Altius Lesson: Chemistry 3
pH of various titrations WB w/ SA : pH ? WA w/ SB : pH ? SA w/ SB : pH ?
- pH < 7 - pH > 7 - pH =7
Conversion between pH and pOH pH + pOH = ? Kw = ? pKw = ?
14 [H3O+] [OH-] pH + pOH
Strong acid titrated with strong base graph looks like?
:)
Relationship between Free energy and Chemical energy equation is?
G = -nFE n is moles of electrons transferred F is faradays constant
pOH equation
pOH= -log[OH-]
Atom and oxidation state Fluorine ? Hydrogen ? Hydrogen w/ a metal ? Oxygen (except peroxides) ? Alkali metals ?
- -1 - +1 - -1 - -2 - +1
For an electrolytic cell Cathode = ?; Anode = ? - Oxidation occurs at the anode or cathode and reduction at the anode or cathode?
- Cathode = (-); Anode = + - oxidation occurs at the anode and reduction at the cathode
Hydrolysis of NaNO2 will result in the reaction of a nitrite ion with water to form? - how will this change the pH? Hydrolysis of NH4Cl will result in reaction of NH4+ with water to form? - how will this change the pH?
- HNO2 and hydroxide ion - increase pH - NH3 and H3O+ - decrease pH
T/F, half reactions always come in pairs, one reduced and one oxidized? Eo for any oxidation half reaction is simply the negative of?
- T - Eo for the associated reduction half reaction
What does it mean if something has a high reduction potential? What does it mean if something has a high oxidation potential?
...
5 Things to remember about electrolytic cells
1) The species with the lower reduction potential will be reduced! 2) The cell potential will always be negative. 3) The sum of the externally applied voltage (Vbattery) and the negative cell potential (-E°cell) must be positive. 4) Cathode = (-); Anode = (+) Note the difference compared to galvanic cells. 5) Oxidation still occurs at the anode and reduction at the cathode.
"One equivalent" definition.
The amount of acid or base necessary to produce or consume one mole of H+ ions.
Oxidation state definition
The apparent charge that an atom takes while in a molecule. The sum of the oxidation states for all of the atoms in a molecule must equal the charge on that molecule.
Titration of SA w/ SB or SB w/ SA Equivalence Point/Stoichiometric Point = ? For a solution of NaOH being titrated with HCl, at the equivalence point [HCl] = ? For titrations involving a SA and a SB, [H+] = [OH-] at ? By definition, if [H+] and [OH-] are exactly equal, pH = ?
[titrant] = [analyte] [H+] =[OH-] pH=7 - midpoint of the nearly vertical section of the graph. - [NaOH] in the flask. Put another way, the moles of HCl in the beaker = the moles of NaOH in the beaker. Because HCl and NaOH are both considered "strong" (i.e., dissociate 100% in water), they will both produce the same amount of ions per mole. - the equivalence point - 7 at the equivalence point.
Galvanic cell diagram Draw a galvanic cell and label the following: anode, cathode, terminals, salt bridge, electron flow, current flow, load (resistor, or voltmeter), electrodes, oxidation half-reaction, and reduction half-reaction.
two half reactions, one reduced and one oxidized
What is most important when creating a buffer at say a pH of 8? What would the pKa be?
- The most important thing is to choose a weak base and its conjugate that has a pKa close to 8. Why should the pka be around 8? Because at the middle of the buffer, the half-equivalence point, pKa equals pH.
A student mixes 2.5 mL of 1.0 M C6H5OH with 5.0 mL of 0.5 M C6H5O-. The pH is measured and found to be 9.9. Based on this information, phenol is best described as a: A) weak base with a pKa of approximately 10. B) strong base with a pKa of approximately 7. C) weak acid with a pKa of approximately 10. D) the conjugate base of a weak acid, whose pKa cannot be determined from the information given.
- The ratios of volume and molarity given result in adding equimolar amounts of phenol (C6H5OH) and its conjugate base. This is a dead giveaway that this is a buffer solution. By adding exactly equimolar amounts the student has recreated the half-equivalence point where [HA] = [A-]. Therefore, pH = pKa, which in this case means the pH is known to be exactly 9.9. Phenol is a weak acid, so combining these two known facts makes answer C correct
What is different about the cathode and anode?
Cathode = (-) and Anode = (+) *opposite as galvanic
What is the hydrogen ion concentration in a solution with a pH of 11.26? a. 5.5 x 10 ^-12 b. 5.5 x 10 ^-11 c. 5.5 x 10 ^-10 d. 1.12 x 10 ^-6
a. 5.5 x 10 ^-12
Biology connection Would you expect a strong oxidizing agent to have a high or low reduction potential?
O2 + 4H+ + 4e- --> 2 H2O E=1.23 V Last step in the electron transport chain, - It would have a high reduction potential. It is favorable for it to take on electrons and oxidize something else in the process. For example, MnO4- and Cr2O7-2 have high reduction potentials.
Amphoteric substance
Substance that can act as either an acid or a base *H2O
What is faradays constant Using the above definition, and combining it with information you should recall from Physics 3, calculate Faraday's constant?
The charge on one mole of electrons 96,485 J - Faraday's constant is the charge on one mole of electrons. One mole of electrons is 6.022 x 1023 electrons. Each electron has a charge of 1.6 x 10-19C. Therefore, the charge on one mole of electrons is: (6 x 1023)(1.6 x 10-19) = 9.6 x 104 C/mol.
