Unit 7: Acids, Bases, and Salts

I. Properties of Acids

A. Taste sour

B. Most acids contain a hydrogen group (dissociates to give off H+ ion, also known as a proton)

C. Reacts with certain metals to produce hydrogen gas/H2(g) (the metal must be above H2 on Table J)

1. Example: Zn + HCl → ZnCl2 + H2 (metal + acid → salt + hydrogen gas)

➢ These reactions are generally not reversible because the hydrogen gas can escape, not allowing for

products to further react with each other.

D. The common acids are found on Table K (note the hydrogen group that all of these acids have)

1. The strong acids (memorize for AP Chem only):

➢ HI : hydroiodic acid

➢ HBr : hydrobromic acid

➢ HCl : hydrochloric acid

➢ HNO3 : nitric acid

➢ H2SO4 : sulfuric acid

➢ HClO4 : perchloric acid

2. For the above acids, note that hydrogen + a halogen (group 17) besides fluorine forms a strong acid.

The other three acids you should commit to memory and familiarize yourself with them. HF is not a

strong acid because the bond between H and F is strong (large difference in electronegativity and a

very short bond: shorter bonds are stronger).

3. Very common weak acid : acetic acid : HC2H3O2 also written has CH3COOH

➢ Dissociates (break apart) in water into C2H3O2

-/CH3COO- anion and H+ cation

II. Properties of Bases

A. Taste bitter

B. Feel slippery

C. Most of the time contains a hydroxide group (however, remember that NH3 (ammonia) is a base) and take

the form of a metal (Group 1 or Group 2) + hydroxide (OH-). Ex. Mg(OH)2

D. The common bases are found on Table L

1. The strong bases (memorize for AP Chem only):

➢ o Examples: LiOH, NaOH, Ca(OH)2

Any group 1 or group 2 metal hydroxides

III. Arrhenius Theory

A. An acid is a substance that yields only H+ ions as the only cations in solution.

1. HBr + H2O → H3O+ + Br-

B. A base is a substance that yields only OH- ions as the anions in solution.

1. KOH + H2O → K+ + OH-

C. Problems with the Arrhenius Model of Acids & Bases

1. Although the Arrhenius model is useful in describing many acids and bases, it does not describe them all

2. For example, NH3 contains no OH- ions, but it is a base

3. Not all compounds that have hydrogen are acids

IV. Hydronium (H3O+) vs. hydrogen ion (H+)

A. They are interchangeable because an acid is defined by the fact that it INCREASES the H+1 concentration

of an aqueous solution. H+ is also known as a proton because a hydrogen-1 atom has 1 proton, 1 electron,

and 0 neutrons. If hydrogen-1 loses an electron, forming an H+ cation, only a proton is left.

B. When H+1 combines with water it becomes H3O+1.

1. HCl + H2O → H3O+ + Cl-

V. Bronsted-Lowry Theory (aka the Alternate Acid-Base Theory or the Other Acid-Base Theory)

A. A Brönsted-Lowry acid: any molecule or ion that can lose an H+ to another species. It is a proton donor.

B. A Brönsted-Lowry base: any molecule or ion that can gain an H+ from another species. It is a proton acceptor.

C. In short, acids are proton (H+) donors and bases are proton (H+) acceptors.

D. Conjugate acids/conjugate bases

1. Take this example: NH3(aq) + H2O(l) → NH4

+(aq) + OH-(aq)

➢ NH3(aq) is the base because it accepts an H+ from H2O(l) which is the acid because it donates an H+ to

NH3(aq).

➢ NH4

+(aq) is the conjugate acid because it is now capable of donating an H+ (to become NH3) and

it is the species that resulted from the base

➢ OH-(aq) is the conjugate base because it is now capable of accepting an H+ (to become H2O) and

it is the species that resulted from the acid.

2. conjugate acid - species produced (the result) when a base accepts a hydrogen ion (H+)

3. conjugate base - species produced (the result) when an acid donates a hydrogen ion (H+)

E. Water is an amphiprotic (aka amphoteric) substance meaning it can act as either an acid or base depending on

whether it accepts (base) or donates (acid) a proton.

1. H2O + HCl → H3O+ + Cl- (water acts like a base)

2. H2O + NH3 → NH4

+ + OH- (water acts like an acid)

VI. Electrolytes

A. Compounds that dissociate (breaks apart) in water in order to form ions that can carry an electric current.

B. Acids, bases, and salts can dissociate into ions in solution. These ions can carry an electric current. The

stronger the acid or base, the more ions will be formed and the solution will be a better conductor of

electricity.

C. Strong Electrolyte = Completely dissociates into ions in water (first beaker in diagram below)

D. Weak Electrolyte = Coexist as ions and compounds in water. Does not fully dissociate into ions, leaving less

ions to carry current compared to strong acids (second beaker in diagram below)

