Chemistry 3
Five main points from the three equilibrium that must be understood
-The equilibrium for the ionization of water is always present in aqueous solutions. -Adding an acid shifts the reaction to the left, increasing the relative [H+] -Adding a base shifts the reaction to the left, increasing the relative [OH-] -Kw, Ka and Kb are used to describe these three equilibriums. -At 25°C, Kw = Ka*Kb; this should make sense because 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.
When is a salt solution basic or acidic?
1) Salts that are from strong bases and strong acids do not hydrolyze. The pH will remain neutral, at 7. Halides and alkaline metals dissociate and do not affect the H+ as the cation does not alter the H+ and the anion does not attract the H+ from water. This is why NaCl is a neutral salt. In General: Salts containing halides (except F-) and an alkaline metal (except Be2+) will dissociate into spectator ions. 2) Salts that are from strong bases and weak acids do hydrolyze giving it the pH greater than 7. The anion in the salt is derived from a weak acid, most likely organic, and will accept the proton from the water in the reaction having the water act as an acid in this case leaving a hydroxide ion (OH-). The cation will be from a strong base, meaning from either the alkaline or alkaline earth metals so like before it will dissociate into an ion and not affect the H+ 3) Salts of weak bases and strong acids do hydrolyze giving it a pH less than 7. This is due to the fact that the anion will become a spectator ion and fail to attract the H+ ion, while the cation from the weak base will donate a proton to the water forming a hydronium ion. 4) Salts from a weak base and weak acid also hydrolyze as the others, but a bit more complex and will require the Ka and Kb to be taken into account.Whichever is the stronger acid or weak will be the dominate factor in determining whether it is acidic or basic. The cation will be the acid, and the anion will be the base and will form either form a hydronium ion or a hydroxide ion depending on which ion reacts more readily with the water.
State the seven clues that can be used to recognize a buffer problems.
1) Watch for equimolar amounts. The maximum buffering strength occurs when the [HA] is equal to [A-]. This is the ratio one would start with if making buffer in the lab. Therefore, be on the look out 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 of each other, such as: NH3 and NH4, CH3COO- and CH3COOH, or HCO3- and CO3(2-). 3) Watch for 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) Watch for 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 that the solution is a buffer. 5) Watch for the half-equivalence point. Remember that 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. Many students memorize this principle, but fail to recognize that what it is really saying is that 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]. One particular problem on the CBT practice exams seems to stump the majority of students. The question stem asks about the ratio of an acid to its conjugate base. Those who recognize it as a buffer problem usually answer it correctly. Those who do not guess.
Talk about the hydrogen half cell.
2H+ + 2e- --> H2 E°= 0.00V The E° value assigned to each half-reaction represent 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 halfreactions are rated and we define its electrical potential as E° = 0.00V. The naught symbol (°) signifies that all potentials are measured under standardized conditions.
What is the difference between faraday and farad?
A Faraday is 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 Farad is 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 x m/s2, we can say 1 Farad instead of saying 1 C2 x s2/m2 x kg. A Farad is the amount of capacitance necessary to hold 1 C of charge on a capacitor with a potential difference of 1 Volt.
What is buffer used for and what is usually contained in buffer solution?
A buffer is a solution that can resist pH change upon the addition of an acidic or basic components. It is able to neutralize small amounts of added acid or base, thus maintaining the pH of the solution relatively stable. This is important for processes and/or reactions which require specific and stable pH ranges. Buffer solutions have a working pH range and capacity which dictate how much acid/base can be neutralized before pH changes, and the amount by which it will change. To effectively maintain a pH range, a buffer must consist of a weak conjugate acid-base pair, meaning either a. a weak acid and its conjugate base, or b. a weak base and its conjugate acid. In a buffer there is an equilibrium between a weak acid and its conjugate base, or between a weak base and its conjugate acid.
A note about galvanic cells.
