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This page contains a text transcript of podcasts created by Dr. Kate Amaral for Chemistry 110 at Penn State Berks. Listen to podcasts »

Converting Atoms to Grams

This is Dr. Kate Amaral from Penn State Berks. This is a podcast for Chem 110. We’re going to be talking about converting atoms to grams. When we do things in chemistry we tend to use grams often because it’s something we can weigh out and measure. The relationship between substances in chemistry tends to actually be in a mole unit rather than a grams unit. And the mole is essentially the chemist’s dozen. If I have a dozen eggs or a dozen pencils, I have 12 of them. It didn’t matter if I’m talking about eggs or pencils. The mole is that same sort of thing. If I have a mole of water or a mole of carbon dioxide, I have the same amount of water as I do of carbon dioxide, and that’s Avogadro’s number which is 6.022X10^23. Remember that number, you’re going to be using it quite a lot over the course of the semester.

So when we start out in grams, it’s often more useful for us to be in moles. And so we’ll have to convert from grams to moles, or if we’re in moles and we want to get to atoms or molecules, we have to be able to make that conversion. So if we’re going to start with the same molecules, I have a lot of molecules of a substance and I want to get to moles so I can use that relationship to relate my substance to some other substance, we would have to use Avogadro’s number.

We were talking about water, H2O. If we wanted to convert 1X10^33 molecules of water to moles, we would use Avogadro’s’ number, for every 6.022X10^23 molecules of water, we have one mole of water. We could also go from moles to molecules and say alright if I have 3.5 moles of water, you can convert to molecules of water by saying for every one mole of water I have 6.022X10^23 molecules of water. That works for compounds going back and forth. If we wanted to just talk about hydrogen atoms in water, though, we would have one extra step. We have one molecule of water; we have two atoms of hydrogen since water’s formula is H2O. That two tells us how many atoms we have in one molecule.

So if we wanted to go to atoms of H from moles of water, so let’s take our 3.5 example, we have 3.5 moles of water, you convert first to molecules. And say for every 6.022X10^23 molecules we have one mole. Then when we are in molecules we can go further and go to atoms of H. For every one mole of water we have two atoms of H. We can also, if we’re just looking at water and we have a mole amount of water and we want to know the mole amount of Hydrogen we have, we can skip that converting to molecule step. And just as we said for every molecule of water, we have two H atoms. We can say for every mole of water, we have two moles of hydrogen. So if we had 3.5 moles of water, we wanted to get two moles of hydrogen, we would say ok for every one mole of water we have two moles of hydrogen.

But it’s really difficult for us in a lab to measure things out it atoms or molecules, instead we measure things out in grams. But often we want to get from grams to atoms or molecules and we can’t go directly there. We have to go through the mole step. So to convert something from grams to moles, we used the molecular weight, which we get off the periodic table. So if we do our water example, let’s say we have four and a half grams of water. We want to know how many molecules of water we have. We have our four and a half grams of water. We’re going to use the molecular weight of water to convert from grams to moles. So water is H2O. So we are going to take two times the weight of hydrogen, which is 1.01 grams per mole. We’re going to add to it one times the weight of oxygen which is 16 grams per mole. So we started off with our four and a half grams of water. We’re going to use 18.02 grams, which gives us one mole of water. And now we can go from moles to molecules just like we did before. For every one mole of water we have 6.022X10^23 molecules of water. And again, if we wanted to go further on to Hydrogen, we could say for every one molecule of water we have two H atoms.

So we can’t go directly from grams to molecules. We have to go through the step of moles no matter which direction we are going in. If we’re going from grams and we want to know molecules, well, we go from grams to moles and moles to molecules. If we’re in molecules, and we want to know how much to weigh out to get that number of molecules we go from molecules to moles, moles to grams. In general, what you want to do, the first rule of thumb is pretty much to convert to moles. If you’ve converted to moles you’re on the right track, it’s just where you go from there.

When Ions Go into Water

This is Dr. Kate Amaral from Penn State Berks. This is a podcast for Chem 110 and today we will be talking about what happens when ions go into water. In general, if we talk about an ionic compound, what keeps the ions together in that ionic compound is the attraction of opposite charges.

We have sodium chloride, everybody’s favorite ionic substance. The sodium plus stays with the chlorine because it’s chlorine minus and the positive and negative attract each other. But if we actually look at sodium chloride, we don’t have just one Na+ and one Chlorine minus. Instead what we have is a Na+ surrounded on all sides by Chlorine minus. And each Chlorine minus is surrounded on all sides by Na+. And those attractions of the positive and the negative keep those sodiums and those chlorines next to one another.

