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Lecture Comments (26)

1 answer

Last reply by: Professor Hovasapian
Fri Jan 13, 2017 7:11 PM

Post by Areez Khaki on December 9, 2016

Hi professor, i dont understand why you would need to balance the reaction solution with the barium chloride example, you added a 2 to KCl to "balence the equation"  what did you mean by that

1 answer

Last reply by: Professor Hovasapian
Sun Feb 22, 2015 7:45 PM

Post by Jason Smith on February 21, 2015

Is it possible for two soluble ionic compounds to mix together and create two non-soluable compounds?

1 answer

Last reply by: Professor Hovasapian
Sun Nov 2, 2014 2:48 AM

Post by David Gonzalez on November 1, 2014

Hi professor, great lecture! Always so thorough and informative in your explanations. I have one question: why does aqueous chemistry always take place in water? Can't the solvent be anything else? Thank you!

1 answer

Last reply by: Professor Hovasapian
Fri Jan 17, 2014 9:24 PM

Post by Jérémie Lessard on January 15, 2014

Hi Professor Hovasapian !

My question refers to the second example (part 26:10). If the ions do not react with each others, can we say that they are in an dynamic equilibrium ? (since there is no precipating species)

Thank you !


0 answers

Post by Professor Hovasapian on December 22, 2013

Hi Burhan.

I hope all is well with you.

The topics you mention run in a straight sequence one after the other for my course, so start with the Kinetics lesson and just move forward. The redox stuff is under Electrochemistry.

I hope that helps.

Best wishes.


1 answer

Last reply by: Professor Hovasapian
Sun Dec 22, 2013 3:51 PM

Post by Burhan Akram on December 22, 2013

Hello Professor Raffi,

I live in B.C, Canada and the grade 12 Chemistry I am studying has the following,

1) Reaction Rates
2) Equilibrium
3) Solubility
4) Acids & Bases
5) Redox Reactions

Could you please tell me which lessons I should watch to help me study efficiently rather than watching lessons which I don't require for this course?

I would really appreciate that,

Thank You


3 answers

Last reply by: Professor Hovasapian
Tue Dec 3, 2013 12:48 AM

Post by Alexis Yates on October 15, 2013

Is there a way to measure the level of dissociation of a compound mathematically of is it something that I should just try to memorize?

1 answer

Last reply by: Professor Hovasapian
Sun Sep 1, 2013 10:51 PM

Post by Stephanie Dahlström on September 1, 2013

You're a great teacher! I've had so much trouble understanding chemistry until I started watching your lectures. You make it really easy to understand. Thank you so much!

5 answers

Last reply by: Professor Hovasapian
Fri Feb 7, 2014 6:50 PM

Post by Antie Chen on April 14, 2013

What's the formula of acetic acid? I consider it should be CH3COOH, and one molecule contains 4 Hydrogen atoms.
What's the formula H2C2H3O2?

2 answers

Last reply by: Thomas Dykstra
Wed Jul 11, 2012 2:43 PM

Post by Pierre-Alexandre Leblanc on April 21, 2012

why cant we fast forward to examples on any videos

Precipitation Reactions

  • When Ionic Compound dissolve, they actually dissociate – break up into free-floating separate ions. We call such compounds soluble.
  • There are guidelines for which ions form compounds that are soluble, and which are not – this is readily available as a chart in your book or on the web.
  • When one solution of dissolved Ionic Compound is mixed with another solution of dissolved Ionic Compound, sometimes a Cation-Anion pair will be such that their combination is not soluble in water.
  • This Insoluble compound forms a solid in solution and, generally, falls to the bottom of the container – it is called a Precipitate.
  • When mixing two soluble ionic compounds, switch Cation-Anion pairs and check the solubility guidelines to decide if a precipitate forms.

Precipitation Reactions

Lecture Slides are screen-captured images of important points in the lecture. Students can download and print out these lecture slide images to do practice problems as well as take notes while watching the lecture.