A functioning Galvanic cell can be created using any two metals, regardless of their reduction potentials. If a galvanic cell is properly set up, it will always produce current. Electrons will automatically flow from the species with the lower reduction potential to the species with the higher reduction potential. However, some MCAT questions read something like this: "Which of the following species can form a spontaneous Galvanic cell with copper, where copper is at the cathode?" That creates an entirely different situation. If the question requires that copper be reduced, then it must be paired with a species with a lower reduction potential.
If the pH of water falls as temperature increases, does this mean that water becomes more acidic at higher temperatures?
A solution is acidic if there is an excess of hydrogen ions over hydroxide ions. In the case of pure water, there are always the same number of hydrogen ions and hydroxide ions. That means that the water remains neutral - even if its pH changes. The problem is that we are all so familiar with 7 being the pH of pure water, that anything else feels really strange. Remember that you calculate the neutral value of pH from Kw. If that changes, then the neutral value for pH changes as well. At 100°C, the pH of pure water is 6.14. That is the neutral point on the pH scale at this higher temperature. A solution with a pH of 7 at this temperature is slightly alkaline because its pH is a bit higher than the neutral value of 6.14. Similarly, you can argue that a solution with a pH of 7 at 0°C is slightly acidic, because its pH is a bit lower than the neutral value of 7.47 at this temperature.
How does the hydrogen ion concentration differ for two solutions, one with a pH = 2 and the other with pH = 4?
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.
What does the sign of E° value tell you?
A species with a positive E° is more likely to gain electrons (i.e., be reduced) than are hydrogen ions. A species with a negative E° is less likely to gain electrons than are hydrogen ions.
What are the differences between strong and weak acids?
A strong acid is one which is virtually 100% ionized in solution. pH is a measure of the concentration of hydrogen ions in a solution. Strong acids like hydrochloric acid at the sort of concentrations you normally use in the lab have a pH around 0 to 1. The lower the pH, the higher the concentration of hydrogen ions in the solution. A weak acid is one which doesn't ionize fully when it is dissolved in water. Most organic acids are weak. Hydrogen fluoride (dissolving in water to produce hydrofluoric acid) is a weak inorganic acid that you may come across elsewhere.
What are the differences between strong and weak bases?
A strong base is something like sodium hydroxide or potassium hydroxide which is fully ionic. You can think of the compound as being 100% split up into metal ions and hydroxide ions in solution. A weak base is one which doesn't convert fully into hydroxide ions in solution.
What is the general form for base dissociation?
A- + H2O <--> OH- + HA Kb = [OH-] [HA]/ [A-]
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?
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.
What are the definitions for Lewis acid and base?
Acids accept a pair of electrons; bases donate a pair of electrons. AlCl3 and BF3 are two common examples of Lewis acids. The electrophiles in all organic chemistry reactions are acting as Lewis acids. NH3, OH- and anything with an electron pair to donate will act as a Lewis base.
What are the definitions for Bronsted-Lowry acid and base?
Acids donate protons H+; bases accept protons H+
What are the definitions for Arrhenius acid and base ?
Acids produce H+ ions in solutions; bases produce OH- ions in solutions.
Describe the concept of acid-base equilibria. (Ka and Kb)
All equilibrium constants (Keq, Ka, Kb. Kw or Ksp) are written via the law of mass action, with pure liquids (l) and solids (s) omitted. You should think of Ka or Kb just as you do Keq. A large Ka (or a small pKa) indicates that at equilibrium there are far more products than reactants. For an acid dissociation, this would mean a lot of dissociation (i.e., a lot of H+ formed) and thus a very strong acid. Similarly, a large Kb (or a small pKb) indicates a very strong base (i.e., a lot of OH formed— either from dissociation of a hydroxide base (i.e., NaOH) or from deprotonating water).
T/F An aqueous solution with a pH of 8 is basic and therefore by definition it does not contain any unreacted H+ ions.
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.
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?
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.
Give the oxidation states of different molecules you must know.