What is the pH of a .1 M solution of acetic acid pKa 4.74? a) 1 b) 2.9 c) 11.1 d) 13
The pka tells us it is a weak acid, also since it is an acid the pH will not be greater than 7, throwing out c and d. A pH of 1 would be a strong acid so it must be b, 2.9
Atom and oxidation state continued Assign an oxidation state to each atom in each of the following molecules: a) (NH4)2SO4, b) FeCO3, c) H2O2, d) NaH, e) SF6.
a) N = -3 ; H = +1 ; S = + 6; O = -2 b) Fe = +2 ; C = +4 ; O = -2 ; c) H = +1 ; O = -1 d) Na = +1 ; H = -1 e) S = +6 ; F = -1.
Lewis Acid/Base definition Two common examples of lewis acids are? The electrophiles in all organic chemistry reactions are acting as? Lewis Acids are all about the what?
Acids accept a pair of electrons Bases donate a pair of electrons - AlCl3 and BF3 - Lewis acids - the accepting an donating of electron pairs
Arrhenius Acid/Base definition
Acids produce H+ ions Bases produce OH- ions in solution
REDOX reaction definition For example: Fe(s) + H2O(l) → H2(g) + FeO(s) In this reaction iron loses two electrons and two hydrogens each gain one electron. Iron is therefore a? What would be the result if you tried to make a buffer out of a strong acid and a strong base?
Any reaction where one or more electrons are transferred from one atom to another. - reducing agent and water is an oxidizing agent. To recognize this, you must be able to calculate the oxidation state of each atom - The pH would change drastically when acid or base was added to it. There would be no weak acid or base to consume the added acid or base.
What do we learn from the above equation? A Faraday is ? A Farad is ? A Farad is the amount of capacitance necessary to hold?
+E = negative G = spontaneous reaction - an obsolete unit of charge equal to the charge on one mole of electrons. In other words, Faraday's constant = 1 Faraday. The Faraday has since been replaced by the Coulomb - a unit of capacitance. It is a "summary" unit similar to a Newton. Just as we can say 1 Newton instead of saying 1 Kg*m/s2, we can say 1 Farad instead of saying 1 C2*s2/m2*kg. - 1 C of charge on a capacitor with a potential difference of 1 Volt.
Atom and oxidation state Alkaline earth metals ? Any elemental atom ? Group V ? Group VI ? Group VII ?
- +2 - 0 - -3 - -2 - -1
Reforming of a weak acid: Reforming of a weak base:
- 1) Na2CO3 ⇄ 2Na+ + CO32- 2) CO32- + H2O ⇄ HCO3- + OH- - 1) NH4NO3 ⇄ NH4+ + NO32- 2) NH4+ + H2O ⇄ NH3 + H3O+
The pH of a solution with H+ concentration 1.0 x 10 ^-4? The pH of a solution with H+ concentration 3.0 x 10 ^-4? a. 3 b. 3.5 c. 4 d 4.5
- 4 - b, 3.5
Acedic acid has a pKa of 4.74, what is the pKb of the acetate ion? What is the pH of a .01 M solution of HBR? What is the pH of a .01 M solution of KOH?
- 9.26 - 2, since HBR is a strong acid it disassociates completely therefore the H+ concentration is also .01M , the -log is simply 2 - 12, remember if the -log of the [OH] is 2 and pH +pOH = 14 the pH is 12
Nernst equation By definition, E° will always equal what for a concentration cell?
E = E° - (0.06/n)*log[lower]/[higher] where n = moles of electrons transferred (e.g., Fe3+(aq) → Fe(s) = 3 and Ag+(aq) → Ag(s) = 1) Zero
Explain the pH scale. What is always present in the solution? The pH scale ranks solutions based not so much on the acids or bases themselves, but on?
- A mathematical ranking system for the acidity or basicity of aqueous solutions. The solution could be water only, water +acid, water +base, or water + acid + base. - Water is always present. The solution being ranked could be water only, water + acid, water + base, or water + base + acid. Notice, however, that water is always there. - how those acids or bases influence the equilibrium for the ionization of water.
Why might a reaction with a very high energy of activation have a fast reaction rate? Why might a reaction with a very high energy of activation have a ΔG less than zero? Why might a reaction with a very high energy of activation have a very large Keq?
- A reaction with a very high Ea can have a fast rate as long as temperature is high - if the temp is high enough it can - There can definitely be more products than reactants, In many reactions the products are highly favored and there will be a lot of them compared to reactants once you reach equilibrium
Bronsted-Lowry Acid/Base definition. Bronsted-Lowry Acid/Base is all about the what?
- Acids donate H+ - Bases accept H+ - the protons
What happens when salts of weak acids/bases are placed in water? NH3 = ? NH4+ = ? NH4NO3 = ? When the salts of weak acids or weak bases dissolve in water one of the ions will undergo?
- Dissolve, one of the ions will undergo hydrolysis to reform the weak acid or base - "weak base" - "conjugate acid" - "salt of a weak base" - hydrolysis to re-form the weak acid or the weak base
Titration definition If a question states: "A strong base is titrated with a strong acid," which one is being added drop-wise and which one is in the beaker? Which solution is referred to as the titrant? Which solution is referred to as the anylate? The base is "titrated with" ? The analyte is? The titrant is?
