5A: acid/base equilibrium, buffers, ions, solubility, titrations
things to remember about ksp problems
** if no common ion effect ** for all sparingly soluble salts X = molar solubility 1. general formula: MX ksp = X^2 2. general formula: MX2 ksp = 4X^3 3. general formula: MX3 ksp = 27X^4
what is pH?
- "power of hydrogen," is a numerical representation of the acidity or basicity of a solution. It can be used to calculate the concentration of hydrogen ions [H+] or hydronium ions [H3O+] in an aqueous solution. Solutions with low pH are the most acidic, and solutions with high pH are most basic.
buffer influence on titration curves
- Buffers make the titration curve "flat" at the region where buffering occurs. On a titration curve, this is the point of inflection. The point of inflection is at pH = pKa (or 14 - pKb) of the buffer. (this is usually seen at half the equivalency point) The area around the point of inflection is the region where the solution has buffering capacity. The pH of this buffering region is typically pKa +/- 1 (or 14 - pKb +/- 1).
Common-ion effect - use in laboratory separations
- In laboratory separations, you can use the common ion effect to selectively crashing out one component in a mixture. For example, if you want to separate AgCl from a mixture of AgCl and Ag2SO4, then you can do so by adding NaCl. This will selectively crash out AgCl by the common ion effect (Cl- being the common ion).
What is the Ksp?
- Solubility product constant - is simply Keq for dissolutions - The higher the value, the more the reaction products dominate in a saturated solution (at equilibrium). ex. AgCl (s) ↔ Ag+ (aq) + Cl- (aq) Ksp for AgCl = [Ag+][Cl-] Ag2SO4 (s) ↔ 2Ag+ (aq) + SO42- (aq) Ksp for Ag2SO4 = [Ag+]^2[SO42-] ** coefficent will become exponent in expression ** solids are not included in ksp expression *** there is never a denominator *** is temperature and pressure dependent
Which of the following best describes a chemical species that is measured to have Kb = 3.2×10^-18 A strong acid B strong base C weak acid D weak base
- VERY low kb - must compare to kw=ka(kb) kw = 1x10^14 M kx=ka (kb) ka = 3.3x10^3 since Ka > 1 (STRONG ACID)
buffer
- a mixture of a weak acid or base with its salt that resists change to pH when small amounts of acid or base is added ex. acid buffer acetic acid______sodium acetate CH3COOH and CH3COO-Na+ weak acid ______salt base buffer ammonia___ammonium chloride NH3 and NH4+Cl- weak base__salt
solute
- a substance that dissolved in a solvent
conjugate acid
- a weak base accepted a proton - STRONGER THAN WEAK ACID
bicarbonate buffer system
- biological acid buffer system in the blood between carbonic acid (H2CO3) and bicarbonate ion (HCO3-) - maintain blood pH CO2 (g) + H2O (l) <-> H2CO3(aq) <-> HCO3- (aq) + H+ (aq)
dynamic equilibrium
- exists once a reversible reaction ceases to change its ratio of reactants/products, but substances move between the chemicals at an equal rate, meaning there is no net change.
How to find pKa from titration curve?
- find the equivalence point on the graph (pH at vertical line) - find 1/2 eq pt (pt at half the titrant vol added) pH at 1/2 eq pt = pKa
How to find Ka from titration curve?
- find the equivalence point on the graph (pH at vertical line) - find 1/2 eq pt (pt at half the titrant vol added) pH at 1/2 eq pt = pKa use Ka = 10^-pKa to find answer ex. eq pt = 9 titrant = 22 ml 1/2 eq pt = 5 (=pKa) titrant 11 ml ka = 10^-5
homogeneous equilibrium
- has everything present in the same phase. The usual examples include reactions where everything is a gas, or everything is present in the same solution.
heterogeneous equilibrium
- has things present in more than one phase. The usual examples include reactions involving solids and gases, or solids and liquids. ** salt dissociations
complex ion formation
- in which a cation (electron acceptor/lewis acid) is bonded to at least one electron pair donor (ligand/ lewis base) and form a stable ionic compound [Metal+] + [Lewis base] → Complex ion [M+] + [L] → [M-Ln+] **The Lewis base can be charged or uncharged. **The Keq for this reaction is called Kf, or the formation constant. ** compounds is held together by a coordinate covalent bond ** have high ksp values
Common-ion effect
- is simply Le Chatelier's principle applied to Ksp reactions - adding the common ion to solution will increase concentration of common ion in solution, cause equilibrium to shift left towards precipitation ex1. AgCl (s) ↔ Ag+ (aq) + Cl- (aq) - if you add Cl- ions to the solution above, then less AgCl would dissolve ex2. if you add NaCl to a saturated solution of AgCl, then some AgCl will crash out of solution. ex3. more AgCl can dissolve in pure water than in water containing Cl- ions.