Any elemental element -- O.S = 0 Flourine -- O.S = -1 Hydrogen -- +1 Hydrogen with a metal -- O.S = -1 Oxygen (except peroxides) -- O.S = -2 Alkali metals -- O.S = +1 Alkali earth metals -- O.S = +2 Group V -- O.S = -3 Group VI -- O.S = -2 Group VII --O.S = -1
Describe oxidation-reduction (Redox) reactions.
Any reaction where one or more electrons are transferred from one atom to another. The atom that loses electron is oxidized and the atom that gains electron is reduced. A reducing agent is an atom or molecule that donates electrons to another atom or molecule and is itself oxidized in the process. An oxidizing agent is an atom or molecule that accepts electrons and is itself reduced in the process. 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 reducing agent and water is an oxidizing agent. To recognize this, you must be able to calculate the oxidation state of each atom.
Acid HX dissociates 80% in water. Would you expect its Ka to be greater than, less than, or equal to one?
Because the acid almost fully dissociates, we know that the ratio of products over reactants would have to be greater than one. In general terms, an acid with Ka greater than one or a pKa less than zero is considered "strong," so this acid would clearly qualify as a strong acid.
Talk about concentration cell and the Nernst equation.
Concentration Cell: A special type of galvanic cell; The same electrodes and solution are used in both beakers. In one beaker the metal is oxidized via its oxidation half-reaction, and in the other beaker it is reduced via its reduction half-reaction. Because the reduction potentials of oxidation and reduction half reactions for the same species only differ by the sign of E°, E°cell = 0.00V. It appears that nothing would happen. However, all E° values are given for standard conditions (the reason for the naught symbol). You do NOT need to know those conditions for the MCAT, but one aspect of standard conditions happens to be 1M concentrations for both solutions. Concentration cells are therefore nonstandard conditions by definition. They have a positive reduction potential E (no naught symbol, signifying nonstandard conditions) if there is a difference in the molarities of the two solutions. The Nernst equation is used to calculate the cell potential based off of the E° of the species and the two concentrations. Yes, you do need to know the Nernst equation—it has shown up on the MCAT at least twice before. o Nernst Equation: 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)
Provide a general overview of the titration process.
Drop by drop mixing of an acid and a base with an indicator. When the indicator changes colour, this is often described as the end point of the titration. The term "equivalence point" means that the solutions have been mixed in exactly the right proportions according to the equation.
What is electrical potentials?
Electric potentials tell us the degree to which a species wants electrons or want to be reduced. These potentials are given in volts and are always presented in what is called a half reaction. Below are a few examples of half reactions and their associated potential in volts: Ag2+(aq) + 2e- Ag(s) E° = 0.80 V Cu+(aq) + 1e- Cu(s) E° = 0.52 V Ni2+(aq) + 2e- Ni(s) E° = -0.23 V Zn2+(aq) + 2e- Zn(s) E° = -0.76 V Notice that in all of these examples an aqueous metal ion is being reduced to form the associated solid metal. This is by far the most common half reaction. The only half reaction you are likely to see on the MCAT that do NOT begin with a metal cations are O2, H2O, H+
Distinguish between equivalence point vs end point.
Equivalence Point vs. End Point: There is always some confusion between these terms. To review, the equivalence point is where [titrant] = [analyte]. The end point is simply the point when the indicator causes the color change. There is no causal relationship whatsoever between the solution reaching the equivalence point and the indicator changing color (i.e., the endpoint). The extent to which they do or do not correlate depends entirely on the indicator chosen.
What are the unique characters of a electrolytic cell?
Essentially, a galvanic cell to which an external voltage is applied, forcing the electrons to flow in the opposite direction. o Oxidation still occurs at the anode and reduction at the cathode. o The species with the lower reduction potential will be reduced! o The cell potential will always be negative. o The sum of the externally applied voltage (Vbattery) and the negative cell potential (-E°cell) must be positive. o Cathode = (-); Anode = (+) Note the difference compared to galvanic cells.
Give examples of weak acids and bases.
