Chapter 19 Acid and Base Chemistry

Properties of Acids Acids effect indicators Blue litmus turns red Acids taste sour (tart) Acids effect indicators Blue litmus turns red Phenolphthalein is clear Acids have a pH lower than 7

Properties of Acids Acids are proton donors (hydrogen ion, H+) Acids react with active metals, produce H2 Acids react with metal carbonates and hydrogen carbonates, produce CO2 Acids neutralize bases, produce salt and H2O Nonmetallic oxides become acids in aqueous solutions

What are acids? Many definitions exist… Acids are compounds that give off hydrogen ions when dissolved in water. They are recognized by having H+1 as the cation. (formula begins with H) The H+1 ion attaches to a water molecules to give rise to H3O+1 ions. HYDRONIUM IONS

Protic Acids Monoprotic acids Diprotic acids Triprotic acids H3PO4 HCl H2SO4 HC2H3O2 H2CO3 HNO3 Mono- can donates 1 hydrogen Di – can donate 2 hydrogens Tri- can donate 3 hydrogen Polyprotic – includes Di- and Tri- protic acids

Acid Nomenclature Rules Naming-Review Binary acids – H+1 and another nonmetal (two elements) Prefix = Hydro __ Suffix = ic HCl = hydrochloric Ternary acids - H+1 and a polyatomic ion Suffix = ic (-ate ion) Suffix = ous (-ite ion)

Properties of Bases Bases feel slippery Bases effect indicators Bases taste bitter Bases feel slippery Bases effect indicators Red litmus turns blue Phenolphthalein turns pink Bases have a pH greater than 7

What are bases? Many definitions exist… BASES are compounds that dissociate to form hydroxide ions when dissolved in water. 0H-1 They are recognized by having 0H-1 as the anion. (formula end with 0H) Bases are proton (hydrogen ion, H+) acceptors Oxides of metallic elements usually form basic solutions

Ions in Solution Pure water is neutral because equal numbers of hydrogen ions (H+1) and hydroxide ions (0H-1) are always present.

Arrhenius Model Ex: HCl + H2O  H3O+ (aq) + Cl – (aq) Svante Arrhenius noticed that aqueous solutions of acids and bases were good conductors (electrolytes). Traditional definition; requires water & focuses on products. ACID- Increases hydrogen ions (H+ or H3O+) in water (ionizes) Ex: HCl + H2O  H3O+ (aq) + Cl – (aq) BASE - Increases hydroxide ions in water (OH-) (dissociates) Ex: NaOH water  Na + (aq) + OH - (aq)

Bronsted-Lowry Model Bronsted –Lowry Acid = Proton Donor Takes into account bases that do not contain a hydroxide group, such as ammonia. Bronsted –Lowry Acid = Proton Donor Bronsted-Lowry Base = Proton Acceptor What is a proton = These donors and acceptors produce partners called _____________ H+ ion lost from an acid Conjugates

Bronsted-Lowry examples HNO3 + H2O  H3O + + NO3 – Acid Base conjugate acid conjugate base HCl + NH3  NH4+ + Cl – Acid Base conjugate acid conjugate base 1. HBr + H2O  H3O + + Br – 2. H2CO3 + H2O  H3O + + HCO3- 3. CN - + H2O  HCN + OH – 4. NH3 + H2O  NH4+ + OH – Substances that can act as both acids and bases are said to be amphoteric.

POGIL Time!

The following substances act as Bronsted acids in water. HF H F + H2O  H3O+ + F- H2SO3 H2SO3 + H2O  H3O+ + HSO3- Nitrous acid, HNO2 + H2O  H3O+ + NO2- Acid Base C.A C.B Acid Base C.A C.B Acid Base C.A C.B

The following substances act as Bronsted bases in water. Carbonate ion CO32- + H2O  HCO3- + OH- Acetate ion C2H3O2- + H2O  HC2H3O2 + OH- Ammonia NH3 + H2O  NH4+ + OH- Base Acid C.A C.B Base Acid C.A C.B Base Acid C.A C.B

HF (aq) + H2O (l) F- (aq) + H3O+(aq) Identify each reactant and each product in these equations as either a Bronsted acid or a Bronsted base. HClO (aq) + H2O (l) ClO- (aq) + H3O + (aq)   HF (aq) + H2O (l) F- (aq) + H3O+(aq) H2SO4 (aq) + SO32- (aq) HSO4 -(aq) + HSO3- (aq) Acid Base C.B C.A Acid Base C.B C.A Acid Base C.B C.A

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Strengths of Acids Strong Acids Weak Acids Sulfuric acid, H2SO4 Phosphoric acid, H3PO4 Hydrochloric acid, HCl Acetic acid, HC2H3O2 Nitric acid, HNO3

HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) Strong Acids HCl is strong because it is very good at transferring an H+ ion to a water molecule. In a 6 M hydrochloric acid solution, 100 % of the HCl molecules react with water to form H3O+ and Cl- ions. HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)

CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2-(aq) Weak Acids Vinegar is a weak acid because it is NOT very good at transferring H+ ions to water. In a 1 M solution, less than 0.5% of the CH3CO2H molecules react with water to form H3O+ and CH3CO2- ions. CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2-(aq) It is considered weak because more than 99 % of the acetic acid molecules remain intact.