If we put a teaspoon of table salt, Sodium Chloride, into water, what happens in water though is that the water has the ability to separate those ions from each other to break that charge attraction. Now, water does that because water is polar. It has a slightly positive end near the hydrogens, H2O, and a slightly negative end near Oxygen. It’s not an ionic substance, those aren’t true positives or negatives, but it is slightly positive by hydrogen and slightly negative by oxygen. If we put a teaspoon of salt in a glass of water, we have a lot more water than we have table salt. And so what happens is, the slightly negative end of water, the oxygen end, will surround the sodium plus ions and manage to pull it away from the chlorines. And for the chlorine minuses, the positive end of the water, the hydrogen end, can surround the chlorine negative and pull it away from the sodiums. And so what you end up with is ions in water surrounded by water molecules. So at one part in the glass you’ll have a sodium plus surrounded by water, the negative end of water. And some place else you’ll have chlorine minus surrounded by hydrogen part of water, the positive part of water.

Now the substances like salt that can completely break apart into its individual component ions, we call strong electrolyte. An electrolyte is basically any substance that can conduct a current.  And since there’re ions in water, there is charge; they’re able to conduct currents. They separate completely. And to find out what ions separate completely we use the table of solubility rules. And this is something you want to look out and learn how to use before you get to the test where you’ll be given one. And so if we look up something like chlorides, they are listed by anion.

We see that all chlorides are soluble unless they contain a cation that’s either, Ag+, Hg2+ or Pb2+. So any chloride we can think of will be soluble, unless it also contains one of those three ions. So NaCl will be completely soluble in water. The ions will completely break apart. It will be a strong electrolyte. We had AgCl on the other hand, which our solubility rules tell us is insoluble. When we put it in water, the charge attraction between the silver and the chloride are so strong that no matter how much water we’re in, that water can’t surround the ions and put it out. It stays as a solid. And there are certain substances that tend to be generally insoluble in water. Most of our carbonates for example, pretty much tend to be solids. They’re not going to separate in water. There are some exceptions to that, of course, if we put it with an alkali metal, which is anything in group one, like lithium, sodium, or potassium. We put potassium carbonate in water; it will split into its component ions and be a strong electrolyte. But if we put anything else, like calcium carbonate in water it’s not going to split into its component ions. It’s not going to be an electrolyte.

So we have strong electrolytes, substances that are completely soluble in water and separate into component ions. We have non-electrolytes, things we put in water and they don’t separate at all. And then we have this category in the middle and that’s weak electrolytes. And there are certain substances that dissolve a little bit in water, but not a hundred percent. There’s a list in your textbook that has your strong electrolytes, we can find them under the solubility rules. Other strong electrolytes would be our strong acids and strong bases that we already need to know. If it’s a weak acid or a weak base, then it’s going to be a weak electrolyte. They’re not going to dissolve one hundred percent in water. So we put it in water and in general it’ll to stay together. Something like Nitrous Acid. HN02 is a weak electrolyte. You put that in water and it kind of stays together. A little bit of it dissociates into H+ and NO2 minus. The majority of it stays together, and so it’s a weak electrolyte. It will carry only a weak current.

Net Ionic Equations

This is Dr. Kate Amaral from Penn State Berks. This is a podcast for Chem 110. Today we’ll be talking about net ionic equations. Now that we know how to determine what ions a substance will separate into in water by looking at the solubility rules, we want to see what happens when we mix different soluble compounds. Sometimes when we put soluble compounds together, they make a new substance which is a solid that can come out of solution. And we call these types of reactions precipitation reactions because the solid looks like it is precipitating down to the bottom of the flask.

So let’s look at two substances. Let’s put silver nitrate into water. So we’re going to look at our solubility rules, they’re listed by nitrates. It says all nitrates are soluble, no exceptions. So silver nitrate is soluble in water. So when silver nitrate is in water, we’ll have Ag+ and NO3 minus floating around all surrounded by water molecules no longer associated with each other. And silver nitrate would be a strong electrolyte. And then let’s put in another beaker sodium chloride. If we put sodium chloride in water, the water’s able to overcome the attractions between the sodium and chlorine and will separate them. So we will have sodium plus floating around and chloride, Cl minus, floating around. To know that for sure check our solubility rules. We look up by the anion, and all chlorides are soluble unless the cation is Silver plus, Hg2+ or Pb2+. Since we threw in sodium chloride, it’s not one of those three. So it’s completely soluble and a strong electrolyte.