Examples of Weak Acids: Anything NOT on the strong acid list. H2O, H2S, NH4+, HF, HCN, H2CO3, H3PO4, acetic acid, benzoic acid, etc. Examples of Weak Bases: Anything NOT on the strong base list. H2O, NH3, R3N, pyridine, Mg(OH)2, etc.
Using the above definition, and combining it with information you should recall from Physics 3, calculate Faraday's constant.
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.
Describe the steps involve in the process of calculating pH of a solution when an acid is added in water.
For Strong Acids/Bases: (e.g., pH of a 1 x 10-3M HCl solution) The pH or pOH = -log[strong acid or base]. We assume a 100 percent dissociation and that the molar concentration of the acid or base is equal to the molar concentration of hydrogen ions or hydroxide ions, respectively. For Weak Acids: (e.g., pH of a 1 x 10^-4M CH3COOH solution) 1) Write out the equilibrium equation (HA <--> H+ + A-) 2) Use x to represent the concentration of each of the two products (or 2x, 3x etc. depending on the coefficients in the balanced equation). 3) Use "[HA] - x" for the concentration of the original acid. 4) If this results in a quadratic equation, assume that x is much smaller than [HA] (in step #3 above) and omit it. 5) Solve for x from Ka = (x)(x)/[HA - x]. 6) Use -log[H+] (i.e., -log[x] to find the pH).
Describe the galvanic cell.
Galvanic cells convert chemical energy into electrical energy. By taking advantage of the difference in reduction potentials between two metals, a current can be spontaneously generated along a wire that connects two metals electrodes submerged in solutions that contain metal ions. Remember that the flow of current is opposite in direction when compare to the flow of electron (i.e, electrons flow to the cathode where reduction happens while current flow to the anode!)
Give the strong bases list.
Group IA hydroxides (NaOH, KOH, etc.), NH2-, H-, Ca(OH)2, Sr(OH)2, Ba(OH)2, Na2O, CaO.
Give the overall reaction and equation you need to know for the ionization of Water.
H2O + H2O <--> H3O+ + OH- (H3O+ is the same as H+) Kw = [H3O+][OH-] = 10^-14 pKw = pH + pOH = 14 pKa + pKb = 14
What is the general form for acid dissociation reaction?
HA + H2O <--> H3O+ + A- (H3O+ is the same as H+) Ka= [H+] [A-]/ [HA]
Give the strong acids list.
HI, HBr, HCl, HNO3, HClO4, HClO3, H2SO4, H3O+ Note: 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." Note: HF is NOT a strong acid, but is often mistakenly labeled as such.
What is the half-equivalent point?
Half-Equivalence Point = midpoint of the nearly horizontal section of the graph. Here the pH = pKa. The half-equivalence point is when half of your starting species has been either protonated (in the case of a base) or deprotonated (in the case of an acid). Say you have an acid to titrate, HA. At the half equivalence point, half of this acid has been deprotonated and half is still in its protonated form.We can also say that [HA] = [A-] at the half-equivalence point. HA will continue to be deprotonated until at the equivalence point the solution contains 100% Aand 0% HA. 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 do NOT have half-equivalence points.
What is the impact of salts on the dissolution of weak acids and bases?
Impact of Salts on the Dissolution of Weak Acids and Weak Bases: -The percent dissociation of benzoic acid (weak acid) decreases in a sodium benzoate solution. -The percent dissociation of ammonium hydroxide (weak base) decreases in an ammoniumchloride solution. Both observations are easily explained by Le Chatelier's principle. In case one, sodium benzoate dissociates to release benzoate ions, which shift the acid dissociation equilibrium for benzoic acid to the left. Similarly, ammonium chloride dissociates to release ammonium ions, which shift the base dissociation equilibrium for ammonium hydroxide to the left.
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
Describe indicators used in titration.
