I wrote this out for my students as a way to review what I consider to be important fundamentals of acid base chemistry.
The proton and the significance of charge
Acid base theory all has to do with just one subatomic particle, the proton--a hydrogen ion. (If we were dealing with Lewis acids and bases it would even be more generalized to the absence of electrons, but we are not--the Bronsted Lowry definition will suit us best.)
I firmly believe that charge is probably the most important thing in chemistry—the force that moves charges around is the force that governs most all chemical phenomena! So you should always pay attention to charges, and think about how they are causing the effect you are looking at.
"Hydrogen" is not the same thing as "hydrogen ion." Ions always have some sort of charge. Atoms never do. If I say "hydrogen" instead of "hydrogen ion" just walk up and slap me. I try not to, but sometimes I slip up during lecture.
Since a hydrogen atom has only one proton and one electron, (all atoms are by definition uncharged) when it loses its one electron it is only a simple subatomic particle, a proton. So, a proton and a hydrogen ion are the same thing, and used interchangeably. Proton is easier to say because it has fewer syllables.
The reason this simple little object is still called hydrogen something is because the number of protons always determines the identity of an atom. Anything with one proton is hydrogen of some sort, anything with two protons is helium, anything with three is some sort of lithium, etc. The number of electrons or neutrons does not change the identity of the thing, just its charge or its weight, respectively. So, the fact that you have just ONE proton makes it hydrogen something. The fact that it has a positive charge makes it an ion.
(If it had an extra electron, which sometimes happens, it gets a different name, hydride, and would be H-. Positive ions derived from atoms are just called “name of atom ion” Negative ions derived from atoms get an IDE ending, on the other hand. We may discuss hydride later in biochemistry, but it is not relevant to this discussion.)
A quick word about electronegativity and polar bonds because it’s so wickedly important
Since electronegativity and polarity is related to charge, it also explains much of chemistry and biology, so you have to pay attention to it in biochemistry.
All atoms are uncharged, but atoms some want electrons more than others. This is not a result of the magnitude of an atom’s charge, which is of course zero, but a distance effect. Like gravity, the electrostatic force is affected by distance. Some atoms have their ball of positive protons in their nucleus closer to the outer edge of their electron shell than others. So, this makes them attract wandering negative charge more easily. These atoms are electronegative.
Memorize that the most electronegative atoms we will regularly encounter are O, N, F, and Cl. Think of electronegative as “electron greedy” if you like. It technically means that these atoms pull on electrons through a bond to another atom. So if the other atom they are attached to is less electronegative, like C or H, they win, and get more than their fair share of electron density. This makes them partially negative, and the electron deprived atom they are attached to is partially positive.
Memorize that polar means “having a negative end and a positive end”. If you are in a polarized discussion, there are two unresolved opinions.
A polar bond has two different ends, a partially negative one and a partially positive one.
(“partially” means somewhere in between 1 and –1. It can’t be that big of a charge, because a covalent bond is two electrons SHARED between two atoms. If one atom loses one of those two electrons, it becomes an ion, and you no longer have a covalent bond there.)
A MORE electronegative atom attached to a LESS electronegative atom creates a polar bond.
Identical atoms joined together create a nonpolar bond.
Just because you have electronegative atoms does not make something polar. If electronegative atoms are attached to the same atom, like O-O, F-F, or Cl-Cl, it is a battle between identically greedy atoms, so neither wins, and neither atom has a partial charge. Electrons slosh back and forth across the bond equally. The bond is NOT polar.
Also, C and H have similar electronegativities, so though a C-H bond technically is slightly polar (with C being a bit more negative than H) it is so slight that we will call it nonpolar.
Notice that polar bonds break more easily than nonpolar ones. Why? Because both electrons in the bond can go with the more electronegative atom.
Just because you have a polar bond does not mean the whole molecule is polar.
