The Chemistry of Acids and Bases

Acids pH below 7 turns litmus paper red taste sour reacts with metals to produce H2(g) generally starts with a hydrogen ion [H+] > [OH-] HCl

generally contains a hydroxide ion Bases pH greater than 7 turns litmus paper blue taste bitter feel slippery generally contains a hydroxide ion [H+] < [OH-] NaOH

Both Acids and Bases an electrolyte

Acidic, Basic, and Neutral Solutions Type of Solution pH Ranges [H+] versus [OH-] Example Acidic Below 7 [H+] > [OH-] Orange Juice Battery Acid Your Stomach Neutral Equals EXACTLY 7 [H+] = [OH-] Distilled Water Basic Above 7 [H+] < [OH-] Bleach Sea Water Blood

Indicators Indicators are compounds that have one color in acidic solutions and another in basic.

Litmus Paper Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

Paper or plastic strips that contain combinations of indicators estimate the pH of a solution by simply dipping a piece of pH paper into it and comparing the resulting color with standards printed on the container Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

Acid Nomenclature Flowchart

BINARY ACIDS HBr (aq) Hydrobromic Acid

Naming Ternary Acids TERNARY ACIDS SO42- H2SO4 H2SO4(aq) Sulfuric acid POLYATOMIC IONS PURE FORMS TERNARY ACIDS SO42- H2SO4 H2SO4(aq) Sulfuric acid SO32- H2SO3 H2SO3(aq) Sulfurous Acid

Naming Bases polyatomic Sodium hydroxide Calcium hydroxide Use the same rules as for ions (name the cation, then name the anion) NaOH  Ca(OH)2  KOH  Sodium hydroxide Calcium hydroxide Potassium hydroxide

Some Common Acids and Bases and their Household Uses.

What are Acids and Bases? There are two common definitions to describe acids and bases: Arrhenius acids and bases Bronsted-Lowry acids and bases These are basically the same although they state different things.

Definitions for Acids & Bases   Arrhenius Brønsted-Lowry Definition for Acids Definition for Bases Key Examples a proton producer in an aqueous solution a proton donor a hydroxide producer in an aqueous solution a proton acceptor Acid – HCl Base - NaOH Acid – HCl Base – NH3 H+ = proton

Arrhenius Acids and Bases Definitions acids in water produce hydronium ions, (H3O+, H+) HNO3(aq) H+(aq) + NO3- 2. Arrhenius Base bases in water produce hydroxide ions, (OH-) KOH(s) K+(aq) + OH-(aq)

Bronsted-Lowry Definitions Acids are proton (H+) donors. Bases are proton (H+) acceptors.  HCl + H2O  Cl– + H3O+ acid conjugate base base conjugate acid Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Bronsted-Lowry Come in Pairs General equation HA(aq) + H2O(l) A-(aq) + H3O+(aq) Acid + Base Conjugate base + Conjugate acid This is an equilibrium. B(aq) + H2O(l) BH+(aq) + OH-(aq) Base + Acid Conjugate acid +Conjugate base

What to Focus On? Arrhenius was the most restrictive definition. This definition required:   the solutions to be aqueous and a base to contain a hydroxide (OH-) ion. Bronsted-Lowry’s definition is the most commonly used. It is helpful to remember: acids tend to “lose“ an H+ ion, while bases tend to “gain“ an H+ ion. Under this definition, ammonia (NH3) is considered a base even though it is NOT an Arrhenius base.