Strengths of Bases Strong Bases Weak Bases Sodium hydroxide (lye), NaOH Weak Bases Potassium hydroxide, KOH Magnesium hydroxide, Mg(OH)2 Calcium hydroxide (lime), Ca(OH)2 not very soluble Ammonia, NH3

NaOH(s)  Na+(aq) + OH-(aq) Strong Bases NaOH is strong because it is very good at dissociating entirely into Na+ ion and OH - NaOH(s)  Na+(aq) + OH-(aq) Even though Ca(OH)2 is only slightly soluble in water, it is a strong base because all of the compound that dissolves dissociates into ions.

NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Weak Bases Ammonia is a weak base because it is NOT very good at transferring OH- ions to water. NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Equilibrium lies to the left because NH3 is weak and the conjugate base, OH- ion, is strong.

Strength Is Not Concentration Although the terms weak and strong are used to compare the strengths of acids and bases, dilute and concentrated are terms used to describe the concentration of solutions. To make the range of concentrations of possible hydronium or hydroxide ion concentrations (10–14 M to 1M) easier to work with, the pH scale was developed by S.P.L. Sørenson.

Neutralization Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl. 2 HCl + Mg(OH)2 MgCl2 + 2 H2O

The Reaction Between Acids & Bases Neutralization reaction: an acid and a base react in aqueous solution to produce a soluble salt and water Salt Ionic compound Cation from base (Mg2+) Anion from acid (Cl-)

Acids Neutralize Bases Bases Neutralize Acids Acids Neutralize Bases Complete each neutralization reaction. HCl (aq) + NaOH (aq)  NaCl (aq) + HOH (l) H2SO4 (aq) + Ca(OH)2(aq)  CaSO4(aq) + 2HOH (l) 2HNO3 (aq) + Mg(OH)2 (aq)  Mg(NO3)2 (aq) + 2HOH (l) 2HBr (aq) + Ba(OH)2 (aq)  BaBr2 (aq) + 2HOH (l) H2CO3 (aq) + Sr(OH)2 (aq)  SrCO3 (aq) + 2HOH (l) HClO4 (aq) + NaOH (aq)  NaClO4 (aq) + HOH (l)

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What is pH? pH is a mathematical scale in which the concentration of hydronium ions in a solution is expressed as a number from 0 to 14. much easier to work with than a range from 1 to 10–14 (each number is 10x the number preceding it) The pH of a solution is the negative logarithm of the hydrogen ion concentration. pH = -log[H+]

What is pOH? pOH is a mathematical scale in which the concentration of hydroxide ions in a solution is expressed as a number from 14 to 0. Runs opposite to the pH scale The pOH of a solution is the negative logarithm of the hydroxide ion concentration. pOH = -log[OH-] pH + pOH = 14.00

pH and pOH

Using your calculator to determine pH Enter the [H+] Take the log of [H+] by pressing the key marked LOG. Change the sign by pressing the +/- key. Ex. Calculate the pH [H+] = 1.0 x 10-7 M [H+] = 1.0 x 10-2 M [H+] = 3.0 x 10-6 M [OH-] = 8.2 x 10-6 M 7.0 2.0 5.5 Determine pOH and subtract from 14.00 8.9

Using your calculator to determine [H+] Enter the [-pH] Take the antilog of [-pH] by pressing the 2nd key and then the key marked LOG. Ex. Calculate the [H+] pH = 6 pH = 3 pH = 4.7 pOH = 2.8 1 x10-6 M 1 x 10-3 M 2.0 x 10-5 M 6.3 x 10-12 M Subtract from 14.00 and determine [H+].

Acid-base Titration Acid-base indicator: chemical dyes whose colors are affected by acidic and basic solutions Titration: method for determining the concentration of a solution by reacting a known volume of the solution with a solution of known concentration

Acid-base Titration End point: the point at which the indicator changes color Equivalence point: stoichiometric point at which the moles of H+ ion from the acid equal moles of OH- from the base

Titration Curves NH3+ HCl Strong CA NaOH + HCl CH3COOH + NH3 CH3COOH + NaOH Strong CB

Calculating Molarity (#H+)MAVA = (#OH-)MBVB What is the concentration (molarity) of a CsOH solution if 30.0 mL are neutralized by 26.4 mL of a 0.250 M HBr solution? HBr + CsOH  CsBr + HOH (1 H+) (0.250 M) (26.4 mL) = (1OH-) MB (30.0 mL) MB = 0.220 M

Buffered Solutions Buffers: solutions that resist changes in pH when limited amounts of acid or base are added a buffer is a mixture of a weak acid and its conjugate base OR a weak base and its conjugate acid the mixture of ions in solution resist changes in pH by reacting with any H+ or OH- ions added to the solution

Buffered Solutions Buffer capacity: amount of acid or base a buffer solution can absorb without significant change in pH ecosystems blood