If we mix our silver nitrate and our sodium chloride together, we have to see if we get a reaction. So what we want to look at is pairing up cations with the opposite anions, the anion that was in the other beaker. So, we had sodium chloride and silver nitrate, so let’s see if our sodium plus our nitrate will make a solid. When we look at our solubility rules, nitrates are soluble. There are no exceptions. So, sodium nitrate will be soluble. They will stay apart in the flask. They won’t associate.  So let’s look at our silver reacting with our chloride. If we look up chlorides, we see that most chlorides are soluble except if it contains silver, mercury, or lead. We have silver around. So when we mix those two together the silver from the silver nitrate will react with the chloride from the sodium chloride and it will form silver chloride, which is a solid which means it will eventually all settle to the bottom of the flask.

So what we have to do now is we have to write a reaction that shows what happened. Well what happened is we started with our silver plus and it was aqueous. It was in a solution of water. And then we added some chloride to it, Cl minus, which was also aqueous in a solution of water. And what we did was we formed a solid, Silver Chloride. And if we look at that, that is what we call a net ionic equation.  It just says specifically what happened. In our reaction, Ag+ Aq reacted with Cl minus Aq and we formed AgCl solid. So a good net ionic equation has the ions that reacted, it has the charges that go with them, and it has the phase label, aqueous, aqueous, and solid.

The sodium and the nitrate, which didn’t do anything, they didn’t take part in our reaction, they didn’t form a solid, they’re still floating around the beaker of water, we call them spectator ions. They don’t do anything. They watch the reaction occur between the silver and chloride. So after we identify what solids we will get we can also identify what the spectator ions will be and we can write a net ionic equation that has the ions that will come together to from the solid.

Significant figures

This is Dr. Kate Amaral from Penn State Berks. This is a podcast for Chem 110. We will be talking about significant figures. Each measurement that we do in a lab, or outside of a lab, has a certain amount of significance. What I mean by that is that we’re accurate to a certain level depending on the device we use to measure that.

If I ask everybody’s age in my classroom and they tell me, 18,18,19,17 and I take an average and say well the average age of student in my classroom is 18 years, 4 months, 3 weeks, 6 days, and 10 minutes, and 5 seconds; well I didn’t measure the age out that far. I can’t do that. I am giving my measurement a level of significance it doesn’t have. If I got my ages in my classroom and I had 18,18,19,17, as far as I could go would be 18. The average age of students in my classroom is 18 years old. That’s as far as I’ve measured it.

So usually the last digit that we measure has some amount of uncertainty to it. So anything after that last uncertain digit, we can’t measure because it’s not significant.  We’re kind of guessing as we get further down. So if we wanted to look at a measurement we can determine what numbers are significant and which one’s not. If it’s a number between 1 and 9 we say ok it’s significant. The zeroes are where the problems occur. If we measured a zero, it’s significant and if we didn’t measure a zero it’s insignificant.  So if I measured something as 1.30 that zero, I measured it, it’s significant. If I measured it as 103, that zero in the middle is also significant. But if I measured it as 130, that last zero is only a place holder to tell us that we’re in the hundreds place and not the tens place. That zero, was not measured, not significant. 130 only has two significant digits.

When we start doing math with that, we have to pay careful attention to what significant figures we have. If we’re adding and subtracting something, we keep our answer to the lowest number of decimal places. So if you’re adding a number with three decimal places to a number with one decimal place, your answer can only have one decimal place. That’s for adding and subtracting. Multiplication and division is a little bit different. For multiplication and division, we keep the lowest number of significant figures. So if I’m multiplying something with five significant figures by something with three significant figures, my answer should have three significant figures. Generally, if I’m only doing one of those rules, if I’m only adding and subtracting, or I’m only multiplying and dividing, when I get to the end is when I want to put it in the right number of significant figures.

If I’m doing both of those things in the same calculation, I have the multiply something and then add something to it and then divide that by something else, at every step of the way I need to make sure I’m at the right number of significant figures so I don’t carry insignificance to the end. But if I’m just doing one, or the other, which is most instances in a generally chemistry classroom, you can wait until the end before you assign significance.

