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Why are alkali and alkali earth metals indicative of a strong acid? Does anyone know why it is so? Or tell me what is indicative of strong acids?

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  • $\begingroup$ Is $HCl$ is a strong acid? I don't see a alkali metal anywhere. $\endgroup$ – most venerable sir May 7 '14 at 23:55
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    $\begingroup$ Alkali and alkaline earth metals are highly electropositive, and thus the electron density is shifted over the hydroxyl group. In essence, it is the -OH that is cleaved and not just the proton. You are actually referring to bases. $\endgroup$ – Jun-Goo Kwak May 7 '14 at 23:56
  • $\begingroup$ What protons get cleaved? That is not in my textbook. $\endgroup$ – most venerable sir May 8 '14 at 0:49
  • $\begingroup$ Ok, then how do we tell if something is a strong acid? How is hydroxyl groups related to bases, not to be confused with hydroxide anions? $\endgroup$ – most venerable sir May 8 '14 at 0:52
  • $\begingroup$ Okay, we're sticking to the Arrhenius theory. Here's the executive summary: Bases and acids often have hydroxyl groups. Whether the entire hydroxyl group is cleaved and leaves as the hydroxide anion or whether a proton is cleaved and leaves as the H+ cation depends (mainly) on the electronegativity of the central atom. Less EN/more electropositive --> probably going to have its hydroxyl group cleaved to form the hydroxide anion. More EN/less electropositive --> probably going to just have its proton cleaved away. $\endgroup$ – Dissenter May 8 '14 at 16:41
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Alkali and alkali earth metals are not indicative of a strong acid; if anything, they're more indicative of a strong base, but this is one of those half-truths/Santa Claus/tooth fairy level "truths" that general chemistry textbooks tell students. Just because something has an "-OH" in it doesn't make it a base or an acid automatically. The textbook is trying to placate readers with a simple rule that is sometimes right so that when you go into the lab and see $\ce{NaOH}$ (sodium hydroxide), you recognize that's a base. $\ce{KOH}$, that's a base too. $\ce{Ca(OH)_2}$, that's a base too.

But what about $\ce{(HO)_3PO}$? What about $\ce{(HO)_2SO_2}$? The two molecules I just mentioned are phosphoric acid and sulfuric acid, respectively (not phosphoric and sulfuric base). The formulas are, in addition, written with respect to structure, to illustrate that both have $\ce{OH}$ groups.

And what about phenol? That's an organic molecule with an OH group hanging off. What about acetic acid, $\ce{H_3CCOOH}$? That has an OH group on the end hanging off (no, it has no O to O bonds). None of these are bases (at least to any significant extent).

One must realize that to be a base - an Arrhenius base (as it seems that your textbook is introducing you to - one must have

a) An electropositive element. An electropositive element would be a group one or group two element and is able to stabilize positive charge no problemo. Example: Sodium, Na. We often find sodium as sodium ion: $\ce{Na^+}$. But do we find any of the electronegative elements - say - fluorine - as $\ce{F+}$? Hell no. $\ce{F-}$, maybe. That's because those electronegative elements are better at stabilizing negative charge while electropositive elements are better at doing the opposite. Remember that chemistry is just one big free energy optimization game; in simpler terms - stability is paramount and is the underlying goal of chemical processes.

b) The electropositive element must be attached to an OH group. Through heterolytic bond clevage, the OH group will leave the electropositive group as the hydroxide ion, $\ce{HO^-}$. The heterolytic bond cleavage (in this case) leaves the electropositive element short of one electron but remember that's perfectly okay - electropositive elements are alright with a positive charge.

