chapter 9 lecture notes: acids, bases and equilibrium 108 chapter 9 lecture notes acids, bases, and...

Chemistry 108 Chapter 9 Lecture Notes Acids, Bases, and Equilibrium

1

Chapter 9 Lecture Notes: Acids, Bases and Equilibrium Educational Goals

1. Given a chemical equation, write the law of mass action. 2. Given the equilibrium constant (Keq) for a reaction, predict whether the reactants

or products are predominant. 3. Use Le Châtelier’s Principle to explain how a chemical reaction at equilibrium

responds when a change is made to the concentration of reactant or product. 4. Know the definitions of Bronsted-Lowry acids and bases. 5. Given the acid form or the base form of a conjugate pair, identify its conjugate. 6. List the properties of acidic and basic solutions. 7. Understand the term "acid strength," and know how acid strength is related to the

acidity constant (Ka) value. 8. Given the [H3O+], be able to calculate the [OH-] (and vice versa). 9. Given the [H3O+], be able to calculate the pH (and vice versa). 10. Given the [H3O+], [OH-], or pH, be able to characterize a solution as being

acidic, basic, or neutral. 11. Given the reactants, predict the products of a neutralization reaction. 12. Given the pH of a solution and the pKa for a particular acid, determine the

relative concentrations of the acid and base forms of the conjugate pair. 13. Define a buffer, and describe how a buffer solution is made.

Equilibrium In earlier chapters, we assumed all reactions converted reactants into products until one of the reactants was completely used up. This is often not the case since many reactions are _________________________. Products can be converted __________________ to reactants!

N2O4(g) 2NO2(g) Colorless Gas Brown Gas

This reaction is_____________________! At Equilibrium: _______________ of forward reaction is equal to the __________ of reverse reaction At __________________________the rate of the forward and reverse reactions are the same and the concentrations (amounts) of reactants and products do not change.


2

Law of Mass Action For any chemical reaction at equilibrium:

aA + bB cC + dD The equilibrium concentrations are given by: Keq =

• Pure _________________or _________________do not appear in the

equilibrium equation

• Equilibrium constant values vary with ________________________. •

Examples: Balance the reaction equation and then write the corresponding equilibrium constant expression.

CO(g) + O2(g) CO2(g) Write the equilibrium constant expression for this reaction.

M

M


3

When an equilibrium constant has a value less than 1, the denominator (related to reactant concentration) of the equilibrium constant expression is greater than the numerator, which means that at equilibrium the reactants are favored (more reactants than products). We say that the equilibrium favors the _____________________ when Keq < 1 When an equilibrium constant has a value greater than 1, the numerator (related to product concentration) of the equilibrium constant expression is greater than the denominator, which means that at equilibrium the products are favored (more products than reactants). We say that the equilibrium favors the _____________________ when Keq > 1

Le Châtelier’s Principle The response to a loss of equilibrium is predicted by: Le Châtelier’s Principle, which states that when a reversible reaction is pushed out of equilibrium, the reaction responds to reestablish equilibrium. Or……think of it like this: The response to a loss of equilibrium is predicted by: Le Châtelier’s Principle, which states that the chemical system will try to _____________what you did!

If the amount of one species is changed, the amounts of other species must change in order for the equilibrium constant to remain ________________________!

M


4

Group Work:

Catalysts do not Effect Equilibrium

• Catalysts increase reaction rates by lowering the activation energy. • When a catalyst is used in a reversible reaction the lowered activation energy

speeds up both the forward and reverse reactions. • The overall result is that a catalyst has no effect on equilibrium or on the value of

Keq.


5

Acids and Bases

Properties of Acids

• Acids will __________________ some metals • Acids will turn a plant pigment (blue litmus) from blue to red • Acids taste _________________

Properties of Bases • Bases feel ____________________ or __________________ • Bases will turn a plant pigment (red litmus) from red to blue • Bases taste _____________________

Bronsted-Lowry Acids and Bases A Bronsted-Lowry acid ________________________H+. A Bronsted-Lowry base ________________________ H+.


6

You Try It: For the forward direction of the reaction, identify the Bronsted-Lowry acid and base. Example:

You try it:

Conjugates Compounds, such as HCN and CN- or NH3 and NH4

+, which differ only in the presence or absence of H+ are called______________________________


7

Amphoteric Compounds .Compounds that can act as acids or as bases are called______________________.

• Examples: • H2O (water)

• H2O + CN-→ OH + HCN (water as an acid) • H2O + HCl → H3O+ + Cl- (water as a base)

• HCO3- (hydrogen carbonate)

• HCO3- + CN-→ CO3

2- + HCN (HCO3- as an acid)

• HCO3- + HCl → H2CO3 + Cl- (HCO3

- as a base) Important note: Another way to write acid-base equations:

• These are equivalent: 1) HCN + H2O → H3O+ + CN- 2) HCN → H+ + CN-

• In #2 above, we write H+ even though we know H+ only exist in water as H3O+ Ionization of Water

Ionization of water obeys the law of mass action just like _________________ other chemical reaction!!