Polarity is actually a vector, which means it has a direction associated with it. A bond is not a molecule. If you have polar bonds in a molecule, you have to look at the whole molecule, and see if it has a net negative end and a net positive end (that’s the definition of polar!) to see if it is polar.
CO2 has polar bonds but is nonpolar as a whole because it is linear. It is negative on the outside but positive in the middle. That is not the definition of polar. HCN is also linear, but one end is not like the other, so it is polar.
Water also has polar bonds, but because the molecule is bent, it has a negative end and a positive end, so it’s polar. Imagine exaggerating the bend, like a hairpin. This will put the partially negative oxygen on one end, and the partially positive hydrogens on the other.
For us, biomolecules that are entirely C and H are ALWAYS NONPOLAR. In chemistry, molecules that are entirely C and H are called hydrocarbons.
biomolecules which are MOSTly C and H are nonpolar.
Nonpolar biomolecules are greases, oils, fats, waxes...things like that. We call them LIPIDS.
Nonpolar is also called hydrophobic, or lipophilic.
The presence of more electronegative atoms like O and N attached less electronegative ones like C or H increases the odds that the molecule is polar, but you have to look at the size and shape of the molecule, too.
Polar is also called hydrophilic.
Definitions of acid and base
An acid is a proton donor.
(Think: An acid is a proton spitter-outer.)
A base is a proton acceptor.
(Think: A base is a proton scarfer-upper.)
Not all molecules with H are acids, of course. The H must be loose. It is most likely to be loose if it is attached to an atom which can support the negative electron that it leaves behind.
Again, think about charge. When an acid loses a proton, the thing leftover not only loses a H but also its charge always decreases by 1. When a base picks up a proton, not only does its formula gain a H, but its charge increases by 1.
A C-H bond is relatively hard to break, because C cannot tolerate a negative (or even a positive) charge. C is right in the center of the trend for gaining or losing electrons (electronegativity.)
An O-H bond is easier to break, because O can sustain a negative charge better than C. But alchohols are not acids, either, nor is water a very good at being an acid. It helps to further spread the abandoned electron out by resonance, on the product, as with carboxylic acids becoming carboxylates.
A base does not need to release OH-
NH3 + H2O <----> NH4+ + OH-
Ammonia makes the solution basic because it increases OH-.
N with three bonds has a pair of nonbonding e-s that is can use to grab onto a proton, SOMETIMES. Not all nitrogen containing molecules are bases. Compounds like amides have an electron withdrawing portion that makes the nonbonding pair unavailable for this activity. Of course, a nitrogen in the form of an alkylammonium ion or quaternary ammonium ion could never be a base, there would be no place to put the proton.
alkylammonium ion, actually an acid
(the Roll book calls it an “amine cation”)
NR4+ (N attached to 4 carbons)
quaternary ammonium ion, neither an acid nor a base.
Conjugate things—acids and bases always come in pairs
An acid always becomes its conjugate base, and a base always becomes it conjugate acid. Adding or removing a proton from a molecule creates a conjugate acid or its conjugate base, respectively.
It helps to think of conjugates as the product of an acid or base reacting, and also remember the definition of acid and base above.
When an acid loses a proton, it the remaining thing (without the proton) is called its conjugate base. When a base does what bases do, the thing it becomes (with a proton) is called its conjugate acid.
Theoretically, in a reverse reaction, the conjugate products would act as acids or bases, but the reaction may or may not be capable of going backwards. This depends on the strength of the reactants, which I talk about later, below.
It is the presence or lack of protons that makes things acidic or basic, not the presence of acids or bases.
Just because you have an "acid" present does not make a solution acidic. Just because you have a "base" present does not make a solution basic.
I could list many situations where acids in solution do not create an acidic solution, and the same for bases. If you want to know if a solution is acidic or basic, don’t look for acids or bases. Look for hydrogen ions or hydroxide ions.
What makes a solution acidic is the presence of lots of protons.
What makes a solution basic is the deficiency of protons, or, if you like, the presence of lots of hydroxide ions.
