Note: this document will print in an appropriately modified format (6 pages)
- The Avogadro constant or Avogadro’s number refers to the number of atoms, molecules, electrons, or ions contained in one mole of a substance. Its value is 6.0221415 × 10 23 mol-1 (the number of atoms or molecules per mole). Amedeo Avogadro, who was an Italian scientist of the early nineteenth century, proposed this constant.
- Updated January 23, 2020 Avogadro's number, or Avogadro's constant, is the number of particles found in one mole of a substance. It is the number of atoms in exactly 12 grams of carbon -12. This experimentally determined value is approximately 6.0221 x 10 23 particles per mole.
Avogadro's Constant is not dimensionless - which is the big difference between the two. The Constant has units of reciprocal quantity. The Constant has units of reciprocal quantity. If that quantity is reciprocal moles, then the numerical value of the Constant and the Number are identical, but the quantity could be something else, for instance. So Avogadro's number is an amount - like atoms in 12 g of carbon - and can mean different things. For example it is also pretty close to the number of atoms in 1 g of hydrogen. This amount is also called a mole. So one mole of carbon weighs 12 g and one mole of hydrogen weighs 1 g.
1 Mole Avogadro
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.
Owing to their tiny size, atoms and molecules cannot be counted by direct observation. But much as we do when 'counting' beans in a jar, we can estimate the number of particles in a sample of an element or compound if we have some idea of the volume occupied by each particle and the volume of the container.
Michael b jordan vote twitter. The latest tweets from @michaelb4jordan. Twitter Erupts In Debate Over Michael B. Jordan’s Suspected Girlfriend Rumors hit the Internet at the top of this year that actor Michael B. Jordan may be off the dating market. The Black Panther.
Once this has been done, we know the number of formula units (to use the most general term for any combination of atoms we wish to define) in any arbitrary weight of the substance. The number will of course depend both on the formula of the substance and on the weight of the sample. But if we consider a weight of substance that is the same as its formula (molecular) weight expressed in grams, we have only one number to know: Avogadro's number, 6.022141527 × 1023, usually designated by NA.
Amadeo Avogadro (1766-1856) never knew his own number!
Avogadro only originated the concept of this number, whose actual value was first estimated by Josef Loschmidt, an Austrian chemistry teacher, in 1895.
You should know it to three significant figures:
NA = 6.02 × 1023
6.02 × 1023 of what? Well, of anything you like: apples, stars in the sky, burritos. But the only practical use for NA is to have a more convenient way of expressing the huge numbers of the tiny particles such as atoms or molecules that we deal with in chemistry. Avogadro's number is a collective number, just like a dozen.
Think of 6.02 × 1023 as the 'chemist's dozen'.
The basic idea (CurtisWang, 4 min) ****
The Mole Explained (utaustinchem, 9 min) ****
The Mole, Avogadro's no., counting by weight (dcaulf, 13½ min) *****
Simple mole calculations (bozemanscience, 5m) ****
Mole and Avogadro's number (Khan, 9½ m) **
The mole concept (IsaacsTeach, 5 min) ****
Mole problems review (LindaHanson, 23 min) ****
Before we get into the use of Avogadro's number in problems, take a moment to convince yourself of the reasoning embodied in the following examples.
Problem Example 1: mass ratio from atomic weightsThe atomic weights of oxygen and of carbon are 16.0 and 12.0, respectively. How much heavier is the oxygen atom in relation to carbon?
Solution: Atomic weights represent the relative masses of different kinds of atoms. This means that the atom of oxygen has a mass that is 16/12 = 4/3 ≈ 1.33 as great as the mass of a carbon atom.
Problem Example 2: Mass of a single atomThe absolute mass of a carbon atom is 12.0 unified atomic mass units (What are these?). How many grams will a single oxygen atom weigh?
Solution: The absolute mass of the carbon atom is 12.0 u,
or 12 × 1.6605 × 10–27 g = 19.9 × 10–27 kg. The mass of the oxygen atom will be 4/3 greater, or 2.66 × 10–26 kg.
Alternatively: (12 g/mol) ÷ (6.022 × 1023 mol–1) × (4/3) = 2.66 × 10–23 g.
Problem Example 3: Relative masses from atomic weightsSuppose that we have N carbon atoms, where N is a number large enough to give us a pile of carbon atoms whose mass is 12.0 grams. How much would the same number, N, of oxygen atoms weigh?
Solution: The mass of an oxygen atom (16 u) is 16/12 = 4/3 that of a carbon atom (12 u), so the collection of N oxygen atoms would have a mass of
4/3 × 12 g = 16.0 g.
Things to understand about Avogadro's number NA
1 Mole Avogadro Sayısı
• It is a number, just as is 'dozen', and thus is dimensionless.
• It is a huge number, far greater in magnitude than we can visualize; see here for some interesting comparisons with other huge numbers.