  • Intro 0:00
  • Precipitation Reactions 0:53
    • Dissociation of ionic Compounds
    • Solubility Guidelines for ionic Compounds: Soluble Ionic Compounds
    • Solubility Guidelines for ionic Compounds: Insoluble ionic Compounds
    • Precipitation Reactions
    • Example 1: Mixing a Solution of BaCl₂ & K₂SO₄
    • Example 2: Mixing a Solution of Mg(NO₃)₂ & KI

Transcription: Precipitation Reactions

Hello, and welcome back to

Welcome back to AP Chemistry.0002

Today, we're going to be talking about precipitation reactions; more generally, over the next couple of lessons, we're going to be talking about aqueous solutions and solution stoichiometry.0005

We'll be talking, particularly, about precipitation reactions today, and then in the next couple of lessons, we'll be talking about acid, base, and oxidation-reduction reactions.0015

Those are the three primary reactions in chemistry; it's either going to be some kind of precipitation, some acid-base, or some oxidation-reduction.0024

These are primarily aqueous because most of the chemistry that you really are going to end up doing is going to be in some sort of a water medium.0032

That is what aqueous means--just chemistry that takes place in the water--wet chemistry, in other words.0041

Let's go ahead and get started and see what we can do.0049

The first thing I want to talk about--we talked about naming ionic compounds; ionic compounds, remember, are compounds that have a cation and an anion, and they come together in the appropriate proportions to form a neutral compound.0054

Well, something very, very interesting happens to a fair number of ionic compounds--in fact, most ionic compounds to some extent--but some more than others.0070

They actually dissolve; and they do more than just dissolve; they actually dissociate, and we're going to discuss that right now, before we get into the precipitation reactions formally.0078

OK, so let's talk about the dissociation of ionic compounds.0088

Dissociation is exactly what the word looks like: it's dis-association.0099

Let's take something like sodium chloride: when they come together, you have this extended crystal of sodium chloride: sodium chloride, sodium chloride, sodium chloride, one after the other, in this long chain--a three-dimensional crystal structure, which is what you have when you look at salt.0103

When you drop it in water, the individual molecules of sodium chloride don't just separate--the ions just completely come apart from each other, so you have chloride ions and sodium ions floating around freely in the solution, each one of those surrounded by water molecules.0117

That is what dissociation means: it means it actually completely comes apart into its constituent elements.0133

Now that I have said that, let's talk about how it actually looks we write it out formally.0141

Let's take NaCl, and I'll put S: S stands for solid--just regular salt; now, you're going to see this written a couple of ways: sometimes (which is often the fashion nowadays), they'll go ahead and just write H2O, which means they are taking solid salt and dropping it in some water; and then you see an arrow and what it is that happens.0148

I do it a little bit differently, and there is a reason why I do it differently, and the reason is I like the actual ionic compound to be what you focus on, not necessarily the water.0171

So, what I do is: I do that, and I put a little H2O on top of the arrow, which means I have dissolved it in water.0184

It could be any kind of amount; so water is irrelevant--it just happens to be the medium in which we're doing our chemistry: aqueous chemistry.0190

Solid salt dissolved in the water: what you end up getting is Na+ + Cl-.0197

Now, sodium chloride happens to be something called a strong electrolyte, and what that means is: it comes apart completely.0202

This NaCl--a crystal of NaCl--doesn't stick together, doesn't clump down at the bottom of the beaker.0210

Yes, when you first drop it in there--yes, you see it at the bottom, but when you stir it around, it dissolves: that means that the sodium and the chloride ions are leaving from the surface of the crystal, and now they're just floating around freely.0216

Once it's fully dissolved, when you look at the solution, you can't actually see the salt; it just looks like a slightly cloudy solution, but it's homogeneous.0227

That is what dissociation means: it comes apart into its constituent compounds.0237

Another example would be something like iron (2) sulfate.0243

Remember iron (2) sulfate?--two oxidation states, so we have to call it that.0249