Ka (acid dissociation constant)
- is the equilibrium constant for chemical reactions involving weak acids in aqueous solution. - The numerical value is used to predict the extent of acid dissociation. A large Ka value indicates a stronger acid (more of the acid dissociates) and small Ka value indicates a weaker acid (less of the acid dissociates). HA (aq) + H2O(l) <-> H3O+(aq) + A- (aq) * water never included in equilibrium (l) * neither are solids (s)
Kb (base dissociation constant)
- is the equilibrium constant for chemical reactions involving weak base in aqueous solution. - The numerical value is used to predict the extent of base dissociation. A large Kb value indicates a stronger base (more of the base dissociates) and small Kb value indicates a weaker base (less of the base dissociates). BOH (aq) + H2O(l) <-> OH-(aq) + B+ (aq) * water never included in equilibrium (l) * neither are solids (s)
Complex ions and solubility / complex ion effect
- is the opposite of the common ion effect - is simply Le Chatelier's principle applied to Ksp reactions - when complex ion forms in solution, it removes common ion from solution and will cause equilibrium to shift right towards dissolution ex1. 1. dissolution AgCl (s) ↔ Ag+ (aq) + Cl- (aq) 2. complex ion formaiton M+ + Cl- ↔ M-Cln complex ion ***When complex ion forms, the Cl- ion is taken out, so more of the AgCl will dissolve. ex2. 1. dissolution AgCl (s) ↔ Ag+ (aq) + Cl- (aq) 2. complex ion formaiton NH3 + Ag+ ↔ Ag-(NH3)n complex ion ***Here, the complex ion formation takes out Ag+, again causing more AgCl to dissolve
hydration
- is where water (solvent) forms a shell around ions in solution -The oxygen atom on water is partially negative, so it surrounds cations. -The hydrogen atoms on water is partially positive, so they surround anions.
solubility rules
- know them in general - do not have to memorize definitely know 1. all salts from group 1 are ALWAYS soluble 1. all nitrate salts are ALWAYS soluble . *******print out rule chart
solubility
- maximum amount of solute that will dissolve in a specific solvent at a given temperature
IP (ion product)
- same as quotient for keq ex. AmBn(s) <-> [mA^n+] (aq) + [nB^m-] (aq) IP = [A^n+]^m[B^m+]^n ** concentrations do not have to be at equilibrium
pka value
- tells you if a given molecule is going to either give a proton or remove a proton at a certain pH - The lower the pKa, the stronger the acid and the greater its ability to donate protons - also a measure of how stable the conjugate base is
pkb
- tells you if a given molecule is going to either give a proton or remove a proton at a certain pH - The lower the pKb, the stronger the base and the greater its ability to take protons
solvent
- the component of a solution that remains in the same phase after mixing or is in greater quantity
solvation / dissolution
- the electrostatic interaction between solute and solvent molecules - breaking intermolecular interactions between solutes and forming new intermolecular interactions between solute-solvent molecules - can be endothermic or exothermic - solvent can be water or any other substance
heat of dissolution
- the enthalpy change during the solvation/ dissolution process H = - (exothermic) H = + (endothermic)
acidity
- the likelihood of a compound to be PROTONATED = increases with higher Ka = increases with LOWER pka = increases with LOWER pH = anything which stabilizes the conjugate base will increase the acidity.
What is the half equivalence point?
