Unit 11: Acids and Bases

Unit Overview… We will learn about Acids and Bases, two important types of compounds in chemistry Learn the distinct properties of Acids and Bases Understand the pH scale, and how we can use indicators to detect pH in an Acid-Base solution Examine important acid-base chemical reactions

Arrhenius’ Theory of Acids and Bases Arrhenius Acid Any substance that dissolves in water to produce hydrogen ions (H+ ) or Hydronium ions (H3O+) Ex: Hydrochloric acid, HCl (a strong acid). HCl (g) H+ (aq) + Cl- (aq) H2O (l) + H+ (aq)  H3O+ (aq) ACIDIC solutions are formed when an acid transfers a proton to water.

Hydronium Ion When an H+ ion interacts with the electrons of the oxygen in a water molecule, an H3O+ ion called the hydronium ion is formed.

Arrhenius’ Theory of Acids and Bases Arrhenius Base Any substance that dissolves in water to produce hydroxide ions (OH- ) Ex: Sodium Hydroxide, NaOH (a strong base) or Ammonia NH3 NaOH (s) Na+ (aq) + OH- (aq) NH3 (l) + H2O (l)  NH4+ (aq) + OH- (aq) *Important Note – most bases end in “OH” but if there is CARBON attached to the compound then it is NOT a base

Properties of Acids They’re electrolytes – chemicals that break up into ions in water, so they conduct electricity in water Strong acid: good conductor Weak acid: poor conductor Sour Taste React with most metals to produce H2 (g) React with bases to form Water and salt React with indicators to change color pH values are from 0 to 6 see Reference Table M for color changes

Properties of Bases Also electrolytes Bitter Taste Feel slippery (like soap) React with acids to form water and salt React with indicators to change color pH values are 8 to 14 See Reference Table M for color changes

Properties of Salts Also electrolytes Ionic Compounds – containing one metal ion (+) and one non-metal ion (-), or a polyatomic ion (except OH-) Ex) NaCl, MgCl2, or Ca3(PO4)2 React with indicators to change color pH values are around 7 (neutral) See Reference Table M for color changes

pH Scale A scale, called the pH scale, has been developed to express the concentration of hydrogen ions [H+] as a number from 0 to 14. A pH of 0 is strongly acidic A pH of 7 is neutral A pH of 14 is strongly basic (High concentration of H+) Equal Concentration of H+ and OH- ions (High concentration of OH-)

pH Scale

Acid – Base Indicators See Table M An indicator is a chemical that changes color when it gains or loses a “proton” (H+ ions). There are many indicators listed on Reference Table M, and the color changes associated with certain pH values are given. See Table M We can use multiple indicators to find the approximate pH of a substance

Table K and Table L These two tables list the most common acids and bases you will see in this course. Table K: The top 4 acids are strong acids The last 2 are weak acids. Table L: The top 3 bases are strong bases The last 1 is a weak base. Table L Table K

Acidity vs. Alkalinity Acidity – is a measure of the strength of an acid related to the concentration of H+ ions. Alkalinity – is a measure of the strength of a base related to the concentration of OH- ions High Acidity = pH less than 7 High Alkalinity = pH more than 7

Strong Acids Strong acids: ionize (dissolve) completely, which means if 100 molecules are placed in water all 100 will break up into ions Good Conductors! Examples: HCl  H+ + Cl- H2SO4  2H+ + SO4-2

Weak Acids Weak acids: ionize (dissolve) only slightly, If 100 molecules are placed in water then only a small percentage will break up into ions. Bad Conductors… Example: acetic acid “vinegar” CH3COOH  CH3COO- + H+ *NOTE: if an organic compound ends in COOH it’s a weak acid!!!

