Acids, Bases, and Acid–Base Equilibria

Acids, Bases, and Acid–Base Equilibria

Strong and weak acids and bases
Buffers Titrations indicators Salt hydrolysis Theory acids and bases Strong and weak acids and bases Ionic product of water The pH scale Ka/Kb/pKa/pKb

Common Household Substances that Contain Acids and Bases

Acids and Bases Acids: Sour taste- Acetic Acid in vinegar
Citric acid in citrus fruits Bases: Bitter Taste - Turnups, Broccoli, Baking Soda

Solubility of Acids and Akalis
Most acids are soluble or react strongly with water Some bases are soluble, some are insoluble Soluble bases are called ALKALIS

Three Models for Acids and Bases

Arrhenius Definition Acids produce hydrogen ions in aqueous solution.
Bases produce hydroxide ions when dissolved in water. Limits to aqueous solutions. Only one kind of base. NH3 ammonia could not be an Arrhenius base.

Bronsted-Lowry Acids and Bases
It’s all about protons (H+) Acid: Proton donor HCl(aq)  H+(aq) + Cl-(aq) H2SO4(aq)  2H+(aq) + SO42-(aq) Base: Proton acceptor NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) OH-(aq)* + H+(aq)  H2O(l) *From any soluble hydroxide or other alkali

Bronsted-Lowry Definitions
And acid is an proton (H+) donor and a base is a proton acceptor. Acids and bases always come in pairs. HCl is an acid. When it dissolves in water it gives its proton to water. HCl(g) + H2O(l) H3O+ + Cl- Water is a base -makes hydronium ion Can also be used on gases

The Brønsted–Lowry Theory
Brønsted–Lowry theory defines acids and bases in terms of proton (H+) transfer. A Brønsted–Lowry acid is a proton donor. A Brønsted–Lowry base is a proton acceptor. The conjugate base of an acid is the acid minus the proton it has donated. The conjugate acid of a base is the base plus the accepted proton.

Ionization of HCl H2O is a base in this reaction because it accepts the H+ Conjugate acid of H2O HCl acts as an acid by donating H+ to H2O Conjugate base of HCl

Ionization of Ammonia

Water Is Amphiprotic H2O acts as an acid when it donates H+, forming the conjugate base ___ H2O acts as a base when it accepts H+, forming the conjugate acid ___ Amphiprotic: Can act as either an acid or as a base

Amphoteric Behave as both an acid and a base. Water autoionizes
2H2O(l) H3O+(aq) + OH-(aq)

Acid in Water HA(aq) + H2O H3O+(aq) + A-(aq) acid base conjugate conjugate acid base

Pairs Competition for H+ between H2O and A- General equation
HA(aq) + H2O(l) H3O+(aq) + A-(aq) Acid + Base Conjugate acid Conjugate base This is an equilibrium. Competition for H+ between H2O and A- The stronger base controls direction. If H2O is a stronger base it takes the H+ Equilibrium moves to right.

Strength of Conjugate Acid–Base Pairs
A stronger acid can donate H+ more readily than a weaker acid. The stronger an acid, the weaker is its conjugate base. The stronger a base, the weaker is its conjugate acid. An acid–base reaction is favored in the direction from the stronger member to the weaker member of each conjugate acid–base pair.

… the weaker the conjugate base. The stronger the acid …
And the stronger the base … … the weaker the conjugate acid.

Strong Acids The “strong” acids—HCl, HBr, HI, HNO3, H2SO4, HClO4—are considered “strong” because they ionize completely in water. The “strong” acids all appear above H3O+ . The strong acids are leveled to the same strength—to that of H3O+— when they are placed in water.

Periodic Trends in Acid Strength
The greater the tendency for HX (general acid) to transfer a proton to H2O, the more the forward reaction is favored and the stronger the acid. A factor that makes it easier for the H+ to leave will increase the strength of the acid. Acid strength is inversely proportional to H— X bond-dissociation energy. Weaker H—X bond => stronger acid. Acid strength is directly proportional to anion radius. Larger X radius => stronger acid.

