We said at the beginning of this series that everyone is capable of being great at chemistry
using motivation and confidence.
And look where that combination has gotten us so far!
We mastered predicting products, met the mole, and calculated yields, excess, and concentrations.
And, now that we’re familiar with all of that and acid-base reactions,
we’ll dig into the details of strong and weak acids and bases, including buffers, and
talk about why they’re important.
We’ll also go through a titration, or the neutralization of an acid and base in reaction,
because it’s a great way to visually show the changes the substances experience.
Titrating acids and bases produces a cool, colorful change, but there’s far more going
on inside that Erlenmeyer flask.
We’ll get into all that detail today, in our final episode.
Hi, I’m Will Comar, and welcome to our last episode of Study Hall Chemistry, presented
by Arizona State University and Crash Course.
Let’s dive in!
Before we splash into any titrations, let’s take a minute to go over some of the essential vocabulary first.
From last episode, you probably remember acids have high H+ concentrations, and bases tend
to have high concentrations of hydroxide, or OH-.
We show those concentrations in two terms.
One is pH, which we covered last time.
It represents the concentration of hydrogen cations, or H+, in a substance.
The other is pOH, a value that represents the concentration of hydroxide anions in that same substance.
Acids and bases have both a pH and a pOH.
In fact, we usually represent the relationship between them as pOH + pH = 14.
If we know the pH, we can find the pOH, and vice versa.
The reason we can do this is because of what these scales actually look like.
The pH scale is a visual representation of the strength of acids and bases on a scale
from 0 to 14, with acids on the low end,
bases on the high end, and neutrality right in the middle, at 7.
The pOH scale is like pH’s mirror image — so when the pH is low, the pOH is high, and they
always add up to 14.
It might seem tricky at first to remember how a pH or pOH value relates to acids and bases,
but we can use two rules to help us out:
One: when acidity increases, the pH decreases.
And two: when substances become more basic, the pOH decreases.
For example, the pH of seawater is at a slightly basic 8.2.
If it becomes more acidic over time, the pH will decrease and move closer to the acid
end of the scale while its pOH increases.
And, adding soap to water makes it more basic, which increases the pH and decreases its pOH.
pH and pOH represent the full spectrum of acids and bases, which means there are both
strong and weak acids and bases.
Strong acids fully dissociate into their ions in solution, as do strong bases.
Weak acids and bases, on the other hand, don’t fully dissociate.
We defined dissociation back in episode 5 when we talked about solubility and ionic
equations, in case you need a refresher.
Strong acids and bases are more well known for playing the blockbuster roles in movies
and dramatic reactions in your average lab experiment.
But even though weak acids and bases tend to fly under the radar, they literally keep us alive!
They play a huge part in creating aqueous buffers, which are solutions with a stable pH.
In other words, they prevent pH from changing significantly when in the presence of a strong acid or base.
For example, our blood contains a buffer made from carbonic acid and a bicarbonate ion that
maintains a narrow pH range.
If we didn’t have that buffer, our blood pH could rise or fall out of that healthy
range, which could be deadly.
Buffers work in pairs, using teamwork to prevent the pH of a solution from changing drastically.
And it’s all got to do with ions.
Buffers require conjugates, which are the parts of substances that either give or receive
a positive ion in reaction.
A conjugate acid is a weak acid that donates, or loses, its H+ ion.
Ammonium is a conjugate acid because when in solution, it donates an H+ ion to become
ammonia, which is a weak base.
Acetate is the conjugate base of acetic acid, since it receives the H+ ion from the weak acid.
Conjugate reactions are reversible, so we could look at them going either direction.
It’s like a seesaw, or a tightrope walk.
If we have a lot of H+ ions, for example, ammonia would take them on to produce more ammonium.
If we have more ammonium, the reaction would donate its H+ ions back to ammonia to maintain balance.
So, when a solution contains either a weak acid and its conjugate base, or a weak base
and its conjugate acid, this is called a buffer solution,
which uses the conjugate ions to resist pH changes.
The conjugate pairs do their balancing act if something new goes into the solution and
there’s extra ions for them to deal with.
Like, if a strong acid gets added, that means we’ve increased the concentration of H+ in our reaction.
This disrupts the equilibrium, and prompts the reaction to form more of the conjugate
base in order to get us back to equilibrium and maintain the solution’s pH.
If we added a strong base instead, that would push the equilibrium the opposite direction,
prompting it to make more of the weak acid in order for balance to be restored.
Buffers are best explained with an example, so let’s talk a bit more about our blood.
The human body is super complex, but simply stated, human blood is at equilibrium between
bicarbonate and carbonic acid.
This is a reversible buffer solution that keeps the pH between 7.35 and 7.45.
The carbonic acid and water react to form bicarbonate and hydronium, which you’ll
remember we can call H+.
In this situation, carbonic acid is the weak acid and bicarbonate is the conjugate base.
Bicarbonate prevents acid from building up in the body, and carbonic acid keeps the pH
from getting too basic.
When we exercise, H+ ions accumulate in our blood, so our kidneys secrete bicarbonate
to bond to it and produce carbonic acid, which we breathe out through our lungs.
