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Chapter 15
Acids and Bases
| Term | Definition |
|---|---|
| General Properties of Acids | Sour taste, the ability to dissolve many metals, the ability to turn blue litmus paper red, and the ability to neutralize bases |
| 1. Common Acids | 1. Hydrochloric Acid(HCl).. Uses: Metal cleaning, food preparation, ore refining, primary component of stomach acid .. 2. Sulfuric Acid(H2SO4).. Uses: Fertilizer/explosives manufacturing, dye/glue production, automobile batteries, electroplating copper |
| 2. Common Acids | 3. Nitric Acid(HNO3).. Uses: Fertilizer/explosives manufacturing, dye/glue production.., 4. Acetic Acid(HC2H3O2).. Uses: Plastic/rubber manufacturing, food preservative, active component of vinegar |
| 3. Common Acids | 5. Citric Acid(H3C6H5O7).. Uses: Present in citrus fruits such as lemons and limes, used to adjust pH in foods and beverages... 6. Carbonic Acid(H2CO3).. Uses: Found in carbonated beverages due to the reaction of carbon dioxide with water |
| 4. Common Acids | 7. Hydrofluoric Acid(HF).. Uses: Metal cleaning, glass frosting etching... 8. Phosphoric Acid(H3PO4).. Uses: Fertilizer manufacture, biological buffering, preservative in beverages |
| Carboxylic Acid | An organic acid containing the functional group - COOH. An H single bonded to an O which is single bonded to a C which is double bonded to an O |
| 1. Common Bases | 1. Sodium Hydroxide(NaOH).. Uses: Petroleum processing, soap and plastic manufacturing... 2. Potassium Hydroxide(KOH).. Uses: Cotton processing, electroplating, soap production, batteries |
| 2. Common Bases | 3. Sodium Bicarbonate(NaHCO3).. Uses: Antacid, ingredient of baking soda, source of CO2... 4. Sodium Carbonate(Na2CO3).. Uses: Manufacture of glass and soap, general cleanser, water softener |
| 3. Common Bases | 5. Ammonia(NH3).. Uses: Detergent, fertilizer and explosives manufacturing, synthetic fiber production |
| General Properties of Bases | Bitter taste, a slippery feel, the ability to turn red litmus paper blue, and the ability to neutralize acids |
| Alkaloids | Organic bases found in plants that are often poisonous. Our aversion to the taste of bases is probably an evolutionary adaption to warn us against alkaloids |
| Arrhenius Definition of Acids and Bases | The definitions of an acid as a substance that produces H^+ ions in aqueous solution and a base as a substance that produces OH^- ions in aqueous solution |
| Hydronium ion | H3O^+, the ion formed from the association of a water molecule with an H^+ ion donated by an acid. H3O^+ and H^+ are often used interchangeable |
| Bronsted-Lowry Definition of Acids and Bases | The definition of an acid as a proton(ion) donor and a base as a proton acceptor |
| Amphoteric | Substances that can act as acids or bases. In other words, can act as a proton donor in one reaction, and a proton acceptor in another. Ex: water |
| Conjugate Acid-Base Pair | Two substances related to each other by the transfer of a proton |
| Conjugate Acid | Any base to which a proton has been added |
| Conjugate Base | Any acid from which a proton has been removed |
| Summarizing the Bronsted-Lowry Definition of an Acid-Base Reaction | 1. A base accepts a proton and becomes a conjugate acid.. 2. An acid donates a proton and becomes a conjugate base |
| Strong Acid | An acid that completely ionizes in solution |
| Weak Acid | An acid that does not completely ionize in solution |
| Monoprotic Acid | An acid that contains only one ionizable proton |
| Diprotic Acid | An acid that contains two ionizable protons |
| Strong Acids | 1. Hydrochloric Acid(HCl).. 2. Hydrobromic Acid(HBr).. 3. Hydroiocid Acid(HI).. 4. Nitric Acid(HNO3).. 5. Perchloric Acid(HClO4).. 6. Sulfuric Acid(H2SO4)(diprotic) |
| Ionic Attraction and Acid Strength | In a strong acid(HA), the attraction between H^+ and A^- is weak, resulting in complete ionization. In a weak acid, the attraction between H^+ and A^- is strong, resulting in only partial ionization |
| Triprotic Acid | An acid that contains three ionizable protons |