Periodic Trends in Acid Strength

The Relationship of Acid Strength and Conjugate Base Strength

Strengths of Acids and Bases
Strong Acids always have very weak conjugate bases Very weak acids always have very strong conjugate bases Acids with mid-range strengths have bases with mid range strengths

Lewis Acids and Bases Acid: electron pair acceptor
Species with an incomplete octet/outer-shell Base: electron pair donor Species with a lone pair For example: Gilbert Lewis

Lewis Acids & Bases A Lewis acid and base can interact by sharing an electron pair.

Lewis Acids & Bases A Lewis acid and base can interact by sharing an electron pair. Formation of hydronium ion is an excellent example. +

Lewis Acids & Bases Other good examples involve metal ions. Such bonds as the H2O ---> Co bond are often called COORDINATE COVALENT BONDS because both electrons are supplied by one of the atoms of the bond.

Lewis Acids & Bases The combination of metal ions (Lewis acids) with Lewis bases such as H2O and NH > COMPLEX IONS All metal ions form complex ions with water — [Cu(NH3)4]2+

:N Lewis Acids and Bases F H B F H F H
Boron triflouride wants more electrons. F H B F :N H F H

:N Lewis Acids and Bases H F H B F F H Most general definition.
Acids are electron pair acceptors. Bases are electron pair donors. H F :N H B F F H

Lewis Acids and Bases F H F B N H F H
Boron triflouride wants more electrons. BF3 is Lewis base NH3 is a Lewis Acid. F H F B N H F H

( ) ( ) Lewis Acids and Bases Al+3 Al H O H +3 H O H 6 + 6
Copyright © Houghton Mifflin Company. All rights reserved. +3 ( ) 6 H Al O H

Types of Acids Polyprotic Acids- more than 1 acidic hydrogen (monoprotic, diprotic, triprotic). Oxyacids - Proton is attached to the oxygen of an ion. Organic acids contain the Carboxyl group - COOH with the H attached to O Generally very weak.

Polyprotic acids Always dissociate stepwise.
The first H+ comes of much easier than the second.

Big Six Strong Acids HCL – Hydrochloric Acid - 1
HNO3 – Nitric Acid - 2 H2SO4 – Sulfuric Acid – 3 HCLO4 – Perchloric Acid - 4 HBr – Hydrobromic acid - 5 HI – Hydroiodic Acid - 6 These are the only acids that completely dissociate in water

Monoprotic Examples

Diprotic Examples

Triprotic Examples

Acetic Acid

Acid Reactions React with the following elements and compounds to form salts Metals Bases Carbonates

Salt Hydrolysis Hydrolysis is the chemical breakdown of a compound due to reaction with water. A neutralization between and acid and a base produces a salt Though salts are products of neutralization, they do not all form neutral aqueous solutions The relative strength of the CA and CB in the salt determine the extant of hydrolysis and so the pH of the solution

Salts Parent Acid and a Parent Base form a salt and water
Parent is used to refer to the relationship between an acid, base and their salt

Acid, Base and Salt Parent acid Parent Base HCl NaOH Salt H2O

Acid + Metal Acid + Metal → Salt + Hydrogen 2HCl + Mg → MgCl2 + H2(g)
Effervescence---reaction with metal gas is given off. Very reactive metals may do this explosively

Acid + Base Reactions between acids and bases are known as neutralization reactions This is always an exothermic reaction. Acid + Base → Salt + Water HCl + NaOH → NaCl + H2O

Acid + Carbonate Acid + Carbonate → Salt + Water + Carbon Dioxide
2HCl + CaCO3 → CaCl2 + H2O + CO2

Relationships KW = [H+][OH-] KW = 1.0 x10-14 14.00 = pH + pOH
[H+],[OH-],pH and pOH Given any one of these we can find the other three.