Protons and hydroxide ions are natural enemies; they cannot both be excessive in the same solution. Either one is high and the other low, or vice versa. In other words, they are “inversely proportional”. You can tell this by this formula:
[H+] x [OH-] = a constant, for any solution
Whenever two variables are multiplied by each other and are equal to a constant, they are inversely proportional. For the right hand side to stay the same (constant), if H+ gets big, OH- must get small, and vice versa.
Of course, H+ and OH- can be equal in concentration, too, giving a neutral solution. Notice “neutral” in chemistry means uncharged. When H+ and OH- meet, they make uncharged water.
OK, there really is not such a thing as a loose H+ in water. But the discussion is simpler if you talk in terms of H+, rather than H3O+, the hydronium ion. I prefer to do this.
A H+ actually combines with one water to make hydronium ion, H3O+. Think of hydronium, if you like, as a proton riding on the back of a water molecule.
To go back and forth between equations using just H+ and H3O+, either add or subtract water to both sides of the equation. Combine H+ and H2O on one side to make H3O+, or do the reverse, when going backwards.
HCl ---> H+ + Cl-
add water to both sides
HCl + H2O ---> H+ + H2O + Cl-
combine H+ and H2O
HCl + H2O ---> H3O+ + Cl- (done.)
What does it mean to be strong?
There is a difference between “strong” and “stronger than...”, so pay attention to that.
Strength has to do with equilibrium.
If you have a “strong” anything as a reactant, there is NO equilibrium.
Strong Acid + whatever -----> products
Strong Base + whatever -----> products
notice the arrows used above. They are NOT equilibrium arrows, which means that when the reaction is done, you have nothing but products.
It helps to remember that strong things destroy themselves completely in water. So they don’t exist in water.
You cannot have an aqueous solution of HCl, for example. It falls apart completely into H+ and Cl-. So solutions labeled “aqueous HCl” are deceptive, but I guess it’s easier to call them that than “aqueous H+ and Cl-”, right?
Most of the acids and bases we deal with in biochemistry are NOT strong. Some are stronger than others, but they are not “strong” strong.
Here are the classic strong acids. Notice they are not organic.
HCl, HBr, HI, H2SO4, HNO3
The reason their reactions cannot go backwards is because:
The stronger a thing, the weaker its conjugate.
So, the “strong” acids have totally incompetent conjugate bases. They aren’t even bases. They can’t react to get the reaction to go backwards.
Technically, Cl- is the conjugate base of HCl. But Cl- is not even a base. So it can’t pick up a proton, and the reaction cannot go backwards after HCl falls apart into H+ and Cl-.
Notice the conjugate, so-called “bases” of the strong acids (Cl-, Br-, I-, NO3-, etc.) are often “spectator ions”, they don’t bond to much, they float around in solution by themselves. They are also more often found in soluble ionic compounds, that is, salts that fall apart into ions in water. They are not very sticky.
The classic biological acids, carboxylic acids, are not strong. (unless you do funky synthetic things to them like attach electron withdrawing atoms like F to stabilize the charge of their conjugate bases.)
Things that are NOT strong engage in equilibrium reactions. This is because their conjugates are likely to have some strength, too.
A moderately strong thing can have a moderately strong conjugate. So the conjugate product can react and the reaction can go backwards, too:
weak + weak <----> weak + weak
if there were any strong things above, there could be no equilibrium.
The conjugate base of a carboxylic acid is a carboxylate, and while neither is “strong”, they can both react, so protons are continuously shuffled back and forth between the two in an equilibrium reaction.
Or, you can imagine a population of carboxylic acids, and some hang on to their protons, and others don’t. If they were “strong” strong, all of them would give up their protons.
The classic “strong” bases are all of the soluble ionic compounds which contain OH-.
NaOH, LiOH, KOH, Ca(OH)2, Ba(OH)2, Sr(OH)2
These also fall apart 100% in water. Other hydroxide containing compounds like Mg(OH) 2 (the white stuff in milk of magnesia) are insoluble and are weak bases.