On the other hand, if we don’t have enough H+, our blood pH increases, so the body converts
carbonic acid to H+ and bicarbonate, which successfully regulates our pH.
We don’t have time here to go into all the details of these interactions, but making
a table to calculate the ICE,
or Initial, Change, and Equilibrium concentrations in a reaction is another way chemists solve
more advanced buffer problems.
We’re going to stick to the basics for today, though.
Let’s move on to strong acid and base reactions.
Last episode, we learned that when strong acids and strong bases combine and there are
no H+ or OH- ions left over,
they’re fully neutralized, since they form neutral water and a salt.
If we want to neutralize an acid or base properly, we also need to know its concentration in
moles per liter, or molarity.
But if we have an unknown concentration of the acid or base in our reaction, we have
to perform a titration.
This is where we slowly add an acid or base with a known concentration to our unknown
solution to determine how much is needed to neutralize it.
From there, we can figure out its concentration.
We add a color indicator chemical to our unknown solution that will change colors once our
solution has neutralized.
We call the point where the color changes and stays that color the endpoint.
At our equivalence point, which is exactly where our titration reaches neutralization,
the solution will have the exact same amount of H+ and OH- ions, and the pH will always be 7.0.
The equivalence point and the endpoint are not the same.
Usually, the color indicator won’t turn until after the solution’s past its equivalence point,
so it’s more precise to use equivalence points for calculating concentrations.
Let’s try a titration with hydrochloric acid and sodium hydroxide, a strong acid and
strong base, on TestTube.
We’ve got 21.00 mL of acid, which is our analyte, the substance with an unknown concentration.
Sodium hydroxide is our titrant, meaning we do know its concentration: 0.450 M.
Our titration will reveal the titration volume, or how much base is needed to neutralize the hydrochloric acid.
First, we’ll add a bit of phenolphthalein, a color indicator,
to the hydrochloric acid and place it under our buret setup.
Then, we’ll pour a bunch of sodium hydroxide into the closed buret — we need enough to
fully complete the titration and not run out too early.
The buret will measure the starting and ending volume of our base.
Next, we’ll open the buret and slowly add drops of sodium hydroxide into our acid, stirring
to let the two solutions mix properly.
At the point where our indicator starts to flash between clear and pink, we have to be
careful not to add too much base too fast.
Usually chemists will go one drop at a time, trying to be as precise as possible.
Once the acid turns pink and stays pink, this is the endpoint, and the end of the titration.
All the acid’s H+ ions have been neutralized by the OH- ions from the base, and we’re
left with some super cool pink water and sodium chloride.
Last thing: we need to subtract our final volume from our initial volume to find out
the amount of NaOH we needed to neutralize the acid.
This is our titration volume.
To get the exact amount needed to reach our equivalence point and not just the endpoint,
chemists usually perform the same titration multiple times,
averaging the titration volumes together to get a more accurate result.
Let’s assume we’ve done this, and for this experiment, we needed 47.00 mL to get
to the equivalence point of our titration.
Now that we know both the volume and molar concentration of our sodium hydroxide solution, and we
know the initial volume of our hydrochloric acid solution,
our next step is to calculate the hydrochloric acid mystery concentration using stoichiometry.
To get to our mystery concentration, we’ll perform three steps.
We’d first need to take our molarity of sodium hydroxide at 0.450 moles per liter, and multiply
by its titration volume to end up with moles.
Since our volumes are in milliliters, we’ll need to convert to liters along the way.
Gotta love the metric system, it makes this a lot easier.
Step two is converting moles of sodium hydroxide to moles of hydrochloric acid with a mole ratio.
In this case, our ratio is one to one, so we end up with 0.02115 moles of hydrochloric acid consumed in the reaction.
Our last step is to take that amount of moles and divide it by our given starting volume
of hydrochloric acid to get our molarity.
We’ll need to convert from milliliters to liters here too.
Thanks to our titration and a little bit of math, we now know that our mystery concentration is 1.01 M HCl.
Sweet!
Chemists use titrations for everything from cleaning up chemical spills in the lab to
ensuring the purity of medicines to evaluating water samples for dangerous pH levels.
Now that we know how to do one, it opens up the door to all kinds of awesome experiences wherever we are.
There’s a larger theme we’ve been pointing to throughout these episodes.
Chemistry in all its forms, from reaction types to the mole to stoichiometry, is an
incredibly helpful tool we can use to make changes — both big and small —
that shape the safety and sustainability of the world we live in.
It unlocks paths to amazing, innovative opportunities, whether that’s winning the Nobel Prize in
Chemistry, inventing brand new materials, or honing the perfect muffin recipe.
The key to getting there lies in taking our time, confidently facing challenges, and practicing.
By titrating in new material bit by bit, we can reveal the determination and drive to
conquer chemistry that’s within each of us.
We can slowly find our own balance between the unknown and our new skills, and neutralize our fears.
We don’t need phenolphthalein to tell us how awesome we are, though.
That’s been there from the start.
Thanks for watching Study Hall: Chemistry, which is produced by Arizona State University
and the Crash Course team at Complexly.
If you liked this video and want to keep learning with us here in Study Hall, be sure to subscribe.
You can learn more about ASU and the videos produced by Crash Course in the links in the description.
See you all!