The pH Scale Graphic: Wikimedia Commons user Slower

Testing Swimming Pool Water
Copyright © Houghton Mifflin Company. All rights reserved.

pH pH= -log[H+] Used because [H+] is usually very small
As pH decreases, [H+] increases exponentially pH is positive and has no units One pH unit represents a 10 X change in [H+] ions

Relationship between H and OH
Acid defined as less then 7 on the pH scale A neutral solution is when H and OH are equal Base solution is greater then 7 on the pH scale

Ionization of Water 2H2O(l) H3O+(aq) + OH-(aq) KW = [H+][OH-]
pKW = pH + pOH KW = 1.0 x10-14 14.00 = pH + pOH KW is known as the ionic product constant of water

Relationship between H and OH
KW = [H+][OH-] KW = 1.0 x10-14 [H+] = KW / [OH-] [OH-] = KW / [H+]

Sample Problem A sample of blood at 298 K has [H+] = 4.60 X mol/L. Calculate the concentration of [OH-] and state whether the blood is acidic, basic, or neutral [OH-] = KW / [H+]

Sample problem [OH-] = 1.00 X 10 -14 / 4.6 X 10-8
[OH-] is greater then [H+] the solution is basic

Chemistry TOK Because many people are so familiar with the value of as the pH of water, it is often more difficult to convince them that water with a pH greater the or less than this is still neutral. Can you think of other examples where entrenched prior knowledge might hinder a fuller understanding of new knowledge? Are we likely to misinterpret experimental data when it does not fit with our expectations based on prior knowledge

Acid dissociation constant Ka
The equilibrium constant for the general equation. HA(aq) + H2O(l) H3O+(aq) + A-(aq) Ka = [H3O+][A-] [HA] H3O+ is often written H+ ignoring the water in equation (it is implied). Also called the acid ionization constant H3O is called the hydronium ion Conjugate base is everything that’s left of the acid molecule after the proton is lost Conjugate acid is formed when the proton is transferred to the bass

Acid dissociation constant Ka
HA(aq) H+(aq) + A-(aq) Ka = [H+][A-] [HA] We can write the expression for any acid. Strong acids dissociate completely. Equilibrium far to right. Conjugate base must be weak.

Weak Acids: Ka and pKa HA(aq)  H+(aq) + A-(aq)
Weak acids dissociate to form an equilibrium HA(aq)  H+(aq) + A-(aq) This has the equilibrium constant (aka acid dissociation constant), Ka The values for Ka are often very small, so we use ‘pKa’ to make them easier to handle: pKa = -log10(Ka)

pKa pKa tells us how acidic (or not) a given hydrogen atom in a molecule is. pH tells us how acidic a solution is. This is also expressed in logarithmic form are follows: pKa = -log10 Ka Because of the minus sign, the lower the pKa, the higher the Ka and the stronger the acid.

Ka and pKa in action In order of decreasing acid strength:
Smaller Ka  weaker acid Smaller pKa  stronger acid Acid Ka pKa Hydronium ion, H3O+ 1.00 0.00 Oxalic acid, HO2CCO2H 5.9x10-2 1.23 Hydrofluoric, HF 7.2x10-4 3.14 Methanoic, CHOOH 1.77x10-4 3.75 Ethanoic, CH3COOH 1.76x10-5 4.75 Phenol, C6H5OH 1.6x10-10 9.80

Back to Pairs Strong acids Ka is large [H+] is equal to [HA]
A- is a weaker base than water Weak acids Ka is small [H+] <<< [HA] A- is a stronger base than water

Weak Bases: Kb and pKb BOH(aq)  B+(aq) + OH-(aq)
Weak bases dissociate to form an equilibrium BOH(aq)  B+(aq) + OH-(aq) This has the equilibrium constant (aka base dissociation constant), Kb The values for Kb are often very small, so we use ‘pKb’ to make them easier to handle: pKb = -log10(Kb)

Kb and pKb in action In order of decreasing base strength:
Smaller Kb  weaker base Smaller pKb  stronger base Base Kb pKb Diethylamine 1.3x10-3 2.89 Ethylamine 5.6x10-4 3.25 Methylamine 4.4x10-4 3.36 Ammonia 1.8x10-5 4.74

Ka × Kb = Kw = 1.00x10-14 pKa + pKb = pKw = 14.0 pH + pOH = pKw = 14.0

Key Points At the point of half-neutralisation: pKa = pH AND pKb = pOH
For a weak acid: AND pKa = -log10(Ka) For a weak base: AND pKb = -log10(Kb) At the point of half-neutralisation: pKa = pH AND pKb = pOH

Percent dissociation = amount dissociated x 100 initial concentration
For a weak acid and weak base percent dissociation increases as acid becomes more dilute.