This happens with an Arrhenius base (water is implied)

$\ce{KOH ->K^+ + HO^-}$

This is what the Arrhenius theory might lead to to believe (which isn't true; and again, water is implied):

$\ce{(HO)_2SO_2 -> 2HO^- + SO_2^2+}$

Rather, this is the proper (first) ionization of sulfuric acid. Remember that sulfur is much more electronegative than one of the group 1 or 2 elements. This, together with the inductive effect of the attached oxygens allows the sulfur to stabilize a negative charge gained through heterolytic cleavage of the O-H bond. Note that I do not mention resonance as a factor (although textbooks might imply this as a factor) because sulfur does not generally access its d-orbitals in bonding, and hence, there is no double bond (and therefore a set of electrons to delocalize) in either sulfuric acid or the hydrogen sulfate ion. Both are best represented as charge-separated species with a +2 formal charge on the sulfur. In other molecules, however, resonance stabilization is a factor in their acidities - e.g. take the acetate ion when considering acetic acid.

$\ce{(HO)_2SO_2 + H2O -> H_3O^+ + HOSO_3^-}$

So in essence the aforementioned two factors (electropositive element and OH group) make an Arrhenius base. Unfortunately the Arrhenius definition of acids and bases is rather limited and if you study chemistry further, you will forget the Arrhenius definition and supplant it with the Bronsted-Lowry definition of acids and bases and the Lewis definition of acids and bases (at the least).

And if you are really looking for what is indicative of a strong base or acid, you just assess conjugate stability. A strong acid has a very stable conjugate base. A strong base has a very stable conjugate acid. A weak acid has a less stable conjugate base. A weak base has a less stable conjugate acid.

Example: $\ce{HCl}$ is the poster child of strong acids. Its conjugate base is $\ce{Cl^-}$, which is basically unreactive in water solution, thereby making $\ce{HCl}$ a strong acid. Remember that IF $\ce{Cl^-}$ were more reactive, then the reverse reaction (un-dissociation, if you will) will be favored, and this isn't going for $\ce{HCl}$'s cause - which is losing its hydrogen proton and keeping it lost.

$\ce{HCl + H_2O -> H_3O^+ + Cl^-}$

This reaction doesn't go backward to any significant extent; i.e.

$\ce{Cl- + H_3O^+ -> not much of a reaction}$

Therefore, $\ce{HCl}$ is a strong acid because it basically ionizes 100%; there is no reverse reaction.

Bottom line for the OP: Take everything your general chemistry textbook has to offer with a grain of salt. Right now your textbook is helping you develop a hammer to take to chemistry problems - crude, but sometimes effective. Later in your studies you will progress to finer tools - the chisel and scalpel, if you will. Here are some common fairy tales offered by general chemistry textbooks:

1) "$\ce{OH^-}$" - the negative formal charge in the hydroxide ion is mainly stabilized by the more electronegative element (as common sense would dictate, and that element would be oxygen). So write $\ce{HO^-}$ and if your teacher asks you what the hell that is, smack him or her down with the science.

2) $\ce{sp^3d^2}$ hybridization and other explanations of hypervalent molecules with d-orbital utilization. This myth started with the misreading of a chemical paper which made use of d-polarization functions (quantum mechanic stuff) as saying that hypervalent molecules access their d-orbitals on a regular basis. D-functions are not d-orbitals. Species such as sulfate ion are best represented with a +2 formal charge on the central atom. Linus Pauling was wrong about Vitamin C and he's wrong about bonding too. So is your textbook (likely, unless it was updated since 1990, the year in which a seminal paper was published which definitely refuted the use of d-orbitals in hypervalent molecules). Note that textbooks may be "updated" yearly but most of the time the only changes are 1) a new photo of a chemical molecule on the cover and 2) a new edition stamped onto the spine.

3) The existence of the $\ce{H^+}$ proton in water. No. This is Coulombically (and comically) unlikely. The high positive charge density of the bare proton means that it is hydrated by water almost immediately in water, creating the hydronium ion. So in writing reactions, write $\ce{H_3O^+}$, not $\ce{H^+}$.

