The Bronsted-Lowry and Lewis Definition of Acids and Bases ...

Hello everyone, I am trying to learn how to figure out how to figure out the most acidic molecule/substance when given a choice. My post has a question but also I'd like your wisdom in validating/refining my understanding of acids/bases as they relate to organic chemistry.
To my understanding so far, Water is more acidic than all alcohols except methanol. The relevant pKa values:
CH3OH 15.54
H2O 15.74
CH3CH2OH 15.9
Tert-Butanol around 19
While I appreciate that all have similar acid strength, there are subtle differences and I would like to understand why (relating back to the structure of the molecule) that causes it to be more acidic. Can anyone explain nice and cleanly why the order of acidity (from strongest to weakest) is methanol, water, ethanol then tert-butanol?
Possible reasons that I have seen on the internet (that I either don't understand or were poorly explained) mention: inductive effect, solvation effects, something about being a better nucleophile, CH3 being an "electron releasing group", ethoxide being a stronger base than hydroxide, something about linearity
My current understanding of deciding acid strength is: The stronger acid has the more stable conjugate base, due to:

increased electronegativity e.g. HF stronger than NH3 because fluorine is more electronegative than nitrogen
larger ion size/atomic size (e.g. HI is a stronger acid than HF because the larger iodine has more room to spread out and stabilise its negative charge)
carbons hybridised with more S character (sp>sp2>sp3) since the s-orbital is closer to the nucleus and hence those electrons experience a stronger force of attraction, resulting in carbon being more electronegative and better at stabilising the conjugate base
resonance stabilisation of conjugate base (more resonance structures = more stable)
inductive effect of electron withdrawal groups (e.g. more electronegative halogen substituent can better pull electrons towards itself, resulting in a more stable conjugate base)
distance between the acidic part and electron withdrawal group (e.g. in 2 constitutional isomers, the molecule with the acidic part and electron withdrawal group closer together will be more acidic)

I am also aware that lower pKa values mean stronger acidity and that the stronger the base, the weaker its conjugate acid.
I am trying to obtain a simplified (first year uni level) conceptual understanding so as to be in the best position to make educated guesses about what is the stronger acid, ultimately to help me predict reactants/products of complicated biochemistry reactions.
Please let me know if there is anything I wrote that I completely misunderstood. My apologies for the long post!

If the pKa of carboxylic acid is 3, then what is the pKa of the resulting carboxylate anion?

Another Uninstall Stroy.

Ammonium NH+ 4 Calcium Ca2+ Iron Fe2+ and Fe3+ Magnesium Mg2+ Potassium K+ Pyridinium C 5H 5NH+ Quaternary ammonium NR+ 4, R being an alkyl group or an aryl group Sodium Na+
Acetate CH 3COO− (acetic acid) Carbonate CO2− 3 (carbonic acid) Chloride Cl− (hydrochloric acid) Citrate HOC(COO− )(CH 2COO− ) 2 (citric acid) Cyanide C≡N− (hydrocyanic acid) Fluoride F− (hydrofluoric acid) Nitrate NO− 3 (nitric acid) Nitrite NO− 2 (nitrous acid) Oxide O2− Phosphate PO3− 4 (phosphoric acid) Sulfate SO2−

4 (sulfuric acid)

