- Welcome to this episode of Real Chemistry.
We're going to be talking about precipitation reactions.
Precipitation reactions are just when you combine
two aqueous chemicals, that is, dissolved chemicals
and get out a solid chemical.
So, if we look here at this reaction,
we're combining lead nitrate which is aqueous,
that's what this aq means, or dissolved,
with potassium iodide, also aqueous or dissolved.
And so, they're both dissolved and that's why
up here they just look like clear solutions.
So it looks like you're adding two clear liquids together
and you get out this yellow.
Right?
So what's going on there?
Well that yellow is a solid and it turns out to be
lead iodide.
When you mix together potassium iodide and lead nitrate,
you get out a solid, lead iodide.
And that's called a precipitation reaction
because a solid is forming from two dissolved chemicals.
Our goal in this video is going to be
to predict the products of a precipitation reaction.
So we'll look at combining two chemical species
that are aqueous and we're going to try to figure out
what compounds form when they're mixed.
Is there a solid?
Now, the reason sodium chloride here
and the other chemicals in the previous slide
were listed as aqueous is because they're soluble.
So both sodium chloride and silver nitrate are soluble.
What we're really looking for is when we combine those
will we get out something that's insoluble or a solid.
You need to be comfortable using solubility tables.
If you're not sure how to use a solubility table,
check out my video on ionic compound solubility.
And that'll help you with that part.
Alright, so how do we do this?
How do we predict what the products will be
if we combine sodium chloride and silver nitrate.
Well, I've broken this process down into five steps.
The very first step we're going to do is
we're going to determine the charges on all of our ions.
Here we'll start off with sodium chloride.
We know that chlorine is minus one.
Chlorine is always a negative one charge.
And sodium is always a plus one charge.
If you're not sure how I did that,
go ahead and check out my video
on predicting the charges of ions.
We can use the periodic table and their position
on that table to know what the charges are.
Over here, we have silver nitrate and
here you basically just have to remember
that nitrate has a charge of negative one.
If you don't know that, you really can't do this problem.
So you have to know nitrate's a negative one.
It becomes pretty important to know the charges
on your polyatomic ions.
And that mean's silver is plus one.
So that's step one.
Now step two, is going to be switch the cations.
What do I mean by that?
Well, if we look over here, we have two cations.
Sodium is a cation that's positively charged,
and silver is a cation that's positively charged.
And what I'm going to do is I'm going to take sodium
and I'm going to put it where silver is
and I'm going to take silver
and I'm going to put it where sodium is.
Why am I doing that?
Well, basically, I'm testing all
of the different types of compounds that can form when
I mix those two things together.
And the only compounds that can form are still pairing
a positive charge with a negative charge.
So, for example, I couldn't have a compound
that was two anions and I couldn't have a compound
that was two cations.
So that's why this is the only switch
that makes sense to do.
We're going to move silver over there
and sodium over there
and we'll get the two compounds we could possibly form.
So what that gives us on the other side then
is silver chloride
and the charges just carry over.
So our chloride is still minus one
and our silver is still plus one.
Then, when we put sodium where silver was,
we get sodium nitrate.
And sodium is still plus one
and nitrate is still minus one.
Okay, now that we switched the cations,
we've just basically written out the products
and the products just have the same ions
that we put in with the cations flipped.
And now, what we're gonna do,
is we're gonna balance the charges.
We're gonna make sure these compounds are neutral.
In this case, you'll notice that silver is plus one
and chlorine is minus one, so that's a neutral compound.
And that sodium is plus one and that nitrate is minus one
so that's a balanced compound.
So this is sort of an easy reaction to start with
because our charges are already balanced,
so we're done with that step.
Now sometimes, let's say that you had a silver that
was plus two, then you'd have to adjust the number
of chlorines to make that neutral.
Alright, now what we're gonna do,
is we're gonna check for solubility.
That's step four.
If it's soluble, we're going to write aqueous.
If it's insoluble, we're going to write solid.
And if both of those compounds are aqueous,
we're going to write no reaction.
The reason for that is that if they're both aqueous,
we don't have a precipitation reaction.
We're not forming a solid.
So that means there's no reaction that occurs.
We just have a mixture of ions that are floating around
in a beaker.
So let's go ahead and check for solubility.
We'll start out with silver chloride.
So we'll start out with silver chloride.
And what we're going to do,
is we're going to go into our solubility tables
and we're going to look for silver or chlorine.
What we see, in the table listing soluble ions
we see chlorine right there.
So chlorine ions, typically are in soluble compounds
except when it's paired with silver in silver plus one.
So now you can see why it's important that we
determine those charges.
If we didn't know it was silver plus one,
then we couldn't tell if this was soluble or insoluble.
So, chlorine compounds typically soluble,
but not when paired with silver.
And that means that our silver chloride is insoluble
so we're going to put an s down there for solid.
A little confusing there.
Don't start to think that s means soluble.
What it means is solid.
That's the phase it's in.
Alright, so we know silver chloride is a solid.
Already we know this is a precipitation reaction.
Our real last question is, do we form two solids or one?
So what's the phase on sodium nitrate?
Here, we're going to go ahead and look
for sodium or nitrate in our tables.
And right at the top of this table on the right hand side
we see nitrate and there's no exceptions.
So anything containing nitrate is soluble.
And that means that we're going to write aqueous down here.
So that's it.
We have a precipitation reaction in this case.
We're forming solid silver chloride
and aqueous sodium nitrate.
The last step, which is not necessary in this case
is balance the chemical reaction.
You'll notice that we have the same number of silvers
chlorines, sodiums and nitrates on both sides
so this reaction is already balanced.
So let's take a look at another example.
Alright, in this case, we have sodium hydroxide
mixing with copper bromide.