Kw = [OH-][H3O+] = 1.0 x 10-14 M2

That means whenever we know [OH-], we can calculate the concentration of__________.

Also, whenever we know [H3O+] we can calculate the concentration of ______________.

Example: What is the concentration of [OH-] in an aqueous solution when [H3O+] = 1.0 x 10-3 M: What is the concentration of [H3O+] in an aqueous solution when [OH-] = 5.0 x 10-11 M


8

The pH Scale

• The “p” in pH means “take ______________________ then multiply by (-1)” • pH means “take logarithm of the ____________--ion concentration, then multiply

by (-1)” • pH = __________________________

Don’t forget about the web site that Dr. Perez made for helping chemistry 108 students with math! • You can link to it from our class webpage.

pH of a Solution Example: What is the pH of an aqueous solution with [H3O+] = 6.3 x 10-1M? Strategy: pH = - log[H3O+] NOTE: pH is one of the few unit-less numbers used in science!


9

You try it: What is the pH of an aqueous solution with [OH-] = 6.3 x 10-1M ? Strategy: pH = - log[H3O+] Kw = [OH-][H3O+] = 1.0 x 10-14 M2 Significant Figures in Log Values (pH) In a log value, the digits to the ___________________of the decimal point are not significant!

Only the digits to the right of the decimal point count as significant figures in pH values.


10

The Antilog What is the [H3O+] of an aqueous solution with pH = 5.3? What is the [H3O+] of an aqueous solution with pH = 12.3? We say a solution is _________________

when [OH-] = [H3O+]. We say a solution is_________________

when there is more [H3O+] than [OH-]

We say a solution is _________________ when there is more [OH-] than [H3O+]. Measuring pH: 1) pH meter 2) Chemical Color Indicators Change Colors according to pH

• Acids will turn a plant pigment (blue litmus) from blue to red. • Bases will turn a plant pigment (red litmus) from red to blue.


11

Acid and Base Strength • The stronger the acid, the more H3O+ it produces and the lower the pH of the

solution that it forms.

• A strong acid such as HCl __________________ almost completely in water.

• HCl(aq) + H2O(l) ↔ Cl-(aq) + H3O+ (aq)

When dealing with acids, the equilibrium constant (Keq) is also known as the _______________ __________________(Ka). The __________________ the Ka for an acid, the ______________________ an acid it is. Sometimes an acids strength is tabulated by pKa (pKa = -logKa). The ___________________ the pKa, the _____________________ the acid.

– pKa for HCl = -7.00 – pKa for HF = 3.18

M

M


12

The stronger the base, the more OH- it produces and the higher the pH of the solution that it forms.

• A _____________ base such as NaOH dissociates (falls apart) almost completely in water.

NaOH(s) ↔ Na+(aq) + OH- (aq)

Neutralizing Acids and Bases

• _______________________takes place when an acid and a base react to form water and a salt (an ionic compound).

• HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

strong acid strong base salt water

• When a strong acid is neutralized by a strong base, the solution has a pH = _____

pH Buffers

Notation→ An interesting relationship exists between pH and an acid and its conjugate base.

• For the acid (HA) and its conjugate base (A-),

– There is more HA when the pH is _______________than the pKa.

– There is more A- when the pH is ________________than the pKa.

– There are equal amounts of HA and A- when the pH = pKa.


13

Henderson-Hasselbalch Equation (you will not need to derive this equation or use it on the exam, just know the rules given on the last three lines of the previous page. I put it here in case you need it in later academic or professional situations)

Strong acids have weak conjugate bases. Strong bases have weak conjugate acids.

Buffers

A ______________________is a solution that resists change in pH when small amounts of acid or base are added. Example: adding acid (H3O+) to acetic acid/acetate buffer

Equilibrium will shift to the left!

Example: adding base (OH-) to acetic acid/acetate buffer

Equilibrium will shift to the right!


14

A buffer is a solution made for fairly high concentration of a weak acid and it’s conjugate base. Buffers are most resistant to pH changes when the pH = pKa of the weak acid and are effective when the pH is within ____________ ___________of the pKa (pH = pKa ± 1).

Buffers in the Blood

• The pH of your blood normally ranges between 7.35 and 7.45. • A blood pH below the normal range is called acidosis, while a blood pH above

this range is called alkalosis, either one of which is potentially fatal.

• Blood is kept in this narrow range • (pH = 7.35 – 7.45) with the help of buffers.

• The most important buffer system in blood is formed from carbonic acid (H2CO3) and its conjugate base, bicarbonate ion (HCO3

-):

chapter 9 lecture notes: acids, bases and equilibrium 108 chapter 9 lecture notes acids, bases, and...

Documents