CHEM 120B Acid / Base Equilibrium

Flashcard 1 (Front): Define an Arrhenius acid. Flashcard 1 (Back): An Arrhenius acid is a compound that produces H+ ions in aqueous solutions1 .

Flashcard 2 (Front): Define an Arrhenius base. Flashcard 2 (Back): An Arrhenius base is a compound that produces OH– ions in aqueous solutions2 .

Flashcard 3 (Front): What is a hydronium ion? Flashcard 3 (Back): A hydronium ion is formed when an H+ ion (proton) reacts with a water molecule, often represented as H3O+2 . It's a simplified representation of more complex ions like H(H2O)n+2 .

Flashcard 4 (Front): Define a Brønsted–Lowry acid. Flashcard 4 (Back): A Brønsted–Lowry acid is a H+ ion donor3 .

Flashcard 5 (Front): Define a Brønsted–Lowry base. Flashcard 5 (Back): A Brønsted–Lowry base is a H+ ion acceptor3 .... For this to occur, the base structure must contain an atom with an unshared (lone) pair of electrons4 ....

Flashcard 6 (Front): What is the acid ionization constant, Ka? Flashcard 6 (Back): The acid ionization constant (Ka) is the equilibrium constant for the reaction of an acid with water3 .... For the reaction HA(aq) + H2O(ℓ) ⇌ A–(aq) + H3O+(aq), Ka = [A–][H3O+]/[HA]3 .

Flashcard 7 (Front): How does Ka relate to acid strength? Flashcard 7 (Back): The larger the Ka value, the stronger the acid6 . Strong acids have Ka >> 1, while weak acids have Ka < 13 .

Flashcard 8 (Front): How do strong acids behave in water? Flashcard 8 (Back): Strong acids are completely ionized in water3 . For example, HNO3(ℓ) + H2O(ℓ) → NO3 –(aq) + H3O+(aq)3 .

Flashcard 9 (Front): How do weak acids behave in water? Flashcard 9 (Back): Weak acids are partially ionized in water6 . For example, HNO2(ℓ) + H2O(ℓ) ⇌ NO2 –(aq) + H3O+(aq)6 .

Flashcard 10 (Front): What is the characteristic functional group of carboxylic acids? Flashcard 10 (Back): Carboxylic acids have a COOH group6 . Only the H atom in the COOH group is acidic7 .

Flashcard 11 (Front): How does the number of oxygen atoms affect the strength of oxoacids? Flashcard 11 (Back): The more oxygen atoms in an oxoacid, the more stable the conjugate base due to delocalization of the negative charge, and therefore the stronger the acid8 . For example, nitric acid is stronger than nitrous acid8 .

Flashcard 12 (Front): How does electronegativity affect acid strength? Flashcard 12 (Back): The source mentions electronegativity and acid strength under "Structure and Acid Strength"9 , implying a relationship, but doesn't explicitly define it. (Further details on this relationship are likely outside the provided sources).

Flashcard 13 (Front): List examples of strong bases. Flashcard 13 (Back): Strong bases are typically hydroxides of Group I and II metal ions, such as MOH(s) → M+(aq) + OH–(aq)4 .

Flashcard 14 (Front): How do weak bases ionize in water? Flashcard 14 (Back): Weak bases partially ionize in water to produce hydroxide ions4 . For example, NH3(aq) + H2O(ℓ) ⇌ NH4 +(aq) + OH–(aq)4 .

Flashcard 15 (Front): What is Kb? Flashcard 15 (Back): Kb is the base ionization constant, which measures the strength of a weak base in water4 . For NH3, Kb = [NH4+][OH–]/[NH3] = 1.76 × 10–54 .

Flashcard 16 (Front): What are amines? Flashcard 16 (Back): Amines are derivatives of ammonia (NH3) that contain a basic nitrogen atom4 .

Flashcard 17 (Front): Define a conjugate acid. Flashcard 17 (Back): A conjugate acid is formed when a base accepts a proton (H+)5 .

Flashcard 18 (Front): Define a conjugate base. Flashcard 18 (Back): A conjugate base is formed when an acid donates a proton (H+)5 .

Flashcard 19 (Front): Give an example of a conjugate acid-base pair. Flashcard 19 (Back): HF(aq) (acid) and F–(aq) (conjugate base)10 , or NH3(aq) (base) and NH4+(aq) (conjugate acid)4 .

Flashcard 20 (Front): What is the leveling effect? Flashcard 20 (Back): In water, H3O+ is the strongest H+ donor that can exist. Strong acids are all completely converted to H3O+ ions, so they all have the same strength in water11 .