Examples HCl(aq) + KOH(s) KCl(aq) + H2O(l)   3 Ca(OH)2(aq) + 2 H3PO4(aq)  Ca3(PO4)2(s) + H2O(l) F-(aq) + H2O(l)  HF(aq) + OH-(aq) HCO3-(aq) + H2O(l)  CO32-(aq) + H3O+(aq) NH4+(aq) + CO32-(aq)  NH3(aq) + HCO3-(aq) acid base base acid base acid conjugate acid conjugate base acid base conjugate base conjugate acid acid base conjugate base conjugate acid

Remember Electrolytes? Ionic Covalent C6H12O6 Na+ NaCl Cl-

Acids and bases are both strong or weak electrolytes (conduct electricity) • Electrolytes = dissociate (break apart into ions) when dissolved • Strong = completely Weak = partially Non = not at all Weak Strong H+ HC2H3O2 C2H3O21- H+ H-Cl Cl- Only a few Ions Lots of Ions

Strong Electrolytes WORD DESCRIPTION Completely breaks apart into its ions Are good conductors of electricity Will produce a bright light bulb Examples of Acids and Bases that are Strong Electrolytes Strong Acids Strong Bases H2SO4 NaOH HCl Ba(OH)2 Notice that all of the ions are separated or dissociated. The title is hyperlinked to a NCSSM animation of HCl dissolving in water.

Weak Electrolytes WORD DESCRIPTION Partially breaks apart into its ions Are poor conductors of electricity Will produce a dim light bulb Examples of Acids and Bases that are Weak Electrolytes Weak Acid HC2H3O2 (Vinegar) Weak Base NH3 (Ammonia) Notice that only some of the ions are separated or dissociated. The title is hyperlinked to a NCSSM animation showing HA (a weak acid) dissolving in water.

What makes a strong acid or a strong base? Strong electrolytes make strong acids and bases Strong Acids HCl - hydrochloric acid HBr - hydrobromic acid HI - hydroiodic acid HNO3 - nitric acid H2SO4 - sulfuric acid HClO4 - perchloric acid Strong Bases The hydroxides of the Group I and Group II LiOH - lithium hydroxide NaOH - sodium hydroxide KOH - potassium hydroxide *Ca(OH)2 - calcium hydroxide *Sr(OH)2 - strontium hydroxide *Ba(OH)2 - barium hydroxide The title is hyperlinked to a NCSSM animation of HCl dissolving in water.

pH Concept

pH Scale Pouvoir hydrogéne (hydrogen power) Is a scale to measure the acidity of a sample, Range: 0 -14 1 14 Highly acidic Very basic (not acidic) neutral 7 Neutral = 7.0 Bases 7-14 Acids 0-7

Relationships between pH, [H+], and [OH-] Click on the picture to get to the animation. As pH increases… The [H+] (increases or decreases). The [OH-] (increases or decreases). The solution becomes more (acidic or basic).

Relationships between pH, [H+], and [OH-] What happens as pH decreases? As pH decreases… The [H+] (increases or decreases). The [OH-] (increases or decreases). The solution becomes more (acidic or basic).

The pH Scale The value of pH is unitless. Solutions with a pH less than 7 are acidic and solutions greater than 7 are basic. If a solution is equal to 7 it is neutral. Here is a typical pH scale.

pH of Common Substances

pH is a Logarithmic Scale Logarithm –The number of times a base must be multiplied by itself to reach a given number # of multiples Base # you’re trying to reach

pH Calculations Given Solving for Formula to Use [H+] pH pH = - log[H+] [OH-] pOH pOH = - log[OH-] [H+] is the concentration of H+ ions, in mol/L.

Logarithms Use your calculator! If you have a log button, you’re all set. Each calculator can have its own method for entering logs. If you don’t know what to do your calculator manual should give examples. 1.44939 E -2 9 - 43

Logarithms If your calculator has a ln button - Don’t use it. Its for taking natural logs. This is different than base 10. 1.44939 E -2 9 - 44

Calculating pH If [H+] is written in scientific notation and has a coefficient of 1, then the pH of the solution equals the absolute value of the exponent Ex. 1.0 x 10-4 M pH = 4.0

Calculating pH Problem 1: If [H+] = 3.40 x 10-5 M, what is the pH? Given Unknown Equation [H+] = 3.40 x 10-5 M pH pH = - log[H+] Solve: pH = -log (3.40 x 10-5) pH = 4.47

