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##### Learning Objectives

- Define Avogadro's number and explain why it is important to know.
- Define the mole. Be able to calculate the number of moles in a given mass of a substance, or the mass corresponding to a given number of moles.
- Define molecular weight, formula weight, and molar mass; explain how the latter differs from the first two.
- Be able to find the number of atoms or molecules in a given weight of a substance.

The chemical changes we observe always involve *discrete numbers of atoms* that rearrange themselves into new configurations. These numbers are HUGE— far too large in magnitude for us to count or even visualize, but they are still *numbers*, and we need to have a way to deal with them. We also need a bridge between these numbers, which we are unable to measure directly, and the weights of substances, which we do measure and observe. The *mole concept* provides this bridge, and is central to all of quantitative chemistry.

## Counting Atoms by the the moles

The mole is an amount unit similar to familiar units like pair, dozen, gross, etc. It provides a specific measure of *the number* of atoms or molecules in a bulk sample of matter. A mole is defined as the amount of substance containing the same number of discrete entities (such as atoms, molecules, and ions) as the number of atoms in a sample of pure ^{12}C weighing exactly 12 g. One Latin connotation for the word “mole” is “large mass” or “bulk,” which is consistent with its use as the name for this unit. The mole provides a link between an easily measured macroscopic property, bulk mass, and an extremely important fundamental property, number of atoms, molecules, and so forth.

The number of entities composing a mole has been experimentally determined to be \(6.02214179 \times 10^{23}\), a fundamental constant named **Avogadro’s number** (\(N_A\)) or the Avogadro constant in honor of Italian scientist Amedeo Avogadro. This constant is properly reported with an explicit unit of “per mole,” a conveniently rounded version being \(6.022 \times 10^{23}/\ce{mol}\).

Consistent with its definition as an amount unit, 1 mole of any element contains the same number of atoms as 1 mole of any other element. The masses of 1 mole of different elements, however, are different, since the masses of the individual atoms are drastically different. The molar mass of an element (or compound) is the mass in grams of 1 mole of that substance, a property expressed in units of grams per mole (g/mol) (Figure \(\PageIndex{1}\)).

Because the definitions of both the mole and the atomic mass unit are based on the same reference substance, ^{12}C, the molar mass of any substance is numerically equivalent to its atomic or formula weight in amu. Per the amu definition, a single ^{12}C atom weighs 12 amu (its atomic mass is 12 amu). According to the definition of the mole, 12 g of ^{12}C contains 1 mole of ^{12}C atoms (its molar mass is 12 g/mol). This relationship holds for all elements, since their atomic masses are measured relative to that of the amu-reference substance, ^{12}C. Extending this principle, the molar mass of a compound in grams is likewise numerically equivalent to its formula mass in amu.

Element | Average Atomic Mass (amu) | Molar Mass (g/mol) | Atoms/Mole |
---|---|---|---|

C | 12.01 | 12.01 | \(6.022 \times 10^{23}\) |

H | 1.008 | 1.008 | \(6.022 \times 10^{23}\) |

O | 16.00 | 16.00 | \(6.022 \times 10^{23}\) |

Na | 22.99 | 22.99 | \(6.022 \times 10^{23}\) |

Cl | 33.45 | 35.45 | \(6.022 \times 10^{23}\) |

While atomic mass and molar mass are numerically equivalent, keep in mind that they are vastly different in terms of scale, as represented by the vast difference in the magnitudes of their respective units (amu versus g).

The relationships between formula mass, the mole, and Avogadro’s number can be applied to compute various quantities that describe the composition of substances and compounds. For example, if we know the mass and chemical composition of a substance, we can determine the number of moles and calculate number of atoms or molecules in the sample. Likewise, if we know the number of moles of a substance, we can derive the number of atoms or molecules and calculate the substance’s mass.

##### Example \(\PageIndex{1}\): Deriving Moles from Grams for an Element

According to nutritional guidelines from the US Department of Agriculture, the estimated average requirement for dietary potassium is 4.7 g. What is the estimated average requirement of potassium in moles?

**Solution**

The mass of K is provided, and the corresponding amount of K in moles is requested. Referring to the periodic table, the atomic mass of K is 39.10 amu, and so its molar mass is 39.10 g/mol. The given mass of K (4.7 g) is a bit more than one-tenth the molar mass (39.10 g), so a reasonable “ballpark” estimate of the number of moles would be slightly greater than 0.1 mol.