- Drop by drop mixing of an acid and a base with an indicator - The terminology used in the question infers that the strong base is in the beaker, which makes it the analyte. - the strong acid, meaning the acid is being added dropwise and is therefore the "titrant." - whats in the beaker. whats being added dropwise
List the 8 MCAT strong acids: H3O+ is a borderline strong acid. It does have a pKa < 0, but just barely (pKa = -1.7). Many texts use it as a line of demarcation: acids stronger than hydronium ion are "strong" and acids weaker than hydronium ion are "weak." Is HF a strong acid? Explain why HF is NOT a strong acid and yet the closely related HCl is a strong acid?
- HI, HBr, HCl, HNO3, HCLO4, HCLO3, H2SO4, H3O+ - The reason HF is not a strong acid and HCl is a strong acid is a matter of structure. Looking at the conjugate bases, F- is far less stable than is Cl- due to its smaller size. When a smaller molecule has to bear a full formal charge it experiences a greater charge density and therefore more instability. A larger atom such as chlorine can spread out this charge over a greater area.
HSCN has a pKa of -1.8, HBF4 of .05 and HIO of 10.5, which one would have the strongest conjugate base? What is the pKb of HIO? How do you solve this?
- HIO will have the strongest conjugate base. Large pKa or a pKa larger than 0 is usually very weak, therefore its conjugate will be relatively strong. HI has the largest pKa, making it the weakest acid, its conjugate will therefore be the strongest base. - 3.5, remember pKa + pKb = 14
What is the molecular form of nitric acid? What about nitrous acid? What happens if you add a small amount of a strong base to a buffer solution, will the pH rise a lot or a small amount? If you add a large amount on the other hand, how will the pH change?
- HNO3 - HNO2 - it will not raise very much - it will raise a lot and become overpowered by the strong base
Species X has a reduction potential of 0.88V. Species Y has a reduction potential of 0.23V. If an electrolytic cell is constructed using these two metals, which metal will be used at the cathode?
- In a Galvanic cell, because reduction happens at the cathode, the species with the higher reduction potential would be at the cathode. However, because this is an electrolytic cell we know that the electron flow will be forced in the opposite direction—toward the metal with the lower reduction potential. In this case, that is species Y, so we know that metal Y will be at the cathode.
Summary Card
- Many students become confused as to how the three equilibriums described above(Ionization of Water, Acid Dissociation, and Base Dissociation) relate to one another. The ionization of water is often called the "autoionization of water" precisely because it happens automatically in allwater. So, when an acid or a base is added to water the autoionization equilibrium is already ongoing in that solution—a process we can describe with an equilibrium constant, Kw. If we are at 25°C we know that Kw will be 10-14 and the concentrations of H+ ions and OH- ions will both be 10-7M. After adding an acid or a base we suddenly have two equilibriums present in the same solution: 1) the ionization of water, and 2) the equilibrium for the acid or the base we added. Just as we described the ionization of water with an equilibrium constant, Kw, we can also describe the dissociation of the acid or base we added with Ka or Kb. The important key many students cannot visualize is this: When we add an acid or base to water, the equilibrium of that acid or base will directly impact the equilibrium for the ionization of water according to Le Chatelier's Principle. Looking at the formula, H2O + H2O H3O+ + OH-, we can see that adding an acid will shift the equilibrium to the left. This will use up hydroxide ions. Each hydroxide ion that reacts will also use up one hydronium ion, but remember that we just added extra hydronium ions in the form of the acid. The net result will be more hydronium ions relative to hydroxide ions and therefore a lower pH (Remember that the [OH-] equaled the [H+] before we added the acid). Similarly, if we add a base to neutral water we can see that it will also shift the reaction to the left, but this time we will be using up H3O+ ions. The net result will be more hydroxide ions relative to hydronium ions and therefore a higher pH. Notice that the addition of either an acid or a base shifts the equilibrium for the ionization of water to the left! Of course, we could add both an acid and a base to the same solution of water—creating three equilibriums in the same solution. In that case the acid and base equilibriums would have competing influences on the equilibrium for the ionization of water. If the acid and base were of equal strength, there would be no net effect and the pH would remain neutral.
Weak Acids and Weak Bases, do they disassociate completely? A general rule is that acids with a pKa > 0 or Ka < 1 and bases with pKb >0 or Kb less than one are considered? Anything not on the strong list is concidered? Examples of Weak Acids: Examples of Weak Bases:
- No - weak. - WEAK - H2O, H2S, NH4+, HF, HCN, H2CO3, H3PO4, acetic acid, benzoic acid. - H2O, NH3, R3N, pyridine, Mg(OH)2
What do you do to get the pKa from Ka? If the exponent is positive is the pKa (+) or (-) ? The Ka of DOOH is 2 x 10-5, how do you find the pKa?
