Ch.15

Question 1

Properties of Acids

Answer 1

– Sour taste
– Can dissolve many metals (redox reactions, react with hydrogen and dissolve)
– Turn blue litmus paper red
– Neutralize bases

• Major types of acids
– Binary acids: Have H atom attached to a electronegative non-metal.
– Oxyacids have H atom attached to O, and that O is attached to some thing electronegative mix more acidic.
– Carboxylic acids have COOH group attached to a carbon chain. (H attached to O which is attached to a carbon, which is double bonded to another O)

Question 2

Properties of Bases

Answer 2

– Bitter taste
– Slippery feel
– Turn red litmus paper blue
– Neutralize acids

• Major types of bases
– Metal hydroxides (group one metals are most common because other groups are not very soluble in water, and when we talk about acid base, chemistry we usually talk about aqueous solutions)
– contain OH- in their crystal lattice
– Salts containing other anions that react with H+ (CO32-, NH2-, etc.) (acts as a base)
– Molecular compounds that react with H+ (amines (N accepts H), etc.).

Question 3

Arrhenius definition of acids and bases

Answer 3

• Arrhenius definition:
– An acid is a substance that produces H3O+ (or H+)ions in water
• HCl(g) + H2O → H3O+(aq) + Cl-(aq)

– A base is a substance that produces OH- ions in water • NaOH(s) → Na+(aq) + OH-(aq)
• NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)

Question 4

Brønsted and Lowry definition of acids and bases

Answer 4

An acid is a proton donor

A base is a proton acceptor

– This is a more general form of the Arrhenius definition
– More accurately describes the interaction of acids and bases with water
– Not restricted to aqueous solutions

If something makes H3O out of H2O it have to donate a proton to H2O to make that happen

If something makes OH out of H2O it have to accept a proton from H2O to make that happen

Question 5

H+ vs. H3O+

Answer 5

The ionization energy of hydrogen is quite large (far higher than most metals)

Naked hydrogen is just a free proton — a tiny, very concentrated positive charge

This makes hydrogen very reactive

Thus, H will always Bond to something
— acids are not ionic salts of hydrogen, but covalent molecules _{if you have an acid in an anhydrous environment where there’s no water around, that hydrogen in the acid when mixed in water will donate the hydrogen to the water. = hydrogen is covalently bonded to the rest of the acid for a free acid. Anhydrous acids are covalent molecules that when mixed with water the ionize (H is transferred and get an anion (the conjugate base of that acid}
— aqueous H isn’t just solvated, but the extra hydrogen is covalently bonded to oxygen
— it is the more correct to write H+ (aq) as H3O(aq)
— nonetheless, it is still coming to see H+(aq)

Acid base reactions favour the strongest bond to hydrogen_{usually to make H3O because the bond between H & O is more stable than the bond between H & the acid.}
— strong acids contain weekly bonded, hydrogen atoms (readily, lost and transfer to water)
— strong bases are capable of forming strong bonds to hydrogen
_{form, strong bonds to hydrogen will grab hydrogen off of other things. If the base acts as a proton, acceptor and grabs hydrogen, that means it’s breaking the covalent bond between that hydrogen and whatever it was bound to, which is favourable when the new covalent bond to the base is stronger than the old one}

Question 6

Chem 101

Answer 6

• bonding, covalent/ionic
• Larger electronegative difference leads to ionic bonds
• Electronegativity is related to electron affinity and ionization energy

— lower ionization energy tend to have lower electronegativity
— higher ionization energy tend to have higher electronegativity

Question 7

H+(aq) = H3O+(aq)

Answer 7

From chem 101
_{Solvation —> where are you have a cation in water all of the oxygen of water molecules are next to the Cations, so you have partial negative charge on the oxygen, stabilizing the positive charge on the Cations = favorable, electrostatic interactions.}

But when you put hydrogen in water, that is not good enough. You Can’t have a free proton with a bunch of partial negative charges from the oxygen atoms surrounding it.

It won’t stabilize a free, proton enough because it’s too it’s concentrated of positive charge. —> end up with a covalent bond, being formed between the hydrogen from the acid, and the oxygen from the water.

So if writing an acid in solution, something has acted as an acid, and a proton has been donated (to water) what we really have is H3O^3, and oxygen that’s making three covalent bonds, two hydrogen atoms, and has a positive formal charge.

Question 8

Naked H+ is a free proton info

Answer 8

For H+ If you start with one proton and one electron and you take the electron away, you’re left with one proton.

Which is very small compared to any other cat ion in a chemical compound because Bill most likely have core electrons around its nucleus.

Vast majority of the volume of an atom is the electron cloud, so if you take all of the electron cloud away, do you have something that’s really tiny —> extremely concentrated positive charge, which makes hydrogen far more reactive than the other Cations, if you will encounter.

Three proton is considered ionizing radiation (rips electrons off things), therefore, will not find free hydrogen in any normal chemical situation, will only find in high energy ionization radiation type conditions.

—> so when we talk about hydrogen an ordinary solutions and gas chemistry (as supposed to plasmid chemistry), we really mean that hydrogen is still attached to something.

Because if you have three protons flying around, it’ll find a pair of electrons and stick to them.

Question 9

Neutralization reaction

Answer 9

What a solution containing H3O (solution of acid that donated a proton to H2O) is mixed with a solution containing OH^- (bass that accepted a proton from H2O), the two species will react to form H2O.

This is a neutralization reaction

H3O^+(aq) + OH^- —> 2 H2O(l)

By the Arrhenius definition, this is the one and only acid base reaction.
—H3O reacts with OH and that’s it

The brønsted and lowry definition allows us to generalize acid base reactions to many other situation.
— any situation in which hydrogen is being transferred, can be considered an acid base reaction.

Question 10

According to the bronsted and lowry definition, acid base reaction’s, always include

Answer 10

Both an acid (the proton donor (but only if there’s a proton acceptor)) and a base (the proton acceptor)

The products are also an acid and a base, as proton transfers are reversible (although sometimes K is very large)

_{equilibrium, constant (k) is very large, strongly favours products)}HCl(aq)_acid + H2O(l) _base → H3O+(aq) _{conjugate acid} + Cl-(aq) _{conjugate base}

NH3(aq) _base + H2O(l) _acid ⇌ NH4+(aq) _{conjugate acid} + OH-(aq) _{conjugate base}

Note, that water is acting as an acid in one reaction and as a base in the other.
Such substances can act as both acids and bases are said to be amphoteric

Ex: lake water, because there’s two hydrogens bonded to water, so we can donate hydrogen Atoms an act as an acid. But oxygen in water also has a couple lone pairs that could potentially be shared with another proton and act as an hydrogen acceptor (a base)

Question 11

Polyprotic Acids and Bases

Answer 11

• Some acids can only donate one proton:

• HNO3(aq)_monoprotic+ H2O(l) → H3O+(aq) + NO3-(aq)

• These are called monoprotic acids (only one proton can be donated or accepted)

• Some acids can donate multiple protons:

• H2SO4(aq) + H2O(l) → H3O+(aq) + HSO4-(aq)_{1st conjugate base of H2SO4}

• HSO4-(aq) + H2O(l) ⇌ H3O+(aq) + SO42-(aq)_{2nd conjugate base of H2SO4}

• These are called polyprotic acids
• Likewise, there are also monobasic bases and polybasic bases

Question 12

Polyprotic Acids and Bases examples

Answer 12

• For the following species write the formula of the conjugate acid and/or base:
– Note that some of these molecules may be amphoteric, and some may be polyprotic acids or polybasic bases

• H2O Conjugate acid: H3O+, Conjugate base: OH- _{H2O access an acid or as a base}
• CO3^2- Conjugate acid: HCO3-, no conjugate base _{has no hydrogen atoms and is an anion, not really going to have a way to generate hydrogen by reaction with H2O, therefore will only act as a base}
• NH3 Conjugate acid: NH4+, Conjugate base NH2- _{has hydrogen so can donate protons also has a lone pair so can except protons}
• H2PO^4- Conjugate acid: H3PO4, Conjugate base HPO42- _{amphoteric, hydrogen atoms can donate, act as an acid, has negative charge, and oxygen atoms that can accept a proton, act as a base.}

• In each reaction, identify the Bronsted-Lowry acid and base: - H2SO4(aq) + H2O(l) → HSO4-(aq) + H3O+(aq)
acid base

• HCO3-(aq) + H2O(l) ⇌ H2CO3(aq) + OH-(aq)base Acid

Question 13

Acid Strength and Ka

Answer 13

• We can classify acids as strong acids or weak acids by looking at the extent of ionization in an aqueous solution

• Strong acids acids transfer a proton to water (almost) quantitatively: _{SA —> dissolved in H2O is qualitatively ionized = Majority of acid is going to transfer protons to H2O and form conjugate base + H3O. Equilibrium constant will be very large}

• HA(aq) + H2O(l) → A-(aq) + H3O+(aq)

• In other words, the equilibrium constant for the above reaction is very large for a strong acid

• For weak acids, the reaction with water is an equilibrium with an intermediate or small K:

• HA(aq) + H2O(l) ⇌ A-(aq) + H3O+(aq)
  _{most of HA will still be there. Only a portion of it will react with H2O. When dissolving a weak acid in H2O reaction will favour reactants}

• The equilibrium constant for reaction of an acid with water is designated Ka, the acid dissociation equilibrium

Larger ka is —> stronger acid is
Smaller ka is —> weaker acid is

Question 14

Acid and Base Strength and Ka

Answer 14

The strength of an acid deals with the nature of the bond to hydrogen in the acid as compared to the bond to hydrogen in H3O+.