Flashcard 21 (Front): What is the autoionization of water? Flashcard 21 (Back): Water can act as both an acid and a base, undergoing self-ionization: 2 H2O(ℓ) ⇌ H3O+(aq) + OH–(aq)12 . The equilibrium constant for this is Kw11 .

Flashcard 22 (Front): What is the value of Kw at 25°C? Flashcard 22 (Back): At 25°C, the ion product of water, Kw = [H3O+][OH–] = 1.0 × 10–1411 . In pure water at 25°C, [H3O+] = [OH–] = 1.0 × 10–7 M11 .

Flashcard 23 (Front): Define pH. Flashcard 23 (Back): pH is a measure of the acidity or basicity of a solution and is defined as pH = −log[H3O+]11 .

Flashcard 24 (Front): What is a neutral pH at 25°C? Flashcard 24 (Back): A neutral solution at 25°C has pH = 711 , because [H3O+] = 1.0 × 10–7 M11 .

Flashcard 25 (Front): How does pH relate to acidity and basicity? Flashcard 25 (Back): pH < 7 is acidic, pH > 7 is basic13 .

Flashcard 26 (Front): How can you determine [H3O+] if you know the pH? Flashcard 26 (Back): [H3O+] = 10−pH13 ....

Flashcard 27 (Front): Define percent ionization for a weak acid. Flashcard 27 (Back): Percent Ionization = ([H3O+]equilibrium / [HA]initial) × 100%13 .

Flashcard 28 (Front): How does the percent ionization of a weak acid change with increasing acid concentration? Flashcard 28 (Back): Percent ionization decreases as [acid] increases13 . This is explained by Le Châtelier’s principle13 .

Flashcard 29 (Front): What happens to the [H3O+] in a more dilute weak acid solution compared to the initial acid concentration? Flashcard 29 (Back): The result will be a larger [H3O+] in the dilute solution compared to the initial acid concentration due to a larger percent ionization15 .

Flashcard 30 (Front): What are polyprotic acids? Flashcard 30 (Back): Polyprotic acids are acids that have more than one ionizable H atom per molecule, such as H2SO416 . They have multiple Ka values (Ka1, Ka2, etc.)16 .

Flashcard 31 (Front): In a polyprotic acid, how do the Ka values generally compare? Flashcard 31 (Back): Generally, Ka1 > Ka2 > Ka317 because it becomes more difficult to remove a H+ ion from a negatively charged anion17 .

Flashcard 32 (Front): For a polyprotic acid, which ionization step typically most significantly affects the pH? Flashcard 32 (Back): Typically, only the first H+ significantly affects pH (Ka1)17 .

Flashcard 33 (Front): What can an anion be thought of as in terms of acids and bases? Flashcard 33 (Back): Every anion can be thought of as the conjugate base of an acid17 . Therefore, every anion can potentially be a base18 .

Flashcard 34 (Front): How does the strength of an acid relate to the strength of its conjugate base? Flashcard 34 (Back): The stronger the acid, the weaker the conjugate base18 .

Flashcard 35 (Front): What is the pH behavior of an anion that is the conjugate base of a strong acid? Flashcard 35 (Back): An anion that is the conjugate base of a strong acid is pH neutral18 . Example: Cl– (conjugate base of strong acid HCl)18 .

Flashcard 36 (Front): What is the pH behavior of an anion that is the conjugate base of a weak acid? Flashcard 36 (Back): An anion that is the conjugate base of a weak acid is basic18 . Example: F– (conjugate base of weak acid HF)18 .

Flashcard 37 (Front): What can some cations be thought of as in terms of acids and bases? Flashcard 37 (Back): Some cations can be thought of as the conjugate acid of a weak base18 . Others are counterions of a strong base18 . Therefore, some cations can potentially be acidic18 .

Flashcard 38 (Front): How does the strength of a base relate to the strength of its conjugate acid? Flashcard 38 (Back): The stronger the base, the weaker the conjugate acid19 .

Flashcard 39 (Front): What is the pH behavior of a cation that is the counterion of a strong base? Flashcard 39 (Back): A cation that is the counterion of a strong base is pH neutral19 . Example: Na+ (counterion of strong base NaOH)19 .

Flashcard 40 (Front): What is the pH behavior of a cation that is the conjugate acid of a weak base? Flashcard 40 (Back): A cation that is the conjugate acid of a weak base is acidic19 . Example: NH4+ (conjugate acid of weak base NH3)19 .

Flashcard 41 (Front): What is the pH behavior of alkali metal and alkaline earth metal cations? Flashcard 41 (Back): Alkali metal cations and alkaline earth metal cations are pH neutral19 .

