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Chemistry 101: General Chemistry14 chapters | 132 lessons | 11 flashcard sets

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Lesson Transcript

Instructor:
*Elizabeth (Nikki) Wyman*

Nikki has a master's degree in teaching chemistry and has taught high school chemistry, biology and astronomy.

Learn the definition of a hydrate and an anhydrate in this lesson. Discover how, when given experimental data, you can determine the formula of a hydrate by following simple steps that include finding the moles of hydrate and anhydrate and comparing the two to write the formula.

'Ah, I need to hydrate!' Everyone knows what it means to hydrate. To hydrate means to drink, but it can also mean to 'combine chemically with water.' Although we use the word hydrate in our everyday lives, it's actually a chemistry term.

A **hydrate** is a compound that contains water with a definite mass in the form of H2O. Hydrates are often in the form of a crystal that can be heated, and the water can be 'burned off' by turning it into steam. This usually causes the hydrate to lose its crystalline structure. The substance that is left over after the hydrate has lost its water is called an **anhydrate**. By measuring the compound before heating and after, the amount of water in the original hydrate can be determined and the formula discovered.

Unknown hydrates are written with their base form; then an *'n*' is placed before the H2O. The *'n*' before the H2O means there is a number there, but we don't know what it is yet, such as in MgSO4 *n*H2O for a magnesium sulfate hydrate or Na2CO3 *n*H2O for a sodium carbonate hydrate. Luckily for us, this is an easy determination. Finding *'n*' is fun!

Here are the steps to finding the formula of a hydrate:

- Determine the mass of the water that has left the compound. This allows us to determine the mass of water that was in the hydrate and the mass of the anhydrate. We do this by subtracting the mass of the anhydrate from the mass of the hydrate. This equals the mass of water.
- Convert the mass of water to moles. To do this, we divide mass of water by the molar mass of water to get moles of water. Remember that the units for molar mass are g/mol. When we divide mass (in g) by molar mass (g/mol), grams will cancel out and we will be left with moles.
- Convert the mass of anhydrate that is left over to moles. To do this, we divide the mass of anhydrate by the molar mass of anhydrate to get the moles of anhydrate.
- Find the water-to-anhydrate mole ratio. Generally, you will have more waters than anhydrate, so divide the moles of water by the moles of anhydrate. This gives you your mole ratio (a note about this step: if your calculations give you a number that is very close to a whole integer, it's generally safe to round to the nearest whole one. If your number has a decimal that is close to .33, .5, or .66, you need to find the lowest common multiple of that number that is a whole number and apply it to the entire formula).
- Use the mole ratio to write the formula.

Let's see how these steps work with a sample problem:

A 210.4 g hydrate of Epsom salt, MgSO4 *n*H2O was heated up, the water was released, and the final anhydrate mass was 120.4 g. What is the formula of this hydrate?

- Determine the mass of the water that has left the compound. Take the mass of the hydrate and subtract the mass of anhydrate from that to get the mass of water. 210.4 g MgSO4
*n*H2O - 120.4 g MgSO4 = 90 g H2O. - Convert the mass of water to moles. Mass of water / molar mass of water = moles of water. 90 g H2O / (18 g/mol H2O) = 5 moles H2O.
- Convert the mass of anhydrate that is left over to moles. Mass of anhydrate / molar mass of anhydrate = moles of anhydrate. 120.4 g MgSO4 / (120.4 g/mol MgSO4) = 1 mole MgSO4.
- Find the water-to-anhydrate mole ratio. Divide moles of water by moles of anhydrate to get the mole ratio. 5 moles H2O / 1 mole MgSO4 = 5:1.
- Use the mole ratio to write the formula. Since there are 5 moles of H2O for every 1 mole of MgSO4, the formula is MgSO4 5H2O.

Are you ready for another example? Feel free to pause the video at any point to do your own calculations:

A hydrate of sodium carbonate, Na2CO3 *n*H2O, originally contains 17.70 g. After heating, its final weight is 15.10 g. What is its formula?