- the point in the titration curve where the pH = pKa - the best buffering region on the curve - nearly a complete horizontal line - the point where the [acid] = [congjuage base] or [base]=[conjugate acid] - Thus the concentration of protonated and deprotonated acid is split 50/50, or 50% of the acid has been deprotonated
titration
- the procedure used to determine the concentration of a known reactant in solution 3 types 1. acid-base 2. oxidation-reduction (redox) 3. complexometric
precipitation
- when solute comes out of solution, spontaneously, due to a reaction or due to a drop in temperature -when IP> ksp
supersaturated
- when solution has max amount of solute that can be dissolved and an increase in temperature allows for even more to dissolve
saturated
- when solution has max amount of solute that can be dissolved and is now in equilibrium with undissolved solute state AT A GIVEN TEMP - when IP = Ksp
unsaturated
- when solution has yet to reach max amount of solute that can be dissolved and is not yet in equilibrium with undissolved solute state AT A GIVEN TEMP - when IP < Ksp
For pure water, pH =
-log[10-7] = 7 * only at temp 25°C
exothermic solvation
-when newly formed solute-solvent interactions are more favorable than old solute-solute interactions - occurs at low temp (heat energy is not required to fuel reaction)
endothermic solvation
-when newly formed solute-solvent interactions are weaker than old solute-solute interactions - occurs at high temp (requires heat input)
what is the normality of 1M HCl?
1 N HCl
acid equivalent
1 mol of H+/ H3O+ ions
base equivalent
1 mol of OH- ions
polyvalent / polyprotic
1 mol of some acids and bases yields to more than 1 acid/base equivalent ex. 1 mol of H2SO4 produces 2 acid equivalents H2SO42- + H2O ←→ H3O+ + HSO4- HSO4- + H2O ←→ H3O+ + SO42-
titration curves exist for the following combinations
1. STRONG ACID + STRONG BASE 2. STRONG ACID + WEAK BASE 3. WEAK ACID + STRONG BASE NOT FOR 4. WEAK ACID + WEAK BASE * because its not accurate, lack sharp curve change, indicators are gradual change not fast
salt formation - 4 types of reactions
1. STRONG ACID + STRONG BASE HCl + NaOH ←→ NaCl + H2O 2. STRONG ACID + WEAK BASE HCl + NH3 ←→ NH4Cl (no water) 3. WEAK ACID + STRONG BASE HClO + NaOH ←→ NaClO + H2O 4. WEAK ACID + WEAK BASE HClO + NH3 ←→ NH4ClO (no water)
equivalence point range for the following 3 titration combinations
1. STRONG ACID + STRONG BASE eqv pt pH = 7 2. STRONG ACID + WEAK BASE eqv pt pH < 7 3. WEAK ACID + STRONG BASE eqv pt pH > 7
Solubility and pH
1. acids are more soluble in excess bases (B- will steal H+ from HA) acid dissociation HA → [H+] + [A-] base formation [B-] + [H+] -> BH+ ***Putting HA in a base will take out the H+, resulting in a shift in eq towards RIGHT -> more HA will dissolve according to Le Chatelier's principle. 1. bases are more soluble in excess acids (A- will steal H+ from BH+) base dissocation BH+ -> B + H+ acid formation [H+] + [A-] → HA ***Putting BH+ in an acid will add more H+, resulting in a shift in eq towards RIGHT -> more BH+ will dissolve according to Le Chatelier's principle. ** redo
5 factors that influence acidity of alpha hydrogens
1. charges (as + charge increases = increase acidity) 2. periodic table - periodic table row (acidity increases across row -> LEFT to RIGHT) - periodic table column (acidity increases going DOWN a column) 3. resonance (increases the stability of compounds = increase acidity) - alcohol < phenol - stabilizing conjugate base = increase acidity 4. inductive effects - EWG and EDG EWG = F > Cl > Br > I (increase acidity with increase electronegativity, and the effect decreases with distance from atom) EDG = alkyl (decrease acidity) 5. orbital (the higher s -character of the bond to the H then increasing acidity) sp> Sp2> sp3 50 % S > 33 % S > 25 % S
properties of strong acids
1. completely ionized in aqueous solutions 2. high solvation 3. will decrease pH below 7 4. strong conductor of electricity 5. completely DISSOLVES in water
properties of strong bases
1. completely ionized in aqueous solutions 2. high solvation 3. will increase pH above 7 4. strong conductor of electricity 5. completely DISSOLVES in water
what effects spontaneity of dissolution?
