## Particles In A Mole

Multiplying by the number of particles in a mole, or dividing by them, doesn’t change the overall ratio. The ratio obtained from the coefficients in a balanced chemical equation is called the mole ratio. What is the mole ratio for the reaction in Model 1? 1 mole of any substance contains 6.022 × 10 23 particles. 6.022 × 10 23 is known as the Avogadro Number or Avogadro Constant and is given the symbol N A(1) N. Portable and easy to use, Particles To Moles study sets help you review the information and examples you need to succeed, in the time you have available. Use your time efficiently and maximize your retention of key facts and definitions with study sets created by other students studying Particles To Moles. Mole Conversions Mole: the amount of a substance that contains 6.02 x 1023 respective particles of that substance Avogadro’s number: 6.02 x 1023 Molar Mass: the mass of one mole of an element CONVERSION FACTORS: 1 mole = 6.02 x 1023 atoms 1 mole = atomic mass (g) Try: 1. How many atoms are in 6.5 moles of zinc? The number of atoms or other particles in a mole is the same for all substances. The mole is related to the mass of an element in the following way: one mole of carbon-12 atoms has 6.02214076 × 10 23 atoms and a mass of 12 grams.

### The Number Of Particles In A Mole Is Known As

### What is Stoichiometry?

Stoichiometry is at the heart of the production of many things you use in your daily life. Soap, tires, fertilizer, gasoline, deodorant, and chocolate bars are just a few commodities you use that are chemically engineered, or produced through chemical reactions. Ryobi 40v battery. Chemically engineered commodities all rely on stoichiometry for their production.

But what is stoichiometry? Stoichiometry is the calculation of quantities in chemical equations. Given a chemical reaction, stoichiometry tells us what quantity of each reactant we need in order to get enough of our desired product. Because of its real-life applications in chemical engineering as well as research, stoichiometry is one of the most important and fundamental topics in chemistry.

### Introduction to the Mole

Which weighs more, 100 pounds of feathers or 100 pounds of bowling balls? You've probably heard this riddle before. Obviously they both weigh the same since I told you I have 100 pounds of each. But if I have 100 pounds of bowling balls and 100 pounds of feathers, do I have more feathers or more bowling balls?The quantities of feathers and bowling balls would not be equal. An individual feather weighs a lot less than a bowling ball. It would take only about 10 bowling balls to get 100 pounds whereas it would take a LOT more feathers.

When you measure quantities in moles, however, the situation is exactly opposite from when you measure according to weight. A mole is defined as the *amount* of a substance. More specifically, there are 6.02×10^{23}particles in a mole of substance. Therefore, if you had 1 mole of feathers and 1 mole of bowling balls, you would have 6.02×10^{23}feathers and 6.02×10^{23} bowling balls. Now suppose you were asked the question, 'Which weighs more, 100 moles of feathers or 100 moles of bowling balls?' The answer this time would be the bowling balls. Although there is an equal number of both feathers and bowling balls, an individual bowling ball weighs much more than an individual feather, and so an equal number of bowling balls would weigh more than an equal number of feathers.

Now, let's return to the number 6.02×10^{23}. This number is known as Avogadro's number and you should definitely commit it to memory. You are probably wondering why it's so large, and it does indeed look intimidating without the exponential notation:

6.02×10^{23} = 60, 200, 000, 000, 000, 000, 000, 000, 000 |

Although you will never have a mole of bowling balls, you will soon be calculating moles of compounds, molecules, atoms, and ions. These representative particles are extremely and incredibly small. That is why there are so many particles in a mole of substance. When you appreciate just how small these particles are, 6.02×10

^{23}stops seeming like such a crazy number.

