Chem 2 Exam 4

Arrhenius Acids

Produces H⁺ ions in aq solutions

Arrhenius Bases

Produces OH^- ions in aq solutions

Arrhenius acid and base

Simplest, but most restrictive definition

when Arrhenius acids and bases react, they neutralize each other and make water and salt

It is restricted to aq solutions

Caveats:

-many substances that don’t produce H⁺ ions are still acidic, Ex. CO₂

-many substances that don’t produce OH^- ions are still basic, Ex. NH₃, Na₂CO_3, Na₂O

Brønsted-Lowry acid and base 

Based on the reactions in water 

Acids and bases will always occur together in pairs , called conjugate acid-base pairs

The original base becomes a conjugate acid

The original acid becomes a conjugate base

Brønsted-Lowry acid

Donates a proton (hydrogen ion, or H⁺) to another substance 

Brønsted-Lowry base

Accepts a proton (hydrogen ion, or H⁺) from another substance

Lewis acid and base

Based on electron donors and acceptors

most expansive definition

Lewis acids and bases must occur in pairs. The Lewis acid accepts the lone pair donated by the Lewis base

All Arrhenius and Brønsted-lowery acids and bases are also considered Lewis acids and bases

Most expansive and can be used in non-aqueous solutions

Lewis acid

Accepts an electron pair

In order to accept the electron pair it must have a hydrogen atom to kick off, incomplete octet, able to expand its octet, or a higher order bond (e.g. double bond)

Ex. HCl, BF₃, AlF₃, CO₂, CU²⁺

Lewis Base

Donates an electron pair

Must have a lone pair to donate

Ex. OH^-, CN^-, CH₃COO^-, NH₃, H₂O, CO

Amphoterism

A species that can act as both an acid and a base

aka: amphiprotic species, ex. water

When an acid donates an H⁺, water will act as a base and accept the H⁺ to form H₃O⁺ (hydronium). Conversely, when a base accepts a H⁺ from water, it would result in an OH^- (hydroxide)

Autoionization of water

Since water can act as both an acid and base it can react with itself and this is called autoionization

K_w = [H₃O⁺][OH^-] = 1.0 × 10^-14

-For aq solutions at 25˚C , temp dependent

-K_w = equilibrium constant for autoionization, ion product constant of water

Acidity and Basicity

Acidic: [H₃O⁺] > [OH^-]

Basic: [H₃O⁺] < [OH^-]

Neutral: [H₃O⁺] = [OH^-]

pH scale for hydronium

pH = -log[H₃O⁺] or

[H₃O⁺] = 10^-pH

Sig figs

For a log the # of sig figs inside the log becomes the number of decimals in the result

When undoing a log the number of decimals becomes the number of sig figs in the result

pH scale and pOH scale

A change in [H⁺] by a factor of 10 causes the pH to change by 1

The scales for pH and pOH are flipped

Ex, pH > 7 = basic but pOH > 7 = acidic

pOH scale for hydroxide

pOH = -log[OH^-] or

[OH^-] = 10^-pOH

Converting between pH and pOH

pH + pOH = pK_w, so

pH + pOH = 14

*only at 25˚C, K_w = 1.0 × 10^-14

Summary of pH 

Solution made from an acid is always acidic 

Solution made from a base is always basic 

Solution of two or more acids would be led by the stronger acid 

Solution of two or more bases would be led by the stronger base 

Acid strength

Strong acids completely ionize in aq solutions

-strong electrolytes basically completely dissociate, conjugate base is very weak

Weak acids only partially ionize in aq solutions

-weak electrolytes and don’t fully dissociate, conjugate base will be much stronger

Types of acids

If not a strong acid then it is a weak acid

Strong acids: HCl, HBr, HI, HNO₃, HClO₃, HClO₄, H₂SO₄

Types of acids cont…

Most Brønsted-Lowry acids are monoprotic, meaning they only donate one proton

Ex. HCl, HNO₃, etc.