And we're just going to go through these same steps.
The first thing we're going to do, is
we're going to determine the charges and
here again, you need to know your polyatomic ions.
Hydroxide is one minus.
That's this whole guy here.
That's hydroxide.
And that means that our sodium is plus one.
Our bromine is minus one.
We know that from the periodic table and
we have two of them.
So since there's two bromines at negative one each,
that gives us a total negative charge of negative two.
And that means that we need a plus two for our copper.
So now, we've determined the charges.
Now the next step is to switch those cations.
So we're going to take sodium and put it where copper is
and we're going to take copper and put it where sodium is.
So now we're going to do step two.
So now we've switched those cations,
we're just going to write out our products
using these same ions but flipped.
What cation is written where?
We would have copper hydroxide because
we're taking that copper
and we're writing it where we saw sodium.
The charges just carry over, so copper is still two plus
and hydroxide is still one minus.
I recommend writing those charges as you go along
because they're going to be necessary in just a second.
Now, what we're going to write is the sodium
where the copper was.
So NaBr.
And bromine is still minus one and sodium is still plus one.
Notice I don't write Br2.
I'm just going to not specify how many bromines I have.
The reason that is is what determines how many
bromines and how many sodiums I have is
if those charges on that ionic compound are balanced.
And that's step three.
So balance charges.
Notice that sodium and bromine have a charge
of plus one and minus one
and that means that sodium bromide is already balanced.
On the other hand, copper two plus is not balanced
with our hydroxide.
Hydroxide is just negative one
and our copper is two plus,
so what we're going to need is
we're going to need two hydroxides.
The way you can think about that
is you can think about it in terms of crossing over.
If you're not sure how to do this step
go ahead and watch Naming Ionic Compounds Part One.
We go through this process carefully
of writing balanced ionic compounds.
Basically you can think about it as crossing over
where you take this two down here
and this one down there and that becomes
the number of elements of each type that we have.
I'm going to erase those arrows.
So that means we have one copper but two hydroxides.
Whenever we have multiple polyatomic ions,
we have to use our parenthesis because that's saying
we have two whole hydroxides.
Alright, so that is step three.
Now, step four, check for solubility.
First, we're going to look at copper hydroxide.
And when we go and take a look through our tables,
we'll notice that the very top of our table
about insoluble compounds, we see hydroxide.
So hydroxide is typically insoluble.
Let's see if copper, which it's paired with,
is one of the exceptions.
And notice you don't see copper anywhere in there.
So in this case, our hydroxide
is making our compound insoluble.
So copper hydroxide is not soluble,
and that means it forms a solid.
So once again, we have a precipitation reaction.
We have copper hydroxide which forms a solid.
Now let's check out sodium bromide.
Again, we'll see sodium in our tables.
We see it right here
under the things that are typically soluble
and there's no exceptions
so that means that any sodium containing compound
is always soluble
so that means our sodium bromide is soluble
and that means we write aqueous.
So, we're done now except for balancing
the chemical reaction.
That's the last step.
Now I'm going to leave you to balance this
chemical reaction and check and make sure it's balanced.
If you have any troubles balancing this reaction,
go check out my video on balancing chemical reactions.
Alright, one more example problem.
Here, we're going to combine iron chloride
with sodium nitrate.
And we're going to follow the same exact steps.
Step one, determine the charges.
So we see here that we have chlorine.
Chlorine we know is one minus.
There's three of them, meaning we have a total
of negative three from our chlorine.
Meaning our iron must have a total of three plus
to make it neutral.
Then our nitrate, again, negative one and
our sodium plus one.
Now let's switch the cations.
So when we switch the cations, we're just switching
the place of those two things that come up front.
So we're going to put iron where we see sodium
and sodium where we see iron.
So I'm going to write the first possible product,
which is sodium chloride.
And again, I'm not going to write how many chlorines I have.
What's going to determine how many
sodium and chlorines I have there is the charges.
So I have sodium which is still one plus
and I have chlorine which is still one minus.
You can see right away that's balanced.
And now, we have to write our iron where we saw our sodium.
So I'm going to write iron nitrate.
And again, that iron charge just carries over.
We see that we have a three plus from our iron
and a one minus from our nitrate.
Now what we're going to do is balance those charges.
Step three, balance the charges.
So we can see right away that sodium chloride
has a plus one and a minus one and that's balanced.
Our iron nitrate however, not balanced.
So what we're going to need to do is
go ahead and cross over that three and that one
to create a balanced compound.
And what that's going to do is take the plus three
that's with iron and that's going to go way down
around the nitrate and it's going to give us
three nitrate groups.
So that's the sort of trick way
to get that compound balanced.
But remember, what it's really saying is that
if you have three things at negative one
and one thing at plus three, you have a balanced compound.
Alright, now last thing we're going to do.
Check for solubility.
Alright, sodium chloride, right away we see
that sodium is in our things that are typically soluble
with no exceptions.
That means that it's aqueous.
Now we're going to check out our iron nitrate.
Our iron nitrate is also aqueous.
Because nitrate is right up there
in our things that are typically soluble
and that means that this guy is also aqueous.
I kind of ran out of room but I'll write the aq there.
And if they're both aqueous,
remember, that means no solids formed.
No solids formed in this reaction,
so it's not a precipitation reaction.
So, instead of writing down these phases,
and leaving this as our final answer,
our final answer is no reaction.
We have no reaction here,
because when we combine these two sets of ions
we just get these sets of ions floating around separately.
There's no solid that's ever formed
so it's not a precipitation reaction.
Thanks for watching this episode of Real Chemistry.
We've learned about precipitation reactions,
how you can combine two aqueous chemicals
and get a solid chemical.
And we've talked about the rules you need
to predict those products.
If you have any questions, please ask them below.
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