1. enthalpy (exothermic or endothermic) 2. entropy (increase or decrease depend on pressure, volume, number of molecules, arrangement of molecules and temp change) 3. gibbs free energy change overall
How is the acid-base equivalence point found?
1. graphical method - plotting pH of unknown solution as a function of adding titrant using pH meter 2. estimate using indicator watching for color change
units of concentration
1. molarity 2. molality 3. normality 4. mole fraction 5. percent mass *** MUST PRACTICE MATH
properties of weak acids
1. partially ionized in aqueous solutions 2. high solvation 3. will decrease pH below 7 4. weak conductor of electricity 5. completely DISSOLVES in water *dissolve and ionize are not the same thing
properties of weak bases
1. partially ionized in aqueous solutions 2. high solvation 3. will increase pH above 7 4. weak conductor of electricity 5. completely DISSOLVES in water *dissolve and ionize are not the same thing
litmus paper rules
1. red lithmus paper - turn blue in base 2. blue lithmus paper - turn red in acid
what is the normality of 1M H2SO4 ?
2 N H2SO4
what is the normality of 1M H3PO4 ?
3 N H3PO4
Normality (N) =
= (Number of equivalents solute)/(liter of solution) - the morality of the stuff of interest (H+, OH, ions)
What is the pH equation?
= -log[H+]
what is the pOH equation?
= -log[OH-]
Molarity
= M = mol solute/L solution (NOT SOLVENT)
what is the Kw equation
= [H+][OH-] = 1*10^-14 at 25°C (298 K) OR Kw = kaxkb
what is the pKa equation?
=-log[Ka]
what is the pKb equation?
=-log[Kb]
acidity of the functional groups is:
Alkene < alkyne < alcohol < carboxylic acid
solute dissociation equation
AmBn(s) <-> [mA^n+] (aq) + [nB^m-] (aq) ex. Ag2SO4 (s) ↔ [2Ag+] (aq) + [SO4^2-] (aq) ** the first step of any solution problem in test day is to write out the balanced dissociation equation
Suppose a nanotechnological innovation allows every single charged ion to be precisely identified and removed from a small volume of water. Which of the following describes Ka at the end of the process, assuming that the filtered water is given adequate time to re-equilibrate? Please choose from one of the following options. A 0 B 10^-14 C 1 D 1x10^7
B 10^-14 - The autoionization of water occurs naturally due to attractive forces between constituents of water molecules. - Even of all ions are removed from a sample of water (including hydronium ions), a short time later the water will re-ionize until it approaches an equilibrium concentration of ions. In equilibrium, Ka = Kw = 10^{-14}
Hydronium ion
H3O+ *****H+ never exist as a proton in water, it always exists as the hydronium ion
How to calculate pH or pOH at equivalence point? HF = 0.3 M Ka = 7.2x 10 ^ -4 NaOH added until eqv reached **** unlikely to be on MCAT
HF - weak acid NaOH - strong base pH will be above 7 need ICE table to do the math the conjugate base or acid will determine the pH in this case the conjugate base F- will be important F- + H2O <-> HF + OH- Need kb kw = kb x ka kb = 1x10^-14/ 7.2x 10 ^ -4 * use fraction table memorized coefficient 1/7 = 0.14 exponents -14 - (-4) = -10 0.14 x10 ^ -10 = 1.4 x 10^-11 = kb (convert to standard notation) kb = [OH][HF]/ F- = x^2/F- = x^2 / 0.3 x^2 = (1.4x10^-11) (3.0x10^-1) x^2= 4.2x10^-12 x = 2x10^6 = [OH-] pOH = 6-.2 = 5.8 pH = 14 - 5.8 = 8.2 (aproximation) real asn is 8.31
when trying to find pH or pOH of weak acid or base must use
ICE table to find [H+] or [OH-] ex. ka = [A-][H+]/ [HA] = x^2/ [HA-x] * x is negligible so [HA-x] is now [HA] = x^2 / [HA]
IP and Ksp
IP < Ksp = unsaturated (solution will continue to dissolve) IP = Ksp = saturated (solution at equilibirum) IP > Ksp = supersaturated (solution will start to percipirate)
how to convert from Ka to Kb or vice versa?