### How Many Particles In A Mole

Because atoms and molecules are extremely small, there are a great many of them in any macroscopic sample. The 1 cm^{3} of mercury referred to in the introduction to moles would contain 4.080 x 10^{22} mercury atoms, for example, and the 3.47 cm^{3} of bromine would contain twice as many (8.160 x 10^{22}) bromine atoms. The very large numbers involved in counting microscopic particles are inconvenient to think about or to write down. Therefore chemists have chosen to count atoms and molecules using a unit called the mole. One **mole** (abbreviated mol) is 6.022 x 10^{23} of the microscopic particles which make up the substance in question. Thus 6.022 x 10^{23} Br atoms is referred to as 1 mol Br. The 8.160 x 10^{22} atoms in the sample we have been discussing would be

[dfrac {8.160cdot10^{22}} {6.022cdot10^{23}text{ mol Br}} = text {0.1355 mol Br}]

The idea of using a large number as a unit with which to measure how many objects we have is not unique to chemists. Eggs, doughnuts, and many other things are sold by the dozen—a unit of twelve items. Smaller objects, such as pencils, may be ordered in units of 144, that is, by the gross, and paper is packaged in reams, each of which contains 500 sheets. A chemist who refers to 0.1355 mol Br is very much like a bookstore manager who orders 2½ dozen sweat shirts, 20 gross of pencils, or 62 reams of paper.

There is a difference in degree, however, because the chemist’s unit, 6.022 x 10^{23}, is so large. A stack of paper containing a mole of sheets would extend more than a million times the distance from the earth to the sun, and 6.022 x 10^{23} grains of sand would cover all the land in the world to a depth of nearly 2 ft. Obviously there are a great many particles in a mole of anything.

Why have chemists chosen such an unusual number as 6.022 x 10^{23} as the unit with which to count the number of atoms or molecules? Surely some nice round number would be easier to remember. The answer is that *the number of grams in the mass of 1 mol of atoms of any element is the atomic weight of that element*. For example, 1 mol of mercury atoms not only contains 6.022 x 10^{23} atoms, but its mass of 200.59 g is conveniently obtained by adding the unit gram to the Table of Atomic Weights. Some other examples are

[begin{align} &text{1 mol H contains 6.022} times 10^{23} text{H atoms;} & text{its mass is 1.008 g.} &text{1 mol C contains 6.022} times 10^{23} text{C atoms;} &text{its mass is 12.01 g.} &text{1 mol O contains 6.022} times 10^{23} text{O atoms;} &text{its mass is 16.00 g.} &text{1 mol Br contains 6.022} times 10^{23} text{Br atoms;} &text{its mass is 79.90 g.} end{align} ]

Here and in subsequent calculations atomic weights are rounded to two decimal places, unless, as in the case of H, fewer than four significant figures would remain.

The mass of a mole of *molecules* can also be obtained from atomic weights. Just as a dozen eggs will have a dozen whites and a dozen yolks, a mole of CO molecules will contain a mole of C atoms and a mole of O atoms.

The mass of a mole of CO is thus

[ text{Mass of 1 mol C + mass of 1 mol O = mass of 1 mol CO}]

[ text{12.01 g + 16.00 g = 28.01 g}]

The molecular weight of CO (28.01) expressed in grams is the mass of a mole of CO. Some other examples are in Table (PageIndex{1}).

Molecule | Molecular Weight | Mass of 1 Mol of Molecules |
---|---|---|

Br_{2} | 2(79.90) = 159.80 | 159.80 g |

O_{2} | 2(16.00) = 32.00 | 32.00 g |

H_{2}O | 2(1.008) + 16 = 18.02 | 18.02 g |

HgBr_{2} | 200.59 + 2(79.90) = 360.39 | 360.39 g |

Hg_{2}Br_{2} | 2(200.59) + 2(79.90) = 560.98 | 560.98 g |

It is important to specify to what kind of particle a mole refers. A mole of Br atoms, for example, has only half as many atoms (and half as great a mass) as a mole of Br_{2} molecules. It is best not to talk about a mole of bromine without specifying whether you mean 1 mol Br or 1 mol Br_{2}.

## Contributors and Attributions

Ed Vitz (Kutztown University), John W. Moore (UW-Madison), Justin Shorb (Hope College), Xavier Prat-Resina (University of Minnesota Rochester), Tim Wendorff, and Adam Hahn.