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 \rightarrow 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):

   ➢ 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 \rightarrow H3O+ + Br-

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

1. KOH + H2O \rightarrow 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 \rightarrow 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) \rightarrow 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 \rightarrow H3O+ + Cl- (water acts like a base)

2. H2O + NH3 \rightarrow 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) \rightarrow NaCl(aq) + H2O(l)

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

3. Neutralization is a double-replacement reaction.

4. Practice neutralization reactions and balancing:

   ➢ HNO3 + KOH \rightarrow

   ➢ HCl + Zn \rightarrow

   ➢ H2SO4 + LiOH \rightarrow

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.

   ➢ Example: Phenolphthalein will show colorless in an acid (pH<7), 1 10 50 60 then start to turn pink between pH values of 8.2 – if I adding base, and it will become fully once the my solution is greater than when enough base added. When using this table every value before first number in range for a specific indicator be initial listed color. Every after last final The two color starts change.

F. can calculated : because depends on concentration H+ solution.

1. Example: What whose x 10-3 M?

2. 3.0 10-10

3. acid rain that has 5.6? 10-6

G. Because strong acids completely ionize (dissociate) Their also ions one mole HCl yield Cl- ions. Therefore, 0.01M would have [H+]=“0.01M” [Cl-]=“0.01M.”

IX. Titrations

A. Titration process by which molarity an unknown or determined slowly neutralizing with known molarity. Note*- you need know volumes solutions. neutralized each other, moles (H+) equals (OH-). This called titration’s endpoint.

B. Equation (Table T) MaVa=MbVb

➢ b=“base,” M → solution/liters solution, V=volume Note* - Volume milliliters liters just consistent same unit both volume base. A 30mL sample 10mL 1.5 NaOH solution? 0.5

C. Modified (ma)MaVa=(mb)MbVb Same variables are original titration equation except:

o ma=“moles” releases mb=“moles” OH-