Acids, Bases, and Salts

Acid/Base Theory Theory Acid Definition Base Definition Arrhenius Substance that releases a H+ ion in water Substance that releases an OH- ion in water Bronsted-Lowry Substance that donates a proton (H+) Substance that accepts a proton (H+) - H+ ion called a proton because no electron, just one proton left (show on board)

Classifying acids, bases, salts donates a proton (H+) Bases accepts a proton (H+) Salts dissolve in water; bound ionically Is not classified as an acid or base using the definitions above

Acids, Bases and Salts Examples: HCl NaOH NaCl H2SO4 Ca(OH)2 H2O

Multiple Hydrogen Atoms Monoprotic acids have one Hydrogen Ex: HCl, HBr, HI Diprotic acids have two hydrogen Ex: H2SO3, H2SO4, H2CO3 Triprotic acids have three hydrogen Ex: H3PO4, H3PO3

HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq) Bases without OH- Example: HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq) Is HCl an acid or a base? Why? Is H2O an acid or a base? Why? - make sure they know what the hydronium ion is

H2O (l) + NH3(l) → NH4+(aq) + OH-(aq) Bases without OH- Example: H2O (l) + NH3(l) → NH4+(aq) + OH-(aq) Is H2O an acid or a base? Why? Is NH3 an acid or a base? Why?

Autoionization of Water H2O(l) + H2O(l) → H3O+(aq) + OH-(aq) Is H2O(l) the acid or base? Kw=?

Naming Acids Hydrogen bonded with an element: Examples HBr H2S hydro__________ic acid Examples HBr H2S hydrobromic acid hydrosulfuric acid

Naming Acids Hydrogen bonded with a polyatomic ion: Examples ______________ic acid Examples HNO3 H2SO4 nitric acid sulfuric acid

Naming Acids Acids with a different number of oxygens: Formula Name HNO3 nitric acid HNO2 nitrous acid HNO hyponitrous acid HNO4 pernitric acid

Naming Acids Example - naming acids: HClO3 HClO2 HClO HClO4 - homework: acids, bases, salts ws

Electrolytes Acids, Bases, and Salts dissociate in water Hydration Simulation Acids, Bases, and Salts dissociate in water (break apart into ions) Ions conduct electricity Substances can be: strong electrolytes weak electrolytes nonelectrolytes

Strong vs Weak Strong Acids HCl HNO3 HI HClO3 HClO4 H2SO4 Strong Bases Group 1A metal with hydroxides Heavy group 2A metals(Ca, Sr, Ba) with hydroxide.

Completely breaks into ions Partially breaks into ions Electrolytes Type of Electrolyte Solubility Ions Conductor of electricity? Strong electrolyte Dissolves in water Completely breaks into ions Very good conductor Weak electrolyte Partially breaks into ions Poor Nonelectrolyte No ions No conduction since ions have charges, they have the ability to “transport” electrons, thus they can conduct electricity demo: light bulb connected with wires to graphite electrodes

Other Properties of Acids and Bases Metals in Acids Recall the zinc in HCl demo What was produced? Taste Acids taste __________ Bases taste _________

Determining Acidity Acidity measured using the pH scale - pH scale starts at 0 and ends at 14. closer to 0, acidic. closer to 14, basic. 7 is neutral solution.

pH = -log [H+] or pH = -log [H3O+] Determining Acidity pH measure of the H+ or H3O+ ions in solution pH = -log [H+] or pH = -log [H3O+] pOH measure of the OH- ions in solution pOH = -log [OH-] pH + pOH = 14 pH: 0 to 14, closer to zero means more acidic pOH: 0 to 14, closer to zero means more basic Scales are exact opposites of one another

Calculating pH and pOH Example Find the pH and the pOH for the following: A 3.0 x 10-4 M solution of HCl? show on board (balanced equation, M of H+, pH equation) lower pH means stronger acids, means strong electrolyte (same for high pH and strong bases)

Calculating pH and pOH Example Find the pH and the pOH for the following: A 1.9 x 10-9 M solution of NaOH? show on board (balanced equation, M of H+, pH equation) lower pH means stronger acids, means strong electrolyte (same for high pH and strong bases)

Calculating pH and pOH Example Find the pH and the pOH for the following: a) 5.4 x 10-2 M solution of KOH? b) 1.0 M solution of HBr? show on board (balanced equation, M of H+, pH equation) lower pH means stronger acids, means strong electrolyte (same for high pH and strong bases) pOH = 1.3 pH = 12.7 pOH = 14 pH = 0

Neutralization Reactions General equation: Acid + Base → Salt + Water Example: HCl + NaOH → HCl + OH- → Cl- + H2O NaCl + H2O another type of double replacement reaction remember, salts are electrolytes and dissolve in water: form ions Acid Base Conjugate Acid Conjugate Base

Neutralization Reactions Example: List which is the acid, the base, the conjugate acid, and the conjugate base. HNO3 + OH- → H2O + NO3-

Titrations Titration Neutralization reaction occurs The reaction between an acid and a base used to find the concentration of the unknown acid or base. Neutralization reaction occurs An indicator is used to find the “end point” Equivalence point (i.e. the “end point”) when the number of moles of the acid equals the number of moles of the base titration demo: prep for lab tomorrow show how calculations can be done to find unknown M of base titration lab

Titrations Use the equation below when calculating the unknown concentration of the acid or the base MAVA = MBVB MA = molarity (concentration) of acid VA = volume of acid used during titration MB = molarity (concentration) of base VB = volume of base used during titration

Titrations Example: It takes 43.81 mL of NaOH to reach the end point with 23.52 mL of 0.5230 M HCl. What is the concentration (molarity) of NaOH?

Titrations Example: It takes 52.40 mL of NaOH to reach the end point with 15.38 mL of 1.2 M HCl. What is the concentration (molarity) of NaOH?

Indicators Indicator pH range for color change 0.5-1.25 3.25-4.25 5-6 Methyl violet 0.5-1.25 Bromphenol blue 3.25-4.25 Methyl red 5-6 Litmus 6-7.5 Phenolphthalein 8-10 Alizarin yellow 10.5-12 indicator is chosen depending on what pH it changes color and on how dramatic that color change is (the more dramatic, the easier it is to see) phenolphthalein is the most commonly used in chemistry

Natural Indicators Red Cabbage

Natural Indicators Hydrangea Acidic Soil Alkaline (Basic) Soil or Neutral Soil