Others can be polyprotic, have multiple protons to donate

Not an indicator of acid strength just percent of ionization in water

Types of bases

If not a strong base then it is a weak base

Strong bases: LiOH, NaOH, KOH, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂

Stronger bases have weaker conjugate acids, weaker bases have stronger conjugate acids

Acid equilibrium constant

aka acid dissociation constant

The larger K_a the stronger the acid

K_a > 1 = strong acid

K_a < 1 = weak acid

Base equilibrium constant

aka base dissociation constant

The larger the K_b the stronger the base and vice versa

K_b > 1 = strong base

K_b < 1 = weak base

pK_a and pK_b

The smaller the pK_a and larger K_a, the stronger the acid

The smaller the pK_b and the larger the K_b, the stronger the base

pK_a = -log(K_a)

pK_b = -log(K_b)

Relationship between K_a and K_b

K_a • K_b = K_w

and

pK_a + pK_b = pK_w

so in standard conditions of 25˚C, pK_w = 14

pK_a + pK_b = 14

Strong acid solutions

The concentration of a strong acid is also equal to [H₃O⁺]

Can use this to find pH, pOH, and [OH^-]

As long as the concentration of the pure acid is sufficiently large (>1.0 × 10^-6), the autoionization of water can be ignored

Weak acid solutions 

Weak acids don’t dissociate completely so we can only do one dissociation at a time

RICE table to determine concentration of hydronium (and everything else)

Strong base solutions

All hydroxide ions will completely dissociate in aq solutions

Monobasic strong bases will only produce one [OH^-] so [OH^-] = base

Dibasic strong bases will produce 2 [OH^-] so [OH^-] = 2 x base

Also can neglect the autoionization of water unless solution is very dilute (<10^-6)

Weak base solutions 

Weak bases do not completely dissociate in water 

Unlike strong bases, weak bases produce OH^- by accepting a proton from water

To get to pH remember to switch from hydroxide to hydronium or from pOH to pH 

Percent dissociation

Degree of dissociation 'α', how much a weak acid/base has dissociated

Weak acid:

α = (conc. ionized acid)/ (initial conc. of weak acid) = [H₃O⁺]_equil/[HA]_init

Weak base:

α = (conc. ionized base)/ (initial conc. of weak base) = [OH^-]_equil/[B]_init

Percent dissociation: α x 100%

Percent dissociation cont…

Trends for acids and bases: (ex. in acids but same for bases) 

The equilibrium conc. of H₃O⁺ of weak acid increases with increasing initial conc. of the acid

-think le chatelier’s principle 

The percent dissociation of a weak acid increases with decreasing acid concentration 

Salts

Are ionic compounds that contain cations and anions and all are strong electrolytes

When dissolved they undergo hydrolysis meaning they can either produce a hydronium or hydroxide

Can yield basic, acidic, or neutral solutions

Neutral salt solutions

Salts that contain only strong bases and acids are neutral and don’t produce a hydronium or hydroxide

Cations from strong bases: Li^-, Na^-, K^-, Mg^2-, Ca^2-, Ba^2-

Anions from monoprotic acids: Cl^-, Br^-, I^-, NO₃^-, ClO₄^-, ClO₃^-

Hydrogen Sulfate ion

H₂SO₄

It is not neutral because sulfuric acid is polyprotic

Basic salt solutions

Salts derived from conjugates of strong bases (cation) and weak acids (anion) hydrolyze to produce a basic solution

-The conjugate acid of a strong base does not react (neutral)

-The conjugate base of a weak acid produces a basic solution

Acidic salt solutions 

Salts that are derived from the conjugates of strong acids (anions) and weak bases (cations) hydrolyze to produce an acidic solution

-The conjugate base of a strong acid does not react (neutral)

-The conjugate acid of a weak base produces an acidic solution

Metal Cations as weak acids

Cations of small, highly charged metals are weakly acidic

Ex. Al³⁺ , Cr³⁺

Weak/Weak salt solutions

If both ions are conjugates of a weak acid or weak base then compare K_a and K_b values

