Acids and Bases

H+ water Acetic acid Citric acid Tannic acid Formic acid Acids H+ Acid – a compound that produces ions when dissolved in Examples: Vinegar – Lemon juice – Tea – Ant venom – water Acetic acid Citric acid Tannic acid Formic acid

Sour Corrosive red Aqueous electrolytes bases H2O salt Properties of Acids Sour taste Turns litmus paper Reacts with metals to form gas solutions of acids are (must be mixed with water!) Reacts with to form and Corrosive red H2 Aqueous Litmus paper is an indicator. Electrolytes can conduct electricity. “Salt” is not always NaCl! It can be other ionic compounds. electrolytes bases H2O salt

Properties of Acids Sugar, corn syrup, modified corn starch, citric acid, tartaric acid, natural and artificial flavors, yellow 5, yellow 6, red 40, blue 1 What ingredients make these… so sour?

“hydro-” root “-ic” hydro chlor ic acid Naming Binary Acids H + one other element Begin with Use the of the element name Add the suffix HCl “hydro-” root “-ic” hydro chlor ic acid

Hydrobromic acid Hydroiodic acid Hydrofluoric acid Naming Binary Acids Hydrobromic acid HBr  HI  HF  Hydroiodic acid Hydrofluoric acid

polyatomic ending -ate -ic -ous -ite Nitrate Nitric acid Naming Ternary Acids H + polyatomic ion Begin with ion without the Add suffix if there was an Add suffix if there was an HNO3  polyatomic ending -ate -ic -ous -ite Nitrate Nitric acid

Chloric acid Phosphorous acid Carbonic acid Naming Ternary Acids Chloric acid HClO3  H3PO3  H2CO3  Phosphorous acid Carbonic acid

2 Strong ionize H+ HCl HBr HI O H 3 – 1 = 2 4 – 2 = 2 Strength of Acids Strong ionize acids – completely in water (create a lot of ) 3 binary acids Ternary acids Strong if # of atoms - # of atoms ≥ H2SO4 HNO3 H+ HCl HBr HI 2 O H 3 – 1 = 2 4 – 2 = 2

1 2 – 1 = 1 3 – 3 = 0 Weak slightly aqueous O H Strength of Acids acids – ionize only in solution Binary acids – all others not listed above Ternary acids Weak if # of atoms - # of atoms ≥ H3PO3 HNO2 aqueous 1 O H 2 – 1 = 1 3 – 3 = 0

OH- (hydroxide) water Magnesium hydroxide Sodium hydroxide Bases OH- Base – a compound that produces ions when dissolved in Examples: Milk of Magnesia – neutralizes stomach acid Drain cleaner– (hydroxide) water Magnesium hydroxide Sodium hydroxide

Bitter Slippery blue Aqueous electrolytes acids H2O salt Properties of Bases Bitter taste Turns litmus paper solutions of bases are (must be mixed with water!) Reacts with to form and Slippery blue Aqueous electrolytes Litmus paper is an indicator. Electrolytes can conduct electricity. “Salt” is not always NaCl! It can be other ionic compounds. acids H2O salt

polyatomic Sodium hydroxide Calcium hydroxide Potassium hydroxide Naming Bases polyatomic Use the same rules as for ions (name the cation, then name the anion) NaOH  Ca(OH)2  KOH  Sodium hydroxide Calcium hydroxide Potassium hydroxide

Strong ionize OH- 1 2 Be Weak ionize slightly Strength of Bases Strong ionize bases – completely in water (create a lot of ions). All hydroxides with groups and metals (except ). bases - only All bases not listed above as strong. OH- 1 2 Be Weak ionize slightly

Arrhenius H H+ / H3O+ HCl + H2O  H3O+ + Cl- Arrhenius OH OH- Arrhenius Theory Arrhenius H An acid must contain a and ionize in water to produce An base must contain a and dissociates in water to produce H+ / H3O+ HCl + H2O  H3O+ + Cl- Arrhenius OH OH- NaOH  Na+ + OH-