4) The 90 degree bond angle. No, methane ($\ce{CH_4}$ does not have any 90 degree bond angles. 90 degree bond angles are actually relatively rare in chemistry. So no, a methane molecule does not look like a set of cardinal axes with H's stuck on the ends where North, East, South, and West are supposed to be. It's a tetrahedron; draw it like a tripod or draw it with projections.

5) $\ce{NaCl_(aq)}$ and other "aqueous" species which actually exist in their dissociated states in water. There is no unionized sodium chloride in water. Proof - put some table salt in water. Do you see anything in the water? No, all the crystals dissolved. So don't write $\ce{NaCl_(aq)}$. Your teacher may ask you to write the net ionic equation for reactions; in my opinion, this is the only correct equation one can right.

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  • $\begingroup$ Overall, this is a pretty solid answer for a beginner to acid-base theory. It's simple and precise while touching on the right topics that may be investigated further by the reader. If I can suggest a slight couple of fixes, it would probably be best to balance the charges in the equation which produces hydroxyl ions out of sulphuric acid (even if the reaction itself is unreal), and I do believe that right after describing the actual ionisation reaction, you mean to say Arrhenius acid instead of base. $\endgroup$ – Nicolau Saker Neto May 8 '14 at 2:20
  • $\begingroup$ Thank you for the feedback! Very excited to hear that my answer was decent! I have balanced the charges in the imaginary ionization of sulfuric acid. About Arrhenius acid - I did mean to say Arrhenius base - in a reference to the two things I listed previously about what made an Arrhenius base. It is true however that it is confusing I mention Arrhenius base right after talking about an acid. Point taken and edit made :). I've tried to be as thorough as possible and as chemically correct and perhaps even a bit "fun" here as possible precisely because I know am speaking to a new chem student :) $\endgroup$ – Dissenter May 8 '14 at 2:39
  • $\begingroup$ Is this college/ AP level chemistry? Man, I have never been taught 30 percent of this stuff. $\endgroup$ – most venerable sir May 28 '14 at 0:43
  • $\begingroup$ What haven't you covered? I learned this in my second semester of college. $\endgroup$ – Dissenter May 28 '14 at 3:18
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Exactly, and to expand on what Nicolau mentions, we can discuss the leveling effect. A chemist (can't remember who) found that some acids - e.g. hydrochloric acid and nitric acid - have the same strength in water. How can this be? They're clearly different, yet they have the same strength in water. This is because these strong acids are "leveled" in water solvent. Take-home:

The strongest acid that can exist in any solvent is the conjugate acid of the solvent. The strongest base that can exist in any solvent is the conjugate base of the solvent.

So for water, the conjugate acid (remember that conjugate means inverse; acid means proton-losing, so inverse/opposite of proton losing is proton gaining) is the hydronium ion, which is formed precisely by tacking a proton to water (this is a Lewis acid/base or coordination chemistry reaction).

And the conjugate base of water is the hydroxide ion. So the strongest base that can exist in water solution is hydroxide ion - anything stronger will be leveled to hydroxide ion. Take the following leveling reactions as examples:

$\ce{HCl_(l) + H_2O -> H_3O^+ +Cl^-}$

Here pure, liquid hydrochloric acid is leveled to its conjugate base and hydronium ion.

$\ce{S^2- + H_2O ->HS^- + HO^-}$

Here, a strong base, sulfide ion, is being leveled to its conjugate acid and hydroxide ion.

So I agree, don't try to memorize any list of strong acids or bases - such a list is just about useless and often the list only works for situations in which water is a solvent.