Problem #1: A M solution of a generic weak acid (HA) has a pH of Determine the Ka. Notice that a generic weak acid is used, symbolized by the formula HA. Solution 1) Write the dissociation equation for the acid: HA <===> H+ + A¯ 2) Write the equilibrium expression: K a = ( [H+] [A¯] ) / [HA]

3) Our task now is to determine the three concentrations on the right-hand side of the equilibrium expression since the Ka is our unknown. We will use the pH to calculate the [H+]. We know pH = -log [H+], therefore [H+] = 10¯pH [H+] = 10¯3.26 = x 10¯4 M b) From the dissociation equation, we know there is a 1:1 molar ratio between [H+] and [A¯]. Therefore: [A¯] = x 10¯4 M

c) the final value, [HA] is given in the problem
c) the final value, [HA] is given in the problem. In the example being discussed, M is the value we want. Ka = [( x 10¯4) ( x 10¯4)] / 0.120 Ka =( x 10¯4)2 / 0.120 Ka = ( x 10¯4)2 / 0.120 Ka = 2.52 x 10¯6

Kb example Problem #1: Codeine (C18H21NO3) is a weak organic base. A 5.0 x 10¯3 M solution of codeine has a pH of Calculate the value of Kb for this substance. Comment: note that I will use B to symbolize the weak base. . Solution: 1) Write the ionizaton equation for the base. B + H2O <===> HB+ + OH¯ 2) Write the equilibrium expression: Kb = ( [HB+] [OH¯] ) / [B]

3) Our task now is to determine the three concentrations
on the right-hand side of the equilibrium expression since the Kb is our unknown. We will use the pH to calculate the [OH¯]. We know pH = -log [H+], therefore [H+] = 10¯pH [H+] = 10¯9.95 = x 10¯10 M Knowing the [H+] allows us to get the [OH¯]. To do this, we use Kw = [H+] [OH¯], so we have this: 1.00 x 10¯14 = (1.122 x 10¯10) (x) x = [OH¯] = x 10¯5 M

b) From the ionization equation, we know here is a 1:1 molar ratio between [HB+] and [OH¯]. Therefore: [HB+] = x 10¯5 M c) the final value, [B] is given in the problem. In the example, 5.0 x 10¯3 M is the value we want. use 5.0 x 10¯3 M Kb = [( x 10¯5) ( x 10¯5)] / 5.0 x 10¯3 Kb = 1.59 x 10¯6

Buffers and pH curves

Buffers and pH Curves Buffers: a solution that resists changes in pH on the addition of small amounts of an acid or a base Water is vulnerable to significant fluctuations pH. Is not a good buffer Buffers are mainly found in biological systems, change in enzyme activity in all biochemical reactions

Examples of reaction that use buffers
Human blood Ocean chemistry Electrophoresis Fermentation Dye industry Calibration

How Buffers work Two main types of buffers
Acidic buffers—maintain a pH at less than 7 Basic buffers—maintain a pH more than 7

Composition of acidic buffers
Made by mixing a weak acid with a solution of its salt of a strong Alkali Example: CH3COOH and NaCH3COO Weak Acid Salt weak acid with strong Alkali

Example (cont) The following equilibrium exists in solution
CH3COOH ↔ CH3COO- + H+ NaCH3COO → Na+ + CH3COO-

Example (cont) Addition of acid (H+)
H+ will combine with the base removing most of the added H+ CH3COO H+ → CH3COOH

Addition of a Base OH- will combine with the acid forming the base and water CH3COOH + OH- ↔ CH3COO- + H2O

Summary The H+ and OH- are used in the reaction
They do not persist in the solution So pH is largely unchanged

Making Buffer solutions
Covered fully in biochemistry It depends on The pKa and pKb of its acid or base The ratio of the initial concentrations of acid and salt or base and salt used in the preparation Copyright © Houghton Mifflin Company. All rights reserved.