The larger value will determine the behavior of the solution

Finding pH of a salt

Determine whether a salt will produce an acidic or basic solution based on its ions

React the stronger ion with water to determine the conc. of hydronium or hydroxide

Acid strength and molecular structure

Acid strength determined by how easily it can donate a proton

-if easy to donate = strong acid

-if hard to donate = weak acid

Based on structure we can determine relative strength of an acid

-How easily can electrons be pulled from H atom

-How stable will conjugate base be

Binary acids 

Have acidic hydrogen directly attached to non-metal atom 

-Ex. HF, HCl, HBr, H₂S

Factors that affect acid strength 

-The more polar a bond, more acidic 

-The stronger the bond the weaker the acid 

-The more stable the conjugate base, the stronger the acid 

Bond polarity in binary acids

In a polar bond the H must have the partial positive in order to be acidic

Bond polarity increases as we go to the right and up on the periodic table

*More specific for when comparing in a row

Bond strength in binary acids

The stronger the bond, the more tightly hydrogen is held, making it harder to dissociate, the weaker the acid

Larger atoms will be further from the H atom, so the bond will be weaker and the acid will be stronger

*More specific for when comparing in a period

Conjugate base stability 

The greater the stability of the conjugate base the stronger the acid 

Stability increases to the right and down a periodic table 

Acidity increases to the right and down the periodic table 

Oxyacids

Have an acidic hydrogen bonded to an oxygen atom, with that oxygen atom bonded to another, non hydrogen central atom

Acid strength is related to the strength and polarity of the H-O bond

-the electronegativity of the central atom

-# of oxygen bonded to the central atom

-oxidation state of the central atom

Electronegativity of central atoms on oxyacids

The more electronegative the central atom:

-the stronger the oxyacid

-weaker and more polarized the H-O bond

-more easily electron density is drawn away from H atom

Acidity of oxyacids decreases down a group

Number of oxygens in oxyacids 

The more oxygens attached to the atom, the stronger the oxyacid 

-weakens and polarizes the H-O bond

Increasing the number of oxygens helps stabilize the conjugate base by spreading out the negative charge 

Oxidation states in oxyacids

The larger the oxidation state of the central atom, the stronger the oxyacid

Acidity of oxyacids increases to the right across a period

Other factors affecting acid/base strength

A compound becomes more acidic as the positive charge increases

-easier to accept electron pair

A compound becomes more basic as the negative charge increases

-easier to donate electron pair

Direction of Acid-Base reactions

Depends on the strength of the acids and bases involved

Can always assume a strong acid will donate a proton and a strong base will accept a proton

For weak acids and bases check K_a and K_b values

Stronger acid + stronger base → weaker acid + weaker base

Direction of Acid-Base reactions cont…

pK_a is an indication of acid strength 

-can use pK_a to calculate the equilibrium constant for the reaction between any acid-base pair 

K_eq = 10^∆pKa

∆pK_a = pKa (product) - pKa (reactant) 

The larger the difference in pK_as, the further the equilibrium lies toward the acid with the higher pK_a (i.e. the weaker acid)  

Common ion effect

The shift in the position of the equilibrium upon the addition of a substance that provides more of an ion that is already involved in the equilibrium

Le chatelier’s principle applies here

Buffer Solutions

A buffer is a solution that resists change in pH with the addition of either an acid or base

aka buffer system

They must contain either:

-significant amount of weak acid and its conjugate base

-significant amount of a weak base and its conjugate acid

Three ways to make buffer solutions 

Mixing a weak acid or a weak base and its conjugate salt 

Mixing two salts that provide a conjugate acid-base pair 

Reacting some of a weal base with a strong acid or reacting some of a weak acid with a strong base 

All three of these methods will leave both a weak acid and its conjugate base (and vice versa) forming a buffer solution