- Simply take the negative log of the Ka to get the pKa - (-) - t is worthwile, however, to work thru how you get 4.8 as the pKa for DOOH. The Ka of DOOH is 2 x 10-5. When the mantissa is 1.0, the pKa (and this applies exactly the same to pH/pOH/etc.) will be exactly the negative of the exponent.However, when the mantissa increases we begin to move toward the next largest exponent. There are ten possible units because once the mantissa reaches ten we have arrived at the next largest exponent. In this case, the 2 tells us that we are 2/10ths of the way between a pka of 4 and a pka of 5. We know we are closer to the exponent of 5 because we're just barely different than 1.0, which would have been a pKa of 5. Thus, we subtract the 2/10ths from 5 and get 4.8
The conjugate base of a strong acid is usually a weak base. Therefore, it can also be said that the Kb of the conjugate base of a strong acid compares in what way to the Kb of a strong base? A) The Kb of the conjugate will exceed the Kb of the strong base. B) The Kb of the conjugate will be approximately equal to the Kb of the strong base. C) The Kb of the conjugate will be less than the Kb of the strong base. D) The Kb of the conjugate cannot be determined without also knowing the Ka.
- This is an excellent illustration of one of the ways in which the MCAT will require you to think conceptually. There is no memorized rule that will allow you to quickly answer. You must stay calm, re-read with careful scrutiny and produce an answer based on your conceptual mastery of each term. The entire stem can basically be rephrased to ask, "How do the Kb values of a strong base and a weak base (the conjugate of the strong acid) compare?" Because Kb times Ka is equal to a constant, if one is high, the other must be low, and vice versa. A strong base must have a high Kb and thus has a low Ka. The weaker base would have a lower Kb, also allowing it to have a relatively higher Ka. Answer C is thus correct.
Titrant and anylate How many equivalents of base can be neutralized by one equivalent of H2SO4?
- Titrant is the known concentration of acid/base that is added to the anylate which is the substance being investigated. - Two equivalents of base can be neutralized by one equivalent of H2SO4 because each sulfuric acid produces two equivalents of hydrogen ions in solution.
All equilibrium constants (Keq, Ka, Kb. Kw or Ksp) are written via? You should think of Ka or Kb just as you do Keq. A large Ka (or a small pKa) indicates? A large Kb (or a small pKb) indicates?
- Via the law of mass action with pure liquids and solids omitted. - that at equilibrium there are far more products than reactants. - very strong base (i.e., a lot of OHformed—either from dissociation of a hydroxide base (i.e., NaOH) or from deprotonating water).
Acid Disassociation Ka= ? Because the acid almost fully dissociates, what do we know about the ratio of products over reactants? In general terms, an acid with Ka greater than one or a pKa less than zero is considered?
- [H+] [A-] / HA - the ratio of products over reactants would have to be greater than one. - "strong," so this acid would clearly qualify as a strong acid.
Multiply the base disassociation equation times the Acid Disassociation equation and what do you get? Acid A has a Kb of 1.0 x 10-9 and Acid B has a Kb of 1.0 x 10-10. Which acid will create the largest decrease in pH when added in equimolar amounts to pure water?
- [H+] [OH-] =Kw - Acid B will give the largest drop in pH. The largest decrease in pH will be caused by addition of the most acidic of the two species. You could simply recognize from the equation Kw = Ka*Kb that Kb and Ka are inversely related, and therefore recognize that the smaller Kb represents the stronger acid (because Ka and acid strength are directly related). You could also use the above equation, plugging in 1 x 10-14 for Kw, and solve for Ka in both cases. Either method leads to the conclusion that the acid with a Kb of 1 x 10-10 is the stronger acid and will therefore lower pH to the greater extent.
Base Disassociation - Kb= ? Ka*Kb = Kw = 10^-14 (at 25°C); because? T/F An aqueous solution with a pH of 8 is basic and therefore by definition it does not contain any unreacted H+ ions.
- [OH-][BH] / [B-] - because ([H+][A-]/[HA])*([OH-][HA]/[A-]) = [H+][OH-] = Kw - An aqueous solution with a pH of 8 is basic, but that does NOT mean that it does not contain any hydrogen ions. In fact, the presence of hydrogen ions is easily verified by solving the formula pH = -log[H+] for [H+]. There are 1.0 x 10-8 moles of hydrogen ions per liter of this solution. It is classified as basic because it has fewer hydrogen ions than are found in neutral water and more hydroxide ions than are found in neutral water.
A buffer solution contains? The term "equivalence point" means that the solutions have been mixed in exactly the right proportions according?
- a weak acid and base, often the conjugates of each other. In a buffer there is an equilibrium between a weak acid and its conjugate base, or between a weak base and its conjugate acid. - to the equation
For which of the following titrations will the [OH-] = [H+] at the equivalence point? For which titrations will [titrant] = [analyte] at the equivalence point? a) SA with SB, b) WB with SA, c) WA with SB, d) WA with WB.
- a) At the equivalence point of the titration of a strong acid with a strong base the [OH-] will equal the [H+] AND the [analyte] will equal the [titrant] in the flask; b) At the equivalence point of the titration of a strong acid with a weak base the [OH-] does NOT equal the [H+], but the [analyte] will equal the [titrant]; c) At the equivalence point of a titration between a weak acid and a strong base the [OH-] will NOT equal [H+], but the [analyte] will equal the [titrant]; d) At the equivalence point of a titration between a weak acid and a weak base the [OH-] will NOT equal [H+], but the [analyte] will equal the [titrant] (remember WA/WB titrations are rarely attempted or useful). The pattern is that the hydroxide and hydrogen ions will be equal at the equivalence point for any "strong/strong" titration, but NOT for any other titrations. The concentration of the analyte will equal the concentration of the titrant at the equivalence point for all titrations.
Hydrolysis of which of the following salts in solution will increase the pH of the solution? a) NaNO2, b) NH4Cl, c) NaF, d) NaClO2, e) CH3COONa, f) NaCl. How do you solve?