Factors influencing Buffers
Dilution: Ka and Kb are not changed by dilution The Ratio of the acid, base and salt concentrations does not change by dilution Diluting a buffer does alter the amount of acid or base that it can absorb without changing the pH Buffering Capacity– the amount of acid or base it can absorb Copyright © Houghton Mifflin Company. All rights reserved.

Factors influencing Buffers
Temperature Does affect the values of Ka and Kb Does affect the pH of the buffer Constant temp required when working with buffers Copyright © Houghton Mifflin Company. All rights reserved.

Titrations

Titrations Millimole (mmol) = 1/1000 mol Molarity = mmol/mL = mol/L
Makes calculations easier because we will rarely add liters of solution. Adding a solution of known concentration until the substance being tested is consumed. This is called the equivalence point. Graph of pH vs. mL is a titration curve.

Key Titration Terms Titrant - solution of known concentration used in titration. Analyte - substance being analyzed. Equivalence point or Stoichiometric points - enough titrant added to react exactly with the analyte. Point of Inflection –the pH jump at the equivalence point Endpoint or Change point- the indicator changes color so you can tell the equivalence point has been reached. 4–97

Half-equivalence point:
Anatomy of an acid-base titration curve: In this case weak-acid/strong-base Equivalence point: i.e. The point of inflection (where the gradient starts to drop again). Represents the point at which the acid has just been neutralised. Half-equivalence point: i.e. Half the volume added at the equivalence point The pH of this point is equal to the pKa of the weak acid. Buffer Region: A lot of base has to be added to result in only a small change of pH pH Intercept: Higher pH if a weaker acid used

Strong acid with Strong Base
Do the stoichiometry. mL x M = mmol There is no equilibrium . They both dissociate completely. The reaction is H+ + OH-  HOH Use [H+] or [OH-] to figure pH or pOH The titration of 50.0 mL of M HNO3 with M NaOH

7 pH mL of Base added Strong acid with strong Base Equivalence at pH 7
Copyright © Houghton Mifflin Company. All rights reserved. mL of Base added

The pH Curve for the Titration of 50. 0 mL of 0. 200 M HNO3 with 0
The pH Curve for the Titration of 50.0 mL of M HNO3 with M NaOH

The pH Curve for the Titration of 100. 0 mL of 0. 50 M NaOH with 1
The pH Curve for the Titration of mL of 0.50 M NaOH with 1.0 M HCI

Strong acid /Strong base Titrations
Problem: how much of .100 M NaOH is required to neutralize .200 M HNO3 STEP 1: No NaOH has been added H OH H2O [H+ ] = M pH = -log [H+ ] = Copyright © Houghton Mifflin Company. All rights reserved.

7 <7 pH mL of acid added Weak base with strong acid
Equivalence at pH <7 When the base is neutralized it makes a weak acid 7 Copyright © Houghton Mifflin Company. All rights reserved. <7 pH mL of acid added

Summary of important graphs

Indicators Signal change in pH Are weak acids and weak bases
The undissociated and the dissociated form have different colors The indicator HIn is a weak acid that is at equalibrium HIn H+ + In color A Color B

Indicators Increasing [H+] the equilibrium will shift to the left and favor HIn Decreasing [H+] the equilibrium will shift to the Right and favor In- Low pH color A will dominate High pH color B will dominate

INDICATORS Indicators signal change in PH Universal paper indicator: pH pH

pH Scale Indicators Range and Color Changes of Some
Common Acid-Base Indicators Indicators pH Scale 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Methyl orange red – yellow Methyl red red yellow Bromthymol blue yellow blue Neutral red red yellow Phenolphthalein colorless red colorless beyond 13.0 Bromthymol blue indicator would be used in titrating a strong acid with a strong base. Phenolpthalein indicator would be used in titrating a weak acid with a strong base. Methyl orange indicator would be used in titrating a strong acid with a weak base.