| Some Weak Acids | 1. Hydrofluoric Acid(HF).. 2. Acetic Acid(HC2H3O2).. 3. Formic Acid(HCHO2).. 4. Sulfurous Acid(H2SO3)(diprotic).. 5. Carbonic Acid(H2CO3)(diprotic).. 6. Phosphoric Acid(H3PO4)(triprotic) |
| Acid Ionization Constant(Ka) | The equilibrium constant for the ionization reaction of a weak acid; used to compare the relative strengths of weak acids... |
| Acid Ionization Constant(Ka) Equation | For the two equivalent reactions.. HA(aq) + H2O(l) <----> H3O^+ (aq) + A^- (aq)... HA(aq) <----> H^+ (aq) + A^- (aq)... Ka = ([H3O^+][A^-]) / [HA] = ([H^+][A^-]) / [HA] |
| Autoionization | The process by which water acts as an acid and a base with itself... H2O(l) + H2O(l) <----> H3O^+ (aq) + OH^- (aq)... or... H2O(l) <----> H^+ (aq) + OH^- (aq) |
| Ion Product Constant for Water(Kw)(Dissociation Constant for Water) | The equilibrium constant for the autoionization of water... Kw = [H3O^+][OH^-] = [H^+][OH^-] = 1.0 * 10^-14 @ 25 degrees celsius |
| Neutral | The state of a solution where the concentrations of H3O^+ and OH^- are equal. Pure water is neutral |
| Acidic Solution | A solution containing an acid that creates additional H3O^+ ions, causing the concentration of H3O^+ to increase and that of OH^- to decrease. The ion product constant still applies... [H3O^+] * [OH^-] will always equal 1.0 * 10^-14 @ 25 degrees celsius |
| Basic Solution | A solution that contains a base that creates additional OH^- ions, causing [OH^-] to increase and [H3O^+] to decrease but again the ion product constant still applies |
| 1. Summarizing Kw | 1. A neutral solution contains [H3O^+] = [OH^-] = 1.0 * 10^-7 M (at 25 degrees celsius)... 2. An acidic solution contains [H3O^+] > [OH^-]... 3. A basic solution contains [OH^-] > [H3O^+]... |
| 2. Summarizing Kw | 4. In all aqueous solution both H3O^+ and OH^- are present, with [H3O^+][OH^-] = Kw = 1.0 * 10^-14 (at 25 degrees celsius) |
| The pH(Power of Hydrogen) Scale | A compact way to specify the acidity of a solution... pH = -log[H3O^+]... In general, at 25 degrees celsius.. 1. pH < 7 The solution is acidic.. 2. pH >7 The solution is basic.. 3. pH = 7 The solution is neutral |
| The pOH(Power of Hydroxide) Scale | The pOH scale is analogous to the pH scale but is defined with respect to [OH^-] instead of [H3O^+]... pOH = -log[OH^-]... 1. pOH < 7 The solution is basic.. 2. pOH > 7 The solution is acidic.. 3. pOH = 7 The solution is neutral |
| Relationship Between pH and pOH | pH + pOH = 14.00.. The sum of pH and pOH is always equal to 14.00 at 25 degrees celsius. Therefore, a solution with a pH of 3 has a pOH of 11. |
| Percent Ionization | The concentration of ionized acid in a solution divided by the initial concentration of acid multiplied by 100... Percent Ionization = [(concentration of ionized acid) / (initial concentration of acid)] * 100 = ([H3O^+]equil / [HA]init) x 100 |
| Relationship Between Initial Concentration of Weak Acid and Percent Ioniization | 1. The equilibrium H3O^+ concentration of a weak acid increases with increasing initial concentration of the acid... 2. The percent ionization of a weak acid decreases with increasing concentration of the acid |
| Strong Base | A base that completely dissociates in solution. Most strong bases are group 1A and 2A metal hydroxides. Unlike diprotic acids, which ionize in two steps, bases containing two OH^- ions dissociate in one step |
| Common Strong Bases | 1. Lithium Hydroxide(LiOH).. 2. Sodium Hydroxide(NaOH).. 3. Potassium Hydroxide(KOH).. 4. Strontium Hydroxide(Sr(OH)2).. 5. Calcium Hydroxide(Ca(OH)2).. 6. Barium Hydroxide(Ba(OH)2) |
| Weak Base | A base that only partially ionizes in water. Unlike strong bases that contain OH^- and dissociate in water, the most common weak bases produce OH^- by accepting a proton from water, ionizing water to form OH^- |
| Base Ionization Constant(Kb) | The equilibrium constant for the ionization reaction of a weak base; used to compare the relative strengths of weak bases. |