Flashcard 42 (Front): How is the Ka of an acid related to the Kb of its conjugate base? Flashcard 42 (Back): The relationship is Ka × Kb = Kw20 .

Flashcard 43 (Front): Define ocean acidification. Flashcard 43 (Back): Ocean acidification is the decrease in the pH of the ocean due to the absorption of increased atmospheric CO221 . CO2 dissolves in water to form carbonic acid (H2CO3), which then forms H3O+21 .

Flashcard 44 (Front): How does ocean acidification affect marine organisms with CaCO3 exoskeletons? Flashcard 44 (Back): Increasing H3O+ concentrations reduce the concentration of CO3 2–22 , which is needed by marine organisms to form their CaCO3 exoskeletons (Ca2+(aq) + CO3 2–(aq) ⇌ CaCO3(s))21 .

Flashcard 45 (Front): What are buffer solutions? Flashcard 45 (Back): Buffer solutions are solutions that resist changes in pH when an acid or base is added22 .

Flashcard 46 (Front): What are the components of a buffer solution? Flashcard 46 (Back): Buffer solutions must contain either significant amounts of a weak acid and its conjugate base, or significant amounts of a weak base and its conjugate acid22 .

Flashcard 47 (Front): How does an acidic buffer solution neutralize added base? Flashcard 47 (Back): An added strong base is neutralized by the weak acid component of the buffer23 . Example: HA(aq) + OH–(aq) → A–(aq) + H2O(l)24 .

Flashcard 48 (Front): How does an acidic buffer solution neutralize added acid? Flashcard 48 (Back): An added strong acid is neutralized by the conjugate base component of the buffer25 . Example: H+(aq) + A–(aq) → HA(aq)26 .

Flashcard 49 (Front): What is the common ion effect? Flashcard 49 (Back): The addition of a salt containing an ion common to an equilibrium shifts the position of the equilibrium. For a weak acid equilibrium (HA(aq) + H2O(l) ⇌ A–(aq) + H3O+(aq)), adding a salt with the conjugate base (A–) shifts the equilibrium to the left, lowering [H3O+] and increasing pH26 ....

Flashcard 50 (Front): Write the Henderson–Hasselbalch equation for an acidic buffer. Flashcard 50 (Back): pH = pKa + log ([base]/[acid]) or pH = pKa + log ([A–]/[HA])28 ....

Flashcard 51 (Front): When is the Henderson–Hasselbalch equation applicable? Flashcard 51 (Back): The Henderson–Hasselbalch equation is applicable when the "x is small" approximation is valid, meaning the initial concentrations of the weak acid and its conjugate base are not very dilute and Ka is fairly small28 .... Generally, initial concentrations should be over 100 to 1000 times larger than Ka30 .

Flashcard 52 (Front): How do you calculate the pH of a buffer solution after adding a strong acid or base? Flashcard 52 (Back): It requires two parts: 1. A stoichiometry calculation for the reaction of the added acid or base with the buffer components31 . 2. An equilibrium calculation of [H3O+] using the new concentrations, often with the Henderson–Hasselbalch equation31 .

Flashcard 53 (Front): What is buffer range? Flashcard 53 (Back): Buffer range is the pH range within which a given buffer can provide pH protection, typically pKa ± 132 .

Flashcard 54 (Front): What is buffer capacity? Flashcard 54 (Back): Buffer capacity is the quantity of acid or base that a buffer can neutralize while maintaining pH within a desired range32 . It is proportional to the component concentrations32 .

Flashcard 55 (Front): How does increasing the concentrations of a buffer's components affect its resistance to pH change? Flashcard 55 (Back): The greater the initial concentrations of HA and A–, the less the change in pH as acid or base is added to the buffer solution32 .

Flashcard 56 (Front): What is pKa? Flashcard 56 (Back): pKa = -log[Ka]33 .

Flashcard 57 (Front): What is pKb? Flashcard 57 (Back): pKb = -log[Kb]33 .

Flashcard 58 (Front): What is the relationship between pKa and pKb for a conjugate acid-base pair? Flashcard 58 (Back): pKa + pKb = 1433 .

Flashcard 59 (Front): What is the pH at the midpoint of a weak acid-strong base titration? Flashcard 59 (Back): At the midpoint of the titration, [HA] = [A–], and therefore pH = pKa34 ....

Flashcard 60 (Front): Define titration. Flashcard 60 (Back): Titration is a process where a solution of unknown concentration (titrant) is slowly added to a solution of known concentration until the reaction is complete36 .

Flashcard 61 (Front): What is the equivalence point in an acid-base titration? Flashcard 61 (Back): The equivalence point is the point in a titration where moles of H3O+ = moles of OH–36 .