A base and an acid, e.g., NH3 + HCl → NH4Cl A metal and an acid, e.g., Mg + H2SO4 → MgSO4 + H2 A metal and a non-metal, e.g., Ca + Cl2 → CaCl2 A base and an acid anhydride, e.g., 2 NaOH + Cl2O → 2 NaClO + H2O An acid and a basic anhydride, e.g., 2 HNO3 + Na2O → 2 NaNO3 + H2O Salts can also form if solutions of different salts are mixed, their ions recombine, and the new salt is insoluble and precipitates (see: solubility equilibrium), for example: {\displaystyle {\ce {{Pb(NO3)2}+Na2SO4->PbSO4(v)+2NaNO3}}} {\displaystyle {\ce {{Pb(NO3)2}+Na2SO4->PbSO4(v)+2NaNO3}}}
Color
Potassium dichromate, a bright orange salt used as a pigment Salts can appear to be clear and transparent (sodium chloride), opaque, and even metallic and lustrous (iron disulfide). In many cases, the apparent opacity or transparency are only related to the difference in size of the individual monocrystals. Since light reflects from the grain boundaries (boundaries between crystallites), larger crystals tend to be transparent, while the polycrystalline aggregates look like white powders.
Salts exist in many different colors, for example:
yellow (sodium chromate) orange (potassium dichromate) red (cobalt nitrate) mauve (cobalt chloride hexahydrate) blue (copper sulfate pentahydrate, ferric hexacyanoferrate) purple (potassium permanganate) green (nickel chloride hexahydrate) colorless (sodium chloride, magnesium sulfate heptahydrate)—may appear white when powdered or in small pieces Most minerals and inorganic pigments, as well as many synthetic organic dyes, are salts. The color of the specific salt is due to the electronic structure in the d-orbitals of transition elements or in the conjugated organic dye framework.
Taste Different salts can elicit all five basic tastes, e.g., salty (sodium chloride), sweet (lead diacetate, which will cause lead poisoning if ingested), sour (potassium bitartrate), bitter (magnesium sulfate), and umami or savory (monosodium glutamate).
Odor Salts of strong acids and strong bases ("strong salts") are non-volatile and odorless, whereas salts of either weak acids or weak bases ("weak salts") may smell like the conjugate acid (e.g., acetates like acetic acid (vinegar) and cyanides like hydrogen cyanide (almonds)) or the conjugate base (e.g., ammonium salts like ammonia) of the component ions. That slow, partial decomposition is usually accelerated by the presence of water, since hydrolysis is the other half of the reversible reaction equation of formation of weak salts.
Solubility See also: Solubility § Solubility of ionic compounds in water Many ionic compounds can be dissolved in water or other similar solvents. The exact combination of ions involved makes each compound have a unique solubility in any solvent. The solubility is dependent on how well each ion interacts with the solvent, so there are certain patterns. For example, all salts of sodium, potassium and ammonium are soluble in water, as are all nitrates and many sulfates – barium sulfate, calcium sulfate (sparingly soluble) and lead(II) sulfate are examples of exceptions. However, ions that bind tightly to each other and form highly stable lattices are less soluble, because it is harder for these structures to break apart for the compounds to dissolve. For example, most carbonate salts are not soluble in water, such as lead carbonate and barium carbonate. Some soluble carbonate salts are: sodium carbonate, potassium carbonate and ammonium carbonate.
Conductivity Solid salts do not conduct electricity. However, liquid salts do. Moreover, solutions of salts also conduct electricity.

In 10th grade I asked a question about finding the pH of a solution with multiple weak acids, and the teacher said the explanation was too complicated for me. Five years and half a chemistry degree later, I still have no idea how to do it.

There are various different cases for what happens when you put acids and/or bases in solutions. Everyone learns how to calculate the pH with pH=-log10([H+]) and then courses cover the use of ICE tables for weak acids/bases. Here are almost all the cases I can think of, sorted by difficulty:

(Easiest)

Trivial:

Strong acid + strong base in stoichiometric quantities
Weak acid + weak base in stoichiometric quantities, where acid's Ka and base's Kb have the same value

If the acid and the base both dissociate completely, or don't but have equal strength, and they are in equal amounts (after multiplying their concentrations if they have multiple H+/OH-), the resulting solution will always be neutral.
Example: 0.05M Ca(OH)2 + 0.10M HCl -> 0.05M CaCl2 => pH=pKw/2=7.00

Very easy:

Strong acid in any quantity
Strong base in any quantity

The acid/base dissociates to produce an anion/cation and leaves hydronium/hydroxide, so the pH is simply -1 times the logarithm of the acid's/base's initial concentration.
Example: 0.10M NaOH => pH=-log10(0.10)=13.00

Easy, with another step:

Strong acid + strong acid in any quantities
Strong base + strong base in any quantities
Strong acid + strong base in non-stoichiometric quantities

After you add the quantities together this is simply the same thing as the last group of cases. If you have two strong acids (or two strong bases), add their initial concentrations. If you have a strong acid and a strong base, subtract the concentration of one from the concentration of the other.
Example: 0.15M HCl + 0.85M HBr => pH=-log10(1.00)=0.00

Average:

Weak acid in any quantity
Weak base in any quantity

Requires an ICE table. The weak acid/base partially dissociates, leaving some HA/BOH and some H+/OH-. You'll have to do a calculation with the Ka/Kb to predict the final pH.
Example: 0.1M CH3O2H -> (x)M H+ + (0.1-x)M CH3O2-, pKa=4.76 => some handwaving => pH=2.88

Average, with an extra step:

Strong acid + weak base in stoichiometric quantities
Weak acid + strong base in stoichiometric quantities

Same as the previous cases, but first, invert the problem: assume the strong acid/base reacts with the weak base/acid and produces water, some irrelevant anions/cations, and the conjugate acid/base of the weak reactant. Then find the conjugate's equilibrium constant: pKa=14-pKb (or for the weak acid, pKb=14-pKa).
Example: 0.1M NaOH + 0.1M HF -> 0.1M Na+ + 0.1M F- => 0.1M F- -> (x)M HF + (0.1-x)M F-, pKb=14.00-3.20=10.80 => use ICE table => pH=11.88