| Base Ionization Constant(Kb) Equation | For the general reaction in which a weak base ionizes water, we define Kb as follows: B(aq) + H2O(l) <---> BH^+(aq) + OH^-(aq)... Kb = ([BH^+][OH^-]) / [B] |
| Common Weak Bases | 1. Carbonate Ion(CO3^2-).. 2. Methylamine(CH3NH2).. 3. Ethylamine(C2H5NH2).. 4. Ammonia(NH3).. 5. Bicarbonate Ion(HCO3^-).. 6. Pyridine(C5H5N).. 7. Aniline(C6H5NH2) |
| Lone Pairs in Weak Bases | Many weak bases have a nitrogen atom with a lone pair that acts as the proton acceptor |
| Anions as Weak Bases | We can think of any anion as the conjugate base of an acid. 1. Cl^- is conjugate base of HCl.. 2. F^- is conjugate base of HF.. 3. NO3^- is the conjugate base of HNO3.. 4. C2H3O2^- is the conjugate base of HC2H3O2 |
| Anions as Conjugate Bases | The anion A^- is conjugate base of acid HA. Since every anion can be envisioned as conjugate base of an acid, every anion itself can potentially act as base. However, not every anion does act as base, it depends on strength of corresponding acid |
| Anions as Conjugate Bases Continuation | 1. An anion that is the conjugate base of a weak acid is itself a weak base.. 2. An anion that is the conjugate base of a strong acid is pH-neutral(forms solutions that are neither acidic nor basic) |
| Strength of Conjugate Acid - Base Pairs | The weaker the acid, the stronger the conjugate base. In contrast, the conjugate base of a strong acid, such as Cl^-, does not act as a base |
| Cations as Weak Acids | Cations can in some cases act as weak acids. WE can generally divide cations into three categories: cations that are the conterions of strong bases; cations that are the conjugate acids of weak bases; and cations that are small, highly charged metals. |
| Cations That Are the Counterions of Strong Bases | Strong bases generally contain hydroxide ions and a counterion. Although these counterions interact with water molecules via ion - dipole forces when the strong base dissociates, they do not ionize water and are themselves pH - neutral |
| 1. Cations That Are the Conjugate Acids of Weak Bases | A cation can be formed from any nonionic weak base by adding a proton(H^+) to its formula. The cation will be the conjugate acid of the base. Consider the following cations and their corresponding weak bases... |
| 2. Cations That Are the Conjugate Acids of Weak Bases | 1. NH4^+ is the conjugate acid of NH3(weak base).. 2. C2H5NH3^+ is the conjugate acid of C2H5NH2(weak base).. 3. CH3NH3^+ is the conjugate acid of CH3NH2(weak base)... Any of these cations, with the general formula BH^+, will act as a weak acid. |
| 3. Cations That Are the Conjugate Acids of Weak Bases | A cation that is the conjugate acid of a weak base is a weak acid |
| 1. Cations That Are Small, Highly Charged Metals | Small, highly charged metal cations such as Al^3+ and Fe^3+ form weakly acidic solutions. For examply, when Al^3+ is dissolved in water, it becomes hydrated according to the equation: Al^3+(aq) + 6H2O(l) --> Al(H2O)6^3+(aq)... |
| 2. Cations That Are Small, Highly Charged Metals | Which then acts as a Bronsted-Lowry acid: Al(H2O)6^3+(aq) + H2O(aq) <---> Al(H2O)5(OH)^2+(aq) + H3O^+(aq)... Neither alkali metal cations nor alkaline earth metals cations ionize water in this way, bu the cation of many other metals do. |
| 3. Cations That Are Small, Highly Charged Metals | The smaller and more highly charged the cation, the more acidic its behavior |
| 1. Classifying Salt Solutions as Acidic, Basic, or Neutral | Four possibilities... 1. Salts in which neither the cation nor the anion acts as an acid or a base form pH-neutral solutions.. 2. Salts in which the cation does no act as an acid and the anion acts as a base form basic solutions |
| 2. Classifying Salt Solutions as Acidic, Basic, or Neutral | 3. Salts in which cation acts as acid and anion does not act as base form acidic solutions.. 4. Salts in which cation acts as acid and anion acts as base form solutions in which pH depends on relative strengths of the acid and the base |