Disadvantages of Arrhenius Theory OH- Only compounds with can be classified as a base. What about ammonia, ? Can only be applied to reactions that occur in Would classify some compounds as acids, such as NH3 water incorrectly CH4

Arrhenius Acids and Bases Classify each of the following as an Arrhenius acid (A – acid) or base (A – base). Ca(OH)2  HBr  H2SO4  LiOH  A - base A - acid A - acid A - base

Bronsted-Lowry Theory acid A Bronsted – Lowry is any substance that can a HCl + H2O  H3O+ + Cl- donate H+ base accept H+ B-L base conjugate base B-L acid conjugate acid Donates H+ Accepts H+

Bronsted-Lowry Theory Let’s look at the reverse reaction. Cl- + H3O+  H2O + HCl B-L base B-L acid conjugate acid Donates H+ conjugate base Accepts H+

Bronsted-Lowry Theory Conjugate acid – formed when a accepts a H+ from an acid. base – a that remains after an acid gives up a H+. Conjugate acid – base pair – 2 substances related to each other by the of a single H+. base Conjugate particle accepting/ donating

Types of Acids Type # of H+ donated Example Monoprotic 1 HNO3 Diprotic Defined by how many H+ they can donate. Type # of H+ donated Example Monoprotic 1 HNO3 Diprotic 2 H2SO4 Triprotic 3 H3PO3

Bronsted-Lowry Theory Identify the acid, base, conjugate acid, and conjugate base. conjugate acid B-L acid B-L base HNO3 + H2O  H3O+ + NO3- Donates H+ conjugate base Accepts H+

Bronsted-Lowry Theory Give the formula and name of the conjugate base of the following B-L acids. (After the B-L acid donates a H+) HI  HCO3-  I- Iodide ion Since we take away a +, make the ion more – CO32- carbonate ion

Bronsted-Lowry Theory Give the formula and name of the conjugate acids of the following B-L bases. (After the B-L base accepts a H+) H2PO4-  ClO3-  H3PO4 phosphoric acid Since we add a +, make the ion more + HClO3 chloric acid

self-ionization conjugate acid B-L acid B-L base conjugate base Acidity/Basicity Water can sometimes act as a B-L acid and sometimes as a B-L base. The of water: H2O + H2O  H3O+ + OH- self-ionization conjugate acid B-L acid B-L base conjugate base

[H+] [OH-] concentration 1 x 10-14 Inverse Acidity/Basicity This reaction occurs to a very small extent: = = 1 x 10-7 M [ ] means [H+] x [OH-] = relationship [H+] [OH-] concentration 1 x 10-14 Inverse

pH = 0 pH = 7 pH = 14 acid base neutral Acidity/Basicity pH = 0 pH = 7 pH = 14 [H+] [OH-] acid base neutral

Neutral [H+] = [OH-] Acidic > Basic < Acidity/Basicity Neutral [H+] = [OH-] Acidic > Basic <

pH [H+] are often small, so the pH scale is easier to use to represent acidity and basicity. pH range is from to log 102 = log 10-3 = 14 pH = -log [H+] 2 -3

neutral [H+] [OH-] 1 x 10-7 - (-7) = 7 pH neutral In water, a solution, = = pH = [H+] [OH-] 1 x 10-7 pH = negative log [H+] So… take the exponent and change the sign! - (-7) = 7

pH is < 7 Solution is acidic pH is = 7 Solution is neutral Solution is basic

pH = - (exponent) = -(-5) = 5 If [H3O+] = 1.0 x 10 –5 M, what is the pH? Is the solution basic, neutral, or acidic? Same as [H+] pH = - (exponent) = -(-5) = 5 Because < 7

pH = - (exponent) = -(-12) = 12 If [H3O+] = 1.0 x 10 –12 M, what is the pH? Is the solution basic, neutral, or acidic? pH = - (exponent) = -(-12) = 12 Because > 7

pH pH = - (exponent) 2 = - (exponent) 2 = - (-2) [H3O+] = 1 x 10-2 Given that a solution has a pH of 2.0, determine the [H3O+]. pH = - (exponent) 2 = - (exponent) 2 = - (-2) [H3O+] = 1 x 10-2

