Acids and Bases

Properties of Acids In aqueous solutions, they conduct electricity They are ______________ Metals above H2 on Table J will react with acids to produce hydrogen gas and a salt Will Mg react with HCl? _________ Will Cu react with HCl? _________ Acids cause color changes in acid-base indicators Acids react with bases to form a salt + H2O _________ sour taste pH levels __________ Common examples:___________________________ Table K

Properties of Bases In aqueous solutions, bases conduct electricity  They are also ______________ Bases cause color changes in acid-base indicators. Bases react with acids to form a salt + H2O!! ____________________________ bitter taste pH levels ___________ Common examples: ____________________ Table L

Properties of Salts Like acids and bases, salts are also _________________ Salts are formed as a product of neutralization reactions Neutral  most have a pH of _______ Salts are ionic compounds ( _______+ _______) Examples:

The Arrhenius Theory Svante Arrhenius Based his ideas on the fact that aqueous solutions of acids and bases are electrolytes said that the properties of acids and bases are because of their H+ and OH- ions

Arrhenius Acids An acid is a substance that has H and releases (yields) H+ ions as the only positive ions in solution Adding an acid to an aqueous solution ___________ the concentration of H+ ions in the solution The H+ ions attach to H2O, forming hydronium ions _________ Examples: ___________________ Can be monoprotic or diprotic________

Arrhenius Bases A base is a substance that has OH and releases (yields) OH- ions as the only negative ions in an aqueous solution Adding a base to an aqueous solution __________ the concentration of OH- ions in the solution Examples: __________________

Naming Acids, Bases and Salts Bases and salts are ionic, and are named the usual way. Binary acids: hydro________ic acid Ternary acids: _________-ic acid (if poly is –ate) _______________ous acid (if poly is –ite)

The Bronsted-Lowry Theory Arrhenius’ theory had some limitations, so they expanded the definitions of acids and bases An acid is an H+ donor A base is an H+ acceptor H+ is simply a proton which interacts strongly with nonbonding e- on the oxygen of a water molecule Interaction results in a hydronium ion H3O+

Bronsted-Lowry Acids and bases can be defined in terms of ability to donate protons Acid: proton donor Base: proton acceptor

Bronsted- Lowry HCl(g) + H2O(l)  H3O+(aq) + Cl- (aq) EXAMPLE HCl: _______ H2O: _______

Comparing the two theories NH3(aq) + H2O(l)  NH4+(aq) + OH- (aq) According to Arrhenius, NH3 is a base because adding it to water increases the hydroxide ion concentration Bronsted-Lowry classifies it as a base because it accepts a proton from H2O, and the water acts as an acid because it donates a proton

Conjugate Acid-Base Pairs In any acid-base equilibrium, both the forward and reverse reactions involve proton transfers Example: HX(aq) + H2O(l)  X- + H3O+ (aq) Forward reaction, HX donates a proton, it is an acid and H2O is the base Reverse reaction, H3O+ is the acid and X- is the base An acid and base that differ only in presence or absence of a proton is called a Conjugate acid-base pair

Conjugate Acid-Base Pairs HNO2(aq) + H2O(l)  NO2- (aq) + H3O+(aq) What are the conjugate acid base pairs in the above reaction?

Amphoteric Substances Substances that can act as an acid or base Ex: H2O With ammonia it’s an acid With HCl it’s a base

Auto-ionization of Water Because water is amphoteric, it can actually donate a proton to another water and vice-versa This is called auto-ionization No individual molecule remains ionized for long, and at room temperature, only about 1 in 109 molecules is ionized at any given moment

Auto-ionization of Water Because it is an equilibrium process, we can write an equilibrium expression Keq= [H3O+][OH-] [H2O]2 This becomes the ion product constant for water (Kw) Kw at room temperature = 1 x 10-14 [H+] = 1x10-7 [OH-] = 1x10-7

Auto-ionization of Water A solution in which [H+] = [OH-] is considered neutral [H+] > [OH-] acid [H+] < [OH-] base When the concentration of one increases, the concentration of the other decreases so their product is always 1 x 10-14

Calculations Calculate the [H+] in a solution in which [OH-] is 0.010 M [H+] = 1 x 10-14 = 1 x 10-14 = 1.0 x 10-12M [OH-] .010 M Is the solution acidic or basic?

Calculations Calculate the [H+] in a solution in which the [OH-] is 2.0 x 10-9 [H+] = 1 x 10-14 = 1 x 10-14 = 5.0 x 10-6 M [OH-] 2.0 x 10-9 Is the solution acidic or basic?