| Salts in Which Neither the Cation nor the Anion Acts as an Acid or a Base Form pH-Neutral Solutions | A salt in which the cation is the counterion of a strong base in which the anion is the conjugate base of strong acid will from neutral solutions. Ex: NaCl(sodium chloride), Ca(NO3)2(calcium nitrate), KBr(potassium bromide) |
| Salts in Which the Cations Does not Act as an Acid and the Anion Acts as a Base Form Basic Solutions | A salt in which the cation is the counterion of a strong base in which the anion is the conjugate base of a weak acid will form basic solutions. Ex: NaF(sodium flouride), Ca(C2H3O2)2(calcium acetate), KNO2(potassium nitrite) |
| Salts in Which the Cation Acts as an Acid and the Anion does not Act as a Base Form Acidic Solutions | A salt in which cation is either conju. acid of weak base or small, highly charged metal ion and in which anion is conju. base of strong acid will form acidic solutions. Ex: FeCl3(iron(III) chloride), Al(NO3)3(aluminum nitrate), NH4Br(ammonium bromide) |
| Salts in Which the Cations Acts as an Acid and the Anion Acts as a Base Form Solutions in Which the pH Depends on the Relative Strengths of the Acid and the Base. Ex: FeF3(iron(III) fluoride), Al(C2H3O2)3(aluminum acetate), NH4NO2(ammonium nitrite) | A salt in which the cation is either the conjugate acid of a weak base or a small, highly charged metal ion in which the anion is the conjugate base of a weak acid will form a solution in which the pH depends on the relative strengths of the acid and base |
| Polyprotic Acid | An acid that contains more than one ionizable proton and releases them in successive steps, each with its own Ka... H2SO3(aq) <---> H^+(aq) + HSO3^-(aq) Ka = 1.6*10^-2... HSO3^- <---> H^+(aq) + SO3^-2(aq) Ka2 = 6.4 *10^-8 |
| Binary Acids | Consider the bond between a hydrogen atom and some other generic element(called Y)... H - Y... The factors affecting the ease with which this hydrogen will be donated(and therefore be acidic) are the polarity of the bond and the strength of the bond |
| Bond Polarity in Binary Acids | The H-Y bond must be polarized with the hydrogen atom as the positive pole in order for HY to be acidic... δ+H - Yδ- |
| Bond Strength in Binary Acids | The strength of the H-Y bond also affects the strength of the corresponding acid. The stronger the bond, the weaker the acid --- the more tightly the hydrogen atom is held, the less likely it is to come off. |
| Acidity of the Group 6A and 7A Hydrides | From left to right, the hydrides become more acidic because the H-Y bond becomes more polar. From top to bottom, these hydrides become more acidic because the H-Y bond becomes weaker |
| The Combined Effect of Bond Polarity and Bond Strength for Binary Acids | The hydrides become more acidic from left to right as the H-Y bond becomes more polar. The hydrides also become more acidic from top to bottom as the H-Y bdon becomes weaker |
| Oxyacids | Oxyacids contain a hydrogen atom bonded to an oxygen atom. The oxygen atom is in turn bonded to another atom(which we will call Y)... H-O-Y---... Y may or may not be bonded to yet other atoms. |
| Factors Affecting Acidity of Oxyacids | The factors affecting the ease with which the hydrogen in an oxyacid will be donated( and therefore be acidic) are the electronegativity of the element Y and the number of oxygen atoms attached to the element Y |
| The Electronegativity of Y in Oxyacids | The more electronegative the element Y, the more it weakens and polarized the H-O bond and the more acidic the oxyacid is |
| The Number of Oxygen Atoms Bonded to Y | Oxyacids may have additional oxygen atoms bonded to element Y. Additional oxygen atoms are electronegative so they draw electron density away from Y, which draws electron density away from the H-O bond, weakening and polarizing it and increasing acidity |
| Lewis Acids and Bases | Lewis Acid: Electron pair acceptor... Lewis Base: Elecctron pair donor |
| Important Property of Lewis Acids | A Lewis acid has an empty orbital( or can rearrange electrons to create an empty orbital) that can accept an electron pair |
Created by:
TimChemistry2
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