Indicators: Indicators are weak acids that change color as they dissociate from HA into H+ and A-. To set up a titration, you must know beforehand the approximate pH of your equivalence point; you then select an indicator that will change color at that approximate pH. The dissociation of the indicator and the acid/base reaction we are analyzing run simultaneously in the same beaker but are otherwise unrelated. The amount of indicator is so small compared to titrant and analyte that we can assume it has no impact on pH.
Would you expect a strong oxidizing agent to have a high or low reduction potential?
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.
Explain why Ka x Kb = 10^-14 at 25 degree Celsius
Ka x Kb = Kw = 10^-14 (at 25°C); because ([H+][A-]/[HA]) x ([OH-][HA]/[A-]) = [H+][OH-] = Kw
Summary about the relationship of ionization of water, acid dissociation and base dissociation. (Part 1)
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 all water. 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.
A biology connection with the cell potential concept.
O2(g) + 4H+(aq) + 4e- --> 2H2O(l) E° = 1.23 V; This is the last step of the Electron Transport Chain.
What is the meaning of the term one equivalent?
One "equivalent" = the amount of acid or base necessary to produce or consume one mole of [H+] ions.
What is the purpose of the salt bridge within the galvanic cell?
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
Describe the unique properties of a galvanic cell.
Reduction always happen at the cathode while oxidation always happen at the anode. This is true for all electrochemical cells. Cathode is + while anode is - This is true ONLY for galvanic cells NOT electrolytic cells. Cell potential is always positive. This is true ONLY for galvanic cells NOT electrolytic cells.
general overview of aqueous solution of salts.
Salts, when placed in water, will often react with the water to produce H3O+ or OH-. This is known as a hydrolysis reaction. Based on how strong the ion acts as an acid or base, it will produce varying pH levels. When water and salts react, there are many possibilities due to the varying structures of salts. A salt can be made of either a weak acid and strong base, strong acid and weak base, a strong acid and strong base, or a weak acid and weak base. The reactants are composed of the salt and the water and on the product side remains the conjugate base (from the acid of the reaction side) or the conjugate acid (from the base of the reaction side).
Show how the equation for pKw= pH + pOH = 14 is derived from the equation for Kw.
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.
What are amphoteric substances?
Substances that can act as either an acid or a base (e.g., H2O).
Important note about acids and bases convention on MCAT.
The MCAT will always tell you if the substances in question are acting as Lewis acids and base. If not mention otherwise, always assume that the substances are acting as Bronsted Lowry acids and bases. Arrhenius definition is rarely used on the MCAT.
What is the cell potential and how would you calculate it?
The cell potential, E°cell, is the sum of the electrical potentials for the two half-reactions that make up an electrochemical cell. Remember the following: Half-reactions always come in pairs—one reduction half-reaction plus one oxidation half-reaction. This makes logical sense because in order for one species to be reduced, another species must be oxidized to make those electrons available. Although listed individually in a table, it always requires two species (i.e., two half-reactions) to make a cell. Usually, only the reduction half-reactions are given in tables. The oxidation half-reaction is the reverse of the reduction half-reaction. E° for any oxidation half-reaction is simply the negative of E° for the associated reduction half-reaction. You CANNOT add two E° values directly off of a half-reaction table. These are all reduction half reactions and you need one of each—one reduction and one oxidation. Therefore, you must reverse the half-reaction of the species with the lowest reduction potential and take the negative of its E° value. Only after changing the sign can you add these two together to get the E°cell. DO NOT USE STOICHIOMETRY! One mole of Cu2+ has the same reduction potential as two moles of Cu2+.
How will temperature effect Kw value?
The formation of hydrogen ions (hydroxonium ions) and hydroxide ions from water is an endothermic process. The forward reaction absorbs heat. According to Le Chatelier's Principle, if you make a change to the conditions of a reaction in dynamic equilibrium, the position of equilibrium moves to counter the change you have made. Therefore, if you increase the temperature of the water, the equilibrium will move to lower the temperature again. It will do that by absorbing the extra heat. That means that the forward reaction will be favoured, and more hydrogen ions and hydroxide ions will be formed. The effect of that is to increase the value of Kw as temperature increases.
