Chapter 11: Acids and Bases

Arrhenius Acids

Produce hydrogen ions (H⁺) when they dissociate in water

• Are also electrolytes (bc they make H+ in water)

• Sour taste

• My sting to the touch

Arrhenius Bases

Produce hydroxide ions (OH⁻) in water

• Are electrolytes (bc they make OH⁻ in water)

• Chalky, bitter taste

• Soapy, slippery to the touch

What color do Arrhenius bases vs acids turn litmus paper?

Bases: Blue

Acids: Red

What color do Arrhenius bases vs acids turn using phenolphthalein indicators?

Bases: Pink

Acids: Colorless

Naming Acids (hydrogen ion and a nonmetal)

• What is the exception to the nonmetal rule?

Exception: CN⁻ also counts although its a polyatomic

HCN, is called hydrocyanic acid

Use the prefix Hydro- and end with -ic acids

• Use the name of the element in between (first syllable prob)

HCl = Hydrochloric acid

Naming Acids (hydrogen ion and a polyatomic)

Change the end of the polyatomic name

Polyatomic ends in ‘ate‘ = ic acid

• HClO₃⁻; take out the H+ you get ClO₃²⁻ which is chlorate so HClO₃⁻ is chloric acid

Polyatomic ends in ‘ite‘ = ous acid

• HClO₂; take out the H+ and you get ClO₂⁻ which is chlorite so HClO₂ is chlorous acid

Naming Bases

• What’s the exception?

Exception: NH₃ which is ammonia

Typical bases are named as hydroxides

• NaOH → sodium hydroxide

• Ba(OH)₂ → barium hydroxide

Br∅nwsted-Lowry Thory

• An acid is a substance that donates H⁺ (donor)

• A base is a substance that accepts H⁺ (recepient)

How does NH₃ act as a Br∅nwsted-Lowry base

• Why?

When combined with water, NH₃ acts as a base that accepts H⁺ from H₂O

• The nitrogen in NH₃ has a stronger attraction for the H⁺ than the oxygen, so water acts as the donating acid

What is H⁺ also referred to as?

• H₃O⁺ = hydronium ion

• A proton

Why is an hydrogen ion also called a proton?

Hydrogen atoms which have only one electron lose that one electron to become an ion leaving only one proton behind

Conjugate Acid-Base Pairs

There are two in any acid-base reaction

• Each pair is related by the loss and gain of H⁺

• One pair occurs in the forward direction and another in the reverse direction

HA + B ⇌ A⁻ + BH⁺

• HA and A⁻ is a pair

• B and BH⁺ is a pair

Amphoteric Substances

Substances that can act as both acids and bases

• Water is an example

How is water an amphoteric substance?

It can act as both a base and an acid

• Becomes OH⁻

  • If it acts as a acid

  • When it reacts with a stronger base

• H₃O⁺

  • If it acts as a base

  • When it reacts with a stronger acid

What are Amphoteric Substances also called?

Amphiprotic

Strong and Weak Acids

• Dissociation

• Strong acids ionize 100% in an aqueous solution

• Weak acids dissociate only slightly in water to form a few ions in an aqueous solution

Strong Acids

• HA(strong acid) Ions dissociate 100%

• Produce large concentrations of H₃O⁺ and the anion (A⁻)

HI(aq) + H₂O(l) → H₃O⁺(aq) + O⁻(aq)

Weak Acids

• Only a few molecules dissociate; most stay in the undissociated molecular form of the acid

• The concentrations of H₃O⁺ and the anion (A⁻) are low

HF(aq) + H₂O(l) ⇌ H₃O⁺(aq) + F⁻(aq)

Diprotic Acids

Some strong/weak acids have two H⁺ which dissociate one at a time

• Reacts twice

Diprotic Acid: Carbonic Acid

Weak diprotic acid

• First reaction:

  • H₂CO₃(aq) + H₂O(l) ⇌ H₃O⁺(aq) + HCO₃⁻(aq)

• Second Reaction:

  • HCO₃⁻(aq) + H₂O(l) ⇌ H₃O⁺(aq) + ClO₃²⁻(aq)