The molar amount of a substance may be calculated by dividing its mass (g) by its molar mass (g/mol):

The factor-label method supports this mathematical approach since the unit “g” cancels and the answer has units of “mol:”

\[ \mathrm{4.7\; \cancel{g} K \left ( \dfrac{mol\; K}{39.10\;\cancel{g}}\right)=0.12\;mol\; K} \nonumber\]

The calculated magnitude (0.12 mol K) is consistent with our ballpark expectation, since it is a bit greater than 0.1 mol.

##### Exercise \(\PageIndex{1}\): Beryllium

Beryllium is a light metal used to fabricate transparent X-ray windows for medical imaging instruments. How many moles of Be are in a thin-foil window weighing 3.24 g?

**Answer**-
0.360 mol

##### Example \(\PageIndex{2}\): Deriving Grams from Moles for an Element

A liter of air contains \(9.2 \times 10^{−4}\) mol argon. What is the mass of Ar in a liter of air?

**Solution **

The molar amount of Ar is provided and must be used to derive the corresponding mass in grams. Since the amount of Ar is less than 1 mole, the mass will be less than the mass of 1 mole of Ar, approximately 40 g. The molar amount in question is approximately one-one thousandth (~10^{−3}) of a mole, and so the corresponding mass should be roughly one-one thousandth of the molar mass (~0.04 g):

In this case, logic dictates (and the factor-label method supports) multiplying the provided amount (mol) by the molar mass (g/mol):

\[\mathrm{9.2 \times10^{-4}\; \cancel{mol} \; Ar \left( \dfrac{39.95\;g}{\cancel{mol}\;Ar} \right)=0.037\;g\; Ar} \nonumber\]

The result is in agreement with our expectations, around 0.04 g Ar.

##### Exercise \(\PageIndex{2}\)

What is the mass of 2.561 mol of gold?

**Answer**-
504.4 g

##### Example \(\PageIndex{3}\): Deriving Number of Atoms from Mass for an Element

Copper is commonly used to fabricate electrical wire (Figure \(\PageIndex{2}\)). How many copper atoms are in 5.00 g of copper wire?

**Solution **

The number of Cu atoms in the wire may be conveniently derived from its mass by a two-step computation: first calculating the molar amount of Cu, and then using Avogadro’s number (*N _{A}*) to convert this molar amount to number of Cu atoms:

Considering that the provided sample mass (5.00 g) is a little less than one-tenth the mass of 1 mole of Cu (~64 g), a reasonable estimate for the number of atoms in the sample would be on the order of one-tenth *N _{A}*, or approximately 10

^{22}Cu atoms. Carrying out the two-step computation yields:

\[\mathrm{5.00\:\cancel{g}\:Cu\left(\dfrac{\cancel{mol}\:Cu}{63.55\:\cancel{g}}\right)\left(\dfrac{6.022\times10^{23}\:atoms}{\cancel{mol}}\right)=4.74\times10^{22}\:atoms\: of\: copper}\]

The factor-label method yields the desired cancellation of units, and the computed result is on the order of 10^{22} as expected.

##### Exercise \(\PageIndex{3}\)

A prospector panning for gold in a river collects 15.00 g of pure gold. How many Au atoms are in this quantity of gold?

**Answer**-
\(4.586 \times 10^{22}\; Au\) atoms

## Summary

The formula mass of a substance is the sum of the average atomic masses of each atom represented in the chemical formula and is expressed in atomic mass units. The formula mass of a covalent compound is also called the molecular mass. A convenient amount unit for expressing very large numbers of atoms or molecules is the mole. Experimental measurements have determined the number of entities composing 1 mole of substance to be 6.022 × 10^{23}, a quantity called Avogadro’s number. The mass in grams of 1 mole of substance is its molar mass. Due to the use of the same reference substance in defining the atomic mass unit and the mole, the formula mass (amu) and molar mass (g/mol) for any substance are numerically equivalent (for example, one H_{2}O molecule weighs approximately18 amu and 1 mole of H_{2}O molecules weighs approximately 18 g).