- a) Hydrolysis of NaNO2 will result in the reaction of a nitrite ion with water to form HNO2 and hydroxide ion, increasing pH; b) hydrolysis of NH4Cl will result in reaction of NH4+ with water to form NH3 and H3O+, DECREASING pH; c) hydrolysis of NaF will result in fluoride ion reacting with water to form HF and OH-, increasing pH; d) hydrolysis of NaSO2 will result in reaction of ClO2-with water to form HClO2 and OH-, increasing pH; e) hydrolysis of CH3COONa will result in reaction of acetate with water to form CH3COOH and OH-, increasing pH. In summary, all of the options will increase pH except for option b) Look for the one that makes H3O+ instead of OH-
To review, the equivalence point is? The End point is? Three common indicators used in the lab are methyl orange, litmus, and phenylphthalein. These indicators change color at a pH of 3.7, 6.5, and 9.3, respectively. Which indicator should be chosen to identify the equivalence point for the titration of a weak acid with a strong base?
- where [titrant] = [analyte] - where the indicator causes the color to change. There is no causal relationship between the two. - An indicator should be chosen that will change color at a pH as close to the equivalence point as possible. A titration of a weak acid with a strong base will have an equivalence point greater than seven on the pH scale. The first two indicators would change color before the equivalence point was reached, so phenylphthalein would be the appropriate choice.
T/F? a) Species for which E° is negative cannot be spontaneously reduced, but are often oxidized; b) The hydrogen half-cell has no affinity for electrons as demonstrated by its electrical potential, E° = 0.00V; c) based on the half-reactions given in the table above, the potential for Cu(s) to be reduced by one electron is -0.52V.
- a) This statement is false. A species with a negative reduction potential can be spontaneously reduced as long as it is paired with another species that has a more negative reduction potential. This is easily proven. Suppose species A has a potential of -1.5V and species B has a potential of -1.8V. Species B will be oxidized, so we reverse the sign and add it to the potential for species A: -1.5 + 1.8 = 0.3V. With a positive cell potential we know this pairing would proceed spontaneously in a Galvanic cell. b) This statement is also false. Hydrogen ions do have an affinity for electrons and can be reduced. The potential of 0.00V was arbitrarily assigned to the hydrogen half-cell to facilitate an assignment of standardized potentials. This is a good example of the need for students to think—to actually think—about the logic of a statement and try to evaluate it based on other things they know. If one thinks about this only in terms of "reduction potentials" it may not be obvious that this is a false statement. However, if one asks: "Is it logical that hydrogen ions cannot be reduced? Is it logical that a hydrogen ion, with its positive formal charge, has no affinity for electrons? Have I ever seen an H+ ion get reduced? A reduced hydrogen ion would be a hydrogen without a charge: or nearly every hydrogen one would encounter in chemistry outside of acids. Notice that this little demonstration of the need to "think" involved asking oneself a series of questions—or in other words—the Socratic Method! c) We included this question because it seems to create confusion. Reduction half-reactions can be reversed to give oxidation potentials. In other words, the half-reaction runs in the opposite direction. Notice, however, that the reverse of one of these half-reactions involves the LOSS of one or more electrons as the metal forms the associated metal cation. For some reason, it is common for students to think that reversing the sign of the reduction potential gives the voltage associated with reduction of the solid metal. For the MCAT, just remember that cations (Cu+, Fe2+, etc.) get reduced to form solid metals (Cu(s), Fe(s), etc.), and solid metals get oxidized to form cations, but solid metals are NOT reduced.
H2O + H2O ↔ H3O+ + OH- The equilibrium for the ionization of water is always present in? Adding an acid shifts the reaction to? Adding a base shifts the reaction to? Kw, Ka and Kb are used to describe? At 25°C, Kw = Ka*Kb; this should make sense for two reasons
- aqueous - the left, increasing the relative [H+] - the left, increasing the relative [OH- - the three equilibriums - 1) we demonstrated above that this is mathematically true, and 2) if we always remain at 25°C the Kw for the ionization of water should never change—per our rule that temperature is the only thing that changes Keq.
Ionization of water H2O + H2O --> H3O+ + OH- -H3O+ is the same as? -Kw=[H3O+] [OH-] = ? -pKw=pH + pOH = ? -pKa + pKb = ? Show how the equation above for pKw is derived from the equation for Kw.
- as H+ - 10^-14 - 14 - 14 - Starting with Kw = [H3O+][OH-] = 10^-14, we take the negative log of all terms, yielding: -logKw = -log[H3O+] + -log][OH-] = -log(10-14). The middle two terms come from the log rule that states: logAB = logA + logB. The first term can be replaced with pKa by definition, the second term with pH by definition, the third term with pOH by definition, and the fourth term by 14 because 14 is the -log of 10^-14. This leaves: pKw = pH + pOH = 14
In a Galvanic cell, reduction happens at the ? The species with the higher reduction potential would be at? In an electrolytic cell which metal is used at the cathode, this will have a higher or lower reduction potential?
- cathode - cathode - lower
Example of the previous All strong acids and strong bases dissociate 100% in water (making them good?
- electrolytes
Weak acids will have a higher pKa or lower?