VII. Neutralizations and salts

A. Neutralization Reaction - a reaction in which an acid and a base react in an aqueous solution to produce a

salt and water: Acid + Base → Salt + Water

1. Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

2. Example: H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2 H2O(l)

3. Neutralization is a double-replacement reaction.

4. Practice neutralization reactions and balancing:

➢ HNO3 + KOH →

➢ HCl + Zn →

➢ H2SO4 + LiOH →

B. Salts

1. All salts are ionic compounds (remember that polyatomic ions from Table E can also form salts)

2. The salt created from a neutralization can be used to predict whether the solution was acidic, basic, or neutral

3. If the salt is formed from a reaction of a strong acid with a weak base, the salt is considered acidic

➢ NH4Cl (the salt) comes from the reaction of HCl (strong acid) and NH4OH (weak base) therefore

this salt is considered acidic

4. If the salt is formed from a reaction of a weak acid with a strong base, the salt is considered basic

➢ NaC2H3O2 (the salt) comes from the reaction of NaOH (strong base) and HC2H3O2 (weak acid) therefore this

salt is considered basic

5. If the salt is formed from a reaction of a strong acid with a strong base, the salt is considered neutral

➢ K2SO4 (the salt) comes from the reaction of KOH (strong base) and H2SO4 (strong acid) therefore

this salt is considered neutral

6. Note* - In all three reactions above, water is also a product along with the salt and is not shown above

VIII. pH

A. pH stands for “pouvoir hydrogene” or hydrogen power = = how great is the concentration of hydrogen ions

(H3O+, H+) in the solution

B. The pH scale is a mathematical scale in which the concentration of hydronium (hydrogen) ions in a solution

is expressed as a number from 0 to 14.

1. pH of 7 is neutral. (H+ concentration = OH- concentration)

➢ concentration can be expressed as square brackets, “[ ]” around the solute

2. pH less than 7 is acidic. ([H+] > [OH-])

3. pH greater than 7 is basic. ([H+] < [OH-])

C. Each change in the pH value going by whole numbers is a change in 10x more or less acidic/basic.

1. Example: going from a pH of 5 to a pH of 3, there are 100x more H+ ions and 100x less OH- ions in the

solution than before

2. Example: going from a pH of 7 to 10, there are 1000x more OH- ions and 1000x less H+ ions in the solution

than before

D. Note* - Adding a stronger acid to a solution will LOWER the pH value. Adding a stronger base to a solution

will RAISE the pH value.

1. Example: What would happen to the pH of a solution containing 6.0 M HCl if 4.0 M HCl was added to it? The

pH value would decrease

E. Indicators - Acids and bases change the colors of indicators.

1. Indicators – changes colors according to what the approximate pH of the solution is.

2. Table M – shows the list of common indicators that chemists use.

➢ o Example: Phenolphthalein will show colorless in an acid (pH<7), then start to turn pink

between pH values of 8.2 – 10 if I start adding base, and then it will become fully pink

once the pH of my solution is greater than 10 when enough base is added.

When using this table every pH value before the first number in the range for a specific indicator

will be the initial listed color. Every pH value after the last number in the range for a specific

indicator will be the final listed color. The range between the two values is when the color starts to

change.

F. pH can be calculated using : pH = -log[H+] because the pH of a solution depends on the concentration of H+ in the

solution.

1. Example: What is the pH of a solution whose H+ concentration is 1 x 10-3 M? pH = 3

2. Example: What is the pH of a solution whose H+ concentration is 3.0 x 10-10 M? pH = 9.5

3. Example: What is the H+ concentration in acid rain that has a pH of 5.6? H+ = 2.5 x 10-6

G. Because strong acids completely ionize (dissociate) in solution. Their concentration will also be the concentration

of ions in solution.

1. Example: one mole of HCl will yield one mole of H+ ions and one mole of Cl- ions. Therefore, a 0.01M

solution of HCl would have [H+] = 0.01M and [Cl-] = 0.01M.

IX. Titrations

A. Titration – The process by which the molarity of an unknown acid or base can be determined by slowly neutralizing

it with a solution of a known molarity.

1. Note*- you will also need to know the volumes of the known and unknown solutions.

2. When the acid and base have neutralized each other, the number of moles of acid (H+) equals the number

of moles of base (OH-). This is called the titration’s endpoint.

B. Titration Equation (Table T)

1. MaVa = MbVb

➢ a = acid, b = base, M = molarity → moles of solution/liters of solution, V = volume

2. Note* - Volume can be in milliliters or liters – just be consistent using the same unit for both the volume

of the acid and the volume of the base.

3. Example: A 30mL sample of HCl is completely neutralized by 10mL of a 1.5 M NaOH solution. What

is the molarity of the HCl solution? 0.5 M HCl

C. Modified Titration Equation

1. (ma)MaVa = (mb)MbVb

➢ Same variables are original titration equation except:

o ma = moles of H+ ions the acid releases

o mb = moles of OH- ions the base releases

2. Example: How many milliliters of 1.0M NaOH will neutralize 50 milliliters of 0.6M H2SO4?

➢ 60 mL of NaOH

X. Naming Acids and Bases (not on the Regents!)

A. Rules of naming acids (name the original compound first then try to name the acid)

1. hydrogen ________ide ending: “ic acid”

➢ When H combines with one element, such as halogens (HCl) put “hydro” at the beginning of the

name, like hydro-chlor-, then, put “ic acid”: hydro-chlor-ic acid.

➢ Example: H2S = hydrogen sulfide (original compound) → hydrosulfuric acid (acid)

2. hydrogen ________ate ending:

➢ Drop hydrogen and -ate ending, and add “-ic acid.”

➢ Example: hydrogen sulfate (H2SO4) → sulfuric acid (H2SO4 (aq))

3. hydrogen ________ite ending:

➢ Drop hydrogen and -ite ending, and add “-ous acid.”

B. Rules for naming bases

1. Name metal part without any changes then name the hydroxide part without any changes

➢ NaOH = sodium hydroxide

➢ LiOH = lithium hydroxide

➢ Ca(OH)2 = calcium hydroxide

➢ KOH = potassium hydroxide