Chapter 18 notes-Acids and bases

Naming Acids Review Acids Examples: Compounds that form H+ in water. Formulas usually begin with ‘H’. Examples: HCl – hydrochloric acid HNO3 – nitric acid

Relationship to –ate anion Oxyanions -ate has MORE oxygen atoms -ite has LESS oxygen atoms per- is greatest oxygen atoms hypo- is least oxygen atoms Relationship to –ate anion General Name Name Formula Acid Name Acid Formula 1 more oxygen per(root)ate perchlorate ClO4- Perchloric acid HClO4 most common (root)ate chlorate ClO3- Chloric acid HClO3 1 less oxygen (root)ite chlorite ClO2- Chlorous acid HClO2 2 less oxygen Hypo(root)ite hypochlorite ClO- Hypochlorous acid HClO No oxygen (root)ide chloride Cl- Hydrochloric acid HCl

Naming Acids Practice HBr hydrobromic acid H2CO3 carbonic acid H2SO4 sulfuric acid HClO3 chlorous acid Phosphoric acid H3PO4 Persulfuric acid H2SO5 Hydrofluoric acid HF Hyponitrous acid HNO

properties of acids Produce H+1 (as H3O+1) ions in water hydronium ion is a hydrogen ion attached to a water molecule Taste sour pH is less than 7 Good conductors of electricity = electrolytes React with metals to produce H2 Zn + 2HCl  ZnCl2 + H2 React with metal carbonates to produce CO2 NaHCO3 + HC2H3O2  NaC2H3O2 + H2O + CO2 React with bases to form a salt and water HCl + NaOH  NaCl + H2O Turns blue litmus paper to red “Blue to Red A-CID”

Examples of acids

Examples of acids

Examples of acids

Properties of Bases Produce OH-1 ions in water Taste bitter, chalky Hydroxide ion Taste bitter, chalky Feel soapy, slippery pH greater than 7 Good conductors of electricity = electrolytes React with acids to form salts and water NaOH + HCl  NaCl + H2O Turns red litmus paper to blue “Basic Blue”

Examples of Bases

Examples of Bases

Hydronium vs. hydroxide ions All water solutions contain hydrogen ions (H+ or H3O+) and hydroxide ions (OH-) The relative amounts determine whether a solution is acidic, basic or neutral Type of solution (H+) (OH-) acidic More Less basic neutral Same

Hydronium vs. hydroxide ions Pure water produces equal numbers of hydrogen ions (H3O+) and hydroxide ions (OH-) The hydronium ion is a hydrogen ion which has a water molecule attached to it by a covalent bond. H+ and H3O+ can be used interchangeably H2O + H2O H3O+ + OH-

Arrhenius Model Swedish chemist Svante Arrhenius – 1883 Acids – produce H+ ions (or hydronium ions H3O+) when dissolved in water Bases – produce OH- ions when dissolved in water Problem: some bases don’t have hydroxide ions! Examples: Ammonia (NH3) Sodium carbonate (Na2CO3)

Arrhenius Model Examples What happens when hydrochloric acid dissolves in water? Is this an acid or base? Acid: ionizes to produce H+ What happens when sodium hydroxide dissolves in water? Is this an acid or base? Base: ionizes to produce OH- HCl  H+ + Cl- NaOH  Na+ + OH-

Brønsted-Lowry Model Danish chemist Johannes Brønsted and English chemist Thomas Lowry - 1923 Model focuses on ‘hydrogen ion’ a.k.a. proton Acids – proton donor a.k.a hydrogen ion donor Bases – proton acceptor a.k.a hydrogen ion acceptor Theory helps to explain why ammonia is a base!

Brønsted-Lowry Model The forward reaction & the reverse reaction is a reaction with an acid and a base. The reverse reactions = conjugate acids and bases Conjugate Acids – produced when a base accepts a hydrogen ion Conjugate Bases – produced when an acid donates a hydrogen ion Conjugate Acid-base pair- consists of 2 substances related to each other by the donating and accepting of a single hydrogen ion HCl + H2O  H3O+ + Cl- A B CA CB

Brønsted-Lowry Model Analogy: Father and son playing catch with a ball The Ball is the hydrogen ion Forward Reaction: Father has the ball  he is an acid because he has the ball to give away Son does not have the ball  he is a base because he can accept the ball when it is thrown to him Throws the ball to his son: His son is now an acid (conjugate) because he has the ball to give away Father is now a base because he is available to accept the ball Reverse Rection: Son has the ball  he is a conjugate acid Father does not have the ball  he is a conjugate base and will accept the ball when it is thrown to him

Brønsted-Lowry Acid HF + H2O  H3O+ + F- What are the conjugate acid-base pairs for the following reaction? Draw the arrows on the reaction to demonstrate this. HF + H2O  H3O+ + F-

Brønsted-Lowry Base NH3 + H2O  NH4+ + OH- What are the conjugate acid-base pairs for the following reaction? Draw the arrows on the reaction to demonstrate this. NH3 + H2O  NH4+ + OH-

Brønsted-Lowry Acid & Base When HF dissolves in water, water acts as a base When NH3 dissolves in water, water acts as an acid Water is acting as an acid and a base!! Anything that can act as both an acid and a base are said to be amphoteric.