Slightly harder, multi-step:

Strong acid that also has another weak dissociation
Strong base that also has another weak ionization
Strong acid + weak acid
Strong base + weak base
Strong acid + weak base in non-stoichiometric quantities
Strong base + weak acid in non-stoichiometric quantities

Assume the first H+/OH- dissociates completely and record the concentration. Then do an ICE table with those ions in the initial concentration, and the second ionization as an independent acid/base.
Example: 0.001M H2SO4 -> 0.001M H+ + 0.001M HSO4- -> (0.001+x)M H+ + (0.001-x)M HSO4- + (x)M SO42- => too lazy to calculate, queried wolframalpha => pH=2.70

Very difficult:

Weak acid + weak base, in non-stoichiometric quantities and/or with non-equivalent Ka and Kb
Weak acid + weak acid, with different Ka values
Weak base + weak base, with different Kb values

I don't think it's possible to do ICE tables for multiple different reactions simultaneously, where the products of each one influence the reactants of the other. I think you can approximate the solution with a series of small-step calculations, where you assume each reaction progresses a little bit and calculate the change in pH, then repeat the process until it settles around a fixed point. Something similar to the method of successive approximations, but for two processes running parallel to each other and whose products interact with each other. Perhaps you could find an exact solution with Markov chains?
Example: 0.1M CH3CO2H + 0.1M NH3; 0.1M HCN + 0.1M HCO2H; 0.1M CH3NH2 + 0.1M C5H5N
(I checked these in WolframAlpha and it doesn't know how to find the pH.)

Even harder, multi-step:

Weak acid with multiple dissociations
Weak base with multiple proton acceptances

You can't just use an ICE table for each dissociation/ionization separately because the products of one reaction with water change the concentrations of the reactants in the other aqueous reaction. For acids, the equilibrium for H2A + 2H2O <-> HA- + H3O+ + H2O <-> A2- + 2H3O+ could be literally anywhere between the first reactants and the last products (you don't know where without knowing pKa1 and pKa2) and furthermore, it could have a lot of H2A with very little HA- and A2- (where pKa1 and pKa2 are small), or lots of HA- and not much H2A and A2- (where pKa1 is large and pKa2 is small), or lots of A2- and not much H2A and HA- (if pKa1 and pKa2 are very large), or various other cases.
Examples: 0.1M H2CO3, pKa1=6.35, pKa2=10.33; 0.1M N2H4, pKb1=4.75, pKb2=15.08

Apparently not impossible:

Multiple weak acids/bases with multiple ionizations

Example: 0.1M H2CO3 and 0.1M N2H4 in the same solution

(Hardest)

I can think of others, but they're just combinations of the above categories. Anyway, how do you solve the last two sets of cases?

Edit: my title is wrong. I actually asked about a solution with a weak acid plus a weak base, in unequal amounts. But the two problems are essentially equivalent.

Simple Equilibrium/Organic Chem/Acid-Base Question

Hi all! Trying to solidify some orgo acid-base concepts, and want to make sure I'm understanding/explaining them well. I chose a simple example-- in this given rxn of ammonia & water:

 NH*_3_* + H*_2_O* - NH*_4_*'+' + OH*'*

NH3 is our base, H2O is our acid; NH4+ is our conjugate acid, OH- is our conjugate base. (Using Bronsted-Lowry definitions).
The Keq for this reaction would be 10^-6.3, (calculated by subtracting the pKa for acid from the pKa for the conjugate acid), which means that the reaction would favor the left side, i.e. the reactants. Correct?
The strength of an acid is determined (or can be determined) by the stability of its conjugate base-- which, if we look at the above rxn, makes perfect sense, since H2O and and NH3 are both far more stable molecules than NH4+ and OH- (namely, H2O is more stable than OH-, so, when acting as acids, the reaction should favor H2O), correct?