Notice again their conjugate so-called “acids” are not even acids, they don’t do much of anything in solution. Na+, Li+, etc., tend to be spectator ions floating about uselessly. They are not sticky.
Biochemical bases are usually weak. This is because their conjugate acids can actually work and we will get an equilibrium reaction. Amines (bases) and their conjugate alkylammonium ions (acids) can both engage in a wimpy sort of exchange of protons, so we get equilibrium reactions with them.
A brief word about equilibrium
Most of our biochemical reactions will be equilibrium reactions, so it is important to understand equilibrium.
It can be confusing, because you have two competing processes going on (forward and reverse reactions.)
Equilibrium means the rates of both processes are equal.
It does not mean the concentrations of products and reactants are equal.
Imagine an up elevator and a down elevator in a store. There are people on the bottom floor (reactant people) and people on the top floor (product people).
Equilibrium means that, every time one person goes up, one person goes down. You could have fifty people on the bottom and ten on top, but the concentrations of reactants and products stay constant.
Adding a catalyst speeds up both elevators equally, so does not affect the equilibrium.
At equilibrium, we have both products and reactants present in the same solution simultaneously.
Since the degree of equilibrium is related to an acid’s or bases’ strength, we can quantify that strength with the equilibrium constant, which we will call Ka or Kb.
Remember the equilibrium constant? (Kc)
first you need an equilibrium, and not one of those reactions make nothing but products. Then you measure the concentrations of products and reactants.
K = products’ concentration divided by reactants’ concentration
(and each substance is also raised to the power of its coefficient in the equation.)
The important thing to remember is that, because products are on top of this equation, the bigger the K, the more products you have, compared to reactants. So the more the reaction has “gone to the right”.
If we measure the K value for an acid, we call it Ka for some reason. But its the same thing as K.
HA <-----> H+ + A-
The bigger the Ka, the more the acid does what acids do, and the more products you will have. So the bigger the Ka, the stronger the acid.
When you measure the degree to which a base does what bases do, (picks up protons) you call the constant Kb.
B- + H2O <-----> BH + OH-
Again, the bigger the Kb, the stronger the base, for the same reason.
Kb’s and Ka’s are still annoyingly small numbers though, like # x 10-5, or # x 10 -3. Notice the more negative the exponent, the smaller the number. So an acid with a Ka of 8 x 10 -5 is stronger than than acid with a Ka of 8 x 10 -9. Don’t pay attention to the number in front, it usually has little significance—just look at the magnitude or exponent.
pKa’s are most often used in biochemistry.
Because it’s so useful to just focus on the exponent, why not do that? “p-something” means take the log of something, then change its sign. So if something is a multiple of 10 this is easy.
The pKa of a Ka which is 1 x 10 -5is 5
The pKa of a Ka which is 1 x 10 -6is 6, etc.
If it is not a multiple of 10, there will not be a 1 out in front, and you need to whip out your scientific calculator or log tables.
The pKa of a Ka which is 5 x 10 -6= ? Well, somewhere near 6, but you don’t know exactly....use your calculator, and it is 5.3. Not exactly, 6, is it, but it wasn’t too far off. It is a blatant ballbark sloppy estimate to just look at the exponent and change the sign, but without a calculator at hand, what else can you do? I think it’s helpful to glance at the exponent, change the sign, and get your sloppy estimate, even though you know it will be naughty to use it officially. It gives you an idea, at least.
Since you change the sign with a p-something, you invert the scale.
So, the bigger the pKa, the weaker the acid.
There are also pKb’s which are just the negative log of a bases’ Kb. Again, the the bigger the pKb, the weaker the base.
But instead of using the Kb or pKb or a base, the Ka or pKa of its conjugate acid is more often used. So you have to do a reversal again! This is because
The stronger a base, the weaker its conjugate acid. The weaker a base, the stronger its conjugate acid. (That was discussed above.)