• The stronger an acid is, the weaker its bond to hydrogen is and the weaker its conjugate base is. (Means won’t be able to reverse rxn)
_{transfer of hydrogen from acid to H2O, more favorite. This reaction is, the stronger the acid will be.}

– Remember that if the K for a reaction is very large, the K for the reverse reaction will be very small
– Thus, the conjugate base of a strong acid is not really a base at all in aqueous solution

• The opposite applies to strong bases – these form strong bonds to hydrogen and their conjugate acids are not really acids at all in aqueous solution
_{bass, except proton from H2O forming OH- and the conjugate acid of the base. Stronger the bond to hydrogen in the base added, the more favourable that reaction will be. Strong Bond to H = strong base, weak bond to H = weak base. (Will have a harder time removing H from H2O)}

• The conjugate bases of weak acids are weak bases

• The larger the Ka, the stronger the acid is and the weaker the conjugate base is

• The smaller the Ka, the weaker the acid is and the stronger the conjugate base is

Strong acid
H+ <——> A-
- weak attraction
- Complete ionization

H+ <—> A-
- Strong attraction
- Partial ionization

_{weak acid dissociates to a limited extent in H2O, the reaction between HA and H2O to form A- and H3O+ proceeds to a limited extent, that means if start with A- (the reverse reaction) can also proceed in a limited extent. So, weak acid —> weak conjugate base. Weak base —> weak conjugate acid}

Question 15

Strong Acids

Answer 15

Need to do

Question 16

Weak acids

Answer 16

Most acids are weak acids

Weak acids do not dissociate completely in water

– in solutions of most weak acids, the majority of acid molecules are not disassociated.

HCN (aq) + H2O (l) ⇌ H3O+ (aq) + CN-(aq)

Qc= [H3O+][CN-] / [HCN]

at equilibrium Qc=Kc«1

Reacting quotient = equilibrium, constant and for a weak acid most of the time will be less than 1.

Dissociation reaction doesn’t favour products as much
— at equilibrium, you will have significant amount of products and reactants
— an acid will produce H3O+ in aqueous solution, but not all of the acid added, will undergo that reaction.

Question 17

General Trends in Acidity

Answer 17

• Higher oxidation number = stronger oxyacid
– H2SO4>H2SO3; HNO3>HNO2

• A cation is a stronger acid than a neutral molecule; neutral molecule is a stronger acid than an anion

– H3O+>H2O>OH-; NH4+>NH3>NH2-

– Trend is reversed for base strength

Question 18

General Trends in Acidity, dealing with oxyacids

Answer 18

When dealing with oxyacid, generally dealing with something where you have a Central non-metal atom with oxygen around it, and some of the oxygen have hydrogen attached to them (H or acidic protons)

Compare oxyacid

—> look at the oxidation state of the central atom.
Ex: H2SO4 (large Ka) vs. HSO4- (small ka) has more oxygen attached to the central atom, putting it in a higher oxidation state, H2SO4 is a stronger acid than HSO4-

—> look at the nature of the central atom.
— if have structures that are identical, minus the central atom, and not changing the number of hydrogen.
Ex: H3PO4 —> H3SO4. The more electronegative atom is going to lead to a stronger acid; sulfur in this case.

—> look at the charge.
— two structures are identical, but one has an additional hydrogen compared to the other.
Ex: H2O and H3O+, H3O it’s a better acid, then H2O. Has positive charge so somethings going to lose H+, it’s going to lose its positive charge. If it already has a positive charge, it’ll be the more favourable process. If starting with a negative charge, that’ll be more difficult to lose the additional positive charge.

Question 19

General Trends in Acidity, dealing with binary acids

Answer 19

—> just looking at the halogens (HCl, HBr, HF), what she’s looking at side effect, larger the central atom is the stronger the acid is (looking down the columns of the periodic table)

—> looking across the rows in the periodic table, if compare H2S to HCl, the thing that controls the acidity is the electronegativity of the central atom. Cl more electronegative than S, so Cl is a stronger acid.

More examples:
- Even though F is more electronegative than The Elements below it in a column, HF is a weaker acid than the rest of them because it’s much smaller than the rest of them and forms, stronger bond to hydrogen (applied to other groups as well)

• H2O vs H2S, H2S is a stronger acid, then H2O, because S is bigger, even though it’s less electronegative than H2O.
• Compare across a row, NH4+ to H2O, H2O is a stronger acid.
  — NH3 is a weak base in H2O, amphatric, weaker acid than NH4+.
  — NH2-, normally don’t consider as an acid, considered a strong base. (So can be said it’s a weaker acid than NH3)

Question 20

Acid Ionization Constant, Ka

Answer 20

_{SA —> at equilibrium, favours products, large numerator, small denominator = results in larger ka. Ka > 1000
WA—> at equilibrium, significant amounts of undissolved HA, and significant amount of dissociated A-, ka will be smaller, between 1 and 10^14}
Let’s derive the expression for the equilibrium constant for the ionization reaction of a weak acid:
_{true as heterogeneous equilibrium because there’s large quantities of H2O}
• HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)

• Ka= [H3O+][A−] / [HA]

– Note that, by definition, the activity of water in aqueous solution is 1

• Ka for weak, monoprotic acids can vary from ~1 _{concentration of A- and HA will be similar at equilibrium} to ~10-14_{far less of A- than HA, very little dissociation, very little excess of H3O+
range: barely dissociates to roughly 50%}

• A substance with a Ka less than 10-14 is a weaker acid than water itself, and thus will not act as a proton donor in aqueous solution. _{in other words in order for acid dissociation equilibrium to proceed to any significant extent, acid (HA in this case) has to be a stronger acid in H2O}

• There are no known “moderately strong” acids, with Ka > 1 but less than 1000 (True strong acids have Ka > 1000)
_{will encounter strong acids, where they strongly favourite products, negligible amount of HA left at equilibrium, or reaction will favour reactants where HA and A- present at equilibrium will have more HA than A- or at the most be similar amounts}

_{will not come across acids in which the majority dissociate, but not quite all of it —> weak acid with intermediate ka, we are products are favored, but there’s still significant amount of reactants around.}

Question 21

Acid ionization constants for (ka) for some monoprotic weak acids at 25°C

Answer 21

In general, for typical concentration of weak acids, are going to have more HA than A-.

The weaker the acid gets, the further you go down the table, the more HA and less A- at equilibrium.