Kw = Kb x Ka
dilution equation
M1V1 = M2V2
normality equation to calculate unknown titrand
NaVa = NbVb Va = volume of acid Vb = volume of base Na = normality of acid Nb =normality of base * units for V must be the same but do not have to be Liters
how to multiply and divide scientific notation ? ex. (3.04×10^5)÷(9.89×10^2) = ?
Since all number in scientific notation have base 10 , we can always multiply them and divide them. To multiply two numbers in scientific notation, multiply their coefficients and add their exponents. To divide two numbers in scientific notation, divide their coefficients and subtract their exponents. In either case, the answer must be converted to scientific notation. ex. (3.04×10^5)÷(9.89×10^2) = ? coefficents divide 3.04 ÷ 9.89 = rounding = 3 ÷ 10 = 0.33 exponents subtract 5-2 = 3 0.33 x 10 ^ 3 (not scientific notation yet) 3.3x10^2 the exact answer is 3.07x10^2 however you only need to approximate on the MCAT
most common strong acids
Strong acid Formula Perchloric acid HClO4 Hydroiodic acid HI Hydrobromic acid HBr Sulfuric acid H2SO4 Hydrochloric acid HCl Nitric acid HNO3 Hydronium ion H3O+ or H+ * all are protonated
most common strong bases
Strong bases Formula Lithium hydroxide LiOH Sodium hydroxide NaOH Potassium hydroxide KOH Rubidium hydroxide RbOH Cesium hydroxide CsOH Calcium Hydroxide Ca(OH)2 Strontium hydroxide Sr(OH)2 Barium hydroxide Ba(OH)2
common Weak acids
Weak acid Formula Formic acid HCOOH Acetic acid CH3COOH Hydrofluoric acid HF Hydrocyanic acid HCN Hydrogen sulfide H2S Water H2O benzoic acid C6H5COOH
common weak bases
Weak base Formula Ammonia NH3 Amine NR3 Pyridine C5H5N Ammonium hydroxide NH4OH Water H2O
how to solve ksp question?
What is the solubility of MX2 if given Ksp? MX2 ↔ M2+ + 2X- Ksp = [M2+][X-]2 Ksp = [M2+][2M2+]2 (because for every M2+, there's two times as much X-) Ksp = 4[M2+]3 ***Solve for [M2+]. molar solubility is the same thing as [M2+] because you used Q = Ksp for a saturated solution. If you solved for [X-] instead, divide your results by 2. If you were given solubility and asked to solve Ksp, then know that solubility = [M2+] = [X-]/2
Methanethiol, CH3SH, has a pKa of 10.3 and methanol, CH3OH, has a pKa of 15.5. Which is a stronger acid? Which is a stronger base, CH3S- or CH3O-? A Methanethiol is the stronger acid, its conjugate base is the stronger base B Methanol is the stronger acid, its conjugate base is the stronger base C Methanethiol is the stronger acid, methanol's conjugate base is the stronger base D Methanol is the stronger acid, methanethiol's conjugate base is the stronger base
Why? A lower pKa indicates a stronger acid, therefore, methanethiol is the stronger acid. Because methanol is a weaker acid, its conjugate base is stronger than the conjugate base of methanethiol, hence, the methoxide ion is the stronger base.
Which of the following is least acidic? A HF B HI C HBr D HCl
Why? Comparing the respective, deprotonated halide ions, F- would be the least stable since it is the smallest in size (I- would be the largest). Thus HF would be the least acidic. acidity of halogen compounds is based on electronegativity HF is small and is less able to delocalize its e- then a larger halogen compound like IF
Knowing that the Ka of triflic acid (monoprotic) is 8.0 x 10^14 and will almost completely dissociate in solution, what is the pH of a 1M solution of triflic acid? A 8.0 x 1014 B 1/(8.0 x 1014) C 0 D 1
Why? Ka = [H+][A-]/[HA]. Therefore, if the Ka of an acid is very large (as is 8.0 x 1014), then nearly all of it will dissociate since the ratio of A-:HA is so high. Thus, in a 1M solution of triflic acid, there will be 1M of H+. Recall that pH = -log[H+]. Thus, -log(1) = 0.