- higher
Whats the best way to find which one has the lowest pH? A. NH4ClO4 B. NH4CH3COO C. NaOH D. NaH
- if you know your strong bases you can see that C and D are both basic. If you know how salts work choice B is a salt made of the conjugate acid of a weak base and the conjugate base of a weak acid. In solution, the conjugates will revert to the associated weak acid and weak base as they exchange protons with water. For the pair in answer A, however, the ammonium will react with water to reform the weak base ammonia plus hydronium ions (slightly decreasing pH)
An atom that is reduced looses or gains electrons?
- it gains electrons. It also is an oxidizing agent.
Summary page The ionization of water is often called the "autoionization of water" precisely because? When we add an acid or base to water, the equilibrium of that acid or base will? Notice that the addition of either an acid or a base shifts the equilibrium for the ionization of water to the?
- it happens automatically in all water - directly impact the equilibrium for the ionization of water according to Le Chatelier's Principle - left
An atom that is oxidized looses or gains an electron?
- it looses an electron. Which makes it a reducing agent.
Titration of a WA w/ SB or WB w/ SA Equivalence Point = ? Because the acid and base are not both strong (i.e., dissociate 100% in water) the [H+] ? This also tells us that the pH?
- midpoint of the nearly vertical section [titrant] = [analyte] - [H+] DOES NOT EQUAL [OH-] - DOES NOT EQUAL 7 - This is why it is very important that you clearly master these two unique types of titrations and keep their behavior clearly separated in your mind. An MCAT question will simply ask: "What is the pH of the solution at the equivalence point?" or an answer choice will say "The concentration of hydroxide ions equals the concentration of hydrogen ions at point B." You will need to remember that the first thing you must do on such a question is decide which type of titration is being performed.
If you know the pKa of a substance can you subtract from 14 to find the pKb?
- no, subtracting from 14 will give the pKb of its conjugate
At the half-equivalence point the concentrations of HA and A- are equal. Therefore, the ratio of [A-]/[HA] must be equal to? When we plug the answer into the Henderson Hasselback equation we get: ? The log of one is? Show how pH = pKa
- one. - pH = pKa + log(1). - zero, Since the log of 1 is zero, the log term term falls out, demonstrating that pH = pKa at the half-equivalence point.
Rusting, is a reduction or oxidation process? If something rusts easily it must have a high or low oxidation potential? What about its reduction potential? Will something that rusts easily have a low reduction potential?
- oxidation - high - low
If Keq is greater than 1, is Gibbs Free energy negative? Gibb's is measured as the difference between? If the absolute value of the free energy of the products exceeds that of the reactants, will Gibbs be negative?
- yes - products and reactants - yes
Calculate pH for Strong Acid/Base What must we assume to make this calculation?
- pH or pOH = -log [strong acid/base] - We calculate pH and pOH by taking the negative log of the concentration of H+ ions and OH- ions respectively. Therefore, if we take the -log of the concentration of the acid or base directly we are assuming that the molar concentration of the acid or base is equal to the molar concentration of hydrogen ions or hydroxide ions, respectively. This cannot be exactly true because some of both of these ions (10-7M) are already present in water before addition of the acid or base. The assumption is usually safe because the molar concentration of the strong acid or base is usually many magnitudes larger than 10-7. If the difference were smaller this would weaken the validity of our assumption. Finally, we are also assuming that the acid and base dissociate 100%. If not, even if the first assumption were true, we could not take the -log of the concentration of the acid or base directly. Say, for example, that only 75% of the acid dissociated in solution—this would be a significant difference in concentration compared to 100% dissociation. Assuming that very strong acids dissociate 100% is usually a safe assumption, although ion-pairing and other factors due reduce the effective concentration of ions. This seems like a very likely MCAT topic because they tend to ask many questions that require a comparison between how we assume things to be for calculation purposes and what they are actually like in reality (real vs. ideal fluids, real vs. ideal gases, ignoring air resistance, etc.).
pH equation pH of pure H2O at 25°C = ? pH > 7 = ? pH < 7 ?
- pH=-log [H+] - 7 - Basic - acidic
Cell potential for a Galvanic cell is always positive or negative? A functioning Galvanic cell can be created using ? If a galvanic cell is properly set up, it will always produce? Electrons will automatically flow from the species with?
- positive - any two metals, regardless of their reduction potentials - current - the lower reduction potential to the species with the higher reduction potential
Blue litmus paper turns what color in acid? In a base what color does it change?
- red - stays blue in base
For an electrolytic cell, the species with the lower reduction potential will be reduced or oxidized? - The cell potential will always be positive or negative? - The sum of the externally applied voltage (Vbattery) and the negative cell potential (-E°cell) must be positive or negative?
- reduced - negative - positive
The nearly horizontal area surrounding the half-equivalence point is called? Adding a relatively large amount of titrant at this point in the titration will have little or a lot effect on pH?
- the "buffer region." - little. That is the very reason the graph is flat in this region—the volume of titrant added increases (x-axis) with little or no increase in pH (y-axis).
The "conjugate base" of an acid is? The "conjugate acid" of a base is?
- the acid minus its hydrogen - the base plus a hydrogen (e.g., NH3 = base; NH4+ = conjugate acid) *you could call HCl the acid and Cl- the conjugate base or call Cl- the base and HCl the conjugate acid.
Reduction always happens at? Oxidation always happens at ? Cathode = ? Anode = ?
- the cathode - the anode - (+) - (-) True of galvanic cells only, NOT electrolytic cells!
