Solutions Unit Honors Chemistry

Naming Acids Review: Ex) HCl hydrochloric acid A. Binary – H +one anion Prefix “hydro”+ anion name +“ic”acid Ex) HCl hydrochloric acid Ex) H3P hydrophosphoric acid B. Tertiary – H + polyatomic anion no Prefix “hydro” (oxo) end “ate” = “ic” acid end “ite” = “ous” acid Ex) H2SO4 sulfuric acid Ex) H2SO3 sulfurous acid

Properties of Acids and Bases: Taste Touch Reactions with Metals Electrical Conductivity Acid sour looks like water, burns, stings Yes-produces H2 gas electrolyte in solution Base (alkali) bitter looks like water, feels slippery No Reaction

Indicators: Turn 1 color in an acid and another color in a base. Litmus Paper: Blue and Red An aciD turns blue litmus paper reD A Base turns red litmus paper Blue. Phenolphthalein: colorless in an acid and pink in a base pH paper: range of colors from acidic to basic pH meter: measures the concentration of H+ in solution

Reactions Neutralization: A reaction between an acid and base. When an acid and base neutralize, water and a salt (ionic solid) form. Acid + Base → Salt + Water Ex) HCl + NaOH → NaCl + HOH

Arrhenius Definition (1884): A. An acid dissociates in water to produce more hydrogen ions, H+. HCl  H+1 + Cl-1 B. A base dissociates in water to produce more hydroxide ions, OH-. NaOH  Na+1 + OH-1 C. Problems with Definition: • Restricts acids and bases to water solutions. • Oversimplifies what happens when acids dissolve in water. • Does not include certain compounds that have characteristic properties of acids & bases. Ex) NH3 (ammonia) doesn’t fit

Bronsted-Lowry Definition (1923): A. An acid is a substance that can donate hydrogen ions. Ex) HCl → H+ + Cl- Hydrogen ion is the equivalent of a proton. Acids are often called proton donors. Monoprotic (HCl), diprotic (H2SO4) , triprotic (H3PO4) B. A base is a substance that can accept hydrogen ions. Ex) NH3 + H+ → NH4+ Bases are often called proton acceptors. C. Advantages of Bronsted-Lowry Definition •Acids and bases are defined independently of how they behave in water. •Focuses solely on hydrogen ions.

Hydronium Ion: Hydronium Ion – H3O+ This is a complex ion that forms in water. H+1 + H2O  H3O+1 To more accurately portray the Bronsted-Lowry, the hydronium ion is used instead of the hydrogen ion.

STRONG Acid/Base versus WEAK Acid/Base Strength refers to the % of molecules that form IONS. A strong acid or base will completely ionize (>95% as ions). This is represented by a single () arrow. HNO3 + H2O  H3O+ + NO3- A weak acid or base will partially ionize (<5% as ions). This is represented by a double (↔) arrow. HOCl + H2O ↔ H3O+ + ClO-

HF < HCl < HBr < HI increasing strength 7 Strong Acids HNO3 H2SO4 HClO3 HClO4 HCl HBr HI 8 Strong Bases LiOH NaOH KOH RbOH CsOH Ca(OH)2 Sr(OH)2 Ba(OH)2

Strength vs. Concentration Strength refers to the percent of molecules that form ions Concentration refers to the amount of solute dissolved in a solvent. Usually expressed in molarity.

Ionization of Acids & Bases H2SO4  2 H+ + SO4-2 Sulfuric acid H3PO3  Phosphorous acid Ca(OH)2  Calcium hydroxide 3 H+ + PO3-3 Ca+2 + 2 OH-1

Conjugate Acid-Base Pairs: A pair of compounds that differ by only one hydrogen ion Acid donates a proton to become a conjugate base. Base accepts proton to become a conjugate acid. A strong acid will have a weak conjugate base. A strong base will have a weak conjugate acid.