Rank the following compounds in order of the most acidic to the least acidic. I. p-chlorobenzoic acid II. benzoic acid III. p-methylbenzoic acid IV. p-nitrobenzoic acid A I > II > III > IV B IV > I > II > III C IV > II > I > III D III > IV > II > I
Why? The acidity of each of these compounds can be determined by the strength of the electron withdrawing properties of each functional group. The more electron-withdrawing a functional group, the more acidic the compound will become (since electron density will be drawn further from the hydrogen, making it more easily donated, hence more acidic). The nitro (NO2) group is the most electron-withdrawing, therefore 'IV' is the most acidic. The methyl (CH3) group is the only one listed that is electron-donating, hence III is the least acidic. Chlorine is electron withdrawing, more so than hydrogen, hence 'I' is more acidic than 'II'
How to convert pH to [H+]?
[H+] = 10^-pH
How to convert 10^-n concentration of H+ to pH? * short cut no calculators in MCAT
[H+]=10^-n pH = n pOH = 14 - n *works for Ka to pKa and Kb to pKb * must convert all values to standard notation *if [H*] = 10^n then pH = - n (negative value)
percent composition by mass
[mass of element/ mass of total compound] x 100%
mole fraction
[moles of A/total mole of all species] ***sum of all fractions will always equal 1
conjugate bases
a weak acid donated a proton - STRONGER THAN WEAK BASE
Brønsted-Lowry base
accepts an H+
lewis acid
accepts e- pair
pH less than 7 pOH greater then 7
acidic
substituents effects on basicity
as the e- donating effect of substituents INCREASE = increase in basicity common e- donating groups (alkyl) C6H > C4H > C3H> CO2-> CH3
substituents effects on acidity
as the e- withdrawing effect of substituents INCREASE = increase in acidity common e- withdrawing groups (high electronegativity) NO2 > F > Cl >Br > I
In a titration curve where will the solution be most well buffered?
at the half equivalence point
pH greater than 7 pOH less than 7
basic
Below what Ka is an acid considered a weak acid?
below 1.0
Below what Kb is an base considered a weak base ?
below 1.0
amphiprotic
can either donate or accept a proton, thus acting either as an acid or a base. Water, amino acids, hydrogen carbonate ions and hydrogen sulfate ions are common examples of amphiprotic species. ALL contain a hydrogen atom.
calculating pH and pOH of strong acids and bases assumes
complete dissociation of ions and so concentration of H+/OH- are equal to acid * use short cuts to calculate logs
strong acids and bases
completely dissociate in aqueous solutions
'concentration' and 'strength' when referring to acids (or bases)
concentration - number of moles of a dissolved substance per litre of aqueous solution. It is a measure of the dilution of the acid (or base). A more concentrated solution contains more moles of solute. strength - refers to the degree of ionisation or dissociation of the acid or base.
Brønsted-Lowry acid
donates an H+
lewis base
donates e- pair
strong acid and strong base titration curve
eqv pt = pH 7
weak base and strong acid titration curve
eqv pt pH < 7
weak acid and strong base titration curve
eqv pt pH > 7
How to find the square root of scientific notation number ? * with out a calculator on MCAT square root of 4 x 10^ 6 * needed for ICE tables
ex. square root of 4 x 10^ 6 coefficient square root 4 = 2 exponent square root (basically divide by 2) = 3 Answer = 2x 10 ^ 3
how does temperature effect Ksp?
gas solutes - Ksp decreases with increasing temp non-gas solutes - ksp increases with increasing temp
solution
homogenous mixutres of two or more substanes that combine to form a single phase usually the liquid phase
how does pressure effect Ksp?
increasing pressure will increase ksp for gas solutes ** no effect on solid or liquid solutes
pH and pOH are
inversely related
ionization vs dissociation vs dissolution
ionization - is the conversion of a neutral molecule or atom to an ion. ex. Na atom to Na+ dissociation - is a process in which a compound separates into two or more parts. ex. water dissociates into hydronium and hydroxide ions. 2H2O -> H3O+ + OH- dissolution - is the process in which the particles of a substance move into the solvent. *** NOTE: The process may or may not involve dissociation ex. dissolving of sugar in water (not ionizing - just solvating) ex. Nacl (s) -> Na+ + Cl- (dissolving AND dissociating)
acid-base equivalence point
is the point at which chemically equivalent quantities of acid and base have been mixed - is reached when the number of acid equivalents present in the original solution equals the number of added base equivalents. ex. it is NOT always pH 7
End Point
is the point at which the indicator changes color - this is NOT the equivalence point but RIGHT after it when the pH is slightly higher or lower then eqv point
Polyvalence/Normality
is the relative acidity/basicity of an aqueous solution, which is determined by the relative concentrations of acid/base equivalents.