Hydrolysis of a salt terminology. This terminology can be a bit tricky: The "salt of a weak acid" refers to? HCO3- = ? CO32- = ? Na2 CO3 = ? Similarly, the "salt of a weak base" refers to?
- the conjugate base of that weak acid combined with a cation to form a salt (see example below). - weak acid - conjugate base - salt of a weak acid - the conjugate acid of that weak base combined with an anion to form a salt (example: NH3 = "weak base"; NH4+ = "conjugate acid"; NH4NO3 = "salt of a weak base")
Electrical potentials tell us? These potentials are given in volts and are always presented in what is called?
- the degree to which a species "wants electrons," or "wants to be reduced." - "half-reaction." o Ag2+(aq) + 2e- → Ag(s) E° = 0.80 V o Cu+(aq) + 1e- → Cu(s) E° = 0.52 V o Ni2+(aq) + 2e- → Ni(s) E° = -0.23 V o Zn2+(aq) + 2e- → Zn(s) E° = -0.76 V
Kw is ? What is pKw? What is the p for?
- the ionic product for water, Kw is essentially just an equilibrium constant - p stand for power, indicating that pKw is the -log10 of Kw
A student is using a coffee cup calorimeter to determine the enthalpy of a reaction. When assembling the apparatus, one of the coffee cups is accidentally torn along one side. The student checks to make sure the reactants aren't leaking out and decides to continue with the experiment. The value he calculates for the heat of reaction will be lower than the true value, why?
- to calculate heat you would need the total change in temperature. If there is a hole in the cup, more heat would escape than normal, and the change in temperature would appear less than had the hole not been there. For example, if the hole was not there the temp change may have been 40 degrees, but since some heat escaped from the hole the now measured temperature change could be 20 degrees which would distort the real each exchange of the reaction.
Impact of salts on the dissolution of Weak Acids and Bases. The percent disassociation of benzoic acid (weak acid) decreases in a sodium benzoate solution. Sodium benzoate dissociates to release benzoate ions, which shift the acid dissociation equilibrium in what direction? The percent disassociation of ammonium hydroxide (weak base) decreases in an ammonium chloride solution. Ammonium chloride dissociates to release ammonium ions, which shift in what direction?
- to the left. - to the left.
Indicators are ? To set up a titration, you must know beforehand? The dissociation of the indicator and the acid/base reaction we are analyzing run simultaneously in the same beaker but are ? The amount of indicator is so small compared to titrant and analyte that we can assume ?
- weak acids that change color as they dissociate from HA into H+ and A- - the approximate pH of your equivalence point; you then select an indicator that will change color at that approximate pH - otherwise unrelated. - it has no impact on pH
4 things to remember about cell potential
1) Half reactions always come in pairs; one reduced + 1 oxidized 2) Eo for any oxidation half reaction is simply the negative of Eo for the associated reduction half reaction 3) You CANNOT add 2 Eo values off of a table. Reverse the species with the lowest reduction potential and take the negative of its Eo value, now add to get Ecell. 4) DO NOT USE STOICHIOMETRY. One mole of Cu2+ is same E.
7 Buffer problem clues
1) The maximum buffering strength occurs when the [HA] is equal to the [A-]. This is the ratio one would start with if making a buffer in the lab. Therefore, be on the lookout for equimolar amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid. 2) Watch for conjugates. To be a buffer, the two equimolar substances are often conjugates ofeach other, such as: NH3 and NH4, CH3COO- and CH3COOH, or HCO3- and CO32-. 3) WEAK acids or bases. The equimolar pair must be a weak base or a weak acid and its conjugate. Strong acid or strong base conjugate pairs do NOT form buffers. 4) Resistance to pH change. Any time an acid or base is added to a solution and the pH does not change "very much," or "changes slightly," this should be a dead giveaway. 5) Half-equivalence point. Only solutions of weak conjugate acid/base pairs have a buffer region, and therefore they are the only solutions that have a halfequivalence point. 6) Watch for pH = pKa. pH = pKa at the midpoint of the buffer region. This is another unmistakable clue that you are dealing with a buffer. 7) Watch for the ratio of [HA]/[A-] or [A-]/[HA]. Ratio of an acid to its conjugate base. Recognize it as a buffer problem
Calculate pH for Weak Acid, 6 steps
1) Write out the equilibrium reaction (HA-> H+ + A-) 2) Use x to represent concentrations 3) use [HA] - x for original acid 4) If a quadratic formula results, assume x is much smaller and omit 5)Solve for X using Ka=(x)(x) / [HA-x] 6) Use -log[H+] to find pH
Steps to determine acidity or basicity of a salt:
1. Write the neutralization reaction in reverse (salt and water makes base and acid). 2. Break up the salt into its ions. 3. To write the base, add an OH- to the cation (the positive ion) for each + charge. 4. To write the acid, Add H+ to the anion (the negative ion) for each negative charge. 5. If you form a strong base, the salt is basic. (NaOH, KOH, LiOH, Mg(OH)2 ...etc.) If you form a strong acid it is acidic. (HCl, HNO3, HClO4, HBr...) If you form both a strong acid and a strong base it is neutral.
Hydrogen half cell The E° values assigned to each half-reaction represent ?