Summary about the relationship of ionization of water, acid dissociation and base dissociation. (Part 2)
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.
Describe the buffer region and the Henderson Hasselbach equation.
The nearly horizontal area surrounding the half-equivalence point is called the buffer region. Adding relatively large amount of titrant at this point in the titration will have little effect on pH. In fat that is the very reason that graph is flat in this region--the volume of titrant added increase (x-axis) with little or no increase in pH (y-axis) pH= pKa + log [A-]/[HA] or pH=pKa - log [HA]/[A-] using this equation you can figure out the ratio of the acid and its conjugate base (e.g, if the ratio is a 100 then you know that there will be 1 molecule of the conjugate base per 100 molecules of the acid)
Provide an overview of the pH scale.
The pH scale is a mathematical ranking system for the acidity or basicity of aqueous solutions. The solution being ranked could be water only, water + acid, water + base, or water + base + acid. Notice, however, that water is always there. In fact, as you'll see in the next section, the pH scale ranks solutions based not so much on the acids or bases themselves, but on how those acids or bases influence the equilibrium for the ionization of water. The pH scale is logarithmic, meaning that an increase or decrease of an integer value changes the concentration by a tenfold. For example, a pH of 3 is ten times more acidic than a pH of 4. Likewise, a pH of 3 is one hundred times more acidic than a pH of 5. Similarly a pH of 11 is ten times more basic than a pH of 10.
Explain why HF is NOT a strong acid and yet the closely related HCl is a strong acid?
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.
Why is HF not a strong acid?
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.
Talk in details about the concept of acids strength.
The strength of an acid refers to its ability or tendency to lose a proton (H+). A strong acid is one that completely ionizes (dissociates) in a solution. In water, one mole of a strong acid HA dissolves yielding one mole of H+ (as hydronium ion H3O+) and one mole of the conjugate base, A−. Essentially none of the non-ionized acid HA remains. In contrast, a weak acid only partially dissociates. Stronger acids have a larger acid dissociation constant (Ka) and a smaller logarithmic constant (pKa = - log Ka) than weaker acids. The stronger an acid is, the more easily it loses a proton, H+. Two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths also depend on the stability of the conjugate base.When all other factors are kept constant, acids become stronger as the X--H bond becomes more polar. The second-row nonmetal hydrides, for example, become more acidic as the difference between the electronegativity of the X and H atoms increases. HF is the strongest of these four acids, and CH4 is one of the weakest Brnsted acids known.At first glance, we might expect that HF, HCl, HBr, and HI would become weaker acids as we go down this column of the periodic table because the X-H bond becomes less polar. Experimentally, we find the opposite trend. These acids actually become stronger as we go down this column.This occurs because the size of the X atom influences the acidity of the X-H bond. Acids become stronger as the X-H bond becomes weaker, and bonds generally become weaker as the atoms get larger as shown in the figure below
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 terminology used in the question infers that the strong base is in the beaker, which makes it the analyte. The base is "titrated with" the strong acid, meaning the acid is being added dropwise and is therefore the "titrant."
Talk about oxidation state.
This is the apparent charge that atoms takes on while in a molecule. The sum of the oxidation states for all of the atom in a molecule MUST equal the charge on that molecule, or equal zero if the molecule is neutral. Oxidation states have an assignment order.
Talk about the hydrolysis of salt and the term " the salt of a weak acid".
This terminology can be a bit tricky: The "salt of a weak acid" refers to the conjugate base of that weak acid combined with a cation to form a salt (see example below). HCO3- = "weak acid"; CO3(2-) = "conjugate base"; Na2CO3 = "salt of a weak acid" Similarly, the "salt of a weak base" refers to the conjugate acid of that weak base combined with an anion to form a salt (see example below). NH3 = "weak base"; NH4+ = "conjugate acid"; NH4NO3 = "salt of a weak base" When the salts of weak acids or weak bases dissolve in water one of the ions will undergo hydrolysis to re-form the weak acid or the weak base: Reforming of a weak acid: 1) Na2CO3 <--> 2Na+ + CO3(2-) 2) CO3(2-) + H2O <--> HCO3- + OH- Reforming of a weak base: 1) NH4NO3 <--> NH4+ + NO3(2-) 2) NH4+ + H2O <--> NH3 + H3O+
Talk about the titration between a strong acid and a strong base or a strong base and a strong acid.