Chlorous acid (HClO2)
Ka = 1.1x10^-2
_{related to perchloric acid, but it’s in a lower oxidation state. 2 O instead of 4 O}
HClO2 + H2O<—>H3O^+ + ClO^- _{significant amount of HClO2 Will dissociate into H3O+ and ClO2^-. But most will remain in HClO2 form for most conditions. For the association equilibrium, notice two things are being multiplied on the top(numerator), and only one thing on the bottom(denominator), so the more dilute the acid gets, the more strongly dissociation is favoured}

Nitrous acid (HNO2)
Ka = 4.6x10^-4
_{ka is smaller than the one above it, so if you have the same concentration of HNO2 and HClO2, going to have less H3O+ in HNO2 solution been in the HClO2 solution}

Hydrofluoric acid (HF)
Ka = 3.5x10^-4

Formic acid (HCHO2)
Ka = 1.8x10^-4
_{the ones with 10^-4. Are still awake or I should go down but somewhat similar because they’re all 10 to the power of -4. So only slightly less than each other} 
Benzoic acid (HC7H5O2)
Ka = 6.5x10^-5

Acetic acid (HC2H3O2)
Ka = 181x10^-5
_{ones with x10^-5. Are in the typical range for carboxylic acids, which are mediums, strength acids}

Hypochlorous acid (HClO)
Ka = 2.9x10^-8
_{removing oxygen had a significant effect because went from 10^-5 to 10^-8}

Hydrocyanic acid (HCN)
Ka = 4.9x10^-10

Phenol (HC6H5O)
Ka = 1.3x10^-10
_{ones with 10^-10. Very weak acids, only a little dissociation}

Question 22

Base Solutions

Answer 22

• Most strong bases are Group 1_{very soluble} or Group 2 _{heavy metals, slightly less soluble} metal hydroxides with the generic formulas MOH and M(OH)2, respectively

• These produce OH- directly when dissolved in water (b/c they’re ionic compounds)_{OH is a part of the crystal structure of the solid}
NaOH(aq) → Na+(aq) + OH-(aq)

• Some other strong bases and all weak bases accept a proton from water to form OH-.
B(aq) + H2O(l) ⇌ BH+(aq) + OH-(aq)
_{when deal with base equilibria, this is the reaction we are interested in. Metal-OH’s that are strong bases have a very large Kb’s (basically producing OH directly), so not calculating those but equal Librium calculations with weak bases}

_{doesn’t really happen with any metal OH because not very soluble in H2O, not soluble means OH well stay in a solid phase}

Question 23

Base Ionization Constant, Kb

Answer 23

• The equilibrium constant for the ionization reaction of a weak base is Kb

• B(aq) + H2O(l) _{activity of 1. Same concept as with acids} ⇌ BH+(aq) + OH-(aq)

• Kb= [BH+][OH−] / [B] (original base concentration left)

• Kb for weak, monoprotic bases can vary from ~1 to ~10^-14

• The smaller the value of Kb, the weaker the base

• A base with a Kb of less than 10-14 is a weaker base than H2O, and will not accept protons from water to a significant extent.

_{Larger the kb, stronger base you have & the more [BH+] & [OH-] will produce. Smaller kb, weaker base have}

Question 24

Autoionization of Water and pH

Answer 24

So what is so special about a Ka or a Kb of 10-14?

• Pure water also acts as an acid and a base with itself: autoionization (b/c it’s amphoteric)

• The equilibrium constant for this reaction is designated Kw

H2O(l) + H2O(l) ⇌ H3O+(aq) + OH-(aq) _{reaction not really favourable, equal Librium strongly favours reactants, so equilibrium constant is very small}
Kw = [H3O+][OH-] = [H+][OH-] = 1.0 x 10-14 at 25°C

• In pure water the concentration of H+ and OH- are equal so we can calculate their concentrations as the √Kw = 1.0 x 10-7 M
_{can change the temp, heat water up, favour products a little bit more because its an endo reaction, but still have a really small k}

• Even when the concentrations are not equal, the equilibrium constant is still 1.0 x 10-14, so if you know the concentration of either H3O+(aq) or OH-(aq) you can calculate the concentration of the other

• When Ka or Kb is less than ~10-14, the excess H+ or OH- from the reaction of the acid or base with water is negligible compared to the H+ and OH- already present.
_{b/c already H3O and OH in pure H2O, if I had a very weak acid or weak base to H2O and the acid or base is so weak it can’t significantly change the [] of H3O+ or OH- that’s already there, it won’t be noticeably acidic or basic in H2O. Even if theoretically you can act as a proton, donor or acceptor
hence why a ka or kb less than x10^-14 indicate somethings not significantly acidic or basic because you want to be able to change [] of H3O or OH in pure H2O}

• This also applies to very dilute solutions (< 10-6 M) of stronger acids or bases
_{also applies to very dilute solutions. Dissolve less than 10^-6 in H2O, amount of H3O and OH from the acid or base added, is going to start to become small relative to the amount of H3O+ and OH- that’s already there. And if you want to calculate [], going to need to take autoionization into account}

Question 25

Acidic and Basic Solutions

Answer 25

• When [OH-] > [H+] the solution is basic
_{I had a base to pure H2O, face, proton acceptor, accept H+ from H2O, or from H3O+. Result in [OH-] being greater than [H3O+]}

• When [H+] > [OH-] the solution is acidic
_{added acid, donate H+ to H2O, make more H3O, or remove some H-. [H3O+] greater than [OH-]}

• When [H+]=[OH-] the solution is neutral
_{[] are equal, like in pure H2O, or when we dissolve things in H2O that aren’t acids or bases}

• Because knowing one concentration enables the other to be calculated, specifying one concentration also specifies the ratio

• Kw = [OH-][H+] = 1.0x10-14
_{can use this equilibrium. Constant show us what one of those concentrations are going to be if know the other one.
— know [H3O], can calculate [OH]
— know [OH], can calculate [H3O]
= b/c Eq constant (kw), still happens even after we dissolve an acid or base in aqueous solution. Kw will still be the same, even if you change [H3O] or [OH] by adding an acid or base}

• Acidic solutions must have [H+] > 10-7 M (greater than) and [OH-] < 10-7 M (less than)

• Basic solutions must have [H+] < 10-7 M (less than) and [OH-] > 10-7 M (greater than)

• It is thus possible to specify the acidity or basicity of a solution using either [OH-] or [H+]
— b/c If know 1, know the other.

Question 26

pH and pOH

Answer 26

• pH is a convenient logarithmic scale for reporting [H+]:
pH = –log[H+]
_{remember H+ = H3O+. Higher [H+] is, the smaller the value of the negative log is going to be. —> b/c the higher the [H3O] is, the greater the log, but we’re taking the negative of the log}

• pOH is an analogous logarithmic scale for reporting [OH-]:
pOH = –log[OH-]

• Because the [H+] (and the [OH-]) in pure water is 1.0 x 10-7 M, the pH (and the pOH) of pure water (or a solution of a substance that is neither acidic nor basic) is 7
_{if have a neutral solution, will have x10^-7 mol/l (M) of H3O, and the negative log of that is 7. So ph is neutral. Can do the same thing with [OH], neutral solution x10^-7 mol/l (M) of OH, the negative log of 10^-7 is 7. So pOH = 7}

• pH + pOH = 14

• pH < 7 (pOH > 7): the solution is acidic ([H+] > 10-7 M, [OH-] < 10-7M)
_{start making solution acidic by adding an acid. Acid increases [] oh H+, results in pH getting lower. (if go from 10^-7 to 10^-6, now - log is 6 and at pH 6}

• pH = pOH = 7: the solution is neutral ([H+] = [OH-] = 10-7 M)

• pH > 7 (pOH < 7) : the solution is basic ([H+] < 10-7 M, [OH-] > 10-7M)
_{if start making solution basic by adding a base. Base increases [OH] resulting in pOH to get lower. Add base = pH increase: pOH decrease.}

_{if add a whole bunch of strong acid to pure H2O to make an aqueous solution so our [H3O] is 1 mol/L, that’ll bring pH down to zero, that means the [OH] is going to have to be 10^-14 mol/L —> pOH = 14}

_{if add 1 mol/L of a strong base to pure H2O to make an aqueous solution. Have 1 mol/L of OH, means going to have 10^-14 mol/L of H3O. pOH = 0, pH = 14}

_{typically 0 and 14 are considered the limits of PH & POH scales, but it is possible to have solutions of strong acids or strong bases that have concentrations greater than 1 mol/L. But run into solubility limit soon after that. So really, the limits are close to PH = -1 (which is what you get from HCl at approximately 10 mol/L, which is close to the solubility limit) and pH = 15 (if you have a solution that contains approximately 10 mol/L of NaOH or KOH, which is close to the solubility limit}

Question 27

Finding the pH, [H3O+] or [OH-]

Answer 27

• Strong acids and strong bases dissociate quantitatively
(Able to take assumptions that v)
• Every mole of strong acid produces 1 mole of H3O+(aq), every mole of strong base produces 1 mole of OH-(aq)