Monoprotic & polyprotic Acids Monoprotic-an acid that can donate only 1 hydrogen ion Examples: HF, HCl, HBr, HNO3, HClO4, HC2H3O2 Mono-1; protic-proton Polyprotic-an acid that can donate more than 1 hydrogen ion Examples: H3PO4, H2SO4, H2CO3, H3BO3 Poly-more than one; protic-proton All polyprotic acids ionize in steps: Step 1: H3PO4 + H2O H3O+ + H2PO4- Step 2: H2PO4- + H2O H3O+ + HPO4-2 Step 3: HPO4-2 + H2O H3O+ + PO4-3

Polyprotic Acids Write the ionization steps for the polyprotic acid: H2SO4

Lewis Acids & Bases Even more general model of acids and bases Based on Lewis structures Lewis acid-an ion or molecule with a vacant atomic orbital that can accept (share) an electron pair (electron pair acceptor) Lewis base-an ion or molecule with a lone electron pair that it can donate (share) an electron pair (electron pair donator)

Strengths of acids Acids that completely ionize in the solution are called strong acids. Produce the maximum amount of ions  good conductors of electricity Examples: HCl, HI, HBr, HClo4, HNO3, H2SO4 Acids that ionizes only partially in the solution are called weak acids. Do not conduct electricity as well as a strong acid. Examples: HF, HC2H3O2, H2S, H2CO3, HClO

Ionization of Acids HCl + H2O  H3O+ + Cl- HF + H2O H3O+ + F- Strong Acids-ionize completely: HCl is a strong acid  HCl is 100% ionized Demonstrated with a single headed forward arrow () Weak Acids-do not ionize completely: HF is weak acid  HF is not 100% ionized Demonstrated with a reverse reaction arrow ( ) The amount that a weak acid ionizes can be found using the acid ionization constant. Value of Ka tells whether reactants or products are favored Small Ka = weakest acids HF + H2O H3O+ + F-

Strength of Bases Bases that completely ionize in the solution are called strong bases. Produce the maximum amount of ions  good conductors of electricity Examples: Metal hydroxides: NaOH, KOH, RbOH, CsOH, Ca(OH)2, Ba(OH)2 Bases that ionizes only partially in the solution are called weak bases. Do not conduct electricity as well as a strong base. Examples: CH3NH2, C2H5NH2, NH3, C6H5NH2

Ionization of Bases NaOH  Na+ + OH- NH3 + H2O NH4+ + OH- Strong Bases-ionize completely: NaOH is a strong acid NaOH is 100% ionized Demonstrated with a single headed forward arrow () Weak Bases-do not ionize completely: NH3 is weak base  NH3 is not 100% ionized Demonstrated with a reverse reaction arrow ( ) The amount that a weak base ionizes can be found using the base ionization constant. Value of Kb tells whether reactants or products are favored Small Kb = weakest base NH3 + H2O NH4+ + OH-

Self-Ionization of Water Water produces both hydronium and hydroxide ions through self-ionization: (Brackets [ … ] are used to symbolize concentration in molarity) Kw = ionization constant of water [H3O+] = 1 x 10-7 [OH-] = 1 x 10-7 H2O(l) + H2O(l) H3O+(aq) + OH-(aq) Kw = [H3O+] [OH-] (Kw = 1 x 10-14 )

pH SCALE A means of expressing acidity or basicity of a solution Ranges from 0-14: 0 to <7 : ACID 7: NEUTRAL >7 to 14: BASE

pH & pOH pH = -log [H+] pOH = -log [OH-] Express the concentration of hydrogen ions or hydroxide ions using a pH scale based on common logarithms pH is a more convenient means of quantitatively describing acidity of a solution. pH stands for – power of hydrogen Positive hydrogen balance Potential of hydrogen Sometimes it is convenient to express the basicity of a solution on a pOH scale. pOH: 0-7  base pOH: 7 neutral pOH: 7-14  acid pH = -log [H+] pOH = -log [OH-]

Measuring pH Indicator paper Indicators pH Meter Litmus paper  tells if an acid or base Wide range indicator paper  tells specific pH Indicators Phenolphthalein  indicates when a base is present (pink) Bromothymol Blue  yellow if acidic, blue if basic pH Meter Measures specific pH of solution by placing electrodes into the solution

Indicators

EQUATIONS pH = -log [H+] pOH = -log [OH-] [H+] = 10-pH [OH-] = 10-pOH pH + pOH = 14 [H+] [OH-] = 1 x 10-14

pH + pOH practice What is the pH of a solution with a [H3O+] = 1.0 x 10-7? 2. What is the pOH of a solution with a [OH-] = 3.21 x 10-4?

pH + pOH practice 1. What is the [H3O+] of a solution with a pH = 12.3? 2. What is the [OH-] of a solution with a pOH = 4.25?

Connection between hydronium and hydroxide pH + pOH = 14 [H+] [OH-] = 1 x 10-14 Neutral solutions: [H3O+] = [OH-] & pH = 7 Acidic solutions: [H3O+] > [OH-] & pH < 7 Basic solutions: [H3O+] < [OH-] & pH > 7

Practice