Trends for Atoms and Ions

This is Dr. Kate Amaral from Penn State Berks. This is a podcast for Chem 110. And today we will be talking about trends for atoms and ions. When we look at a periodic table, the atoms are arranged by atomic number. It goes Carbon, Nitrogen, Oxygen, (6,7,8) and that tells us the number of protons. But from the periodic table, we can also get other useful information such as size. So if we’re looking at a periodic table, as we go down a group, we go from period one to period two to period three to period four, what happens is we’re increasing our N level, from our Quantum number, so basically the distance from the nucleus is increasing as we add another N level from two to three to four. And as we increase the N level, the size of an atom gets bigger, and bigger, and bigger. So something in period three is bigger than something in period two. Something in period six is bigger than something in period five, or four, or three, or two.

The other part of the size trend that we have to look at is going across the period going from Carbon to Nitrogen to Oxygen, to Fluorine. If we look at that, they’re all at the same N level, N=2, so they are all roughly the same distance from the nucleus. All their orbitals are roughly the same distance from their nucleus. But if we look at Fluorine vs. Lithium, Fluorine has nine protons and Lithium only 3.  So what keeps the electrons close to the nucleus is the attraction between the positive protons in the nucleus and the negative electrons in the oribitals. And since they’re the same N level, the orbitals are the same rough distance from the nucleus but Fluorine has nine positive charges to attract its electrons while Lithium only has three. So since Fluorine has more positive charges to attract its electrons, it’s much smaller than Lithium even though Fluorine has more electrons than Lithium. So our general trend for size is that size increases as we go down a group. So Francium would be our biggest element, and size decreases as we go from left to right, so Helium would be our smallest element.

The next trend we’re going to look at is the ionization energy. Ionization energy is the ability of an atom to lose an electron. It’s lithium gas going to lithium plus (which is also gas) plus an electron. And as things get bigger, the attraction between the nucleus and electrons gets weaker and weaker. And so bigger things can really give up electron quite easily while small things that are holding its electrons tight, they don’t give them up very easily. So our trend for ionization energy, which is the ability to give up an electron, is the opposite of our size trend. Francium, our biggest element, can easily give up an electron so it has the lowest ionization energy. Helium, which is our smallest one, holds on tight to its electron so it’s extremely difficult to give up its electron. So it has very high ionization energy. The noble gas column is sort of an exception to our general trend.  If we look across a row, let’s say Nitrogen to Oxygen to Fluorine, Oxygen’s ionization energy is a little bit higher than Nitrogen’s, and Fluorine’s is just a little bit higher than Oxygen’s. When we get that noble gas column, our noble gases are un-reactive unless it’s much, much higher. They really don’t want to give up an electron. They’re going to hang out to the as tight as possible.

The next trend is electron affinity. Electron Affinity, what happens is an atom gains an electron. So this is lithium gas, plus an electron, going to lithium minus gas. So again, something that’s small has a strong attraction for its electrons; it’s really easy to give an electron to it. But something that’s big, where there’s not so much of an attraction; it’s kind of difficult to give it another electron. So again, this trend is opposite of our size trend. Francium, our biggest thing has a very low electron affinity. It doesn’t attract electrons well. But Fluorine, for most purposes are smallest atom, with its very tight attraction for its electrons, attracts electrons very easily and so it has a high electron affinity. Again, like with ionization energy, our noble gases are different. They’re un-reactive and they don’t want an electron just like they didn’t want to give up an electron. So they have extremely high ionization energy and an extremely low electron affinity.

Iso-electronic Series

This is Dr. Kate Amaral from Penn State Berks. This is a podcast for Chem 110. We will be talking about iso-electronic series. As we’ve talked about the size trend on the periodic table, in general the size trend goes towards Fluorine; things get smaller as we go towards Fluorine. So as we go down a group from period two to period three to period four, we get larger and larger. Also, as we go across a group we get smaller and smaller if we’re going left to right. Even though Fluorine has nine electrons and Lithium only has three, they’re in the same N level so they have roughly the same distance from the orbitals to the nucleus. But, Fluorine has nine protons to attract all of its electrons, where Lithium only has three. So those nine protons attract the electrons tighter and Fluorine is smaller than Lithium. So in general, we increase towards Cesium and decrease towards Fluorine with Helium really being the smallest element on the periodic table.

