Unformatted text preview: ACIDS AND BASES BASES Goals
Compare/contrast Arrhenius and BronstedLowry definitions of acids and bases. Determine the strengths of acids and bases based on the degree of dissociation Relate dissociation and ionization constants to equilibrium constants Use the ion product constant for water to calculate the concentration of hydrogen and hydroxide ions. Acids and Bases Acids Examples
HCl: stomach acid Citric acid: citrus acid (lemon juice) Carbonic acid H2CO3: soda drinks Ammonia NH3: cleaner Sodium hydroxide NaOH: drain Sodium cleaner cleaner Magnesium hydroxide Mg(OH)2: milk Magnesium of magnesia of Properties Properties
PROPERTY Taste Touch Reaction w/ Reaction metal metal Electrical Electrical conductivity conductivity Indicators ACID Sour Sharp sting Reacts Reacts vigorously vigorously Yes BASE Bitter Slippery, Slippery, smooth smooth Does not react Yes Litmus: blue to Litmus: red to Litmus: Litmus: red blue red blue Neutralization Neutralization
Neutralization reaction: Mix the proper Neutralization amounts of acid and base together and the result is a solution that has neither the properties of the base or acid (it has been neutralized). neutralized). One product that forms is an ionic One compound (salt). (salt). Types
Strong acid + strong base = neutral Strong acid + weak base = weak acid Weak acid + strong base = weak base Weak acid + weak base = neutral The Arrhenius Definition The
Acid: A substance that dissociates in Acid: water to produce hydrogen ions (H+). water Base: A substance that dissociates Base: in water to produce hydroxide ions (OH-). (OH Acid – Base Neutralization: Acid + Base Water + Salt Examples: Hydrochloric acid + Sodium hydroxide Nitric acid + Potassium hydroxide The Arrhenius Definition Limitations
Water solutions only. Water Oversimplifies how acids dissolve in Oversimplifies water. water. Does not include some compounds Does that have the characteristics of a base. base. The Bronsted-Lowry Definition Definition
What is a H+ ion? Acid: Can donate H+ ion. Proton donor. donor. Base: Can accept H+ ion. Proton acceptor. acceptor. Monoprotic acid: An acid that can Monoprotic donate only one H+ ion. donate Diprotic acid: Two H+ ions. Triprotic acids: Three H+ ions. The Hydronium Ion (H3O ) The
1. 2. 3. 4. 5. 6. H+ + H2O H3O+ HCl + H2O H3O+ + ClCl Bronsted-Lowry perspective, when one Bronsted-Lowry compound acts as an acid (gives up a H+ compound ion), another compound must act as the base (accept the H+ ion), therefore HCl is base the acid and H2O is the base. the NH3 + H2O NH4+ + OHOH NH3 is the base and H2O is the acid. A substance such as water that can act substance as either an acid or a base is described as amphoteric. as Conjugate Acid-Base Pairs Conjugate
NH3 + H2O NH4+ + OHOH 2. What is the acid / base in the forward What direction? What is the acid / base in the reverse direction? reverse 3. When an acid loses an H+ ion it becomes its conjugate base and when a base gains an H+ it becomes its conjugate gains acid. acid. 4. Examples.
1. Strengths of Acids and Bases Bases
1M acetic acid vs. 1M HCl Strong acids readily transfer H+ ions to water to form H3O+ ions and react water completely (dissociates) to form ions. completely Substances that have a strong affinity for Substances H+ ions are called strong bases. The OHion is also a strong base. ion There is an inverse relationship between There the strengths of conjugate acid-base pairs. pairs.
Stronger acid = weaker conjugate base Stronger base = weaker conjugate acid Acid Dissociation Constant (Ka) (K
A measure of the strength of an acid. The larger the Ka the more the acid reacts with water to produce H3O+ reacts ions, therefore, the stronger the acid. ions, All the weak acids have Ka values less than one. less Ka = [H3O+] [Cl-] [HCl] [HCl] Base Dissociation Constant (Kb) (K
A measure of the strength of a base. The stronger the base, the larger the The concentration of hydroxide ions in the solution, therefore, the larger the Kb. solution, The dissociation constant of a weak base The is less than one. is Kb = [NH4+] [OH-] [NH3] [NH Ionization of H2O Ionization
H2O is not pure---it contains small is quantities of two important ions (H3O+ and quantities OH-). OH Water is amphoteric. Self-ionization of water. Water acts as Self-ionization both an acid and base in the same equation. equation. H2O + H2O H3O+ + OHNote the long and short arrows. In pure water at 25oC, both H3O+ ions and OH- ions are found at concentrations of 1.0 OH
-7 -7 Ion-Product Constant (Kw) Ion-Product
Kw = [H3O+] [OH-] Kw 7 7 = [ 1.0 x 10--7 ][ 1.0 x 10--7 ] ][ 1.0 = 1.0 x 10-14 1.0 Used to calculate concentrations of Used both H3O+ and OH- ions in a solution both OH and to determine if the solution is acidic or basic. acidic Sample problems. Ion-Product Constant (Kw) Ion-Product
Acidic: [H3O+] > 1x10-7 M > [OH-] Neutral: [H3O+] = 1x10-7 M = [OH-] Neutral: Basic: [H3O+] < 1x10-7 M < [OH-] Basic: Goals (part II)
Calculate pH Use pH to determine the concentration of hydrogen and hydroxide ions Define a buffer solution Explain the purpose and process of an acid/base titration. The pH Scale The
pH = -log [H3O+] pH = -log [H+] pOH = -log [OH-] Strong acid = 1, weak acid = 5, neutral = 7, weak base = 10, strong base = 14. base Practice problems. Measuring pH: acid-base indicators Measuring or pH meter. or Buffers Buffers
Control the pH of a solution with a buffer. A buffer is a mixture that is able to release buffer or absorb H+ ions, keeping a solution’s pH or constant. constant. Human body pH is 7.35 to 7.45. The amount of acid or base that a buffer The can neutralize is called the buffer capacity. can ...
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