Both are reversible reactions due to the acid’s weak nature

Note that HCO₃⁻(bicarbonate) is amphoteric because it can become ClO₃²⁻(conjugate base) or H₂CO₃(conjugate acid)

Diprotic Acid: Sulfuric Acid

Strong diprotic acid

• First reaction:

  • H₂SO₄(aq) + H₂O(l) → H₃O⁺(aq) + HSO₄⁻(aq)

• Second Reaction:

  • HSO₄⁻(aq) + H₂O(l) ⇌ H₃O⁺(aq) + SO₄²⁻(aq)

Only the second reaction is reversible due to the acid’s strong nature, but the HSO₄⁻ formed after the first reaction is weak

Note that HSO₄⁻ is amphoteric because it can become SO₄²⁻(conjugate base) or H₂SO₄(conjugate acid)

Strong Bases

Strong electrolytes

• Formed from metals of groups 1A and 2A

• Include LiOH, NaOH, KOH, Ba(OH)₂, Sr(OH)₂, Ca(OH)₂

• Dissociate completely in water

  • However, they have low solubility

KOH(s) → K⁺(aq) + OH⁻(aq)

Weak Bases

Weak electrolytes

• Poor acceptors of H⁺ ions; not readily available to accept the H⁺

• Produce very few ions in solution

• Include ammonia NH₃

NH₃(g) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

• Conjugate acid-base pairs exist

• ⇌ = equilibrium, doesn’t dissociate completely

What is the rule with acids and their conjugate base?

The stronger an acid is, the weaker its conjugate base is

What are the strong acids?

HCI = hydrochloric acid

HBr = hydrobromic acid

HI = hydroiodic acid

H₂SO₄ = sulfuric acid

HClO₄ = perchloric acid

HNO₃ = nitric acid

What are the strong bases?

• What group are they in?

LiOH = Grp 1

NaOH = Grp 1

KOH = Grp 1

RbOH = Grp 1

CsOH = Grp 1

Ca(OH)₂ = Grp 2

Sr(OH)₂ = Grp 2

Ba(OH)₂ = Grp 2

In an acid-base reaction, are both the acids and bases of the same strength?

Nope, one base will be stronger and one acid will be stronger

How do you determine the direction of an reaction?

By comparing the relative strengths of the acids and bases in the reaction

How does a reaction move?

Strong → Weak

What are the directions of these reactions and what type of reaction are they?

• Reactants: H₂SO₄ and H₂O Products: H₃O⁺ and HSO₄⁻

• Reactants: CO₃²⁻ and H₂O Products: OH⁻ and HCO₃⁻

Reactants: H₂SO₄ and H₂O Products: H₃O⁺ and HSO₄⁻

• H₂SO₄ and H₂O = stronger acid and base

• H₃O⁺ and HSO₄⁻ = weaker acid and base

  • The reaction is normal →

Reactants: CO₃²⁻ and H₂O Products: OH⁻ and HCO₃⁻

• CO₃²⁻ and H₂O = weaker acid and base

• OH⁻ and HCO₃⁻ = stronger acid and base

  • The reaction is reversible ⇌

Acid and base dissociation constant

Ka (a is a subscript) = acids

Products in brackets(multiplied) over the reactants(multiplied) in brackets

• Only includes gasses and aqueous solutions

  • Not solids or liquids

HCHO₂(aq) + H₂O(l) ⇌ H₃O⁺(aq) + CHO₂⁻(aq)

• Ka = [H₃O⁺][CHO₂⁻]/[HCHO₂]

Kb (b is a subscript) = base

• Same idea as acids

What is the acid dissociation constant used for?

Used to distinguish strong acids from weak acids

• Stronger acids have larger Ka values

Dissociation of Water

The equilibrium reach between the conjugate acid-base paired of water produces both H₃O⁺ and OH⁻

• Reversible reaction

• H₂O(l) + H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)

What is the water dissociation constant?

• Formula

• Values in what temp?

Kw = [H₃O⁺][OH⁻]

At 25°C, H₃O⁺ and OH⁻ are both 1.0 x 10⁻⁷ M

• Therefore, [H₃O⁺][OH⁻] would equal 1.0 x 10⁻¹⁴ M

How can you use [H₃O⁺] and [OH⁻] to determine the acidity of a solution?