pOH = -log [OH-] 1 x 10-14 pH + pOH = 14 Similar to pH, there is also pOH. Because [H+] x [OH-] = pOH = -log [OH-] 1 x 10-14 pH + pOH = 14

pH/pOH [H+] 1 pH 2 pOH 14 10 [OH-] 10-14 10-8 Ex. battery acid stomach acid tomatoes milk 10-2 10-4 10-6 4 6 12 8 10-12 10-10

pH/pOH [H+] 10-12 pH 8 14 pOH 4 [OH-] Ex. seawater detergent ammonia oven cleaner 10-8 10-10 10-14 10 12 2 6 10-6 10-4 10-2 1

pOH = - (exponent) = -(-10) = 10 pH/pOH If [OH-] = 1.0 x 10 –10 M, what is the pOH? What is the pH? Is the solution basic, neutral, or acidic? pOH = - (exponent) = -(-10) = 10 pH + pOH = 14 pH + 10 = 14 pH = 4

pH = - (exponent) = -(-6) = 6 pH/pOH What is the pH and the pOH for 1.0 x 10 –6 M HF? pH pOH pH = - (exponent) = -(-6) = 6 pH + pOH = 14 6 + pOH = 14 pOH = 8

pH/pOH pH + pOH = 14 9 + pOH = 14 pOH = 5 pOH = - (exponent) Given that a solution has a pH of 9.0, determine the [OH -] and the pOH. pOH [OH -] pH + pOH = 14 9 + pOH = 14 pOH = 5 pOH = - (exponent) 5 = - (exponent) 5 = - (-5) [OH-] = 1 x 10-5

Review: Acids  Conjugate Bases Acids LOSE H+ to become conjugate bases. This is a H atom. When a H+ is lost from an acid, this (-) electron remains. The (+) proton is taken with the H. - + o

Review: Acids  Conjugate Bases What is the conjugate base for the acid HBr? HBr  H+ + Br- H Br Proton is kept by H. Electron is left by H. + - H + Br Conjugate base

Review: Acids  Conjugate Bases What is the conjugate base for the acid HNO2? HNO2  H+ + NO2- NO2 H Proton is kept by H. Electron is left by H. + NO2 - H + Conjugate base

Review: Acids  Conjugate Bases What is the conjugate base for the acid HSO3-? HSO3-  H+ + SO32- - SO3 H Electron is left by H (added to the one that was there already). Proton is kept by H. + SO3 2- H + Conjugate base

Review: Bases  Conjugate Acids Bases GAIN H+ to become conjugate bases. What is the conjugate acid for CN-? CN- + H+  HCN - + CN H + The + and the – cancel out in the final molecule. CN H

Review: Bases  Conjugate Acids What is the conjugate acid for NH3? NH3 + H+  NH4+ + NH3 H + There is no – on the NH3 to cancel the + from H, so the final molecule is positive. + NH3 H

-log[H+] 10-pH -log[OH-] 10-pOH 14 pH/pOH cont. Review: pH = [H+] = pOH = [OH-] = pH + pOH = -log[H+] 10-pH -log[OH-] 10-pOH 14

Neutralization Reactions What happens with you mix an acid with a base? A ____________________________________________ reaction Neutralization Reactions What happens with you mix an acid with a base? A reaction HCl + NaOH  + Products are always a ( and ) and This is called a reaction double replacement H2O NaCl salt nonmetal metal water neutralization

Neutralization Reactions Write the balanced chemical equation for the neutralization reaction between: nitric acid and potassium hydroxide + + HNO3 KOH  H2O KNO3 Must use criss-cross to make salt Must balance equation This equation is already balanced

Neutralization Reactions Write the balanced chemical equation for the neutralization reaction between: sulfuric acid and magnesium hydroxide 2 + + H2SO4 Mg(OH)2 H2O MgSO4  Must balance equation

Neutralization Reactions acid + base  salt + water

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