Calculating pH Problem 2: If [H+] = 1 X 10-10, what is the pH? Given Unknown Equation [H+] = 1 X 10-10 pH pH = - log[H+] Solve: pH = - log 1 X 10-10 pH = - (- 10) pH = 10

Calculating pH Problem 3: If [H+] = 1.8 X 10-5, what is the pH? Given Unknown Equation [H+] = 1.8 X 10-5 pH pH = - log[H+] Solve: pH = - log 1.8 X 10-5 pH = - (- 4.74) pH = 4.74

Calculating pOH Solve: Problem 1: pOH = -log (2.3 x 10-12) pOH = 11.6 If [OH-] = 2.30 x 10-12 M, what is the pOH? Given Unknown Equation [OH-] = 2.30 x 10-12 M pOH pOH = - log[OH-] Solve: pOH = -log (2.3 x 10-12) pOH = 11.6

Calculating pOH If [OH-] is written in scientific notation and has a coefficient of 1, then the pOH of the solution equals the absolute value of the exponent Problem 2: If [OH-] = 1.0 x 10-9 M, what is the pH? pOH = 9.0

What’s in a glass of water? distilled

Distilled H2O at the Molecular Level What’s in a glass of distilled water? Water Molecules (H2O) Hydronium Ions (H3O+) Hydroxide Ions (OH-) What’s happens in the glass of water? H2O + H2O ⇆ H3O+ + OH- This is called the self-ionization of water.

pH + pOH = 14 Water Water ionizes- falls apart into ions. H2O ® H+ + OH-. Only a small amount. [H+ ] = [OH-] = 1 x 10-7M A neutral solution. In water Kw = [H+ ] x [OH-] = 1 x 10-14 Kw is called the ion product constant. pH + pOH = 14 Amphoteric a molecule or ion that can react as an acid as well as a base Ex: H2O, NH3

Calculating pOH from pH Problem 1: If the pH is 3.25, what is the pOH? Given pH = 3.25 Unknown pOH ? Equation pH + pOH = 14 Substitute and solve : 3.25 + pOH = 14 3.25 + (- 3.25) +pOH = 14 (- 3.25) pOH = 10.8

Calculating pH from pOH Problem 2: What is the pH of a solution if [OH-] = 4.0 x 10-11 M? Given [OH-] = 4.0 x 10-11 M Unknown pH? Equation pH + pOH = 14 Step 1: Find pOH pOH = -log [OH] pOH= -log[4.0 x 10-11 ] = 10.4 Step 2: Calculate pH pH + pOH= 14; pH = 14 – 10.4 pH = 3.6

Looking at the Math Given Solving for Formula to Use pH [H+] pOH [OH-]

Calculating [H+] from pH If the pH of Coke is 3.12, [H+] = ??? [ [H+] = 10-3.12 = 7.6 x 10-4 M *** to find antilog on your calculator, look for “Shift” or “2nd function” and then the log button Known pH = 3.12 Unknown [H+] ? Analysis [H+] = 10 -pH Substitute and solve :

Calculating [H+] from pH The pH of an unknown solution is 6.00. What is its [H+]? Known pH = 6.00 Unknown [H+] ? Analysis [H+] = 10 -pH Substitute and solve : [H+] = 1x 10 -6 M

Calculating [H+] from pH A solution has a pH of 8.5. What is the Molarity of hydrogen ions in the solution? Known pH = 8.5 Unknown [H+] ? Analysis [H+] = 10 -pH Substitute and solve : [H+] = 10-8.5 3.16 X 10-9 M

Acid-Base Reactions or Neutralization Reactions acid + base  water + salt HBr(aq) + NaOH(aq)  H2SO4(aq) + KOH(aq)  H3PO4(aq) + Ba(OH)2(aq)  H2O + NaBr H2O + K2SO4 H2O + Ba3(PO4)2 * Double replacement reactions