There is another inversely proportional equation for this concept, for an acid base conjugate pair
Ka x Kb = constant (or you could use pKa and pKb if you like)
So it can be really confusing to hear the pka of a base! They are really talking about the pKa of the conjugate acid of the base.
If the base is weaker, its conjugate acid will be stronger, and it will have a high Ka and a low pKa.
If the base is stronger, its conjugate acid will be weaker, and it will have a low Ka and a high pKa.
A word about pH
The pH is the negative log of the hydrogen ion molarity.
pH = -log [H+]
or you can think, if you like [H+] = 10 -pH
If OH- and H+ are equal, using
[H+] x [OH-] = 1 x 10 -14(this is the constant for which most of our aqueous solutions hold true)
and substituting n=[H+] = [OH-]
n 2= 1 x 10 -14so (taking the square root of both sides to solve for n) n = 1 x 10 -7. So the pH is 7, which is neutral.
Below this the solution has more H+ than OH- and it is acidic (the scale is upside down because of the sign change that “p” does)
Above this the solution has more OH- than H+ and is basic.
If you don’t want your patients to die remember that a decrease of 1 pH unit is a drastic change (times ten) of H+ concentration, because it is a log scale.
Again, without a calculator, do ballpark stuff, but don’t tell anyone your thoughts out loud, and don’t use your estimate on real patients unless you want them to die.
if pH of normal blood is around 7.4, this means [H+] = 10 -7.4M
which is somewheres between 10 -7 and 10 -8M
But only your calculator can tell you it is really (use the alog or 10x function) = 4 x 10 -8 M H+
pH does not tell you the strength of an acid or a base. Not without additional info, anyway.
I could put a drop of the strong acid HCl in a swimming pool but the pH would remain
unchanged, and you could say the same for a base.
A very concentrated weak thing has a more dramatic effect. Again, acidity and basicity has to do with the absence or excess of protons. You can’t tell an acids’ strength by its pH. You also must know its concentration.
The pH will depend on two things, the concentration of the thing, and its strength.
These are the original comments from my old site, reposted by me, just for completeness sake
that was a really helpful overview and explained a lot
Posted by: sally | December 30, 2006 at 08:50 AM
Thanks for putting that together! It was very clear and helped a lot!
Posted by: Emily | December 18, 2007 at 06:02 PM
That was great! Best explanation of acid/base relationship I've come across. Thank you!
Posted by: david | April 11, 2009 at 11:28 AM
Posted by: neely | May 10, 2009 at 06:38 PM
Posted by: jeelia | October 29, 2009 at 11:18 PM
WOW! very useful! and easy to read! as a yale student taking chemistry right now... i really appreciated this!
Posted by: shell | March 03, 2010 at 08:24 PM
great!!! loved it! i was worried about acid and base but this cleared everything up!
Posted by: ge | September 28, 2010 at 07:40 PM
Hey I really appreciate this very accessible explanation of things. I find that my biggest struggle in math and science is just understanding what all the notation means. So to have a verbal, descriptive, in-plain-American-English explanation of things is so totally helpful. Thanks very much.
Posted by: LF | October 14, 2011 at 09:30 PM
typo on strong acids it's ClO4 not ClO.right?
Posted by: me | October 28, 2011 at 02:08 PM
Acid Solvent for Inorganic clogs
So it can be really confusing to hear the pka of a base! They are really talking about the pKa of the conjugate acid of the base.
Posted by: Acid Solvent for Inorganic clogs | November 28, 2011 at 05:50 AM
Wow! This is incredibly helpful. I was almost always confused when we talked about acids and bases, but this clears up most of my confusion. Thanks!
Posted by: Taylor | January 28, 2012 at 02:52 PM
Finally an explanation of acids and bases without all the jargon! Thank you so much!
Posted by: Mayci | February 28, 2012 at 07:09 PM
Posted by: Holly Phaneuf Erskine | 05/02/2022 at 03:54 PM