- Take the mass of the hydrate and subtract the mass of the anhydrate to get the mass of water. 17.70 g Na2CO3
*n*H2O - 15.10 g Na2CO3 = 2.60 g H2O. - Divide the mass of water by the molar mass of water to get moles of water. 2.60 g H2O / (18.00 g/mol H2O) = 0.144 moles H2O.
- Divide the mass of anhydrate by the molar mass of anhydrate to get moles of anhydrate. 15.1 g Na2CO3 / (106 g/mol Na2CO3) = 0.142 moles Na2CO3.
- Divide moles of water by the moles of anhydrate to get the mole ratio. 0.144 moles H2O / 0.142 moles Na2CO3 = 1.01:1. It is safe to round this to a 1:1 ratio.
- Use the mole ratio to write the formula. Since there is only 1 mole of H2O for every 1 mole of Na2CO3, the mole ratio is 1:1 and the formula is Na2CO3 H2O.

Let's try an example that you might encounter in a lab setting. Here we are in the lab after dehydrating some iron (III) chloride salts. Now it's time for data analysis. We want to determine the formula for a hydrate of iron (III) chloride, FeCl3 *n*H2O. Let's look at our data:

- Mass of the empty dish used for weighing = 2.5 g
- Mass of dish and sample before heating = 7.4 g
- Mass of dish and sample after heating = 5.4 g

What's the formula for our hydrate? To solve this, we can use the same steps we used in the previous example. This time, however, we must subtract the mass of our weighing dish from the mass of our sample before and after heating.

The mass of the hydrate = the mass of the dish and sample before heating - the mass of the empty dish. 7.4 g - 2.5 g = 4.9 g, the mass of our hydrate.

The mass of the anhydrate = the mass of the dish and sample after heating - the mass of the empty dish. 5.4 g - 2.5 g = 2.9 g, the mass of the anhydrate.

The mass of water = the mass of the hydrate - the mass of the anhydrate. 4.9 g - 2.9 g = 2.0 g of water.

Moles of water = the mass of water / the molar mass of water. 2.0 g H2O / (18 g/mol H2O) = 0.11 moles H2O.

Moles of anhydrate = mass of anhydrate / molar mass of anhydrate. 2.9 g FeCl3 / (162 g/mol FeCl3) = 0.018 moles FeCl3.

The mole ratio is all we have left to determine. Mole ratio = moles of water / moles of anhydrate. 0.11 moles H2O / 0.018 moles FeCl3 = 6.1:1. 6.1 is very close to 6, so it is safe to round down. We'll call this 6. Since there are 6 waters to every 1 iron (III) chloride, the ratio is 6:1. The formula for our hydrate is FeCl3 6H2O.

A **hydrate** is a compound that contains water with a definite mass in the form of H2O. An **anhydrate** is a hydrate that has lost its water molecules. Determining the formula for a hydrate means discovering the number of water molecules that the substance contains. The steps to determining the formula for experimental data are easy:

- Determine the mass of water that has been removed from the compound
- Convert the mass of water to moles
- Convert the mass of anhydrate that is left over to moles
- Find the water-to-anhydrate mole ratio
- Use the mole ratio to write the formula

Once you've completed this lesson, take a few moments to:

- Compare hydrate and anhydrate
- Enumerate the steps for determining the formula for a hydrate
- Analyze two related examples

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Chemistry 101: General Chemistry14 chapters | 132 lessons | 11 flashcard sets

- Go to Atom

- Go to Gases

- Go to Solutions

- Chemical Reactions and Balancing Chemical Equations 6:07
- Mole-to-Mole Ratios and Calculations of a Chemical Equation 10:13
- Mass-to-Mass Stoichiometric Calculations 7:04
- Stoichiometry: Calculating Relative Quantities in a Gas or Solution 11:07
- Limiting Reactants & Calculating Excess Reactants 7:31
- Calculating Reaction Yield and Percentage Yield from a Limiting Reactant 7:07
- Calculating Percent Composition and Determining Empirical Formulas 11:04
- Hydrates: Determining the Chemical Formula From Empirical Data 10:01
- Go to Stoichiometry

- Go to Kinetics

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