This is a solid; when we drop it in water, we're going to get a free iron ion floating around in solution, plus SO42-; and notice this polyatomic ion--it stays together.0253

Dissociation of ionic compounds means it dissociates into its individual ions, not into its individual elements.0266

So, a polyatomic ion stays together as a group: carbonate, dichromate--these things stay together.0273

There is also something else that we write--a little (aq) that you see; this (aq) means that it is dissolved; it is in water; it's surrounded by water--solvated; (aq), (aq), to differentiate it from the solid.0278

Again, iron sulfate--iron (2) sulfate--is also a strong electrolyte, and we'll talk about solubility in a minute--the extent to which these things actually do come apart completely.0296

But, it dissociates completely; you are not going to find any iron, any sulfate, in solid form.0307

OK, now let's talk about acid; and we're going to use acetic acid, so I'm going to write H2C2H3O2 (acetic acid is just vinegar).0315

When I drop it in water, something interesting happens with acetic acid; now, acid--for those of you that are familiar with acids (if you are not, not a problem--we're going to be discussing it in the next lesson)--an acid is something that has hydrogens that it gives up.0327

So, when you drop an acid in water, the actual compound in water--it actually gives up one of its hydrogen ions, and it gives it up one at a time.0342

But what happens first is: when you drop acetic acid in water, acetic acid itself stays together as a molecule, and it actually dissolves in water.0351

So, we actually write it like this: H2C2H3O2, and we put a little (aq) here.0361

Well, this (aq) means that the whole molecule is surrounded by water, so you can't actually see the individual molecules of acetic acid anymore.0368

Then, what we do--we have a double arrow here; now that it is an aqueous solution, now it gives up one of its hydrogen ions--it dissociates into its ionic constituents: hydrogen ion plus (if you don't mind, instead of writing C2H3O2, I'm just going to write Ac--that is a common abbreviation for the acetate ion).0377

So, notice, in this particular case, what we have done here is just write the particular ionic compound dissociates into its constituents.0401

This ionic compound dissociates into its constituents; acetic acid happens to be a weak electrolyte; in other words, it doesn't dissociate completely into its hydrogen ion and acetate ion form.0409

In fact, most of it stays in this form--the acetic acid form.0421

But, it's dissolved in water; I take pure acetic acid, and drop it in water, and you can't tell the difference: when you look, it just looks like a solution of water.0425

Well, there is acetic acid in there; it is surrounded by water molecules, but it hasn't dissociated, which is why you see this double arrow.0433

Notice, up here, you see single arrows; double arrow means that it goes both ways.0441

In fact, later in the course, we'll find a way to actually quantify the extent to which this dissociates--meaning we'll attach some math to it and measure how much of this dissociates.0445

As it turns out, acetic acid is about 1 to 2% dissociated, under normal conditions.0456

Just to let you know, you will often see it like this: you will see it from here--this is usually how it's written, but I wanted you to see what it is that actually happens, chemically.0461

Pure acetic acid, dropped in water, is solvated by the water, and then it will dissociate into its constituent ions partially; weak electrolyte/strong electrolyte.0472

Strong electrolyte means completely dissociated in aqueous solution.0482

OK, now for the most part we're going to be concerned with electrolytes that dissociate completely, but every once in a while we will discuss things like acetic acid.0486

All right, having said this, now let's talk about precipitation reactions.0497

Precipitation reactions (rxn is just a shorthand for reaction): If I have a solution of some salt (and again, salt is just a fancy word for ionic compound)--sodium chloride, and I mix it with a solution of iron sulfate, something might happen.0504

As it turns out, there are certain salts that are so tightly bound together that, even when I drop them in water, they don't dissociate.0526

Or, if they do, they dissociate so little that, for all practical purposes, they don't dissociate.0537

We can measure very, very tiny amounts, but really, what happens--it just sinks to the bottom.0544

