To determine the number of moles present in 16 grams of water, we need to use the mole concept. The mole is a unit used to represent the amount of a substance. It allows us to count particles on a macroscopic scale rather than having to count individual atoms or molecules. The mole relates mass and molecular weight. Molecular weight tells us the mass of one mole of a substance. For water, the molecular weight is 18.015 g/mol. This means one mole of water has a mass of 18.015 grams.

## Converting Grams to Moles

We can use the molecular weight of water to convert from grams to moles using the following relationship:

moles = grams / molecular weight

For our example with 16 grams of water:

moles of water = 16 g / 18.015 g/mol = 0.889 moles

So for 16 grams of water, the number of moles present is 0.889 moles. This makes sense because 16 grams is less than the 18.015 grams in one mole of water.

## Using the Mole Ratio

Another way to think about this problem is using the mole ratio between water and individual water molecules. One mole of water contains 6.022 x 1023 water molecules. This is known as Avogadro’s number. Therefore, the mole ratio between water and water molecules is:

1 mole of H2O = 6.022 x 1023 molecules of H2O

If we have 0.889 moles of H2O present in our sample, we can multiply it by this ratio to find the number of water molecules:

0.889 moles H2O x (6.022 x 1023 molecules H2O / 1 mole H2O) = 5.35 x 1023 molecules H2O

So our 16 gram sample contains roughly 5.35 x 1023 water molecules.

## Molar Mass Calculation

As a check, we can calculate the molar mass of water from the atomic weights of hydrogen and oxygen:

Hydrogen: 1.0079 g/mol

Oxygen: 15.999 g/mol

For H2O, there are two hydrogen atoms and one oxygen atom, so the molecular weight is:

(2 x 1.0079) + 15.999 = 18.015 g/mol

This matches the known molecular weight of water, confirming our calculations are correct.

## Importance of the Mole Concept

Being able to interconvert between mass, moles, and number of particles is incredibly important in chemistry. Stoichiometry relies on using mole ratios to determine the amounts of reactants and products involved in chemical reactions. The mole allows us to bridge the macroscopic world that we perceive and the molecular-scale world influencing what we observe. It puts different chemicals on an equal footing for quantitative analysis.

Many key chemistry concepts such as solution concentration, reaction yields, and gas laws depend on the mole. Mastering the mole conversion process is a fundamental skill for all chemists, both in the classroom and the laboratory. With practice, mole calculations become second nature. The more familiar you become with various mole conversion problems, the easier they become over time.

## Mole Conversions for Other Substances

The process we used above can be generalized to calculate moles from grams for any substance, not just water. The steps are:

1. Identify the molecular weight of the substance. This can be looked up on a periodic table or other chemical reference source if unknown.

2. Use the formula: moles = mass (g) / molecular weight (g/mol)

3. Rearrange the ratio to match your known and unknown values. Make sure the units cancel out properly.

As long as you know the mass and the corresponding molecular weight, you can find the moles for any sample of a pure substance.

## Examples of Mole Calculations

Let’s practice mole calculations for some other common substances:

**Carbon dioxide CO2**

Molecular weight = 44.01 g/mol

If we have 22.0 g of CO2, the moles are:

moles CO2 = 22.0 g / 44.01 g/mol = 0.500 moles

**Sodium chloride NaCl**

Molecular weight = 58.44 g/mol

For 29.2 g NaCl:

moles NaCl = 29.2 g / 58.44 g/mol = 0.499 moles

**Sucrose C12H22O11 (table sugar)**

Molecular weight = 342.3 g/mol

For 85.6 g sucrose:

moles sucrose = 85.6 g / 342.3 g/mol = 0.250 moles

You can see how this mole calculation method works for any substance where you know its mass and molecular weight. With some practice, you can become skilled at quickly converting between different units using dimensional analysis and appropriate mole ratios.

## Determining Molecular Weight

For simple molecules, you can determine the molecular weight by summing the atomic weights of the constituent atoms. The examples earlier showed this for calculating the molecular weight of water.

