength is measured by the size of the equilibrium constant when it reacts with H₂O

• HA (aq) ⇋ H⁺ (aq) + A^- (aq)

• The equilibrium constant is the acid ionization constant K_a

• The larger the ionization constant, the stronger the acid or base

Autoionization of water

• Water is amphoteric, it can act as a acid or base

• About two of every one billion water molecules forms ions through autoionization

• All aqueous solutions contain both H+ and OH-

• The concentrations of these are equqal in water

• [H+][OH-] = 1 × 10^-7 M at 25 C

Ion product of water K_w

• The product of [H+] and [OH-] is always the same at room temp = 1.00 × 10^-14 at 25 C

• K_w is the ion product of water or the dissociation constant of water

• [H+] and [OH-] are inversley proportional

Measuring acidity: pH

• The acidity or basicity of a solution is often expressed as pH

• pH = -log[H+]

• pH water = -log[10^-7] = 7

• [H=] = 10^-pH

• pH < 7 is acidic

• pH > 7 is basic

• pH = 7 is neutral

What does pH imply?

• Lower pH = more acidic, higher = more basic

• One pH unit is a factor of 10 difference in acidity

• pH range from 0 to 14

• pH 0 is [H+] = 1 M

• pH 14 is [OH-] = 1 M

• pH has as many decimal places as sig figs the number used to calucltae it had, normally 2

Acid and Basic Solutions

• All aqueous solutions ahve both H+ and OH- ions

• Neutral solutions have equal [H+] and [OH-], each at 1.00 × 10^-7

• Acidic solutions: [H+] > 1.00 × 10^-7

• Basic solutions: [OH-] > 1.00 × 10^-7

Measuring acidity: pH and pOH

pH = -log[H+] = 7

[H+] = 10^-pH

pOH = -log[OH-]

pH + pOH = 14 at 25 C

pK_a

Another way of expressing the strength of an acid or base is through pK

Acid: pK_a = -log(K_a), K_a = 10^-pK_a

Stronger acid = smaller pK_a

Larger K_a = smaller pK_a

pK_b

Another way of expressing the strength of an acid or base is through pK

Base: pK_a = -log(K_b), K_b = 10^-pK_b

The stronger the base, the smaller the pK_b

Larger K_b = smaller pK_b

Percent ionization

The percentage of acid molecules that ionize when dissolved in water

( molarity of ionized acid / initial molarity of weak acid ) x 100

[H+]_equili / [HA]_init x 100 = % ionization

Higher percent ionization = stronger acid

Finding pH of a strong acid

For a monoprotic strong acid, the acid concentration equals the hydronium concentration [H+] = [H_acid]

For a diprotic acid (H₂SO₄) the first ionization is most significant, and second can generally be ignored

Relationship between [H+] and [HA]_inital

• Increasing the initial concentration of an acid results in increased [H+] at equilibrium

• Increasing the initial concentration of an acid results in decreased percent ionization

• The increase in [H+] is slower than increase in acid concentration

Finding the pH of mixtures of acids

• Generally, you can ignore the contribution of the weaker acid to the [H+]_equili

• For a mixture of a strong acid with a weak acid, the complete ionization of the strong acid provides more than enough to shift the weak acid equilibrium to the left so far that the weak acids added [H+] is negilible

Strong bases: K_b > 1

Hydroxide compounds are strong bases

Ex: KOH, NaOH, Ca(OH)₂

For ionic bases (soluble), practically all units are dissociated into OH- ions or accept H+ ions, they are strong electrolytes, and soluble multi-hydroxide bases completely dissociate in water

Weak bases: K_b < 1

In weak bases, only a small fraction of molecules accept H+ ions, weak electrolytes, and most of the weak base molecules dont accept H+ ions from water

Ex: NH₃, Amines, Carbonates

Base ionization constant, K_b

Base strength is measured by the size of the equilibrium constant when it reacts with H₂O

B(aq) + H₂O(l) ⇋ BH+ (aq) + OH- (aq)

The larger the K_b = stronger base

Acid-base properties of ions and salts

• Salts are ionic compounds formed by a neutralization reaction

• Salts that contain the cation of a strong base and the anion that is the conjugate base of a weak acid are basic

• Salts that contain cations that are the conjugate acid of a weak base and an anion of a strong acid are acidic

Anions as weak bases

• Every anion can be thought as the conjugate base of an acid, therefore, every anion can potentially be a base

• The stronger the acid, the weaker the conjuagte base

• An anion that is the conjugate base of a strong acid is pH neutral

• An anion that is the conjugate base of a weak acid is basic

Relationship between K_a of an acid and K_b of its conjugate base

K_a x K_b = [H+][OH-] = K_w (1.00×10^-14)

Cations as weak acids

• Some cations can be thought of as the conjugate acid of a weak base, so some may be acidic

• Others are the counterions of a strong base

• The stronger the base, the weaker the conjuagte acid

• A cation that is the counterion of a strong base is pH neutral

• A cation that is the conjugate acid of a weak base is acidic

Metal cations as weak acids

• Cations of small, highly charged metals are weakly acidic

• Alkali metal cations and alkaline earth metal cations are pH neutral

• Cations are hydrated