Titration of a SA w/ SB or SB w/ SA: - Equivalence Point/Stoichiometric Point = midpoint of the nearly vertical section of the graph. At this point [titrant] = [analyte]. For example, for a solution of NaOH being titrated with HCl, at the equivalence point [HCl] = [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. Thus, for titrations involving a SA and a SB, [H+] = [OH-] at the equivalence point (this is NOT true if a WA or WB is involved, as we'll discuss below). By definition, if [H+] and [OH-] are exactly equal, pH = 7 at the equivalence point.
Talk about the titration of a weak acid and a strong base or a weak base and a strong acid.
Titration of a WA w/ SB or WB w/ SA: - Equivalence Point/Stoichiometric Point = midpoint of the nearly vertical section of the graph. Again, at this point [titrant] = [analyte]. However, because the acid and base are not both "strong" (i.e., dissociate 100% in water) the [H+] does NOT equal [OH-]. This also tells us that the pH does NOT equal 7. Notice that this is different than for SA/SB titrations. 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. --For WB w/ SA: pH < 7 --For WA w/ SB: pH > 7 --For SA w/ SB: pH = 7 (roughly) --For WA w/ WB: pH = 7 (roughly) (assuming equal strength; these titrations are rarely attempted)
Talk in details about the concept of ionization of water.
Water molecules can function as both acids and bases. One water molecule (acting as a base) can accept a hydrogen ion from a second one (acting as an acid). However, the hydroxonium ion is a very strong acid, and the hydroxide ion is a very strong base. As fast as they are formed, they react to produce water again. The net effect is that an equilibrium is set up. Kw is essentially just an equilibrium constant for the reactions shown. Like any other equilibrium constant, the value of Kw varies with temperature. Its value is usually taken to be 1.00 x 10-14 mol2 dm-6 at room temperature around 25 degree Celsius. The concentration of hydroxonium ions and hydroxide ion present in pure water turns out to be 1.00 x 10-7 mol dm-3 at room temperature. This lead to the Kw constant value if you multiple the two values.
Based on the half reaction given in the table above, the potential for Cu(s) to be reduced by on electron (i.e oxidized) is -0.52 V
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
General overview of weak acids/bases.
Weak acids and bases do not dissociate readily in solution. As a general rule, an acid with a pKa greater than zero, or a Ka less than one, can be considered a weak acid. Similarly, a base with a pKb greater than zero, or a Kb less than one, can be considered a weak base.
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.
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 ClO2 will result in reaction of ClO2- with water to form HClO2 and OH-, increasing pH; e) hydrolysis of CH3COONa will result inreaction of acetate with water to form CH3COOH and OH-, increasing pH. In summary, all of the options will increase pH except for option b).
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.
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
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!
How would you calculate pH? Talk briefly about H2O pH level.
pH = -log[H+] pH of pure H2O at 25°C = 7.0, and the [H+] = [OH-]; pH > 7 = basic, and the [H+] < [OH-]; pH < 7 = acidic, and the [H+] >[OH-]; pH = 7 is defined as "neutral"
How would you calculate pOH? And what is the equation that combine both pOH and pH?
pOH = -log[OH-] pOH + pH= 14
Talk about the relationship between the free energy and chemical energy.
ΔG° = -nFE°; where n is the number of moles of electrons transferred in the balanced redox reaction,and F is Faraday's constant. o From the above EQN we learn: + E° = negative ΔG = spontaneous reaction o Faraday's Constant = the charge on one mole of electrons.