For more complex molecules like sucrose with a larger chemical formula, finding the molecular weight is not as straightforward. In these cases, you’ll need to look up the accepted molecular weight in a reference source.

Some common sources for molecular weight data include:

– CRC Handbook of Chemistry and Physics

– Merck Index

– Published literature for the compound

– Chemical manufacturer data and safety sheets

If you can’t locate a reliable source for the molecular weight of a complex compound, you can estimate it by summing the atomic weights of all atoms in the molecular formula. However, this method may not account for interactions between atoms that influence the overall molecular weight.

## Molecular Weight vs. Atomic Weight

It’s easy to confuse molecular weight and atomic weight when you first learn about moles. Just remember:

– Atomic weight is the mass of an individual atom of an element. It is based on the number of protons and neutrons in the nucleus.

– Molecular weight is the mass of a molecule of a compound or element. It sums the atomic weights of all constituent atoms.

So atomic weight refers to the mass of a single atom, while molecular weight deals with the mass of molecules consisting of two or more atoms bonded together. Keeping them straight will help you correctly apply mole calculations.

## Molar Mass Conversion Table

Here is a table for reference with the molecular weights of some common substances:

Substance | Molecular Weight (g/mol) |

Water H2O | 18.015 |

Carbon dioxide CO2 | 44.01 |

Sodium chloride NaCl | 58.44 |

Sucrose C12H22O11 | 342.3 |

Methane CH4 | 16.042 |

Acetone C3H6O | 58.08 |

You can use this table to quickly lookup molecular weights for mole calculations. Just be sure the source of the molecular weight is reliable, and confirm any estimates with reference data.

## Limitations of the Mole Concept

While extremely useful, the mole concept does have some limitations:

– Avogadro’s number is an approximation, albeit a very good one. There is some uncertainty in its experimental determination.

– It only applies to pure substances. Normal laboratory samples may contain impurities that affect the calculated mole values.

– The mole concept assumes complete dissociation of molecular compounds into individual atoms and perfectly gas behavior. Real gases deviate somewhat from ideal behavior.

– It ignores interactions between molecules that may alter the effective molecular weight under different conditions.

However, for most basic chemistry purposes, these limitations introduce minimal error. Using Avogadro’s number and molecular weights provide an excellent approximation for working with macroscopic amounts of substances. Within its working domain, the mole concept remains one of the most powerful and useful tools in chemistry.

## Practice Problems

Here are some practice problems to reinforce your understanding of mole calculations:

1) How many moles are present in 25.0 grams of methane, CH4?

2) What mass of water, H2O, contains 0.655 moles?

3) How many carbon dioxide molecules are present in 1.25 moles of CO2?

4) What is the mass in grams of 0.157 moles of glucose, C6H12O6?

5) How many moles are in a 145.5 g sample of uranium oxide, UO2?

Work these out on your own, showing your method clearly. Then check your answers below:

**Answers:**

1) 1.56 moles

2) 11.8 g

3) 7.5 x 1023 molecules

4) 28.2 g

5) 0.485 moles

With practice, you’ll get better at quickly navigating mole calculations. Don’t worry if it feels tricky at first – stick with it and the concepts will become more familiar. Moles are key to success in chemistry problem-solving.

## Conclusion

In summary, here are the key points about calculating moles from mass:

– The mole represents 6.022 x 1023 particles of a substance (Avogadro’s number).

– Moles provide a link between mass and molecular weight.

– Molecular weight of a compound sums the atomic weights of its elements.

– Use the formula moles = mass (g) / molecular weight (g/mol) to convert between mass and moles.

– Practice mole calculations with different substances to get comfortable with the conversions.

– Pay close attention to units and use dimensional analysis to cancel properly.

– Remember the difference between atomic weight and molecular weight.

– Use reference tables to find molecular weights of compounds.

Mastering mole calculations is critical for success in chemistry. With dimensional analysis and a solid understanding of the mole concept, converting between mass and moles becomes straightforward. Practice makes perfect – so be sure to work plenty of different mole calculation problems.