how to go from pka/pkb to ka/kb ?
ka/kb=10^-pka or pkb
ksp vs kf
ksp - solubility constant - for the dissolution of the original solid to ions kf - formation or stability constant for complex ions
Whether a chemical is an acid or base depends on whether Ka or Kb is bigger ex. Is a chemical with Kb=3.2×10^−18 an acid or base?
kw = Ka x Kb 1.0x10^-14 = ka x 3.2×10^−18 Ka = 3.125x10^3 Ka >1 means strong acid
How to determine titrand and titrant combination in titration curve ?
look at the started point in pH curve pH >> 7 titrand is a strong base pH >7 (slightly) titrand is weak base pH < 7 (slightly) titrand is weak acid pH << 7 titrand is strong acid then look at eqv pt which is the straight line region
Molality
m = mol solute/kg solvent
if concentration of acid or base are way above 10^-7 M then autoionization of water is *Remember large negative exponent means smaller actual value
negligible - do not need to include it in the calculates for finding PH. HOWEVER. If the concentration of acid or base is less then 10^-7 M you must include auto ionization of water in equation
Salt Formation
occurs during acid/base neutralization reactions. Salt may precipitate and fall out of solution depending on the solubility and amount produced.
amphoteric
of a compound, especially a metal oxide or hydroxide) able to react both as a base and as an acid.
weak acids and bases
only partially dissociate in aqueous solutions
What is the mathematical relationship between pH and pOH?
pH + pOH = 14
How to convert Ka to pKa? no calculators in MCAT Ka = nx10^-m ex. Ka = 1.8 x10^-5 * n must be between 1 and 10
pKa = m-0.n * move decimal place over by 1 to the left ex. pka = 5-0.18 = 4.82 (actual is 4.74 - but close approximation) * works for [H+] to pH and [OH-] to pOH
concentrated
ratio to solute to solvent is high ** still considered unsaturated
dilute
ratio to solute to solvent is small ** still considered unsaturated
hydrolysis
reverse salt formation reaction salts reaction with water to create weak acid and base
aqueous solution
solution in which solvent is water symobl (aq)
titrant
solution of known concentration
titrand
solution of unknown concentration
Ka > 1
strong acid
Kb > 1
strong base
large Ka means
stronger the acid (MORE dissociation)
larger Kb means
stronger the base (MORE dissociation)
concentration
the amount of solute dissolved in solent
basicity
the likelihood of a compound to be DEPROTONATED
molar solubility
the molarity of solute in saturated solution *** the same thing as the concentration of dissocated ions
redox titration
the transfer of electrons required to reach equivalence point * indicators used change color at a particular voltage Redox = reduction + oxidation = species A gains electrons + species B lose electrons. Reduction = reduction in charge = decreased oxidation number = gaining electrons. Oxidation = increase in charge = increased oxidation number = losing electrons.
what is Kw?
the water dissociation constant - equilibrium constant * ONLY temp dependent
True or false: all solutions are mixtures but not all mixtures are solutions
true
henderson-hasselbalch equation
used to estimate pH or POH of a butter solution pH = pKa + log [A-]/[HA] *** this is just a rearrangement of the ka or kb expression
autoionization of water
water can react with itself H2O(l) + H2O(l) <-> H3O+(aq) + OH-(aq)
Smaller Ka means
weaker the acid (LESS dissociation)
Smaller Kb means
weaker the base (less base dissociates)
chelation
when a central cation is bonded to the same ligand in multiple places creating an even more complex ion
indicator
which are weak organic acid and bases that have unique properties and display different colors in the protonated and de-protonated states based on pH level. They are used in such a low concentration that does not affect the equivalence point.