Flashcard 62 (Front): What is an indicator in a titration? Flashcard 62 (Back): An indicator is a chemical that changes color with pH, used to determine when the equivalence point has been reached36 .

Flashcard 63 (Front): What is the endpoint in a titration? Flashcard 63 (Back): The endpoint is the point of pH change where the indicator changes color37 . Indicators are chosen so the endpoint coincides with the equivalence point37 .

Flashcard 64 (Front): What is a titration curve? Flashcard 64 (Back): A titration curve is a plot of pH versus the amount of added titrant38 . The inflection point of the curve is the equivalence point38 .

Flashcard 65 (Front): What is the pH at the equivalence point of a strong acid-strong base titration? Flashcard 65 (Back): The equivalence point of a strong acid-strong base titration results in a neutral salt solution, so the pH is 7.0039 .

Flashcard 66 (Front): What is the pH at the equivalence point of a weak acid-strong base titration? Flashcard 66 (Back): The equivalence point of a weak acid-strong base titration results in a basic salt solution, so the pH is > 7.0034 .

Flashcard 67 (Front): What happens at half-neutralization in a weak acid-strong base titration? Flashcard 67 (Back): At half-neutralization, [HA] = [A–], and the pH = pKa of the weak acid35 .

Flashcard 68 (Front): How many equivalence points are expected in the titration of a diprotic acid where Ka1 >> Ka2? Flashcard 68 (Back): If Ka1 >> Ka2, there will be two distinguishable equivalence points in the titration40 . The closer the Ka values, the less distinguishable the equivalence points40 .

Flashcard 69 (Front): Define a Lewis acid. Flashcard 69 (Back): A Lewis acid is a substance that accepts a lone pair of electrons40 .

Flashcard 70 (Front): Define a Lewis base. Flashcard 70 (Back): A Lewis base is a substance that donates a lone pair of electrons40 .

Flashcard 71 (Front): What is a coordinate bond? Flashcard 71 (Back): A coordinate bond is a covalent bond formed when one anion or molecule donates a pair of electrons to another ion or molecule41 .

Flashcard 72 (Front): What is a ligand? Flashcard 72 (Back): A ligand is a Lewis base that is bonded to the central metal ion of a complex ion41 .

Flashcard 73 (Front): What is a complex ion? Flashcard 73 (Back): A complex ion is an ionic species consisting of a metal ion bonded to one or more Lewis bases (ligands)41 .

Flashcard 74 (Front): What is a formation constant (Kf)? Flashcard 74 (Back): A formation constant (Kf) is an equilibrium constant describing the formation of a metal complex from a free metal ion and its ligands41 . A large Kf indicates a stable complex ion42 .

Flashcard 75 (Front): What is solubility? Flashcard 75 (Back): Solubility is the amount of solute that dissolves, typically expressed in g/L42 .

Flashcard 76 (Front): What is molar solubility (S)? Flashcard 76 (Back): Molar solubility (S) is the amount of dissolved solute expressed as mol/L42 .

Flashcard 77 (Front): What is the solubility product constant (Ksp)? Flashcard 77 (Back): The solubility product constant (Ksp) is an equilibrium constant that describes the formation of a saturated solution of a slightly soluble salt42 . For Mg(OH)2(s) ⇌ Mg2+(aq) + 2 OH–(aq), Ksp = [Mg2+][OH–]242 .

Flashcard 78 (Front): How do you calculate molar solubility (S) from Ksp? Flashcard 78 (Back): S is calculated by solving the equilibrium expression for the dissolution of the slightly soluble salt using an ICE table42 ....

Flashcard 79 (Front): How does the common ion effect affect solubility? Flashcard 79 (Back): The presence of a common ion from another source will shift the solubility equilibrium to the left, decreasing the solubility of the sparingly soluble salt43 . Example: Adding SO4 2– to a solution of BaSO4 decreases the solubility of BaSO443 .

Flashcard 80 (Front): How does pH affect the solubility of insoluble salts containing basic anions? Flashcard 80 (Back): For insoluble salts containing basic anions, solubility increases as pH decreases (i.e., in acidic solutions)43 . The added H+ reacts with the basic anion, consuming it and shifting the dissolution equilibrium to the right43 .... Example: The solubility of CaF2 increases in acidic solution because F– is a weak base44 ....

Flashcard 81 (Front): How does pH affect the solubility of salts with pH-neutral anions? Flashcard 81 (Back): The solubility of salts with pH-neutral anions (conjugate bases of strong acids) is not significantly affected by pH45 . Example: The solubility of AgI is not greater in acidic solution because I– is pH-neutral45 .