If [H₃O⁺] and [OH⁻] are equal, then the solution is neutral

If [H₃O⁺] is in a larger quantity, then the solution is acidic

If [OH⁻] is in a larger quantity, then the solution is basic

What is the acidity of pure water?

• Why

Neutral

• In pure water, the ionizing energy of water molecules produces small but equal amounts of H₃O⁺ and OH⁻ ions ; they are both 1.0 x 10⁻⁷ M

What happens when you add an acid to pure water?

The solution becomes acidic

• Increases the [H₃O⁺]; causing it to exceed 1.0 x 10⁻⁷ M

• Decreases the [OH⁻]

• [H₃O⁺] > [OH⁻]

What happens when you add a base to pure water?

The solution becomes basic

• Increases the [OH⁻] ; causing it to exceed 1.0 x 10⁻⁷ M

• Decreases the [H₃O⁺]

• [H₃O⁺] < [OH⁻]

True or False: When acid or base is added, the combined value of [H₃O⁺] and [OH⁻] is still 1.0 x 10⁻¹⁴

True

What is the pH scale used for?

Describing the acidity of solutions

What is the range of pH scale and what do the values mean?

0 - 14

pH = 7 → neutral

pH > 7 → basic

pH < 7 → acidic

What are is the [H₃O⁺] concentration representations of pH?

pH < 7.0 → [H₃O⁺] > 1.0 x 10⁻⁷

pH > 7.0 → [H₃O⁺] < 1.0 x 10⁻⁷

pH = 7.0 → [H₃O⁺] = 1.0 x 10⁻⁷

What are the three methods of pH determination?

• pH meter

• pH paper

• Indicators

What kind of scale is the pH scale?

A logarithmic scale that corresponds to the [H₃O⁺] of aqueous solutions

• A change of one pH unit corresponds to a tenfold change in [H₃O⁺]

• pH decreases as [H₃O⁺] increases

Calculating pH

• How do the sigfigs work?

The negative log(base 10) of the [H₃O⁺]

• -log[H₃O⁺]

• Always positive

If the given value of [H₃O⁺] is 1.0 x 10⁻², we take the sigfig count of the 1.0 which is two, and determine that there are two decimal places

• The pH value is 2, and since we need two decimal places, we get 2.00

How do you calculate [H₃O⁺] from pH?

[H₃O⁺] = 10⁻ph

• Basically, [H₃O⁺] is equal to 10 to the negative pH

Reaction of acids with metals → what does it produce

Produce salt and hydrogen

Reaction of acids with carbonates and hydrogen carbonates → what does it produce

Carbon dioxide gas, salt, and water

Reaction of acids with bases → what does it produce

Salt of the metal and water

Reaction of acids → what’s produces by bicarbonate and carbonate ions

Carbon dioxide gas

What is a salt?

An ionic compound that does not have H⁺ as the cation or OH⁻ as the anion

Neutralization Reactions

An acid reacts with a base to produce salt and water

• The salt formed is the anion from the acid and the cation from the base

• One H⁺ always reacts with one OH⁻

HCI(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

acid base salt water

How do you write a neutralization reaction?

Compltete Molecular Equation: HCI(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Total Ionic Equation: H⁺ + Cl⁻ + Na⁺ + OH⁻ → Na⁺ + Cl⁻ + H₂O(because only aqueous solutions split into ions)

Net Ionic Equation:

H⁺ + Cl⁻ + Na⁺ + OH⁻ → Na⁺ + Cl⁻ + H₂O

H⁺(aq) + OH⁻(aq) → H₂O(l)

Acid-Base Titrations

A known volume of an acid is placed in a flask with an indicator and titrated with a measured volume of base solution to the neutralization endpoint

Titration

A laboratory procedure used to determine the molarity of an acid

• Uses a base to neutralize a measured volume of an acid

• Requires a few drops of indicator such as phenolphthalein to identify the endpoint

Titration Endpoint

When the moles of base = the moles of acid

• The concentration of base is known

• The volume of the base used to reach the endpoint is measured

• The molarity of the acid is calculated using the neutralization equation for the reaction

Indicator: Phenolphthalein

• In relation to titration

Added to identify the endpoint of a titration

• Turns pink when a solution is neutralized

What is the molarity of an HCl solution if 18.5mL of 0.225M NaOH is needed to neutralize 0.0100L of HCl?