 But when we start talking about ions, then things get a little bit more difficult. There are certain ions that we say are iso-electronic. What Iso-electronic means is that they have the same number of electrons.  Iso means same and electronic is our electrons. So for example Cl minus has eighteen electrons, Argon has eighteen electrons, and K plus has eighteen electrons. That would be an iso-electronic serious, they all have the same number of electrons. So when we’re trying to determine the size of an iso-electronic series, we can’t use the general size trend to predict that. Instead, we have to think about what’s happening with the electrons. As we take an electron away from an atom, that atom will hold its electrons that it has left even tighter, the attraction with be even stronger. But if we add an electron to an atom, then the atom can’t attract the electrons as well because now there are more of them and it gets a little bit bigger. So if we’re talking about an iso-electronic series. A positive thing is smaller than a neutral atom is smaller than a negative atom. The more positive we get, the smaller we get. The more negative we get, the bigger we get.

 So if we look at our iso-electronic series of Chlorine minus, Argon, and K plus, if we were to arrange those from smallest to largest, K plus would be smallest, Argon would be in the middle, and Chlorine minus would be the largest. If we were to expand that a little bit, expand our iso-electronic series to five atoms and ions, we could do S2 minus, which has eighteen electrons, keep the Chlorine minus, which has eighteen electrons, Argon with eighteen electrons, K plus with eighteen electrons and Calcium two plus with eighteen electrons. If we had to organize those according to size, Ca2 plus would be smallest because it’s the most positive, then K plus, Argon, then we would get to Cl minus, and then finally Sulfur two minus. Sulfur two minus is the biggest because it is the most negative.

Phase Diagrams

This is Dr. Kate Amaral from Penn State Berks. This is a podcast for Chem 110. We will be talking about Phase Diagrams. Phase diagrams are essentially a graph that can tell us whether a substance will be a solid, liquid, or gas at certain temperatures or pressures. So picture a graph in your mind with temperature on the x axis and pressure on the y axis. If we start off at a very low temperature, we’d probably be solid. As the temperature increases, we’ll become liquid and as we increase even more we’ll eventually become a gas. So if we’re looking at a phase diagram, from left to right we’ll have solids at low temperatures moving on to liquids at intermediate temperatures and gases at high temperatures. If we go up the y axis and think about pressure, as we start at a very low pressure, we will be a gas. Our substance will exist in a gas phase. If we start to compress that and add more and more pressure, eventually we’ll become a liquid and finally as we compress at a very high pressure we become a solid. So going up the y axis, we start off as a gas eventually go to a liquid and finally, a solid. So if we’re looking at a phase diagram, on the left side of the diagram with low temperatures and high pressures will be solid. So we move a little bit to the right, temperature goes up a little bit, pressure decrease a little bit, we’ll be liquid. And then at high temperatures and at low pressures we will exist as a gas.

So there are lines that separate these three phases from each other. There’s a line that separates a solid from a liquid and one that separates a liquid and a gas. And those lines tend to slope slightly to the right. If we cross the line between a solid and a liquid, if we’re going from a solid to a liquid, what we’re doing at that point is melting. We’re melting our solid. If we go in the opposite direction, from a liquid to a solid, we’re freezing our liquid and forming a solid. If we do that at a pressure of one atmosphere, we call that the normal freezing point. So the normal freezing point is at one atmosphere of pressure when we cross the line from liquid to solid. If we were to cross the line between the liquid and the gas, we went from a liquid to a gas, we would be vaporizing.  And if we went back from gas to liquid we would be condensing. Just like water, you can see when you boil water to make pasta that the water vaporizes and becomes steam and if you have a lid on, it’ll condenses on the top of the lid and fall back as a liquid.  So it’s going back and forth across that liquid gas line. Just like with melting, if we go across that liquid gas line at one atmosphere, we call that the normal boiling point. So we go from liquid to a gas at one atmosphere, it’s our normal boiling point.

The only other two points of note on a phase diagram are the triple point and the critical point. The triple point is that perfect temperature and pressure where a substance can exist as a solid, a liquid, and a gas all at once and that’s a very specific temperature and pressure for each substance. The critical point we find at the very top of our liquid gas line. That critical point basically tells us after we pass that point; no matter what temperature we’re at or pressure we’re at we’re still going to be a gas phase. We can’t go back to a liquid phase anymore once we pass that critical point.

So in a phase diagram, when we look at it we have solids, liquids, and gases. The solids, liquids, and gases are separated by lines. If we cross one of those lines, say from a solid to a liquid, we melt. We go from a solid to a liquid phase.  If we go from a liquid phase to a gas phase, we boil. The triple point is that special temperature and pressure where the substance can exist as a solid, liquid, and gas all at once. The critical point is where no matter what we can’t become a liquid or a solid anymore.