• The [H3O+] in a strong acid solution > 10-6 M is thus is equal to the total acid concentration (for a monoprotic acid)
_{because the amount of H3O produced by acid association is going to be very large relative to the amount of H3O produced by autoionization of H2O. Keep in mind as increase [H3O], in order to continue to satisfy the kw (produce of H3O and OH) the [OH] is going to go down. So adding SA two H2O result in significantly more H3O compared to amount produced by autoionization of H2O. But also suppresses autoionization H2O by adding the acid. —> will have 10^-6 mol/L of H3O and 10^-8 mol/L of OH, only one percent of total H3O is from autoionization of H2O.}

• The [OH-] in a strong base solution > 10-6 M is thus is equal to the total base concentration (for a monobasic base)
_{same assumptions, but just about all the OH comes from that association or hydrolysis of the strong base. —> use base concentration to find [OH] and pOH using kw and find pH.}

• For very dilute solutions, the [H3O+] and [OH-] resulting from autoionization of water become significant, and you must use an ICE table
_{typically will be using for weak acids and bases dissociate qualitatively (significant amounts of unassociated and associated acid. In these instances will need to use ka or in to calculate how much you half of everything}

• Weak acids and weak bases do not dissociate quantitatively

• We need to use Ka or Kb to calculate the equilibrium [H3O+] or [OH-]
– This will involve an ICE table
– We need the Ka or Kb value for the acid or base

Question 28

Review - Sig. Figs. & Logs

Answer 28

• In logarithms, only the numbers to the right of the decimal point are significant

• The number of sig. figs. in the original value (the one you’re taking the log of) determines the number of digits after the decimal point

-log 1.00x10^- = 3.000
3 sig fig. 3 digits after decimal

Anti-log -4.26(10^-4.26)_pH = 5.5x10^-5_[H3O]
2 decimal places. 2 sig figs

Question 29

pKa

Answer 29

• We can apply a p-scale to other measures as well
_{pka just like pH, more acidic solution result in lower pH, a shrunk acid is going to result in lower pka. Taking negative log of equilibrium constant, the closer that equilibrium constant get to 1, the closer of that negative log gets to zero. The smaller that equilibrium constant gets, the larger that negative log gets.}

• When applied to acid dissociation constants (Ka), we get pKa

• A higher pKa corresponds to a smaller Ka,

• Smaller pKa = stronger acid (just like lower pH = more acidic)

_{very weak acids —> pka around 14, as pka gets smaller, dealing with stronger and stronger acid}

_{strong, weak acid pka around 0-1}

_{strong acids will have negative pka’s. Because I have very large equilibrium constant much greater than one. So when take negative log of that going to get negative number.}

_{pka greater than 14 is possible, especially wouldn’t start talking about acid-base chemistry in non-aqueous solutions. But anything with a pka greater than 14 is not going to act as an acid in H2O. Ex: NH3 (ammonia) if you look at table with a pka with non aqueous conditions, you’ll see the pka for NH3 is somewhere in the 30s. But if there’s of NH3 in H2O, it’s not an acid. It’s actually a base. It’s amphoteric Aunt is a much stronger base and acid.}

Question 30

pH of Weak Acid Solutions

Answer 30

• In weak acids, the [H3O+] is not equal to [HA]
_{little more difficult because you don’t have complete the association of acid will have equilibrium constant between one and 10^-14, which means in order to determine how much H3O is in solution we need to equal Librium calculations using the equilibrium constants, such as setting up ICE tables and solving for equilibrium concentration}

• In order to find out the [H3O+] (and thus the pH) of weak acid solutions, we have to consider the equilibrium reaction between HA and H2O:

• HA(aq) + H2O(l) ⇌ A-(aq) + H3O+(aq)

You will encounter two general types of problems:

1. Given equilibrium concentrations, find Ka
2. Given Ka (or at least can find ka) and initial concentrations, find equilibrium concentrations

Question 31

Problem solving strategy

Answer 31

• Write the balanced equation and Ka expression; these tell you what to find
_{camel equation leaves to equilibrium expression = prod/react}

•Construct an ICE table , and define x as the unknown change in concentration that occurs
_{Weak acid or base in H2O. Initial conditions, have no dissociation or base hydrolysis that occurred yet, reaction proceeds in the forward direction. So can set up changes in ICE tables and make a polynomial expression for equilibrium concentration and solve the problem from there}

• We can make some assumptions that simplify the calculations:

– [H+] resulting from acid dissociation is large relative to the 10-7 M from autoionization of water _{or if dealing with a base the [OH] from Bayside trellisses is significantly larger than the [OH] from autoionization from H2O (more than 10^-7 mol/L.
so for initial H3O or OH in ICE table can just write zero because 10^-7 it’s not going to be a significant to the x you’ll get later and so we can just ignore that}
• this is almost always true
_{only exception is when you have a very, very low concentrations and that’s a rare problem}

– x (the fraction of the acid that dissociates) is small relative to initial acid concentration [HA], so [HA]-x ≈ [HA]
_{dissociation fraction is going to be small only a small amount of acid dissolved in aqueous solution is actually going to dissociate and produce conjugate base and H+. So [HA] - x it’s pretty much the same as it was before, the fraction that dissociate is negligible}
• This is often true, but there are some exceptions, particularly for larger Ka and/or dilute solutions _{more dilute solutions below 10^-2 mol/L. Strong acid, larger ka is, the larger, as a fraction of the association is going to be less, likely going to be able to make the x is small approximation, when referring to the fraction of dissociation. —> if can’t you end up with a quadratic equation.}

• Substitute the known values into the Ka expression and solve for x.

• Check that the assumptions are justified (<5%) - If not, solve with quadratic equation
_{if assumed OH or H3O from autoionization of H2O is negligible, just remember when doing calculations that either H+ or OH is significantly greater than 10^-7
compare value of X to initial concentration of HA, or value of X to initial concentration of base and make sure it’s less than 5%. If not, go back and solve with the quadratic equation}

Question 32

Percent Ionization / Percent Dissociation

Answer 32

• We can relate the concentration of [H3O+] to the initial concentration of an acid
_{is the fraction of the acid pit into aqueous soln that undergoes the dissociation rxn & produces H3O + conjugate base. So it’s just, [H3O] in acid soln / [HA] + [A-] total [] of acids}

% dissociation = [𝐻3𝑂+] / [𝐻𝐴] +[𝐴−]
_{typically given initial [] before it dissociates is the total of the amount that does dissociate and amount that doesn’t}

• Strong acids have 100% ionization

• Weak acids will have % ionization < %100
– (usually < 1%)_{most sig below 1%, so small % dissociate}

• Equilibrium [H3O+] increases with increasing acid concentration
_{increase [acid], decrease [H3O] (make sense b/c acid dissociation Eq, add reactant(acid) to it, result in increase in [] of product and decrease pH}

• Percent ionization of a weak acid decreases with increasing acid concentration
_{start off with acid soln with some amount of weak acid dissolved in it, add more WA, double amount of acid in soln([HA] + [A]), going to less than double the [H3O], [H3O] will increase when do that but not going to double, fraction of dissociation is going to decrease.}

_{H2O given activity 1 on react side, have 2 prods H3O and A, & only 1 react HA, so as increase [] numerator is squared, denominator is not. So as increase [], decrease % ionization}

_{higher [] of WA, will have higher [H3O], but lower % ionization/dissociation. Double [WA], less than double the [H3O] in the resulting soln}

• Stronger weak acids (larger Ka / lower pKa) have a higher % dissociation than weaker weak acids (smaller Ka, higher pKa) at the same concentration
_{stronger acid is, higher % ionization. Stronger acid is, larger ka or lower pka, the more rxn favs prods. (Now all WA dissociation rxns fav reactants but the larger the ka, the less strongly fav’ed reactants are, and more products going to get and H3O is a prod so stronger acid is, lower pH of soln of the same [] is going to be, b/c going to higher % dissociation and greater fraction of acid put in is going to dissociate}

_{put 2 things together, can come up with some ideas of what’s going to happen when we dissolve acid in aq soln with respect to our ability to use 5% rule and make x is small approximations.
—if have High [] of WA, % dissociation is going to be small and amount of HA is going to be larger relative to the amount of A at Eq and be able to make small x approximations and be able to say [H3O] is going to be square root of the acid’s [] times the ka.
—on the other hand dilute soln of stronger acids, then prod not going to be able to make x is small x approximations and will have to solve quadratic equation}

Question 33

Mixtures of Acids – Strong + Strong

Answer 33

• This is trivial – both acids dissociate 100%
• [H3O+] = [HA1] + [HA2]