Well, what is kind of interesting is: when you take, let's say, a mixture of a salt that is fully dissociated, and you pour in another solution of a salt that is also dissociated, now ions are sort of slamming into each other randomly.0549

Well, some of those ions, when they stick together--they stick together so tightly that they literally fall out of solution as a solid--they just...literally, a solid shows up in the middle of solution and drops to the bottom.0562

We call that a precipitate.0573

Precipitation reactions are reactions where, when you mix things, a solid all of a sudden shows up in the reaction mixture.0575

Sometimes, it will be a liquid, but more often than not, it will be a solid.0585

The way we decide what is actually going to stick together: we use something called the Solubility Guidelines.0588

The Solubility Guidelines--I'm going to refer you to your text, because I think it's important that you find some of these important tables in your text, and be familiar with them--at least where they are.0597

Basically, they are just guidelines to tell you which ionic compounds tend to dissolve in water, and which don't dissolve in water--which just fall to the bottom without dissolving at all.0610

You're going to be referring to this over and over again.0620

Your teacher may ask you to memorize it; maybe not; for the most part, it's just a reference to decide, when you're faced with a reaction of two ionic compounds, is there going to be a precipitate?--and you decide based on solubility.0624

You'll see what I mean in just a minute.0637

So, a Solubility Guideline for ionic compounds looks like the following.0639

I just want you to see what you can actually expect to see in your book as a table.0650

You're going to see something that looks like this: it says (you know what, let me do this in red--it'll be a lot better): Soluble Ionic Compounds, and then it has a column called Important Exceptions.0655

You will see something that says like this: Compounds Containing...and it will list a bunch of ions: nitrate, carbonate, iodide, bromide, things like that.0698

Over here, it will say None, because any ionic compound that is compounded with a nitrate ion--they are all going to be soluble--completely soluble.0714

Then, let's say there is an entry for iodide: this means that all compounds that have an iodide in them (like magnesium iodide, lead iodide, something like that)--in general, they tend to be soluble.0728

Then, the exceptions are: in the case of things that are mixed with iodide, lead2+ ion, mercury22+ ion, and silver.0741

So, you're going to see something like this--a bunch of charts; you're going to see ions, and then you're going to see the exceptions.0753

Then, there is going to be another set, and it's going to say (oops, I have all kinds of stray lines showing up here--let's see...)--then you're going to see another that says Insoluble Ionic Compounds.0759

It's also going to have a column of important exceptions, and it's going to list a bunch of ions that...those ions that have those particular ion it are generally insoluble--they don't dissolve in water.0784

Then, there is going to be some Exceptions.0798

For example, you will see hydroxide; generally, anything that is bound to a hydroxide or multiple hydroxides--for the most part, most of the metals--most of the ionic compounds--that have hydroxide in them are insoluble in water, the exceptions being the alkali metals.0802

So, over in this column, it will say "alkali metal," and you will know that, when you see a hydroxide with an alkali metal, like lithium hydroxide or potassium hydroxide, you will know that it is an exception to this rule and they are completely soluble.0819

Now that we know what this looks like, again, I encourage you to find it in your book; it will usually be in one of the chapters, and it will say Solubility Guidelines for Ionic Compounds, or something like that.0834

Get to know that very, very well.0845

OK, so let's do a problem, and I think everything will come together nicely.0848

Let's look at magnesium nitrate, Mg(NO3)2 (right?--magnesium is 2+; NO3 is 2-; we need two of them to balance the charge) (aq) + sodium hydroxide--again, (aq), and Mg(OH)2; this one...we'll say S, and we have NaNO3, sodium nitrate.0852

OK, when you see an equation like this, the question is, "What is going on here?"0889

This is a shorthand notation for what is going on chemically.0894

Now, we're going to go ahead and describe what is going on chemically, and how you decide what is going on chemically.0898