Indicators pH=pKa + log([HI]/[In-]) = pKa + log(10)
pH=pKa+1 on the way down Choose the indicator with a pKa 1 more than the pH at equivalence point if you are titrating with base. Choose the indicator with a pKa 1 less than the pH at equivalence point if you are titrating with acid. This is called the endpoint range

Indicators Since it is an equilibrium the color change is gradual.
It is noticeable when the ratio of [In-]/[HI] or [HI]/[In-] is 1/10 An indicator changes color at the equivalence point One drop of added solution can produce a dramatic change

Acids and Bases can be distinguished
Acid Base indicators are used to distinguished between acids and bases This usually results in a color change of the indicator Litmus paper is the most common (dye derived from Lichens (algae, cyanobacteria, and fungus) Turns pink in acids and blue in bases

Other Indicators

Strength based on ionization
The reactions of acids and bases are dependent on the fact that they dissociate in solution to produce H+ ions and OH- ions The extant of this dissociation is what define the strength of an acid or a base

Strength based on ionization
If the acid or base dissociates fully it will exist entirely as ions in solution then is a strong acid. The big six completely dissociate

Strong Acids Strong Acids: HA(aq)  H+(aq) + A-(aq) …..the acid fully dissociates into ions For example: HCl(aq)  H+(aq) + Cl-(aq) Includes: hydrochloric, sulfuric, phosphoric, nitric Strong acids have weak conjugate bases

Weak Acids CH3COOH + H2O  CH3COO- + H3O+
HA(aq)  H+(aq) + A-(aq) …..the acid only partially dissociates into ions CH3COOH + H2O  CH3COO- + H3O+ Includes: hydrofluoric, ethanoic, carbonic Weak acids have strong conjugate bases The equilibrium shifts to the left in favor of the reactants

Strong and Weak Bases Strong Bases: BOH(aq)  B+(aq) + OH-(aq) …..the base fully dissociates into ions For example: NaOH(aq)  Na+(aq) + OH-(aq) Includes: group (I) hydroxides, barium hydroxide Strong bases have weak conjugate acids

Weak Bases Weak Bases: BOH(aq)  B+(aq) + OH-(aq) …..the base only partially dissociates into ions For example: NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) Includes: ammonia, amines Weak bases have strong conjugate acids

Simulation

So what? The equilibrium has a profound effect on the properties of the acid/base Compared with strong acids of the same concentration, weak acids: Have lower electrical conductivity React more slowly pH is higher (less acid) Change pH more slowly when diluted However, they neutralise the same volume of alkali Weak bases follow a similar pattern

Common strong acids and bases
HCL Hydrochloric Acid LiOH Lithium Hydroxide Common examples of strong species HNO3 Nitric Acid NaOH Sodium Hydroide H2SO4 Sulfuric Acid KOH Potassium Hydroxide Ba(OH)2 Barium Hydroxide

Common weak acids and bases
CH3COOH and other organic acids Acetic acid or ethanoic acid NH3 Ammonia Common examples of weak species H2CO3 Carbonic Acid C2H5NH2 and other amines Ethylamine H3PO4 Phosphoric Acid

Acid Deposition

Causes of Acid Deposition
All Rain water is naturally acidic due to the presence of carbon dioxide Which dissolves is water forming a weak acid, carbonic acid H2O + CO2  H2CO3

Acid Rain All rain water has a pH of 5.6. slightly acidic
Acid rain refers to rain water that has a pH below 5.6 Contains additional acids The main contributors to acid rain are the oxides of Sulfur and Nitrogen primary pollutant Acid rain is a secondary pollutant

Acid Deposition Is a broader term than just acid rain and includes all processes by which acidic components precipitate out or leave the atmosphere as gases

Two main types of acid deposition
There are 2 main types of acid deposition Wet acid deposition: rain, snow, sleet, hail , fog, mist, dew all fal to the ground as aqueous precipitates Dry acid deposition: acidifying particle fall t othe ground as dust and smoke later dissolve in water to form acids

Sulfur Oxides

Nitrogen Oxides

Impact on materials

Impact on plant life

Impact on animals

Impact on water

Impact on human health

Fighting acid rain

Acids, Bases, and Acid–Base Equilibria

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