Strongly fav prod

Question 34

Mixtures of Acids – Strong + Weak

Answer 34

• Full ionization of a substantial concentration of a strong acid will suppress ionization of the weak acid (Le Chatelier’s principle)

_{WA diss RXN}
• HA(aq) + H2O(l) ⇌ A-(aq) + H3O+(aq) Adding H3O+ is adding product
_{if also have SA and solution adding bunch of H3O to solution by its fav diss, eventually it’s adding products to WA dis equilibrium, forcing equilibrium back towards reactants. Occurs significantly to point that dissociation of WA mix on negligible contribution to the total [H3O]}

• Therefore, [H3O+] = [HA]strong for most strong acid + weak acid mixtures
_{if have relatively small amount of SA and WA is on the stronger side of WA’s. If total amount of SA is it enough to exceed the amount of H3O that the WA generates by a factor of at least 10, then can have significant amount of dissociation of WA along with the SA also contributes to the [H3O]}

• If [HA]strong is less than ~10 times the [H3O+] that would result from the weak acid alone, you will need to use an ICE table
_{but initial [HA] isn’t zero, it’s the concentration of the SA v}

• Set the initial [H3O+] = [HA]strong instead of 0, then proceed as for a weak acid alone

Question 35

Mixtures of Acids – Weak + Weak

Answer 35

• If the concentrations of the acids are similar and the Ka of the two acids is much different, we can ignore the contribution of the weaker acid
– Only necessary to solve the ICE table for the strongest acid
_{stronger of the WA is going to dominate the [H3O] production.}

• If the Ka are similar, or if the weaker acid is present in a much higher concentration than the stronger acid, we must solve the equilibrium problem for both acids simultaneously
_{if mixing, too weak acids with similar ka, or have a lot of a weaker WA and a small amount of a stronger WA}

• If ( [𝐻𝐴]_stronger / [HA]_weaker ) ( ka_stronger / ka_weaker )> ~100, you can ignore the weaker acid

Question 36

Polyprotic Acids

Answer 36

• Contain two or more ionizable protons – H2SO4, H3PO4
_{can consider a mixture of a strong acid and a weak acid that happens to be attached to the same conjugate base}

• Each proton has its own Ka value, and they ionize in successive steps
– Ka1 > Ka2 > Ka3

• Example: H2SO3(aq) – a diprotic acid (two ionizable protons)

— H2SO3(aq) + H2O(l) ⇌ HSO3-(aq) + H3O+(aq) Ka1= 1.6 x 10-2
_{conjugate base (HSO3^- still has acidic proton, therefore amphoteric}
— HSO3-(aq) + H2O(l) ⇌ SO32-(aq) + H3O+(aq) Ka2= 6.4 x 10-8
_{HSO3^- is weaker acid than H2SO3}

• In this case, the difference in Ka is large and the contribution of the second ionization to [H3O+] can be ignored (except for very dilute solutions) _{because the first ka1 large and will suppress it. (Just like with SA/WA mixtures. Association of SA inhibits the association of WA, stronger acid will increase [H3O] in aqueous solution, that [H3O] will act as added reactant to the 2nd dissociation and suppress it.
exceptions is with very dilute solutions, the actual amount of [H3O], produced by the first association is small, so it can’t inhibit the smaller dissociation to a large extent}

• When the successive Ka are closer together both equilibria must be evaluated simultaneously.
_{Ex: H2SO3, small molecule, and removing one of the hydrogens producing a negative charge in Theory close proximity of the second one. In some acids with larger molecules, might have acidic proton at one end, and another acidic proton at the other end of the molecule, so that negative charge, produced by the first association doesn’t have nearly as large of an affect on the second dissociation, so ka’s end up being much closer together}

Question 37

Polyprotic Acids

Answer 37

• For most polyprotic acids H2A:

• Ka1 is so much larger than Ka2 (> 100 times) that essentially all of the H3O+ is produced in the first ionization step

– H2A(aq) + H2O(l) ⇌ HA-(aq) + H3O+(aq)

• Thus, very little of the deprotonated acid (HA-) dissociates any further, so we can assume that [HA-]=[H3O+]
_{2nd ka is still important if you want to know the concentration of the conjugate base (A2-)}

• The concentration of the doubly deprotonated acid ([A2-]) ≈ Ka2
_{remember that ka1 has produced H3O and HA, (intermediate product, and the conjugate base of the original acid, that’s still an acid itself) so that means so long as the extent to which the second association, occurs and changes the two concentrations is very small, those to concentrations will stay about the same and will be able to make ex is small approximations and cancel them out
in other words [] of doubly deprotonated acid of A2- is going to equal the second association equilibrium constant.}

• Ka = [𝐻3𝑂^+]+ [𝐴2^−] / [HA^-] (recall that [HA-]=[H3O+])
_{A2- is the conjugate base. For ka2, HA is the single deprotonated species. The intermediate species, the one that’s already lost one proton and is the acid.}

• Exception – polyprotic carboxylic acids where the COOH groups are well separated often have very similar Ka1, Ka2, Ka3 etc.

_{COOH, relatively large molecules and relatively well separated, so negative charge at one end of the molecule doesn’t really affect the other of the molecule that much, so ka1 and ka2 aren’t that much different
significant amount of 2nd association occurring, happens when ka1 and ka2 are quite similar}

Question 38

H2SO4

Answer 38

• Sulfuric acid differs from most other polyprotic acids (as others are WAs)

• H2SO4 is a strong acid in the first dissociation step, and a weak acid in the second
_{HSO4-, larger ka, stronger WA for Ka2. Means both []’s can be important}

• HSO4- has a relatively large Ka, so in many sulfuric acid solutions (concentration of ~0.1 M or less) the second ionization is significant and must be accounted for.

• Treat it as a strong acid + weak acid mixture where the dissociation of the weak acid cannot be ignored
_{won’t always be the case. Higher [H2SO4], the more H3O there is from the first dissociation and the higher [H3O], means less the second dissociation happens}

— H2SO4(aq) + H2O(l) ⇌ HSO4-(aq) + H3O+(aq) Ka1=&raquo_space;1

— HSO4-(aq) + H2O(l) ⇌ SO42-(aq) + H3O+(aq) Ka2= 1.11 x 10-2
_{stronger WA. If have somewhat more dilute solution with less [H3O] in solution from first dissociation, see significant dissociation of the 2nd proton and significant amounts of A2- in solution}

Question 39

Concentrations of Anions of polyprotic acids

Answer 39

• We have already encountered how [A2-] = Ka2 for solutions of most diprotic acids, by substituting the [HA-] and [H3O+] from the Ka1 calculation into the Ka2 equation.
_{second dissociation (ka2) is often much smaller than the first (ka1), result in [A2-] = ka2}

• A more general solution involves solving successive ICE tables

– First solve for the concentration of all of the species after the first ionization step

– Use these values as your initial values for the second ionization step

– When you do this, you are ignoring any potential feedbacks on the earlier equilibria

– verify that this is assumption is justified (5% rule)
_{simplification of the cancellation of the two terms cancelling in ka equation, the HA- and H3O+, or assuming are equal (are valid so long as the second ionization is relatively small compared to the first) in other words, so long as ka1 and ka2 are significantly different. —> justifies ignoring the feedbacks of acid dissociations}

Question 40

Strong Bases

Answer 40

• Strong bases will dissociate completely in water:
_{often salts containing hydroxide and completely dissociate in water. Typically, for every hydroxide in the crystal lattice of a strong base, will get one hydroxide in aqueous solution}
• [HO–] = [strong base] x (# OH)

• NaOH(aq) → Na+(aq) + OH-(aq)

• Ba(OH)2(aq) → Ba2+(aq) + 2 OH-(aq)

• Most water soluble strong bases only have one OH, so [OH-] = [base]
_{other metals, conform metal hydroxides, but they’re solubility in water is very very low, so doesn’t happen}

Alkaline metal OH-
Lithium hydroxide (LiOH)
Sodium hydroxide (NaOH)
Potassium hydroxide (KOH)

Group 2 heavy metals
Strontium hydroxide [Sr(OH)2]
Calcium hydroxide [Ca(OH)2]
Barium hydroxide [Ba(OH)2]

Question 41

Weak Bases

Answer 41

• In weak bases, only a small fraction of molecules accept protons
_{reaction favours reactants for a week basis, so only a small amount of product is produced}

• most of the weak base molecules do not take protons from water
• much less than 1% ionization in water
  _{1% or less (just like for WAs}

• [HO–] &laquo_space;[weak base]
_{less than weak base []. Just like H3O for WA’s is less than the total WA []}