The pH scale The concentration of [H+] in aqueous solution is usually very small, so the concentration is expressed as pH ( the negative logarithm in base 10) of [H+] pH = -log [H+] Calculate the pH of a solution with [H+] concentration of 1.0 x 10-7 M -log(1.0 x 10-7 M) = -(-7) = 7

Calculations with pH (regents) Examples: Determine the pH given the [H+] concentration 1. If [H+] = 1 x 10-4 pH = ___________ 2. If [H+] = 1 x 10-11 pH = ___________ 3. If [H+] = 1 x 10-7 pH = ___________ 4. What is the pH of a 0.00001M HCl solution? _________ 5. What is the pH of a 0.000000001M solution? _________ 6. What is the pH of a 0.0001M solution? _________

Calculations with pH A sample of lemon juice has [H+] = 3.8 x 10-4 M. What is the pH? What is the pH of a window cleaning solution with an [H+] = 5.3 x 10-9 Calculate the pH of a solution in which [OH-] is 0.010 M Calculate the pH a solution in which the [OH-] is 2.0 x 10-9 *rules for sig figs and logs: #decimal places = # sig figs in original #

Calculations with pH A sample of apple juice has a pH of 3.76. Calculate [H+] pH= -log[H+] =3.76 Thus : log [H+] = -3.76 Use antilog or INV Ans: 1.7 x 10-4 M

Calculations with pH A solution formed by dissolving an antacid tablet has a pH of 9.18. Calculate the [H+] Calculate the [OH-]

Other “p” values pOH = -log[OH-] pH + pOH = -log Kw = 14.00

Calculations with pH Strong acids are acids which completely ionize, and there are relatively few. HCl, HBr, HI,HClO3, HNO3 and HClO4 are STRONG monoprotic acids H2SO4 is a STRONG diprotic acid. Memorize the STRONG acids [H+] concentrations are assumed to be = to the concentration of the solution because complete ionization

Calculations with pH What is the pH of a 0.040M solution of HNO3?

Strength of Acids and Bases The principles of equilibrium apply to weak acids and bases because they only partially dissociate into ions. The dissociation of acetic acid into acetate ions and hydronium ions is shown as: Chapter 8 © 2011 Pearson Education, Inc.

Strength of Acids and Bases The Equilibrium Constant, Ka Weak acids dissociate much less than 100% and have an equilibrium constant called an acid dissociation constant, Ka. The strength of a weak acid can be determined from the Ka value. The larger the Ka value, the stronger the acid (the more protons dissociated). Refer to the table distributed in class Chapter 8 © 2011 Pearson Education, Inc.

Weak acids and bases This table shows Ka values for substances acting as weak acids. Chapter 8 © 2011 Pearson Education, Inc.

Strength of acids and bases Using Ka to calculate pH Using the value of Ka and the initial concentration of the solution, we can determine the [H+] There are four steps to solving a Ka problem- RICE (reaction, initial, change, equilibrium)

Ka Problems Calculate the pH of a 0.20 M solution of HF. We need to find H+ concentration before we find pH Reaction: HF(aq)  H+(aq) + F- (aq) Initial 0.20 X x Change -xM +xM Equilibrium (0.20-x)M xM

Ka Problems Substitute values from table into equilibrium expression Because Ka is very small, we can presume that x is very small compared to M, therefore we say 0.20M-x ~ 0.20 Ka = (x)(x) = 3.5 x 10-4 0.20 Solve x = 8.37 x 10-3 pH = - log [H+] pH= -log 8.37 x 10-3 pH = 2.078

Ka problems The Ka of HCN is 4.9 x 10-10 Calculate the pH of a 0.20 M solution

Ka problems Diprotic and polyprotic acids will have two or more Ka because each proton takes more work. The subsequent Ka values will be smaller.

The Meaning of pH (regents) pH measures the H+ ion concentration in a solution. If….. [H+] = [OH-] ______________________ [H+] > [OH-] ______________________ [H+] < [OH-] ______________________ An increasing pH means the H+ ion concentration is _________ A decreasing pH means the H+ ion concentration is _________ The pH scale is logarithmic Each change of a single pH unit signifies a TENFOLD change in the concentration of H+ ions A solution with a pH of 4 is 10x more acidic than a pH of ___ A solution with a pH of 7 is 100x more acidic than a pH of __

Neutralization Reactions When an acid reacts with a base to produce salt and water It occurs when there are the same amount of H+ ions as OH- ions General formula: acid + base  water + salt

Writing Neutralization Reactions Take the H+ from the acid and the OH- from the base and combine them to form H2O Then combine the other ions to form the salt Remember to look at oxidation numbers!!!! Make sure the equation is balanced When a solution is neutral, the moles of H+ = moles OH- Ex: HCl + KOH  ___________________ Ex: H2SO4 + NaOH  ___________________

Acid-Base Titration A laboratory technique used to achieve neutralization between an acid and a base It is used to find the unknown molarity of an acid or a base by slowly adding measured volumes of an acid or base of known molarity until neutralization occurs End point: The point when the indicator changes color nHMAVA = nOHMBVB

Examples How many milliliters of 4.00M NaOH are required to exactly neutralize 50.0mL of a 2.00M solution of HNO3? If it takes 55mL of 0.1M NaOH solution to neutralize 450mL of an H2SO4 solution of unknown concentration, what’s the molarity of the acid?

Table M Examples You put bromcresol green in a solution  it turns blue Then you put bromthymol blue in the same solution  it turns yellow What is the pH of the solution? _____________ Is the solution acidic or basic? _____________ A solution turns yellow with thymol blue and blue with bromthymol blue. What is the pH of th\e solution? ____________ Is the solution acidic or basic? ____________