Acid (A), Base (B), Conjugate Acid (CA), Conjugate Base (CB) NH3 + H2O ↔ NH4+ + OH- HCl + H2O ↔ Cl- + H3O+ Base and Conjugate Acid are a Conjugate Pair. Acid and Conjugate Base are a Conjugate Pair. B A CA CB A B CB CA

AciDonates & Bases accept H2O + H2O ↔ H3O+ + OH− B A CA CB H2SO4 + OH− ↔ HSO4− + H2O A B CB CA HSO4− + H2O ↔ SO4−2 + H3O+ A B CB CA OH− + H3O+ ↔ H2O + H2O B A CA CB

The Self-ionization of Water & pH 1. Water is amphoteric, it acts as both an acid and a base in the same reaction. Ex) H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq) Keq = equilibrium constant = [H3O+] [OH-] Because reactants and products are at equilibrium, liquid water is not included in the equilibrium expression @ 25C, [H3O+] = 1 x 10-7 M and [OH-] = 1 x 10-7 M Kw = ion product constant or equilibrium constant for water Kw = [H3O+] [OH-] = 1 x 10-14 M2 1.0 x 10-14 M2 = [1.0 x 10-7 M] [1.0x10-7 M] 1.0 x 10-14 = [H3O+] [OH-]

Acids: [H3O+] > 1 x 10-7 M Bases: [OH-] > 1 x 10-7 M Using Kw in calculations: If the concentration of H3O+ in the blood is 4.0 x 10-8 M, what is the concentration of OH­ ions in the blood? Is blood acidic, basic or neutral? Kw = [H3O+] [OH-] 1.0 x 10-14 M2 = [4.0 x 10-8 M] [OH-] 2.5 x 10-7 M = [OH-] slightly basic

The pH scale (1909): the power of Hydrogen Measure of H3O+ in solution. pH = -log[H3O+] Range of pH: 0-14 pH < 7: acid pH > 7: base pH = 7: neutral pOH = -log[OH-] pH + pOH = 14

OH- H+ 14 1 pH [H3O+] [OH-] 14 1x10-14 1x100 13 1x10-13 1x10-1 12 11 1x10-11 1x10-3 10 1x10-10 1x10-4 9 1x10-9 1x10-5 8 1x10-8 1x10-6 7 1x10-7 6 5 4 3 2 1 14 H+ OH- 1

Significant Digits Rule The number of digits AFTER THE DECIMAL POINT in your answer should be equal to the number of significant digits in your original number Ex -log[8.7x10-4M] Calc Answer = 3.0604807474 Sig Fig pH = 3.06

Concentration Percent concentration by mass (mass % Molarity (M) (solute/solution) x 100% = % Concentration Molarity (M) Moles of solute/Liters of solution = mol/L Molality (m) Moles of solute/mass of solvent = mol/kg ppm and ppb Used for very dilute solutions Dilution – a process in which more solvent is added to a solution How is this solution different? Volume, color, molarity How is it the same? Same mass of solute, same moles of solute In Dilution ONLY – M1V1 = M2V2

Dissolution Process Ionic Compounds NaCl(s)  Na+1(aq) + Cl-1(aq) For dissolution to occur, must overcome solute attractions and solvent attractions. Dissociation Reaction: the separation of IONS when an ionic compound dissolves (ions already present) Try calcium chloride electrolyte nonelectrolyte Dissolving NaCl in water hexahydrated for Na+1; most cations have 4-9 H2O molecules 6 is most common Solvation: process of solvent molecules surrounding solute Hydration: solvation with water

V. Solution Stoichiometry A. Many reactants are introduced to a reaction chamber as a solution. B. The most common solution concentration is molarity. molarity = mol/liter C. Examples Excess lead(II) carbonate reacts with 27.5 mL of 3.00M nitric acid. Calculate the mass of lead(II) nitrate formed PbCO3 + 2HNO3  Pb(NO3)2 + H2CO3

2. Calculate the volume, in mL, of a 0 2. Calculate the volume, in mL, of a 0.324 molar solution of sulfuric acid required to react completely with 2.792 g of sodium carbonate according to the equation below. H2SO4 + Na2CO3  Na2SO4 + CO2 + H2O

Energy Changes Heat of solution = Hsoln Endothermic Exothermic Solute particles separating in solid Solvent particles moving apart to allow solute to enter liquid Energy absorbed Exothermic Solvent particles attracted to solvating solute particles Energy released