The first thing we want to do is to make sure that this is actually balanced, so let's see: I have two nitrates over here; I have one magnesium, one magnesium; two hydroxides, one hydroxide--so I'm going to stick a 2 in front of that; now, that gives me 2 sodiums--I'm going to stick a 2 there--2 sodiums; 2 nitrates; 2 nitrates; I'm balanced--2 in front of the sodium hydroxide, 2 in front of the sodium nitrate.0904

Now, notice this aqueous, aqueous, aqueous; there is an (aq) over here, too--my apologies (and let me go back to blue).0931

This aqueous--what this means is that, even though we write it as a molecule, well...if I look under the Solubility Guidelines for magnesium nitrate, anything that has a nitrate is soluble.0938

It means that a solution of this (which--aqueous means that you have dropped some magnesium nitrate in water)--it's completely dissolved; it's a strong electrolyte--it dissolves completely; it's soluble.0950

So, it looks like this: what is floating around in solution is a magnesium 2+ ion and two free nitrate ions.0961

You also have two sodium ions, because sodium hydroxide is also soluble.0972

Hydroxide is generally insoluble, but the exception is the alkali metals; sodium is an alkali metal; the hydroxide of an alkali metal is completely soluble, so you have two sodium ions floating around, and you have two hydroxide ions floating around.0980

That is where the 2 comes from, here.0996

Now, here is what is going on: everything on the left side is what happens before a reaction takes place.0999

This arrow--everything here is after the reaction has taken place; now, let's see what has really happened.1005

Well, what is going to happen here is: this is going to slam into nitrate, but when it slams into nitrate, it's just going to come apart again, because it's soluble.1011

It's going to slam into sodium, but when it slams into sodium, they're both positively charged--they're going to bounce off each other.1020

But, when it slams into a hydroxide, because a hydroxide, under the Solubility Guidelines, is insoluble, it won't bounce off, and it won't dissolve again; the binding strength is so strong that they literally stick together, and they stick together so completely that it literally turns into a solid again.1026

So, this is...essentially, what happens is: you are switching partners.1046

The cation of one goes with the anion of the other, and this cation goes with this anion.1051

So, Mg and OH--so we have to make sure that's appropriate, so it's going to be Mg(OH)2, because hydroxide is -1 and magnesium is +2.1058

Because it is insoluble, that is why we write it as a solid.1069

It literally falls out of solution, down to the bottom of the flask.1075

You will also see this written with an arrow pointing down; that is an old symbol for precipitate.1078

I tend to prefer this symbol, myself.1085

The symbol for a gas is an arrow pointing up, meaning the gas bubbles off.1087

Now, let's look at the other combination--sodium and nitrate: well, sodium nitrate--we know that anything that has a nitrate in it is going to be soluble, so they stay as free ions; they might slam together, but they come apart again, because sodium nitrate is soluble.1092

So, you have two sodium ions and two nitrate ions that are, again, floating around in solution.1107

Now, watch what happens: you know that, in algebra, whenever you have something on the left side of the equality sign which is the same as the right side of the equality sign, they cancel out.1115

The same thing here: I have two nitrates on the left, two nitrates on the right; I can cancel those--they don't really participate in any chemistry.1125

I have two sodiums on the left, two sodiums on the right; now, I'm going to write a third equation, which is basically just what is left over.1133

Mg2+ + 2OH- → Mg(OH)2 as a precipitate.1143

These three equations--the first one is called a molecular equation; you're just giving molecular formulas--you're not really talking about what is going on.1154

The second equation tells you everything that is in solution; everything is right there--that is called the total ionic equation.1162

The total ionic is my personal favorite, because you see everything that is involved--there is nothing hidden.1172

This is it; this is the chemistry; the total ionic--that is where the chemistry takes place.1178

The third equation is called the net ionic; it's the actual chemical reaction that takes place.1183

We cancel the nitrate and the sodium because they don't do anything--they just float around; they are called spectator ions.1189

I'm not sure if the names are altogether that important; what is important is their function, which in this case is nothing--they just sit there and do nothing.1195

This is the chemistry.1202

So, again, the first equation: it is called the molecular, because we give the molecular formulas--the balanced molecular formulas.1203

