Chemistry – April 21, 2017 P3 Challenge – Consider: 2 SO2 (g) + 1 O2 (g)  2 SO3 (g) (Exothermic) at equilibrium. Determine which way the equilibrium will shift when the following disturbances are applied. A) SO3 is removed. B) Temperature is increased . C) Volume is increased. Get out LeChatelier’s Principle worksheet for HMW check

Chemistry – April 21, 2017 Today’s Objective – Acid/Base equilibrium and pH Assignment: Acid/Base pH Worksheet Agenda Acid/base properties Acid/base definitions Acid/base concentrations Dissociation and Ionization Self-ionization of water pH, pOH

Already about acids and bases (U5) Acids react with bases to form water and a salt. Type of Double Replacement reaction called a neutralization. Forming water is a driving force Acidic compounds start with H and have 3 naming rules. There are 6 strong acids: HCl, HBr, HI, H2SO4, HNO3 and HClO4 All other acids are weak. Strong bases are Group 1 and 2 hydroxides.

Acid and base properties Acids properties: Sour taste Dissolves metals Turns litmus paper red Turns phenolphthalein colorless Base properties Bitter taste Slippery feel Turns litmus paper blue Turns phenolphthalein pink

Arrhenius Definitions Proposed by Svante Arrhenius, a Swedish scientist, in 1850 Acid = produces H+ ion in solution Base = produces OH- ion in solution Can label any substance based on its chemical formula. Ex: NaOH Ex: H2SO4 Limited usefulness because it misses some substances that we can label as acids and bases based on their properties: NH3 is a base, NaHCO3 is a base

Brønsted-Lowry Definition Proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923 Acid = proton (H+) donor Base = proton (H+) acceptor Some molecules can both donate and accept protons (Ex: Water) and are called amphoteric. Ex: NH3(aq) + H2O (l)  NH4+ (aq) + OH- (aq) Ex: HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq)

Types of acids Monoprotic acids donate 1 proton Ex: HCl Diprotic acids can donate 2 protons Ex: H2SO4 Triprotic acids can donate 3 protons Ex: H3PO4

Conjugate acids and Conjugate bases The reactions of acids and bases are often reversible creating an equilibrium The “acid product” that forms from the addition of a proton to a base is a conjugate acid. The “base product” that forms from the removal of a proton from an acid is a conjugate base. Any reactant and its conjugate partner form a conjugate acid-base pair. Ex: NH3(aq) + H2O (l)  NH4+ (aq) + OH- (aq)

Dissociation vs Ionization When an ionic compound dissolves in water = dissociation. The ions already exists, they just move apart. Ex: NaOH a strong base added to water is a dissociation. When a molecular compound splits apart and forms ions = ionization Ions must be created. Ex: HC2H3O2 a weak acid, added to water is a partial ionization.

Strong Acid Concentrations 6 Strong acids completely ionize in water and are strong electrolytes. HCl, HBr, HI, HNO3, HClO4, H2SO4 All other acids are weak, only partially ionize. Form an equilibrium. HA (aq)  H+ (aq) + A- (aq) For strong acids, determine ion concentrations to estimate [H+] 1.5 M HCl 3.5 M H2SO4

Strong Base Concentrations Strong bases are the hydroxides of alkali and alkaline earth metals. Strong bases are strong electrolytes and dissolve completely in water. All other bases are weak and only partially dissolve. Determine the [OH-] concentration for 0.45 M Ba(OH)2

Water Because water is amphoteric: can act as an acid or a base. Water can react with itself, acting as both the acid and base: H2O(l) + H2O(l)  H3O+ (aq) + OH- (aq) This reaction is the self-ionization of water, and happens in very small amounts. In pure water at 25◦C, [H+] = [OH-] = 1 x 10-7 M = 0.0000001 M Product of these two concentrations is called the ion product of water Kw = 1 x 10-14 = [H+] [OH-] The ion product of water is true for all acidic or basic solutions. H3O+ = H+

Relative concentrations [H+] [OH-] Kw 1.E-01 1.E-13 1.E-14 1.E-02 1.E-12 1.E-03 1.E-11 1.E-04 1.E-10 1.E-05 1.E-09 1.E-06 1.E-08 1.E-07

pH pH = -log [H+] [H+] Opp of Exp pH 1.E-01 1 1.E-02 2 1.E-03 3 1.E-04 5 1.E-06 6 1.E-07 7 1.E-08 8 1.E-09 9 1.E-10 10 1.E-11 11 1.E-12 12 1.E-13 13 Easier to refer to [H+] based on the exponent. More specifically the opposite of the exponent because they are negative. We call this value pH. Small pH values, have highest [H+] Mathematically, pH = -log [H+]

Acidic, Basic, or Neutral pH pH is by far the most common way to describe the acidity/basicity of a solution Always remember that pH is a log function: each increase by 1 is a factor of 10. Ex: pH = 7 pH = 3.2 pH = 8.2

pOH It is less common than pH, but you can similarly define pOH pOH = -log [OH-] Low pOH values are basic, high pOH values are acidic, 7 is neutral. pH +pOH = 14 Ex: If pOH = 4, pH = Ex: If pOH = 10, pH = Often easiest, to convert pOH to pH when identifying acidity.

pH Calculations Four related quantities: pH, pOH, [H+] and [OH-] pH = -log [H+] , pOH = -log [OH-], Kw = [H+] [OH-] = 1 x 10-14 , pOH + pH = 14 Note: the opposite of pH = -log [H+] is [H+] = 10 –pH Calculate the other three quantities for the given information. Then indicate if it is an acidic or basic solution. [H+] = 4.0 x 10-3 M [OH-] = 1.5 x 10-5 M pH = 2.1 pOH = 3.5

Exit Slip - Homework Exit Slip: An acidic solution has [H+] = 4.5 x 10-4 M. What are a) the pH b) the pOH and c) [OH-] What’s Due? (Pending assignments to complete.) Acid/Base pH worksheet What’s Next? (How to prepare for the next day) Read p496 - 520