Flashcard 82 (Front): What is Q in the context of solubility? Flashcard 82 (Back): Q is the ion product, which has the same form as Ksp but uses current (non-equilibrium) concentrations of the ions46 .

Flashcard 83 (Front): How can Q and Ksp be used to predict precipitate formation? Flashcard 83 (Back):

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If Q > Ksp, the reaction shifts left, and a precipitate forms46 .

•

If Q < Ksp, the reaction shifts right, and a precipitate does not form46 .

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If Q = Ksp, the solution is saturated, and the system is at equilibrium46 .

Flashcard 84 (Front): Describe the changes in the buffer system H2CO3/HCO3– related to ocean acidification. Flashcard 84 (Back): Increased atmospheric CO2 dissolves in the ocean, forming H2CO321 . The H2CO3 then dissociates to form HCO3– and H3O+21 . The increase in H3O+ decreases the pH of the ocean (ocean acidification) and also reduces the concentration of CO3 2– needed by marine organisms22 .

Flashcard 85 (Front): In a buffer solution containing a weak acid (HA) and its conjugate base (A–), which component reacts with added H+ ions? Flashcard 85 (Back): The conjugate base (A–) reacts with added H+ ions to form more of the weak acid (HA): H+(aq) + A–(aq) → HA(aq)26 .

Flashcard 86 (Front): In a buffer solution containing a weak acid (HA) and its conjugate base (A–), which component reacts with added OH– ions? Flashcard 86 (Back): The weak acid (HA) reacts with added OH– ions to form more of the conjugate base (A–) and water: HA(aq) + OH–(aq) → A–(aq) + H2O(l)24 .

Flashcard 87 (Front): What is the relationship between Kw, Ka, and Kb? Flashcard 87 (Back): Kw = Ka × Kb for a conjugate acid-base pair20 ....

Flashcard 88 (Front): What is the pOH of a solution? Flashcard 88 (Back): pOH = −log[OH–]14 ....

Flashcard 89 (Front): What is the relationship between pH and pOH at 25°C? Flashcard 89 (Back): pH + pOH = 14.0014 ....

Flashcard 90 (Front): How can you calculate [OH–] if you know the pOH? Flashcard 90 (Back): [OH–] = 10−pOH14 ....

Flashcard 91 (Front): What determines the pH of the equivalence point in a titration? Flashcard 91 (Back): The pH of the equivalence point depends on the pH of the salt solution formed38 . A neutral salt gives pH = 7, an acidic salt gives pH < 7, and a basic salt gives pH > 738 .

Flashcard 92 (Front): In the titration of a weak acid with a strong base, what is the predominant species at the equivalence point? Flashcard 92 (Back): At the equivalence point, the weak acid has been completely neutralized by the strong base, resulting in a solution containing primarily the conjugate base anion of the weak acid35 .

Flashcard 93 (Front): What type of buffer solution is blood? Flashcard 93 (Back): Blood has a buffer system consisting of a mixture of H2CO3 (weak acid) and HCO3– ion (conjugate base)22 . This is an example of a buffer system in nature22 .

Flashcard 94 (Front): When choosing an acid to make a buffer, what property should be considered relative to the desired pH? Flashcard 94 (Back): Choose an acid whose pKa is closest to the pH of the buffer you want to prepare32 .

Flashcard 95 (Front): What is the relationship between Ka and pKa? Flashcard 95 (Back): pKa = -log(Ka) and Ka = 10-pKa33 .

Flashcard 96 (Front): What is the relationship between Kb and pKb? Flashcard 96 (Back): pKb = -log(Kb) and Kb = 10-pKb33 .

Flashcard 97 (Front): What is the Ka for the first ionization of carbonic acid (H2CO3)? Flashcard 97 (Back): The source47 ... mentions a pKa1 value of 6.37 for the reaction H2CO3(aq) + H2O(l) ⇌ HCO3 –(aq) + H3O+(aq). Therefore, Ka1 = 10^-6.37.

Flashcard 98 (Front): What is the relationship between pKa and buffer range? Flashcard 98 (Back): The effective buffer range is typically considered to be pKa ± 132 .

Flashcard 99 (Front): In the Henderson-Hasselbalch equation, what do [base] and [acid] refer to? Flashcard 99 (Back): [base] refers to the concentration of the conjugate base, and [acid] refers to the concentration of the weak acid in the buffer solution28 .

Flashcard 100 (Front): What is the significance of the condition Ka1 >> Ka2 in the titration of a polyprotic acid? Flashcard 100 (Back): If Ka1 >> Ka2, the first ionization step goes essentially to completion before the second one begins, resulting in two distinct equivalence points in the titration curve40 .