Also this discussion reminds me of a good quiz question my teacher once asked me:

enter image description here

If one draws out the Lewis structure for the solvent, one finds that there is a lone pair on the nitrogen. This parallels ammonia, $\ce{H_3N}$, which also has a lone pair on the central nitrogen atom. The point was for us to draw the parallel between ammonia and liquid methylhydroxylamine. Upon realizing that, we can then realize that $\ce{H_2S}$ is a strong acid in this solvent because even in water solvent, $\ce{H_3N}$ is effectively protonated by $\ce{H_2S}$ - the two undergo a large extent reaction. One can ascertain this reaction quantitatively through calculating a complex K for the following reaction:

$\ce{H_2S + H_3N <=>H_4N^+ + HS^-}$

$K_{above~reaction}=\frac{[HS^-][H_4N^+]}{[H_2S][H_3N]}=\frac{K_a(reactant~acid)}{K_a(product~acid)}=\frac{(1.1*10^{-7})}{(5.6*10^{-10})}=196=reaction~progresses~far~to~the~right$

Or simply by comparing the $\ce{K_a}$ values of the reactant acid and the product acid. Off the top my head (I basically got the $\ce{K_a}$ values of all the acids memorized because of how much we studied B/L A/B chemistry), they are (also mentioned above in the calculation of the K for the reaction):

${K_a(H_4N^+) = 5.6 x 10^{-10}}$

${K_a(H_2S) = 1.1 x 10^{-7}}$

So as we can see, we are going from a stronger acid (not strong acid) to a weaker acid. It makes sense that this equilibrium reaction would be a large extent reaction to the right; strong to weak; waterfalls; high to low, etc. Analogize as you see fit.

OP: Let the chemistry talk to you; don't let the rules dictate everything because in chemistry, the "rules" are mostly illusions.

EDIT: And further expanding on the fact that acidity and basicity are rather fluid concepts, consider this myth: neutral water has a pH of 7.0. No. The only requirement for neutrality is $\ce{[H_3O^+]=[HO^-]}$ in a solution. Consider the following reaction:

$\ce{2H_2O + heat ->H_3O^+ +HO^-}$

The reaction is stoichiometric - i.e. hydronium and hydroxide ions are created in a 1:1 ratio.

At 25 degrees Celsius, the concentration of both ions produced in this auto-ionization reaction is ${1.0 *10^{-7}}$ M.

However, if you heat up water, the reaction goes further to the right (Le Chatlier's principle - disturb a system at equilibrium and the system will try to return to equilibrium; in this case the system will try to consume the extra heat through further auto-ionization).

So what does this mean? It means that hot water (i.e. steam) is dangerous.

Hot water is BOTH acidic and basic (i.e. corrosive) - and highly so. So don't come into contact with super-heated water/steam. Contact with skin may require a visit to your dermatologist or quite possibly your undertaker.

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I'd just like to add a small commentary to Dissenter's nice answer. Not only is it important to analyze the structure of the solute to determine its acidic or basic behavior, but also the structure of the solvent. We are very used to thinking about aqueous solutions of acids and bases, but that's just a particular case, and limiting oneself to water hinders a deeper understanding of acidity and basicity.

A classic example is to consider nitric acid, $\ce{HNO3}$ as a solute. In liquid water as a solvent, nitric acid is a strong acid, and its behavior is represented by the formal reaction:

$$\ce{HNO3 + H2O -> NO_3^- + H3O^+}$$

However, if pure acetic acid $\ce{H3CCOOH}$ is used as a solvent (also known as glacial acetic acid), nitric acid is a weak acid, behaving like so:

$$\ce{HNO3 + H3CCOOH <=> NO_3^- + H3CCOOH_2^+}$$

But the real fun is what happens if pure sulfuric acid $\ce{H2SO4}$ is used as a solvent. Here, nitric acid is a weak base! Instead of the following reaction:

$$\ce{HNO3 + H2SO4 <=> NO_3^- + H3SO4^+}$$

, what happens in reality is:

$$\ce{HNO3 + H2SO4 <=> H_2NO_3^+ + HSO4^-}$$

The take-home is just that acidity and basicity are somewhat more fluid concepts than they might at first appear from Arrhenius theory, so avoid being too strict about classifying and memorizing what substances are acids, bases or neither.

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