Is all of the above correct? I know that my language in terms of chemistry is not perfect-- "should favor" (I don't mean to anthropomorphize molecules), but, just in terms of an college-level introductory Organic Chemistry course, does my understanding of the above make sense/sound good?
Thanks for any help-- appreciated as always! Cheers :)

conjugate acid base nh3+h2o video

The strength of a conjugate acid/base varies inversely with the strength or weakness of its parent acid or base. Any acid or base is technically a conjugate acid or conjugate base also; these terms are simply used to identify species in solution (i.e acetic acid is the conjugate acid of the acetate anion, a base, while acetate is the conjugate base of acetic acid, an acid). In the following reaction, identify the conjugate acid and base: NH3 + H2O <--> NH4+ + OH- Conj. Acid - NH4+ Conj. Base - OH- In the following reaction, identify the acid and the base, using the Bronsted-Lowry definition. Conjugate base: a base that forms when an acid loses a proton. All Arrhenius acids and bases are also Bronsted-Lowry acids and bases, but the opposite isn't true. An Arrhenius acid/base must be a ... Problem 33 Medium Difficulty. In each reaction, identify the Bronsted-Lowry acid, the Bronsted-Lowry base, the conjugate acid, and the conjugate base. a. H2CO3(aq) + H2O(l) H3O+(aq) + HCO3 3. Complete the Brønsted-Lowry equilibria, label the components acid or base, and pair up the conjugate acid-base pairs: a. HSO4 + H2O . b. NH3 + H2O . c. CN + H2O . d. H + H2O . e. HClO4 + H2O . 4. Is the monohydrogenphosphate ion HPO42- amphiprotic? If so, write the formulas of its conjugate acid and its conjugate base. 5. acid, the weaker its conjugate base and vice versa ↑K a ⇔↓K b ↓ pK a ⇔↑ b Example: The K a of HF is 6.8×10-4. What is the K b of F-? K b = K w / K a = = 1.0 ×10-14/6.8-4 K b = 1.5×10-11 18.3 Relative Acid-Base Strength and the Net Direction of Reaction • An acid-base reaction proceeds to a greater extent from the stronger acid ... In order to find the conjugate base of an acid, you take away a H. In order to find the conjugate acid of a base, you add a H. In the equation: NH3 + H2O -> NH4+ + OH-. We can see that from NH3 to... well NH3 is a base that reacts with H2O to get NH4 + OH- NH3+ H2O-->NH4+ + OH- A conjugate base is the species formed when a Bronsted- Lowry base accepts a proton. NH4+ is the conjugate acid of NH3. So the conjugate acid/base designation simply refers to proton transfer, i.e. plus or minus #H^+#....and as always, charge is balanced as well as mass. Can you tell me the conjugate base of #NH_3# ? This does not occur in aqueous solution, but it is conceived to occur in liquid ammonia. คู่กรด-เบส (conjugate acid-base pairs) ... NH3 + H2O NH4 + + OH- วิธีทำา เริ่มต้น 0.1 M สมดุ ล 0.1 - x Ex 4. สารละลายเบส NH3 0.10 M มี ค่า Kb = 1.8 x 10-5 จงหาร้อยละ การแตกตัว - -ปป. -x +x +x+x +x ...

conjugate acid base nh3+h2o top

Equation for Ba(OH)2 + H2O (Barium hydroxide + Water ...

NaOH and H2O reaction to form Na+ and OH- ions in solution.It is EXOTHERMIC because more energy is RELEASED when then ions attach to water molecules than is ... In this video we will look at the equation for CH3COOH + H2O and write the products. When we add CH3COOH to H2O the CH3COOH will dissociate and break into ... In this video we will look at the equation for H2SO4 + H2O and write the products. There are two reactions we need to consider since both of the hydrogens... In this video we will look at the equation for HCl + H2O and write the products. When we add HCl to H2O the HCl will dissociate and break into H+ and Cl-. ... This chemistry video tutorial explains the concept of acids and bases using the arrhenius definition, bronsted - lowry and lewis acid base definition. It al... In this video we will describe the equation NaOH + H2O and write what happens when NaOH is dissolved in water.When NaOH is dissolved in H2O (water) it will d... In this video we will describe the equation Ba(OH)2 + H2O and write what happens when Ba(OH)2 is dissolved in water.When Ba(OH)2 is dissolved in H2O (water) ... Use Bronsted Lowry Acid/Base Theory to identify conjugate acid base pairs.More free chemistry help at www.chemistnate.com This chemistry video tutorial provides a basic introduction into acids and bases. It explains how to identify acids and bases in addition to how they react ... In this video we determine the type of chemical reaction for the equation NaOH + HCl = NaCl + H2O (Sodium hydroxide + Hydrochloric acid yields Sodium chlorid...

The Bronsted-Lowry and Lewis Definition of Acids and Bases ...

conjugate acid base nh3+h2o

conjugate acid base nh3+h2o - win

Acid base problem

[General chemistry] Brønsted–Lowry acid and bases

Does a "more stable" conjugate acid mean it has a higher pKa and thus its base is strong?

Understanding the relative acid strengths of methanol, water, ethanol and tert-butanol