Always start with the volume that the M is given with

18.5 mL NaOH(1 L NaOH/1000 mL NaOH)(0.225 mol NaOH/1 L NaOH)(1 mol HCl/1 mol NaOH) = 0.00416

0.00416 mol HCl/0.0100 L HCl = 0.416 M HCl

Buffers

A buffer solution maintains the pH by neutralizing small amounts of added acid or base

• An acid must be present to react with an OH⁻ added and a base must be present to react with any H₃O⁺ added

• Weak acid-base pairs

Components of a buffer

Contains a combination of acid-base conjugate pairs, a weak acid, and a salt of its conjugate base

• Equal concentrations of weak acid and salt

Function of a weak acid in a buffer

If a small amount of base is added, it is neutralized by the acid which shifts the equilibrium in the direction of the products

Function of a conjugate base in a buffer

When a small amount of acid is added, the additional H₃O⁺ combines with the acid ion, causing the equilibrium to shift in the direction of the reactants

• The acid produced contributes to the available weak acid

What chemical species affect pH?

H₃O⁺/H⁺ and OH⁻

How does a buffer work?

• Added acid (H₃O⁺)

• HA ⇌ A- + H₃O⁺

The added acid combines with the conjugate base and creates more of the original HA acid

• H₃O⁺ doesn’t change = no pH change

How does a buffer work?

• Added base (OH⁻)

• HA ⇌ A- + H₃O⁺

When a base is added, the H⁺ in the HA combines with the OH⁻ base to form water, and what’s left of the HA is A⁻, creating more of the conjugate base

• H₃O⁺ doesn’t change = no pH change

In the body, what pH do buffers tend to maintain?

7.4

Can buffers soak up limitless amounts of acid or base, keeping the pH constant no matter what?

Nope, they have limits too

• Any contamination may also mess it up

Which ones make a buffer?

1. HF and NaF

2. NH₄⁺ and NH₃

3. HBr and Br

4. HCl and HClO

1. HF and NaF = yes

   Whenever Na is in one of these, you simply cancel it out because it always ionizes, so you are left with HF and F⁻ which is weak and an acid-base pair

2. NH₄⁺ and NH₃ = yes

   It’s a weak acid-base pair

3. HBr and Br = no

   It’s an acid-base pair, but it’s strong, not weak

4. HCl and HClO = no

   It’s not an acid-base pair, since they do not differ by only an H⁺

What are the pH values of

• Pure water

• Blood plasma

• 1M NaOH solution

• 1 M HCl solution

• Urine

• Drinking water

There are a ton more, look at the slides

Pure water = 7.0

Blood plasma = 7.4

1M NaOH solution = 14.0

1 M HCl solution = 0.0

Urine = 6.0

Drinking water = 7.2

What is the pH of a 0.260 M solution of KOH?

-log(0.205) = the pOH value which is 0.602

BUT we need the pH → pOH + pH = 14

14 - 0.602 = 13.398, 3 sigfigs bc the og pH had three decimal places = 13.4

The pH of an acidic solution is 2.11, what’s the [H⁺]?

10⁻²¹¹ (exponent = -2.11)

→ make sure it’s in scientific notation 7.8 x 10⁻³ M

Predict whether the reaction has more products or reactants

• H₃PO₄(aq) + H₂O(l) ⇌ H₃O⁺(aq) + H₂PO₄⁻(aq)

• CO₃²⁻(aq) + H₂O(l) ⇌ OH⁻(aq) + HCO₃⁻(aq)

H₃PO₄(aq) + H₂O(l) ⇌ H₃O⁺(aq) + H₂PO₄⁻(aq)

weaker acid and base stronger acid and base

More reactants = strong to weak

CO₃²⁻(aq) + H₂O(l) ⇌ OH⁻(aq) + HCO₃⁻(aq)

weaker acid and base stronger base and acid

More reactants

What is the formula for acetic acid?

• Is it weak or strong

HC₂H₃O₂ or CH₃COOH

• Weak acid

What is an indicator

• Literally

An organic molecule