The second equation is called total ionic (your book might call it complete ionic), because you are listing every ionic species.1217

The third, which is the actual chemistry taking place, is the net ionic.1227

It's called ionic simply because, on the left-hand side, you have ions, and here you have a precipitate--a solid that sinks to the bottom.1234

Notice, it's neutral; charges are balanced; neutral--zero charge on the right side; 2+, 2-, zero charge on the left side; zero net charge.1242

That is what is happening with precipitation reactions.1252

You're going to really mix two soluble compounds, and sometimes the ions--one cation and one anion--will be such that they are actually insoluble in solution; they'll bind really tightly, and they'll drop to the bottom of the flask as a solid.1255

You pour out the water; you filter it, and you collect the solid; you weigh it, and you do all of the other things that you're going to end up doing in a lab.1269

Let's do an examples.1276

Let's move forward: let's call this one Example 1 (the one previous was just a discussion): this time, we're not going to write the right half; we're going to actually figure out what is going on as a product.1282

The question is: what happens when you mix a solution (sol'n is a shorthand for solution) of barium chloride and potassium sulfate--BaCl2 and K2SO4?1298

Well, let's write out the molecular formula and see what happens; we'll switch partners and see what we get.1326

So, we have BaCl2; that is aqueous (well, actually, let's not write--let's just write BaCl2 first); and then we'll write K2SO4 and draw a little arrow; now, let's decide what is happening here.1332

Barium chloride: we look under the solubility guidelines, and we discover that it is soluble.1349

So, this is going to be a Ba2+, and again, full dissociation; that is what solubility means--full dissociation.1354

This Cl2 breaks up into 2 chloride ions, 2Cl- + K2SO4: sulfate may or may not be soluble, but in the case of a sulfate that is put together with an alkali metal, which potassium is, they are completely soluble.1362

So, again, they dissolve completely; that means we have 2 potassium ions in the solution, plus an SO42-.1381

Now, we switch partners; the barium goes with the sulfate.1392

Now, when we put barium and sulfate together, as it turns out, barium sulfate is not soluble; therefore, it has a little arrow pointing down (or you can put an S for solid); that is a precipitate.1397

The other combination is, as far as the molecular formula is concerned, just potassium and chloride, which is KCl.1412

Again, it isn't K2Cl2; we have to put them together the way we did when we named compounds.1421

Potassium is a +1; chlorine is a -1; it's just KCl.1427

Now, we can balance the molecular formula.1431

I actually should have done that first, instead of starting to write the total ionic; forgive me.1433

All you have do is switch partners; so, there are 2 potassiums here; I'm going to put a 2 there; here it's 2 chloride--I'm going to come over here--that is 2 chloride; 1 barium, 1 barium; 1 sulfate, 1 sulfate; good.1438

All I needed was this 2, and the equation is balanced.1449

Now, we already took care of the reactants: barium, chloride, and potassium sulfate; we broke them up into ions.1453

We said that barium sulfate is a precipitate, so it stays together; it does not separate into ions; they actually bind together.1460

Potassium chloride is soluble; therefore, it stays as 2 free potassium ions plus 2 free chloride ions.1469

Now, I take a look at what is on the left and what is on the right; the 2 chloride and the 2 potassium cancel with the 2 chloride and the 2 potassium, and I end up with Ba2+ ion + a sulfate ion; it gives me barium sulfate as a solid that falls to the bottom of the flask.1478

The first is my molecular equation that I got just by switching partners and combining them in the appropriate stoichiometric quantities.1500

The second is a total ionic equation, meaning everything that is happening, based on the Solubility Guidelines.1508

So, these two are soluble; therefore, they are floating around as free solutions; the barium sulfate is not; therefore, it sticks together.1514

The potassium chloride is soluble; therefore, it is floating around as free ions.1520

We cancel what is on the left and the right, and we are left with our net ionic.1527

