10.1 & 10.2 ACIDS AND BASES PROPERTIES OF ACIDS AND BASES Acids pH Electrical conductivity Taste Feel Colour with litmus paper Colour with phenolphthalein Bases PROPERTIES OF ACIDS AND BASES Acids pH Electrical conductivity Taste Feel Colour with litmus paper Colour with phenolphthalein less than 7 Bases greater than 7 PROPERTIES OF ACIDS AND BASES Acids pH Electrical conductivity Taste Feel Colour with litmus paper Colour with phenolphthalein less than 7 Bases greater than 7 conductivity varies PROPERTIES OF ACIDS AND BASES Acids pH less than 7 Electrical conductivity Taste Feel Colour with litmus paper Colour with phenolphthalein Bases greater than 7 conductivity varies sour bitter PROPERTIES OF ACIDS AND BASES Acids pH less than 7 Electrical conductivity Bases greater than 7 conductivity varies Taste sour bitter Feel no special feel slippery Colour with litmus paper Colour with phenolphthalein PROPERTIES OF ACIDS AND BASES Acids pH less than 7 Electrical conductivity Bases greater than 7 conductivity varies Taste sour bitter Feel no special feel slippery Colour with red litmus paper Colour with phenolphthalein blue PROPERTIES OF ACIDS AND BASES Acids pH less than 7 Electrical conductivity Bases greater than 7 conductivity varies Taste sour bitter Feel no special feel slippery Colour with red litmus paper Colour with colourless phenolphthalein blue pink REACTIONS OF ACIDS Acids react with metals to produce hydrogen gas Magnesium + hydrochloric acid magnesium chloride + hydrogen gas Hydrochloric acid also reacts with carbonate compounds to produce carbon dioxide REACTIONS OF BASES Bases react with carbon dioxide to form carbonates Calcium hydroxide + carbon dioxide calcium carbonate + water NAMING ACIDS Binary Acids •composed of two elements: hydrogen and a non-metal ex. HF (aq) •typically a hydrogen attached to a halogen and dissociated in water, hence the “aq” HF F- + H+ •begin with the prefixes “hydro” HCl (aq) = hydrochloric acid HBr (aq) = Oxoacids •formed from a polyatomic ion that contains oxygen, hydrogen and another element. They have NO overall charge Hypo --- ous Hyponitrous acid HNO --- ous Nitrous acid HNO2 HClO HClO2 H2CO H2CO2 H2SO2 H2SO3 H3PO2 H3PO3 --- ic Per --- ic Nitric acid Pernitric acid HNO3 HNO4 Chloric acid HClO3 HClO4 Carbonic acid H2CO3 H2CO4 Sulfuric acid H2SO4 H2SO5 Phosphoric acid H3PO4 H3PO5 NAMING BASES Ionic hydroxides NaOH = Hydroxide bases only show their basic properties when they are in solution; thus, we use (aq) in the formula of a base. ARRHENIUS THEORY OF ACIDS AND BASES • An acid is a substance that ionizes in water to produce one or more hydrogen ions, H+ • Example: H2SO4 (aq) 2 H+ (aq) + SO4 2- (aq) • A base is a substance that dissociates in water to produce one or more hydroxide ions, OH• Example: Ca(OH)2 (aq) Ca 2+ (aq) + 2 OH- (aq) • Ionization vs dissociation • Both occur when a compound breaks apart in water, causing the presence of ions • Dissociation occurs when water molecules pull the + and – ions apart • Ionization involves the formation of ions from uncharged molecules, causing uncharged entities to become ionized. How? Arrhenius didn’t say! • In water, ionic compounds DISSOCIATE and molecular compounds IONIZE. • Thankfully, the chemical equations look the same! BRøNSTED LOWRY THEORY OF ACIDS AND BASES A broader definition of acids and bases BRØNSTED ACID: Substance that donates an H+ (proton) BRØNSTED BASE: Substance that accepts an H+ (proton) Under this definition, a substance can ONLY act as an acid (donating a proton) IF another substance is acting as a base at the same time. BRøNSTED LOWRY THEORY OF ACIDS AND BASES Two molecules or ions that are related by the transfer of a proton are called: A CONJUGATE ACID-BASE PAIR EVERY BRØNSTED ACID has a conjugate base that forms after the acid has donated its H+ (proton). EVERY BRØNSTED BASE has a conjugate acid that forms after the base has accepted a H+ (proton). HNO3 + N2H4 NO3ACID BASE + N2H5+ CONJUGATE CONJUGATE BASE ACID H2PO4 + CO32- HPO42- + HCO3ACID S2- BASE BASE CONJUGATE CONJUGATE BASE ACID + H2O HSACID + OH- CONJUGATE CONJUGATE ACID BASE COMPARING ACIDS Strong Acids Weak Acids COMPARING ACIDS Strong Acids Weak Acids Completely ionize in water Partially ionize in water COMPARING ACIDS Strong Acids Weak Acids Completely ionize in water Partially ionize in water React very quickly React mildly COMPARING ACIDS Strong Acids Weak Acids Completely ionize in water Partially ionize in