I. Measuring Concentration

1. Molarity (M): the number of moles of solute dissolved per liter of solution; also known as molar concentration

2. Molality (m): the ratio of the number of moles of solute dissolved in one kilogram of solvent; also known as molal concentration

3. We will always use molarity in this class.4. M = # of moles of solute / # of liters of solution5. To calculate the mass of a solute when given the

volume and concentration of solution, use the following setup:

6. To calculate the volume of a solution when given concentration and mass of solute, use the following setup:

II. Acids and Bases

1. Acidic solutions taste sour, turn blue litmus paper pink, and produce more hydrogen ions than hydroxide ions.

2. Basic solutions taste bitter, feel slippery, turn pink litmus paper blue, and produce more hydroxide ions than hydrogen ions.

II. Acids and Bases3. Arrhenius model: states that an acid is a

substance that contains hydrogen and ionizes to produce hydrogen ions in aqueous solution (ex: HCl). A base is a substance that contains a hydroxide group and dissociates to produce a hydroxide ion in aqueous solution (ex: NaOH).

4. HCl(g) H+(aq) + Cl-(aq)5. NaOH(s) Na+(aq) + OH-(aq)

II. Acids and Bases

6. Although the Arrhenius model is useful in explaining many acidic and basic solutions, it has some shortcomings.

7. For example, ammonia (NH3) does not contain a hydroxide group, yet it produces hydroxide ions in solution and is a well known base.

II. Acids and Bases

8. Bronsted-Lowry model: states that an acid is a hydrogen-ion donor and a base is a hydrogen-ion acceptor

9. For example, when a molecule of acid, HX, dissolves in water, it donates an H+ ion to a water molecule. The water molecule acts as a base and accepts the H+ ion.HX(aq) + H2O(l) <--> H3O+(aq) + X-(aq)

10. In the example above, the water molecule becomes an acid (H3O+) by accepting the H+ ion.

II. Acids and Bases

11. In the example above, the acid (HX) becomes a base because it can now accept a positive hydrogen ion.

12. Conjugate acid: the species produced when a base accepts a hydrogen ion from an acid

13. Conjugate base: the species that results when an acid donates a hydrogen ion to a base

14. Amphoteric: water and other substances that can act as both acids and bases

III. Naming Acids and Bases

1. Bases are named like other ionic compounds (e.g., NaOH = sodium hydroxide)

2. Two types of acids: binary and oxyacids3. Binary acid contains hydrogen and one

other element.hydro + nonmetal root + ic acid(e.g., HCl = hydrochloric acid)

III. Naming Acids and Bases

4. Oxyacid contains hydrogen and an oxyanion (polyatomic ion with oxygen in it)

a. Name the polyatomic ion.b. Replace “ate” with “ic”, “ite” with “ous.”c. Change nonmetal root for pronunciation.d. Add “acid” to the name.(e.g., H2SO4 = sulfuric acid)

IV. Acid/Base Strength

1. Strong acids/bases: acids/bases that dissociate (ionize) completely in aqueous solutions

2. Weak acids/bases: acids/bases that dissociate (ionize) only partially in aqueous solutions

3. Reactions involving weak acids/bases reach a state of equilibrium where the forward and reverse reactions occur at equal rates. This is represented by a double arrow.

IV. Acid/Base Strength

4. The equilibrium constant, Keq, provides a measure of the degree of ionization of an acid/base.

5. For the following reaction:HCN(aq) + H2O(l) <-> H3O+(aq) + CN-(aq)

Keq =[H3O+][CN-]

[HCN][H2O]

IV. Acid/Base Strength

Keq [H2O] = Ka =

6. The concentration of liquid water in the denominator of the expression is considered to be constant in dilute aqueous solutions, so it can be combined with Keq to give a new equilibrium constant, Ka.

7. Acid ionization constant (Ka): the value of the equilibrium constant expression for the ionization of a weak acid

[H3O+][CN-]

[HCN]

IV. Acid/Base Strength

8. The value of Ka indicates whether reactants or products are favored at equilibrium.

9. The weakest acids have the smallest Ka values because their solutions have the lowest concentrations of ions and the highest concentrations of un-ionized acid molecules.

IV. Acid/Base Strength

10.Base ionization constant (Kb): the value of the equilibrium constant expression for the ionization of a base

11.CH3NH2 (aq) + H2O(l) <-> CH3NH3+(aq)

+ OH-(aq)

12.Kb =[CH3NH3

+] [OH-]

[CH3NH2]

V. pH and pOH

1. Ion Product Constant for Water (Kw): the value of the equilibrium constant expression for the self-ionization of waterH2O(l) <-> H+(aq) + OH-(aq)

2. Keq =

3. Keq [H2O] = [H+][OH-]4. Kw = [H+][OH-]5. In pure water at 298 K, [H+] and [OH-] are both equal

to 1.0 x 10-7 M.

[H+][OH-]

[H2O]

V. pH and pOH

6. Therefore, the value of Kw is 1.0 x 10-14.7. This is also the product of [H+] and [OH-] for other

solutions.8. Because concentrations of H+ ions are often small

numbers expressed in scientific notation, chemists adopted an easier way to express H+ ion concentration using a pH scale based on common logarithms.

9. pH: the negative logarithm of the hydrogen ion concentration of a solution

10. pH = -log [H+]

V. pH and pOH

11. If [H+] equals 1.0 x 10-7 M, pH is calculated as follows:pH = -log (1.0 x 10-7)pH = -(log 1.0 + log 10-7)pH = -[0 + (-7)] = 7.00

12. pH > 7 is basic/alkalinepH = 7 is neutralpH < 7 is acidic

V. pH and pOH

13.Sometimes chemists find it convenient to express the basicity, or alkalinity, of a solution on a pOH scale that mirrors the relationship between pH and [H+].

14.pOH: the negative logarithm of the hydroxide ion concentration of a solution

15.pOH = -log [OH-]

V. pH and pOH

16.pOH > 7 is acidicpOH = 7 is neutralpOH < 7 is basic

17.pH + pOH = 14.00

VI. Neutralization1. Neutralization Reaction: a reaction in

which an acid and a base react in aqueous solution to produce a salt and water.

2. Salt: an ionic compound made up of a cation from a base and an anion from an acid.

3. Neutralization is a type of double replacement reaction.

VI. Neutralization4. Titration: a method for determining the

concentration of a solution by reacting a known volume of the solution with a solution of known concentration.

5. If you wished to find the concentration of an acid solution, you would titrate the acid solution with a solution of base of known concentration.

6. The solution of known concentration is called the standard solution.

VI. Neutralization7. Measured volumes of the standard solution

are added slowly and mixed into the solution in the beaker.

8. This process continues until the reaction reaches the stoichiometric point (equivalence point), which is when the moles of H+ ion from the acid equal the moles of OH- ion from the base.

9. A buret is used to dispense the standard solution.

VI. Neutralization10. For convenience, chemists often use a chemical

dye rather than a pH meter to detect the equivalence point of an acid-base titration.

11. Acid-base indicators: chemical dyes whose colors are affected by acidic and basic solutions

12. End point: the point at which the indicator used in a titration changes color

13. The equation used in a titration is as follows:MAVA = MBVB

M = molarity A = acid V = volume B = base

VI. Neutralization14. Buffer: a solution that resists change in pH

when limited amounts of acid or base are added

15. A buffer is a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.

16. The mixture of ions and molecules in a buffer solution resist changes in pH by reacting with any hydrogen ions or hydroxide ions added to the buffered solution.

VI. Neutralization

17.Buffer capacity: the amount of acid or base a buffer solution can absorb without a significant change in pH