Molecular, total ionic, net ionic: this is a complete description of the chemistry that is going on here.1530

I hope that this is making sense; it's reasonable straightforward--a lot of symbolism, but there is nothing strange going on here.1536

Your intuition should be able to guide you through, as long as you realize that this molecular--when you're dealing with ionic compounds, they may or may not dissolve.1543

What you are checking is solubility--if they are soluble, you write them out as free ions, and then just mix and match.1552

Really, though, you're not even mixing and matching; you're just switching partners.1559

OK, let's see; let's do one and see what happens here.1563

Our second example is going to be: let's mix a solution (well, let's actually spell out the problem); what happens when a solution of magnesium nitrate (let me write it over here--I don't have enough room), which is Mg(NO3)2, and potassium iodide, are mixed?1577

OK, well, let's write out the molecular formula and see what we get.1620

That is Mg(NO3)2 + KI (we want to write the molecular formula, so all we do is switch partners) magnesium and iodide; they come together, and it's going to be MgI2, and then potassium and nitrate are going to come together, so it's going to be KNO3.1626

Now, I want to balance this; that is the first thing I do, once I come up with a molecular by switching partners.1650

There is 1 magnesium, 1 magnesium; there are 2 nitrates--there is 1 nitrate here, so I put a 2, which gives me 2 potassiums; I stick a 2 in front of the K on the left; that gives me two potassium.1656

That gives me 2 iodide, but there is 2 iodide on the right; I'm fully balanced--all I needed was a 2 here and a 2 there.1668

Oops, some random lines are showing up again, which we do not want.1675

Let me put that Mg there; so now, let's check for solubility, because that is what we do next.1683

Well, magnesium nitrate is fully soluble; therefore, we have Mg2+ + 2NO3-.1690

Potassium iodide: fully soluble; therefore, we have 2 potassium ions and 2 iodide ions.1703

Now, let's look over here--magnesium iodide; it turns out magnesium iodide is also soluble; therefore, you have magnesium 2+, plus 2I-.1711

And potassium nitrate: definitely soluble; so 2K+ + 2NO3-.1723

So, notice, in this case: you ended up with--you mixed two solutions, where both things are fully soluble; and when you switched partners, you ended up with, well, something where both of the products are actually also soluble.1732

So, what is going on here--even though you can write a molecular formula and switch partners, there is no chemistry going on here.1747

All that is really happening is a bunch of free ions floating around, slamming into each other, and then coming apart again, because everything is soluble; every combination is soluble.1753

Really, there is no reaction here, even though we can write a molecular formula.1762

That is what is really important to remember: just because you can write a chemical equation doesn't mean there is any chemistry going on here.1767

Chemistry means something changes; magnesium, nitrate, potassium, iodide on the left, floating around, free in solution; magnesium, nitrate, potassium, and iodide, floating around in solution on the right; there is no difference--everything cancels; nothing happens here.1772

There is no net ionic, because nothing actually happens.1787

That is the important lesson.1791

So, take a look at your two ionic compounds; write a molecular formula on the left; switch partners, and write the appropriate chemical formula on the right.1795

Then, go back to the reactants and do the total ionic; see which one of those four compounds that you have is soluble, and which is not.1805

That will be your total ionic equation.1814

Then, you cancel similar species in the total ionic that are on the left and on the right, like we did for both of these examples; and what you are left with is your net ionic equation.1818

Your net ionic equation is the chemistry.1830

Here, in the last example, there is no net ionic; there is no chemistry; so, just because we can write it doesn't mean that something is actually happening.1833

These things don't really exist as compounds; they are soluble; they're just floating around, free, in solution.1841

This, the total ionic, is what is telling you everything that is going on; nothing is hidden--this tells you that nothing is going on.1848

Again, your intuition should be just as good; you know that free ions floating around--they slam into each other; there is nothing here that is mysterious--you just have to follow the logic.1857

OK, so thank you for joining us here at for the discussion of precipitation reactions.1867

We'll see you next time; take care.1872