On most phase diagrams there is a special place where we can go directly from a solid to a gas without going through the liquid phase. And if we do that, if we cross from a solid directly into a gas, we call that sublimation. Dry ice, which is carbon dioxide, solid carbon dioxide, doesn’t pass through a liquid phase. If you leave it out and it’s solid it will evaporate into a gas which is how they do fog a lot. If we go the other way, from gas to a solid, that’s deposition. We’re depositing our solid onto the substance.

The only phase diagram that’s a little bit different than most of these is a phase diagram for water. The line between a solid and a liquid for most substances slopes a little bit to the right. But for water, it actually slopes a little bit to the left. And that means, as we increase pressure we go from a solid back to a liquid which is one of the few substances we can do that with. We take advantage of that in the winter as we ice skate. We start off on a solid surface and we put our ice skate blade on that solid surface and we increase pressure. And as we increase pressure, it goes back to liquid water and actually what we’re doing is gliding on the liquid water. So water is a strange exception here that line between solids and liquids slants slightly to the left and as we increase pressure, we go back to a liquid.

Stoichiometry with Thermo- Chemical Equations

This is Dr. Kate Amaral from Penn State Berks. This is a podcast for Chem 110.Today we’ll be talking about stoiciometry with thermo-chemical equations. Before when we’ve worked with stoichiometry, we’ve worked directly within an equation. We go from compound A to compound C, calculate how much of that we’re going to get or how much of A we’re going to need to make some amount of C. And we’ve been doing that just with the chemical substances in the equation. But every reaction that we do in a lab either gives off heat or requires heat to run. And so, we call that a thermo-chemical equation. A thermo-chemical equation is when an equation is balanced and includes the heat either given off or required by a reaction.

So an example would be thinking of water liquid going to water gas. To do that would be 44 kilojoules of heat to go from water liquid to water gas. That would be a nice balanced thermo-chemical equation. So if we change the coefficients on that, let’s say we need two waters, liquid waters, to go to two gaseous waters, the enthalpy change, that delta h, will double. It will be twice as much as before. So instead of being 44kilojoules, it will be 88 kilojoules. If we do the reaction in the reverse direction, let’s say we are going to take gaseous water and make it liquid water, the delta h, the enthalpy associated with that equation will change its sign. So instead of being positive 44 kilojoules, it will be negative 44 kilojoules. And so those are some ways that we manipulate our thermo-chemical equations. We change the stoichiometirc amounts of the substances we are doing, we reverse them and we need to be able to handle the enthalpy change that goes along with that. So if we multiple a balanced thermo-chemical equation by some factor like 4, 10 or 30, we have to multiple the enthalpy value, the delta h, by that same number, 4, 10 or 30. And if we reverse a thermo-chemical equation, we change the direction of the arrow; we have to change the sign on the enthalpy value that goes with that.

When we actually do stoichiometry with these, we can figure out how much heat will be release by a reaction or how much heat will be required by that reaction. And that’s actually important in a lab because if a lot of heat is given off, we probably want to make sure that it doesn’t get too hot, that the substance doesn’t burn. And the same thing if it requires heat and we don’t give it any, well the reaction’s probably never going to go. So we need to pay careful attention to are thermo-chemical, our enthalpy values, as we do reactions.

 When we do a reaction in a lab, we are actually going to be starting with a gram amount rather than a mole amount. And so, just with stoichiometry where we convert from grams to moles and use the balanced equation to find products, we’re going to do that exact same thing with thermo-chemistry. If we start off with grams of water, we would take water, convert to moles using molar mass just like we’ve always done, then we can figure out how much heat that requires by using our balanced equation. So if we remember for the one where we had two liquid waters going to two gaseous waters, the delta h value was positive 88. So if we start with our grams of liquid water, we can convert to moles of liquid water using molar mass. Then we use our balanced equation. Well for every two moles of liquid water, the reaction requires 88 kilojoules of heat. We can put that right in our dimensional analysis. We can also work backwards with that enthalpy value. If we said, ok we reacted liquid water to make gaseous water and when we did that we had to put in 150 kilojoules of heat. From that, we could figure out how much water we reacted, just as with normal stoichiometry. So we would take our 150 kilojoules of heat, use our balanced equation, for every 88 kilojoules of heat we reacted two moles of liquid waters, and then we could go back to grams by using molar mass.

So with a thermo-chemical equation even though the enthalpy value is separated from the physical equation by a space we can still treat it like part of the equation and use stoichiometry with it.