water React very quickly React mildly Low pH (0-4) Moderate pH (4-7) COMPARING ACIDS Strong Acids Weak Acids Completely ionize in water Partially ionize in water React very quickly React mildly Low pH (0-4) Moderate pH (4-7) Good conductors Poor conductors COMPARING ACIDS Strong Acids Weak Acids Completely ionize in water Partially ionize in water React very quickly React mildly Low pH (0-4) Moderate pH (4-7) Good conductors Poor conductors Examples HCl (stomach acid) HNO3 (used in rocket fuel) H2SO4 (in car batteries) Examples CH3COOH (vinegar) H3PO4 (in pop) COMPARING BASES Strong Bases Weak Bases Completely ionize in water Partially ionize in water React very quickly React mildly High pH (11-14) Moderate pH (7-11) Good conductors Poor conductors Examples NaOH (in oven/drain cleaner) Ca(OH)2 (limewater) Examples NH3 (Ammonia) **STRENGTH and CONCENTRATION of acids/bases are not related! • Diluting A STRONG ACID does not make it a weak acid. • ie: HCl is always a strong acid, regardless of whether its [ ] is 20 mol/L or 0.001 mol/L • It will always completely ionize in water, regardless of concentration!! • Similarly, a WEAK ACID can be very concentrated. STRONG AND WEAK ACIDS Acidic solutions that have high electrical conductivity are considered STRONG ACIDS In these solutions almost all (>99%) the acid ionizes (produces ions) Typically these acids are highly corrosive HCl, HNO3, H2SO4 WEAK ACIDS have low percentage ionization, i.e. they only ionize partially (<50%) Examples of weak acids are H2CO3 and HCH3COO STRONG AND WEAK BASES An ionic hydroxide compound that dissociates completely (>99%) in solution is considered a STRONG BASE Similar to acids, weak bases do not dissociate completely Some weak bases do not contain a hydroxide ion but, when dissolved in solution, produce hydroxide ions Example: ammonia NH3 + H2O NH4+ + OH- MONOPROTIC ACIDS Contain a single hydrogen ion that can dissociate (can only lose one proton). Example: HCl(aq) H+(aq) + Cl-(aq) What are diprotic acids and triprotic acids? EXAMPLE OF A TRIPROTIC ACID H3PO4(aq) + H2O(l) ↔ H+(aq) + H2PO4-(aq) H2PO4-(aq) + H2O(l) ↔ H+(aq) + HPO42-(aq) HPO4(aq) + H2O(l) ↔ H+(aq) + PO4-(aq) *FIRST DISSOCIATION IS STRONGER THAN SECOND DISSOCIATION. pH In pure water, (neutral solution), a very low concentration of water molecules ionize to produce hydronium ions and hydroxide ions • S.P. Sorenson proposed that the acidity of a solution be expressed in terms of a quantity known as pH The pH of solution is related to the quantity of hydronium ions (H3O+) in a given amount of solution, in the mathematical expression pH • pH = “power of hydrogen” • pH scale is measurements by powers of ten • pH of 3 is 10 times more acidic than pH of 4 100 times more __________than basic • pH of 10 is _______ pH of 8 • Indicators are substances that will change different colours in the presence of an acid or base ACID BASE INDICATORS Indicators can be used to determine the whether a solution is acidic or basic We must choose the right indicator! pH of a Solution Recall… • Acids release H+ when they ionize in water • Concentration is the amount of a solute in a given volume of solvent These two concepts have been combined to describe the acidity of a solution (or pH) • pH is the concentration of H3O+ (or H+) in a solution pH of a Solution Also recall (according to Arrhenius)… Acidic solutions have a higher concentration of hydronium ions Basic solutions have a higher concentration of hydroxide ions Thus, the pH scale works by determining the hydronium ion concentration pH Scale Pure water has a hydronium concentration of 1 x 10-7 mol/L Thus, In an acidic solution: [H3O+] > 1 x 10-7 mol/L In a basic solution: [H3O+] < 1 x 10-7 mol/L pH Scale From this: The pH of a solution is defined as the negative of the exponent to the base 10 of the hydrogen ion concentration + [H ] = -pH 10 pH Scale Example: What is the [H+] of a solution with a pH of 3? [H+] = 10-pH [H+] = 10-3 [H+] = 1.0 x 10-3 mol/L pH Scale Example: What is the [H+] of a solution with a pH of 10.33? [H+] = 10-pH [H+] = 10-10.33 [H+] = 4.7 x 10-11 mol/L pH Scale To determine the pH of a solution, the following equation would need to be used: pH = -log[H+] Example: What is the pH of an antacid solution with a hydrogen ion concentration of 5.3 x 10-11? pOH The same formulas can be used to find the concentration of base pOH = -log[OH-] [OH-] = 10-pOH pH + pOH = 14 • PG 469 # 1, 2, 4, - 8, 9, 10 • PG 475 # 1,4-11