• Finding the pH of a weak base solution is similar to finding the pH of a weak acid

NH3 + H2O ⇌ NH4+ + OH−

Question 42

Base Ionization Constant, Kb

Answer 42

• Base strength measured by the size of the equilibrium constant when reacted with H2O

• B(aq) + H2O (l) ⇌ OH− (aq) + HB+ (aq)

• The equilibrium constant is called the base ionization constant,

— Kb – larger Kb = stronger base
– Analogous to Ka for acids

Kb = [HB+] [OH−] / [B]

_{kb closer to one, on the stronger side of WBs, produce a considerable amount of OH. Smaller kb, 10^-14, weaker WB,  produces a lot less OH}

Question 43

Some common weak bases table

Answer 43

Stronger base → higher [OH-] → larger Kb

Carbonate ion (CO3^-)
Kb = 1.8x10^-4

Methylamine (CH3NH2)
Kb = 4.4x10^-4

Ethylamine (C2H5NH2)
Kb = 5.6x10^-5

Ammonia (NH3)
Kb = 1.76x10^-5

Bicarbonate ion (HCO3^-)
Kb = 2.3x10^8
_{amphoteric —> act as acid or base. Conjugate base of carbonic acid, we could based in carbonate ion. With carbonic acid, expect first diss (1st time acts as a acid) will be stronger acid than the second time. So if start other way around with carbonate ion that’s been doubly deprotonated, except in the first proton, more favourable then accepting the second proton}

Pyridine (C5H5N)
Kb = 1.7x10^-9

Alniline (C6H5HN)
Kb = 3.9x10^-10

_{the carpenter ant, bicarbonate ions must occur with positively charged ion such as Na+ that serves to balance the charge, but does not have any part in the ionization reaction. For example, it is the bi carbonate ion that make sodium bicarbonate (NaHCO3) Basic.}

Question 44

Relationship between Ka of an Acid and Kb of Its Conjugate Base

Answer 44

• A key relationship exists between Ka of HA and Kb of A- :

Ha + H2O <-> A- + H3O+

A- + H2O <-> HA + OH-

——————————
2 H2O <-> H3O+ + OH-

_{see that all but H2O, OH, and H3O cancel out, this is autoionization of H2O}

When you add equations, you multiply the k’s.

Ka = [A-][H3O+] / [HA]

Kb = [HA] [OH-] / [A-] _{trying to show it’s the conjugate base of the acid above}

Ka x kb = [A-][H3O+] /[HA] x [A-][OH-] / [HA]

Ka x kb = [H3O+] [OH-] = kw

Kw = 1.00x10^-14

So Ka for any acid and kb for corresponding conjugate base it’s always going to be 10^-14 (kw)

Question 45

Acid–Base Properties of Salts

Answer 45

• Consider what happens when a salt containing a conjugate acid or a conjugate base, such as ammonium chloride (NH4Cl) is dissolved in water:

NH4Cl —> NH4+ + Cl-

NH4 conjugate acid of a weak base

Conj base of WA is a WB
Conj acid of WN is a WA
So many cells contain a weak acid, a weak base —> when dissolved in H2O, can produce acidic, or basic solutions.

NH4 + H2O <-> NH3 + H3O

pH will be?
Use ka to calculate [H3O], then calculate the pH.
Or find Kb NH3 can convert to in by kw/kb = ka.

Ka = [NH3] + [H3O] / [NH4] = 5.68x10^-10

Question 46

Cations as Acids

Answer 46

Three categories
_{more likely to act as assets because positive charge and positive charged things are better at giving positive charged bits away}

_{are exceptions, ex: polyprotić acid wanna have a conjugate base of polyprotić acid that amphoteric, and still an acid itself}

1. Cations that are conjugate acids of strong bases (low charge metal cations)

• Strong bases completely dissociate/ionize
• The counterion has no affinity for hydroxide and is pH-neutral
• Na+, K+, Ca2+, etc.
  _{will not affect ph of soln, act a spectator ions}

2. Cations that are conjugate acids of weak bases

• The conjugate acid of a weak base is a weak acid

• Solutions of salts containing these cations will be acidic
– NH4+ and other ammonium salts

3. Highly charged metal cations

• These also produce acidic solutions – Fe3+, Al3+
_{Cations that have a charge of 3+ or higher. Exceptions: barium (2+), magnesium (2+)}

Question 47

Metal Cations as Weak Acids

Answer 47

• All cations are hydrated in aqueous solution

• Recall that this is an ion-dipole interaction between the cation (+ charged ion) and the oxygen (- dipole of H2O) of the water.
_{it’s an electrostatic attraction
cation surrounded by H2O molecules, aunties H2O molecules, going to have their oxygen Atoms facing in towards the cation. This is why ionic compounds dissolve in water, in the first place.
chem101: ionic salt, have very high lattice energy, really hard to separate the cations and anion’s but when dissolved in water can get all these electrostatic attraction between the positive charged Cation and negative charged end of water molecule (Oxygen), or negative charged an ion the H- of the H2O molecules}

• This interaction transfers some electron density from water to the cation, generating a partial positive charge on the water molecule
_{polls electron density towards self polarizing the H2O brings closer to positive charge Cation resulting an increase dipole of H2O}

• The effect is small for singly charged cations and larger 2+ cations
_{think about electrostatic attraction is it deals with separation of charge how much charge have, and how closely the opposite charges can approach. More closely charges can approach the stronger, the electrostatic attraction.
most doubly charged cations are smaller and have more charge, but most of the time that still doesn’t matter.
singly charged Cations tend to be fairly large and only have one charge so not something have to worry about}

• alkali metal and most alkaline earth metal cations are pH neutral

• The effect is much larger for small, highly charged cations

• The partial positive charge makes it easier to water to donate a proton
_{small highly charged cation polarizes H2O, molecule enough that it is easier to break oxygen hydrogen bonds in H2O molecule}

• This gets rid of positive charge

• Thus, small and/or highly charged cations produce acidic solutions in water

Question 48

Anions as Bases or Acids

Answer 48

Three categories

1. Anions that are conjugate bases of strong acids: Cl-, Br-, ClO4-, etc.

• Strong acids completely dissociate/ionize
• The counterion has no affinity for H3O+ and is pH-neutral
  _{ka for strong acid. Very large equilibrium. Constant forward. Association reaction is very large means equilibrium. Concentration for reverse reaction is small, therefor it’s going to be ignored}

2. Anions that are conjugate bases of weak acids: F-, CH3COO-, CO32-, etc.

• The conjugate base of a weak acid is a weak base

• Solutions of salts containing these cations will be basic

3. Anions that are partially deprotonated polyprotic acids: HSO4-, H2PO4-, HCO3-, etc.
   _{lost and protons, but not all of them}
   • These materials are amphoteric

• Solutions of salts containing these cations can be acidic or basic

• Compare Ka to Kb to determine which.
_{can calculate kb from ka for conjugate acid}

Ex:

_{HSO4-, ka about 10^-2, conj acid SO4 has very large ka, which means kb for HSO4 is going to be unimportant. It’s a much stronger acid than base
H2PO4-, ka of H2 10^-7, ka for conj acid PO4 is 10^-2, ka / kw = kb. Kb will be around 10^-12, acid stronger than base, acidic soln.
HPO4-2, ka 10^-12, kb 10^-8, more basic than acidic}

Question 49

Anions as weak bases

Answer 49

• The stronger the acid, the weaker the conjugate base and vise-versa
_{conjugate will be basic in Solution and therefore won’t contribute to having a basic salt}

• Remember that partially ionized polyprotic acids are amphoteric

– Look at Ka and Kb to decide which is more important

_{if have WA and salt that contains conjugate base of WA, then, potentially for salt to be basic}

_{if mono Protić acid. Anion insult is the conjugate base for the monoprotic WA, contributes to salt being basic}

_{if it’s an intermediate ion in a poly product dissociations series, look at ka and kb and decide which one is larger and if it’ll be more acidic or basic}

_{is salt contain an ion that’s a conjugate base of something you don’t normally considered to be a acid in H2O and the Ka for it is so tiny that it doesn’t really happen in H2O, it’s going to be a strong base.}

Question 50

Acid-Base properties of salt solutions

Answer 50

• When a salt dissolves, one or both of the ions may react with water to affect the pH