2H+ + 2e- --> H2 Eo =0.00V This is the standard against which all other half-reactions are rated and we define its electrical potential as Eo=0V. - the relative reduction potential of that species compared to the potential of two hydrogen ions to gain two electrons to form hydrogen gas. This is called the "Hydrogen Half-Cell." It is the standard against which all other half reactions are rated and we define its electrical potential as E° = 0.00V. The naught symbol (°) signifies that all potentials are measured under standardized conditions.
Strong base titrated with strong acid graph
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Weak acid titrated with a strong base graph
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Weak base titrated with a strong acid graph
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Electrolytic cells definition
A galvanic cell with a battery attached forcing the electrons to flow in opposite direction.
pH scale (numbers) How does the hydrogen ion concentration differ for two solutions, one with a pH = 2 and the other with pH = 4? The solution with a ph of 2 has a factor of how many more H+ ions?
A solution with a pH of 2 has [H+] = 10^-2, while a solution of pH 4 has [H+] = 10^-4, thus the pH 2 solution has 10^2 or 100 times the [H+] of a solution with a pH of 4
Concentration Cell: In one beaker the metal is oxidized via? In the other beaker it is reduced via ? Because the reduction potentials of oxidation and reduction half reactions for the same species only differ by? The Nernst equation is used to ?
A special type of galvanic cell; The same electrodes and solution are used in both beakers. - its oxidation half-reaction - its reduction half-reaction - the sign of E°, E°cell = 0.00V - calculate the cell potential based off of the E˚ of the species and the two concentrations.
Draw each of the following titration curves and label the axes. - Strong Acid (SA) titrated with a Strong Base (SB) - Strong Base (SB) titrated with a Strong Acid (SA) - Weak Acid (WA) titrated with a Strong Base (SB) - Weak Base (WB) titrated with a Strong Acid (SA)
For all of these plots, the x-axis is volume of titrant (usually mL) and the y-axis is pH. Strong Acid with Strong base: A strong acid in the flask will result in the pH being low, around 1 to 2 depending on the [SA]. As strong base is added, the pH will stay low and slowly rise until the equivalence point is reached and then it will go straight vertical for around 6 pH units. The middle of the equivalence region will be pH 7. After the equivalence point, the plot will slowly go to higher pH as more strong base is added. Strong Base with Strong Acid: With strong base present, the pH will start high. As strong acid is added, it will slowly go to lower pH and then sharply drop when the equivalence point is reached. Again, the middle of the pH region will be pH 7. It will then slowly go to lower pH as more strong acid is added. Weak Acid with Strong Base: As with the strong acid titration, this one will start at a pH below 7, but since the weak acid only partially dissociates, the pH will be higher than it is for a strong acid, or usually around 3 to 5. As strong base is added, the pH will increase slowly through the buffer region and then go vertical at the equivalence point. The equivalence point won't cover as many pH units as it did for the strong acid titration and leading up to it and after it, the pH will more quickly rise than it did for the strong acid. The half way pH of the equivalence point will not be 7 as it is for a strong acid, it will be higher than 7. This is because the conjugate base of the weak acid is now present, creating a basic solution. After the equivalence point, the pH will slowly rise as more base is added. Weak Base with Strong Acid: This will look like the strong base/strong acid titration except the pH will not start as high, the equivalence region won't last for as many pH units, and the equivalence point will be at a pH below 7. The reasons for all of these are for like reasons given in the weak acid/strong base titration, except that they need to be applied to the weak base.
Galvanic cell (aka Voltaic cell) definition
Galvanic cells convert chemical energy into electrical energy. By taking advantage of the difference in electrical potential between two metals, a current can be generated spontaneously.
List the 8 strong bases for the MCAT. What is the henderson hassback equation?
Group IA hydroxides (LiOH, NaOH, KOH, RbOH) NH2-, H-, Ca(OH)2, Sr(OH)2, Ba(OH)2, Na2O, CaO *Disassociate 100%
What is the purpose of a salt bridge?
In the cell above notice that over time there will be a buildup of negative charge in the copper vessel due to continual loss of copper cations, and a buildup of positive charge in the zinc vessel due to the continual production of zinc cations. This polarity resists the flow of electrons and would eventually shut down the cell if a salt bridge were not present. Within the salt bridge sodium ions can flow toward the copper vessel and nitrate ions can flow toward the zinc vessel, neutralizing the buildup of charge and allowing electron flow to continue. The metal cations themselves, as well as any other ions in the solutions, can also flow through the salt bridge. In an electrical sense, the salt bridge connects the circuit, allowing continual flow of electrons from electrode to electrode and then back through the salt bridge via ion diffusion.
Half equivalence point= HA will continue to be deprotonated until at the equivalence point the solution contains? This is not true of a titration involving a strong acid and strong base because they both dissociate 100%. Therefore, by definition, almost immediately after adding HA 100% has been changed to A-. Therefore, SA/SB titrations?
midpoint of the nearly horizontal section of the graph. pH=pKa or [HA] = [A-] BUT THEN at the equivalence point there is 100% A- and 0% HA *SA/SB titrations do NOT have half equivalence points - contains 100% A and 0% HA - do NOT have half-equivalence points
Cell Potential (aka EMF or Ecell) definition A species with a positive E˚ is more likely to? A species with a negative E˚ is less likely to ?
the sum of the electrical potentials for the two half reactions that make up an electrochemical cell. - gain electrons (i.e., be reduced) than are hydrogen ions. - gain electrons than are hydrogen ions.