Strong Bases Strong bases: ionize completely Good Conductors! Example: Sodium Hydroxide NaOH  Na+ +OH-

Weak Bases Weak base: ionizes slightly (won’t dissolve completely in water) Bad Conductors… Example: ammonia NH3 + H2O  NH4+ + OH-

Naming Acids

Binary acids –two elements - H+ * - Previous Naming System - New Naming System hydrogen _____ ide  hydro_____ic acid Hydro_____ic acid Hydrochloric acid Hydrofluoric acid Hydrosulfuric acid Hydrobromic acid Hydrogen _____ ide Hydrogen chloride Hydrogen fluoride Hydrogen sulfide Hydrogen bromide

Ternary acids: H+ & polyatomic ion - Previous Naming System - New Naming System Hydrogen _____ate  _________ic acid Hydrogen _____ate Hydrogen sulfate Hydrogen chlorate Hydrogen nitrate Hydrogen phosphate _________ic acid Sulfuric acid chloric acid nitric acid Phosphoric acid

More Ternary Acids – “-ite Ions” - Previous Naming System - New Naming System Hydrogen _____ite  _______ous acid _______ous acid Nitrous acid Sulfurous acid Chlorous acid Hydrogen _____ite Hydrogen nitrite Hydrogen sulfite Hydrogen chlorite

Naming Bases Ca(OH)2  Calcium Hydroxide No change in how we name these… Full name of the metal atom + Hydroxide Ex) Ca(OH)2  Calcium Hydroxide

Acid and Base Reactions You will need to be familiar with 2 types of reactions involving acids & bases. 1st reactions between any acid and certain metals. 2nd reactions combine acids & bases to produce salt and water (neutralization).

Acid-Metal Reactions Most metals will corrode when exposed to certain acids. When this happens H2 gas is produced. Example: Ni + 2HCl  NiCl2 + H2 or 2Al + 6HCl  2AlCl3 + 3H2 Only precious metals won’t corrode like copper, gold, and silver because they aren’t reactive (Low on Table J)

Neutralization Reaction In a neutralization reaction an Acid will react with a Base to produce salt and water. Word equation Formula equation Net ionic equation A net ionic equation has only the ions that take part in the reaction. The ions that don’t change are removed, called spectator ions. Hydrochloric Acid + Sodium Hydroxide yields Water + Sodium Chloride HCl + NaOH  H2O + NaCl H+(aq) + OH-(aq)  H2O(l)

Try This One H2SO4 + 2NaOH 2H2O + Na2SO4 Sulfuric acid + sodium hydroxide yields water + sodium sulfate H2SO4 + 2NaOH 2H2O + Na2SO4 2H+ + SO4-2 + 2Na+ + 2OH- 2H2O + 2Na+ + SO4-2 H+ + OH- H2O

Finding Concentration of H+ in Acids A 1.0 M HCl “monoprotic acid” HCl  H+ + Cl- Produces 1 H+ so the concentration of H+ is 1.0 M Written like this: [H+] = 1.0 M A 1.0 M H2SO4 “diprotic acid” H2SO4  2H+ + SO4-2 Produces 2 H+ so the concentration of H+ is 2 x 1.0 M Written like this: [H+] = 2.0 M

Finding Concentration of OH- in Bases A 1.0 M KOH KOH  K+ + OH- Produces 1 OH- so the concentration of OH- is 1.0 M. Written like this: [OH-] = 1.0 M A 1.0 M Mg(OH)2 Mg(OH)2  Mg+2 + 2OH- Produces 2 OH- so the concentration of OH- is 2(1.0 M). [OH-] = 2.0 M

Titration An Acid-Base titration is a technique used find the concentration of an acid or a base by neutralizing it. During a titration you add set volumes of a base to an acid until it is neutralized. Using the Acid-Base titration formula listed on Table T you can solve for your unknown concentration.

Acid- Base Titrations MaVa = MbVb (Table T) Ma = molarity of H+ Va = volume of acid Mb = molarity of OH- Vb = volume of base

Titrations (Neutralization) Problems Ex. What volume of 0.50M HCl is required to neutralize 100mL of 2.0M NaOH? MaVa = MbVb Ma = 0.5 M Va = ??? mL Mb = 2.0 M Vb = 100 mL 0.5(x) = 2.0(100) X = 400mL HCl

Bronsted-Lowry Acid-Base Theory Acids: chemicals that act as proton donors in water ( H+ + H2O  H3O+) Bases: Chemicals that act as proton acceptors in water (H2O  H+ + OH-) Water: Can act as an acid or a base because it can either donate or accept “protons” depending on the situation

Bronsted-Lowry Conjugate Acids/Bases A conjugate acid - what is left of a Bronsted-Lowry Base after accepting a proton A conjugate base - what is left of aa Bronsted-Lowry Acid after donating a proton Ex) HNO3 + H2O  NO3- + H3O+ Conj. Base Conj. Acid Base Acid