• If the salt cation _{low charge metal} is the counter-ion of a strong base and the anion is the conjugate base of a strong acid, it will form a neutral solution. _{neither will significantly contribute, so will act as base or acid. pH close to 7}

• NaCl Ca(NO3)2 Kbr

• If the salt cation is the counter-ion _{still won’t significantly interact with H2O as an acid} of a strong base and the anion is the conjugate base of a weak acid, it will form a basic solution. _{anion is a WB}

• NaF Ca(C2H3O2)2 KNO2

• If the salt cation is the conjugate acid of a weak base or a small highly charged metal cation, and the anion is the conjugate base of a strong acid, it will form an acidic solution.
_{usually highly charged, polarize H2O, proton, transfer remain bound to metal, no longer free proton —> make solution acidic. Cation is a WA}

• NH4Cl FeCl3 Al(NO3)3

• If the salt consists of a cation that is a weak acid and an anion that is a weak base, or the anion is amphoteric, the overall acidity depends on the relative acid strength (Ka) or base strength (Kb) of the ion(s).
_{same process for when talked about amphoteric anions, —> look at ka of cation acting as acid and kb of anion acting as a base. —> look at their magnitudes and see which one more dominant —> ka larger than kb = acidic, kb larger than ka = basic}

• NH4CN —> NH4+ + CN-
  _{NH4 conjugate acid, cation of weak base, weakly acidic.
  CN- conjugate base, anion of weak acid, weakly basic}
• NH4+ + H2O <-> NH3 + H3O
  — kb of NH3 = 1.76x10^-5
  Ka of NH4 = kw/kb = 1.00x10^-14 / 1.76x10^-5 = 5.7x10^-10z
• CN- + H2O <-> HCN + OH
  — ka of HCN = 6.2x10^-10
  Kb of CN- = 1.00x10^-14 / 6.2x10^-10 = 1.6x10^-5

• Kb is larger than Ka, so the solution is basic. _{Slightly. CN- acting as a base will occur to greater extent then NH4 acting as an acid}

Question 51

Structure – Acidity Relationships

Answer 51

• The strength of an acid is dictated by the strength and polarity of the H-X bond
_{in order to be an acid two things have to occur in order for proton donation to occur. 1. Have to break a bond between the rest of the acid. (Part becomes a conjugate base) and H. 2. Then have to form a bond between H and the base. So to know if an acid will be strong or not, you need to look at this process, and if just looking at the acid need to look at the nature of the bond between the conjugate base and H.}

– An acid-base reaction consists of a bond formation (between the base and the H+ from the acid) and a heterolytic bond cleavage (dissociation of the acid)

– The relative stabilities of the formed and broken bonds control the position of the equilibrium

_{bond energy. Homolytic cleavage. Take a bond rip it apart, and each atom takes one electron with it and then determine how much energy you need to put into that rxn make that happen. —> table of bond energies.
this isn’t what you doing, when breaking bonds with acids and bases, one break bond, leave both electrons on the conjugate base and new electrons in the newborn to H+ are both coming from the base that’s reacting with the acid}

Two important factors that come into play 1. The strength of the bond. 2. Just how staple the conjugate base that now has an extra negative charge compared to the way it used to be.</sub>

• The weaker the H-X bond, the easier it is for the H-X bond to break so the hydrogen can bond to something else

– This makes compounds with weaker H-X bonds stronger acids

– In aqueous solution, “something else” is the oxygen atom of water, forming H3O+

• The breakage of the H-X bond leaves both electrons on X-

– The more stable X- is as an anion, the easier it is to break the H-X bond
_{can look at how polarized it is compared to the rest of the acid. More polar easier for the rest of the acid to hold onto the electrons. So more polar H-X bond is, the more electron density is pulled away from the H and towards the rest of the conjugate base, the more stable to conjugate base is going to be in the stronger. The acid is going to be.} 

– A more polar H-X bond is an indication of a more stable conjugate base

– It is often said that acidity is controlled by the stability of the conjugate base
_{less stable conjugate base is the week at the acid. More stable conjugate base is the stronger the acid.} 

• In a strong acid, breaking the H-X bond to form a new bond between H+ and H2O is strongly favoured

– Ka is very large

– H3O+ is the strongest possible acid in aqueous solution (the levelling effect)
_{even if ka a different strong acids are different, all = acidity} 
• In a weak acid, breaking the H-X bond to form a new bond between H+ and H2O is possible, but not favourable

– Ka is small
_{if you look at exactly how difficult it is to break H-X bond, can come up with some sort of measure of acidity that something is going to be.}

• If breaking the H-X bond to form a new bond between H+ and H2O is strongly disfavored, H-X is not an acid at all (at least not in water)
_{because remaining conjugate base is very unstable, transferring that proton from H-X to H3O becomes very unlikely and have something that’s not acidic}

– Ka is very small

– H3O+ produced by dissociation is insignificant compared to H3O+ produced by autoionization

Question 52

Acid Strength and Molecular Structure across the periodic table

Answer 52

• The molecular structure of an acid determines the nature of its bond to hydrogen, and thus the strength of the acid.

• In binary acids (compounds containing only hydrogen and one other nonmetal), bond strength and bond polarity both decrease down the periodic table, while bond polarity increases as you move to the right

• Across a row of the periodic table, the strength of binary HX compounds as acids increases to the right
(HF > H2O > NH3 > CH4)
_{increasing acid strength as increasing electronegativity of the other element in a binary acid}

• Increasing electronegativity of the central atom makes the negatively charged conjugate base more stable
_{increasing electronegativity of central atom means molecules get progressively more polar, pull more electron density away from H and towards the other atom, and as they do that, H get some more and more of a positive charge, and it becomes easier to dissociate that H and the materials becomes progressively more acidic}

• Bond strength remains roughly constant
_{bond, strength does increase slightly as move towards more electronegative atoms. Such as H-F bond is a bit stronger than C-H bond. But the main thing affecting is the electron negative affect.}

Question 53

Acid Strength and Molecular Structure across the columns of the periodic table

Answer 53

• Down a column of the periodic table, bond strength and bond polarity of H-X compounds both decrease

• Within a group, bond strength is more important than bond polarity

• Acid strength order is HI > HBr > HCl&raquo_space; HF H2Te > H2Se > H2S > H2O

• Despite not being very electronegative, larger atoms can more easily accommodate extra electrons

– Less electron-electron repulsion when there is more space to spread out

_{go down calm on periodic table electronegativity decreases, yet acidity increases. Ran into this before with Hydrohalides. (HF, HI etc).}

_{HF is a weak acid. F is more electronegative than any other halogen. So H-F more polarized then H-Cl, H-Br, H-I bonds. What we’re seeing here is more of a bond strength affect.}

_{bond as go down get progressively weaker. Another way to look at this is as you go down columns, get bigger atoms (size effect). Bigger Atoms don’t have as an effective orbital overlap with H.}

_{With bigger atoms also have an conjugate base that has an extra electron, an extra negative charge. So bigger the Atom, the more the extra negative charge can spread out and the less electron electron repulsion it’s going to cause}

_{so with binary acids as you go down, the periodic table acid strength increases even though the polarity of the bonds get smaller (the electronegativity of the central atom get smaller).}

_{what dealing with is size effect —> picture of the other. Adam is in the binary acid the stronger the acid gets.}

_{removing a hydrogen from H2O or HF, even if those molecules are quite polarized removing a H is difficult because bonds are stronger and leaving behind an anion that’s []ed on a relatively small Atom. Whereas with heavier members, more acidic, because easier to break bonds to H b/c bond is weaker, and remaining anion is bigger and causing less electron—electron repulsion because there’s more room for that extra negative charge to spread out}

Question 54

Strengths of Oxyacids

Answer 54

• Oxyacids are acids in which hydrogen is bonded to an oxygen, which is in turn bonded to another atom or groups of atoms

– Usually a nonmetal, but not always

• Anything that stabilizes the negative charge on the oxygen in the conjugate base increases the acid strength _{b/c strength of acid depends on stability of conj base. More stable conj base is, weaker O-H bind tend to be and easier for acid to transfer a proton and stronger acid gets}

– In other words, anything that makes the O-H bond more polar makes an oxyacid stronger

– One can also consider delocalization – spreading a charge across many atoms is favorable
_{for oxyacids, conj b will have O with lone pair and neg formal charge. Anything that helps spread neg formal charge around and pull e- density away from O, is going to help make stronger oxyacids}

• The polarity of the O-H bond in an oxyacid is determined by two factors:

1. The electronegativity of the central atom Y

-the more electronegative Y is, the more it pulls electrons away from the oxygen, which stabilizes the O- in the conjugate base _{more polarized O-H bond gets and more stable - charge is after that H has been transferred elsewhere (after diss)}

_{this is the opposite trend for binary acids. Bin acids going down Column, makes acid stronger. Tho HI is a stronger acid than HCl. However, with oxyacids, we aren’t dealing with bond between halogen and H anymore. Don’t care about it’s bond strength b/c isn’t one. All care about is EN, in oxyacids bond breaking when acid dis is always a bond between O-H}

2.The number of oxygen atoms bonded to the central atom Y _{only depends on EN, unlike binary acids where size matters}

-Oxygen is very electronegative, and is also effective at delocalizing the negative charge and stabilizing the conjugate base
_{influnces number of resonance structures CS draw that can delocalized that charge and effects the induction effects. O is a very EN atom, so I’d have central nonmetal atom and start sticking O to it, will start pulling e- density away from it and in turn pull e- density from O-}

Question 55

Carboxylic Acids

Answer 55

• Carboxylic acids are organic acids with the general formula R-COOH _{core of the group. Way functions as acid is the same process of oxyacid dissociation}

• Only the H bonded to the oxygen atom is acidic

• Carboxylic acids are generally weak acids

• Acidity is the result of delocalization of the negative charge across both oxygen atoms
_{to determine strength of carb acids, will look at the rest of the carb acid.}

_{e- withdrawing groups, result in stronger acids. e- donating groups, result in weaker acids}

_{alkene groups are e- donating, alkene are attached to carb acid group tend to make them weaker. Not having alkene group like formic acid (only H is attached to carb group) it’s a SA b/cno e- donating alkene group. Alkene group carb acids (the vast majority of COOH) are all similar in strength}

_{can look at what is attached to group, if start putting e- withdrawing things, substitute H with halogen make acid stronger. Branching alkene groups, even more e- donating, make acids weaker}

Question 56

Base Strength

Answer 56

• A base must be able to form a bond to hydrogen
_{hetreolytic cleavage. Base has to supply both e- to form the bond}

• This normally requires a base to have a lone pair of electrons, that can be converted into a bond
_{in our resulting acid
factors affecting basicity are to a large extent the opposite to factors that effect acidity}

• The more electron density on the atom with the lone pair, the more favorable bond formation is and the stronger the base is_{to share that pair of e-‘s with another atom b/c that’s going to result in removal of density}

– Increased electron density means the base can donate that electron pair more effectively (stronger base will be)

• Factors which weaken an acid (the ability to form a strong, relatively non-polar bond to hydrogen) will strengthen the conjugate base
_{e- donating groups, e- that aren’t delocalized but rather located only on that 1 atom. All the things that’d make for a WA and unstable conj base, make for a stronger base}

• In aqueous solution, the acid that will react with the base is H2O

– No base can be stronger than OH- in aqueous solution (levelling effect)
_{once have large Eq constant, the rxn already strongly favours products so doesn’t matter how much larger you make it. I’d put a stronger base in H2O, immediately rxns and produces more OH- (same concept as if H3O}

Question 57

Amines as bases

Answer 57

Many weak bases are amines
_{work as bases b/c of lone pair present on N}

-organic compounds containing nitrogen with a lone pair of electrons
_{lone pair used to form bond to H, results + charges N containing species that makes 4 bonds}

-Basic because they accept a proton from water
_{to figure out which is stronger or weaker, use the effect with acids. If look at ka’s of conj acids, will find the strongest base (that’s the weakest conj acid) of this series is going to be methylamine
the WB (strongest conj acid) is pyridine}

Question 58

Base Strength - Amines

Answer 58

• Substituent groups that are strongly electronegative (electron withdrawing) will remove electron density from the atom with the lone pair, stabilizing the lone pair and reducing its ability to from bonds

– This weakens the base, and strengthens the conjugate acid

• Substituent groups that are electron donating (low electronegativity) will add electron density to the atom with the lone pair, destabilizing the lone pair and increasing its ability to from bonds

– This strengthens the base, and weakens the conjugate acid

Question 59

The Leveling Effect

Answer 59

• Why are all strong acids equally strong in aqueous solution, when based on the trends in acid strength we can say HI is a stronger acid than HBr or HCl?

• In aqueous solutions any strong acid immediately ionized to form H3O+

• In aqueous solutions any strong base immediately dissociates to form OH- (strong bases without OH, also generate OH- in solution)

• Water exerts a leveling effect on any S.A. or S.B. by reacting with it to form products of water autoionization

• We have seen how the strongest acid we can have in aqueous solution is H3O+, and the strongest base is OH-

• So how do we know that HI really is stronger acid than HCl?
_{use a non aq soln to dissolve the acid}

• To really be able to observe the difference in strength of strong acids we need to dissolve them in a solvent that is a weaker base than water, so it accepts protons less readily.
_{so it’s then possible that not all of them will be SA’s anymore, could be a mixture of protonated Acetic acid and halide ion along with undissolved HCL, HI, etc in soln. If this happens, we can measure Eq constants and determine which has the largest Eq constant which is the strongest acid this is how we extend acid/base beh outside aq soln’s}

• What if we dissolve acids in acetic acid?

– Acetic acid is a weaker acid and behaves as the base

HCl(g) + CH3COOH(l) ⇌ Cl-(acet) + CH3COOH2+(acet)

HBr(g) + CH3COOH(l) ⇌ Br-(acet) + CH3COOH2+(acet)

HI(g) + CH3COOH(l) ⇌ I-(acet) + CH3COOH2+(acet)

_{acetic acid stronger acid than H2O but weaker base than H2O}

Question 60

Lewis Acid – Base Theory for acid

Answer 60

• We have seen how the ability of a molecule to act as a proton donor or a proton acceptor depends on electronic factors.

• In Lewis Acid-Base theory, we look at acid base reactions from the point of view of the electrons, rather than hydrogen atoms (protons)

• A Lewis Acid is any species that accepts an electron pair to form a bond – Does not require H (although H+ is an excellent Lewis acid)
_{anything that can share a pair of e-‘s/has a relatively low energy empty orbital. —> lower the E, stronger the acid.}

– Any material with a vacant low energy orbital is a Lewis acid

– The lower the energy of the vacant orbital, the stronger the Lewis acid
_{empty orbital we’re Interested in is the anti bonding orbital between H and the rest of the acid (which becomes the conj base). —> out e- in this orbital, break bonds}

– All protic (Bronstead-Lowry) acids are also Lewis acids

– Other materials like Mg2+, BF3 and even CO2 are Lewis acids
_{metal ions are good Lewis acids}

_{BF3. B/c has 3 e-‘s but 4 valence orbitals, makes relative low and empty orbital. Esp if other thing are highly EN and pulling e- density away from B, (helps lower energy of orbitals)}

• In the case of CO2, the “low energy” vacant orbital is a C-O antibonding orbital.

Question 61

Lewis Acid – Base Theory for base

Answer 61

• A Lewis Base is any species that donates an electron pair to form a bond
_{anything that has a pair of e- willing to share. Higher the E of the pair, the better the Lewis base is. B/c more favourable it going to be to share them}

– The higher the energy of the orbital containing the lone pair, the better the Lewis base

– Anything with a lone pair can be a Lewis base

– All Bronstead-Lowry bases are also Lewis bases

– Other materials that have little affinity for H+, such as Cl- or even N2, can be Lewis bases in reactions with different Lewis acids

• N2 is an extremely weak Lewis base

• A reaction between a Lewis acid an a Lewis base results in the formation of a single species with a new covalent bond, called an adduct

– When the Lewis acid is H+, this is also a Bronstead-Lowry acid base reaction.

_{possible to make new bond and not break old bond —> product called adduct}

Question 62

How Arrhenius, Brønsted-Lowry, and Lewis theories relate

Answer 62

Lewis
Acid: electron pair acceptor
Base: electron pair donor

_{instead in looking at it from the perspective of H, looking at it from the e- pair’s perspective. Once do that, H no longer essential and other kinds of rxns can be thought of as acid and base rxns?/sub>}

Which Encompasses

Brønsted-Lowry
Acid: H+ donor
Base: H+ acceptor

_{so can contain things that don’t have OH- but can still act as bases and same concept with H3O+ and acids}

Which Encompasses

Arrhenius
Acid: H3O+ donor (makes it in H2O)
Base: OH- donor (makes it in H2O)