## Glossary

- Avogadro’s number (
*N*)_{A} - experimentally determined value of the number of entities comprising 1 mole of substance, equal to 6.022 × 10
^{23}mol^{−1}

- formula mass
- sum of the average masses for all atoms represented in a chemical formula; for covalent compounds, this is also the molecular mass

- mole
- amount of substance containing the same number of atoms, molecules, ions, or other entities as the number of atoms in exactly 12 grams of
^{12}C

- molar mass
- mass in grams of 1 mole of a substance

## FAQs

### What is the mass of 2.50 x10 4 molecules of nitrogen gas? ›

Finally, multiply the number of moles by the molar mass to find the mass of the nitrogen gas: (4.15 x 10^-20 moles) x (28.02 g/mol) = **1.16 x 10^-18 grams** So, the mass of 2.50 x 10^4 molecules of nitrogen gas is approximately 1.16 x 10^-18 grams.

**How do you determine the molar mass of an element one must know? ›**

The molar mass of any element can be determined by **finding the atomic mass of the element on the periodic table**. For example, if the atomic mass of sulfer (S) is 32.066 amu, then its molar mass is 32.066 g/mol.

**What does molar mass tell you about an element? ›**

Molar mass: The molar mass of an element is the mass in grams of one mole of the element. One mole of an element contains the amount of the element with a mass equal to the element's atomic mass expressed in grams.

**How do you calculate moles? ›**

If you want to know how many moles of a material you have, **divide the mass of the material by its molar mass**. The molar mass of a substance is the mass in grams of one mole of that substance. This mass is given by the atomic weight of the chemical unit that makes up that substance in atomic mass units (amu).

**What is the mass of 2.5 g atom of nitrogen? ›**

Since we need to calculate the mass present in 2.5 g atom of nitrogen, this means first we need to consider the atomic mass of 1 mol of atom of N. 2.5x(14)g = **35g**.

**What is the mass in grams of 2.5 x10 9 h2o molecules? ›**

mass of 2.5 x 10^{9} molecules of H_{2}O = **7.5 x 10 ^{-}^{14} g**.

**What is the mass of 6.02 x10 23 molecules of H2? ›**

The mole is simply a very large number, 6.022×1023 , that has a special property. If I have 6.022×1023 hydrogen atoms, I have a mass of 1 gram of hydrogen atoms . If I have 6.022×1023 H2 molecules, I have a mass of **2 gram** of hydrogen molecules.

**What is the molar mass of all elements? ›**

Element | Average Atomic Mass (amu) | Molar Mass (g/mol) |
---|---|---|

C | 12.01 | 12.01 |

H | 1.008 | 1.008 |

O | 16.00 | 16.00 |

Na | 22.99 | 22.99 |

**Do we calculate molar mass? ›**

**Molar mass is calculated by adding the atomic masses of a given compound**. The periodic table provides the mass of each individual element, denoted beneath the element's symbol. By adding the atomic masses taken from the periodic table, the molar mass can be determined.

**How do you find molar mass without periodic table? ›**

**How to find the molar mass of a compound?**

- Make use of the chemical formula to determine the number of atoms of each element in the compound.
- Multiply the atomic weight of each element with its number of atoms present in the compound.
- Add up all and assign unit as grams/mole.
- Example.

### Why is knowing the molar mass of an atom so important? ›

Molar mass is of great importance when setting up an experiment. If you are testing principles involving specific amounts of a substance, the molar mass **allows you to figure out how much you should weigh out on your scale**. As an example, consider an experiment calling for 2 moles of pure carbon.

**What do you understand by the term molar mass What is its significance? ›**

Molar mass is **the mass of one mole of any substance**. One mole of any species (atoms, molecules, ions or particles) is that quantity in number having a mass equal to its atomic or molecules mass in grams. Molar mass is represented by "M"

**How does the molar mass relate to the periodic table? ›**

The molar mass of an element is the mass in grams of one mole (6.02 x 10^{23} particles) of the element. So in most cases, **to find the molar mass of an element, you just need to look at its atomic mass (atomic weight) on the periodic table**. Remember — that's the number written under the element symbol and element name.

**What is the exact molar mass? ›**

Since 2019, a mole of any substance has been redefined in the SI as the amount of that substance containing an exactly defined number of particles, 6.02214076×10^{23}. The molar mass of a compound in **g/mol** thus is equal to the mass of this number of molecules of the compound in grams.

**What is molar mass and formula mass? ›**

**Formula mass and molar mass are the same**. However, the atomic mass is different from formula mass and molar mass. When we say "formula mass," we usually mean the mass associated with $1$ mole of a formula; i.e. the formula's Avogadro number. The elemental makeup of substances is represented by chemical formulae.