• Its practical use is limited to counting tiny things like atoms, molecules, 'formula units', electrons, or photons.
• The value of NA can be known only to the precision that the number of atoms in a measurable weight of a substance can be estimated. Because large numbers of atoms cannot be counted directly, a variety of ingenious indirect measurements have been made involving such things as Brownian motion and X-ray scattering.
• The current value was determined by measuring the distances between the atoms of silicon in an ultrapure crystal of this element that was shaped into a perfect sphere. (The measurement was made by X-ray scattering.) When combined with the measured mass of this sphere, it yields Avogadro's number. But there are two problems with this: 1) The silicon sphere is an artifact, rather than being something that occurs in nature, and thus may not be perfectly reproducible. 2) The standard of mass, the kilogram, is not precisely known, and its value appears to be changing. For these reasons, there are proposals to revise the definitions of both NA and the kilogram. See here for more, and stay tuned!
Wikipedia has a good discussion of Avogadro's number
The mole (abbreviated mol) is the the SI measure of quantity of a 'chemical entity', which can be an atom, molecule, formula unit, electron or photon. One mol of anything is just Avogadro's number of that something. Or, if you think like a lawyer, you might prefer the official SI definition:
Some useful mole links
Wikipedia article on the mole
Mystified by the Mole? Stop it!
A short history of Avogadro's Number
National Mole Day site
Avogadro's number NA = 6.02 × 1023, like any pure number, is dimensionless. However, it also defines the mole, so we can also express NA as
6.02 × 1023 mol–1; in this form, it is properly known as Avogadro's constant. This construction emphasizes the role of Avogadro's number as a conversion factor between number of moles and number of 'entities'.
How many moles of nickel atoms are there in 80 nickel atoms?
Solution: (80 atoms) / (6.02E23 atoms mol–1) = 1.33E–22mol
Is this answer reasonable? Yes, because 80 is an extremely small fraction of NA.
Molar mass and its uses (IsaacsTeach, 6½ m) ****
The atomic weight, molecular weight, or formula weight of one mole of the fundamental units (atoms, molecules, or groups of atoms that correspond to the formula of a pure substance) is the ratio of its mass to 1/12 the mass of one mole of C12 atoms, and being a ratio, is dimensionless. But at the same time, this molar mass (as many now prefer to call it) is also the observable mass of one mole (NA) of the substance, so we frequently emphasize this by stating it explicitly as so many grams (or kilograms) per mole: g mol–1.
Don't let this confuse you; it is very important always to bear in mind that the mole is a number and not a mass. But each individual particle has a mass of its own, so a mole of any specific substance will always have a mass unique to that substance.
Problem Example 4: Boron content of boraxBorax is the common name of sodium tetraborate, Na2B4O7. In 20.0 g of borax,
(a) how many moles of boron are present?
(b) how many grams of boron are present?
Solution: The formula weight of Na2B4O7 is (2 × 23.0) + (4 × 10.8) + (7 × 16.0) = 201.2.
a) 20 g of borax contains (20.0 g) ÷ (201 g mol–1) = 0.10 mol of borax, and thus 0.40 mol of B.
b) 0.40 mol of boron has a mass of (0.40 mol) × (10.8 g mol–1) = 4.3 g.
Problem Example 5: Magnesium in chlorophyllThe plant photosynthetic pigment chlorophyll contains 2.68 percent magnesium by weight. How many atoms of Mg will there be in 1.00 g of chlorophyll?
Solution: Each gram of chlorophyll contains 0.0268 g of Mg, atomic weight 24.3.
Number of moles in this weight of Mg: (.0268 g) / (24.2 g mol–1) = 0.00110 mol
Number of atoms: (.00110 mol) × (6.02E23 mol–1) = 6.64E20
Is this answer reasonable? (Always be suspicious of huge-number answers!) Yes, because we would expect to have huge numbers of atoms in any observable quantity of a substance.
Molar volume of a pure substance
This is the volume occupied by one mole of a pure substance. Molar volume depends on the density of a substance and, like density, varies with temperature owing to thermal expansion, and also with the pressure. For solids and liquids, these variables ordinarily have little practical effect, so the values quoted for 1 atm pressure and 25°C are generally useful over a fairly wide range of conditions. This is definitely not the case with gases, whose molar volumes must be calculated for a specific temperature and pressure.
Problem Example 6 : Molar volume of a liquidAvogadro 1 Ubuntu
Methanol, CH3OH, is a liquid having a density of 0.79 g per milliliter. Calculate the molar volume of methanol.
Solution: The molar volume will be the volume occupied by one molar mass (32 g) of the liquid. Expressing the density in liters instead of mL, we have
VM = (32 g mol–1) / (790 g L–1) = 0.0405 L mol–1
Problem Example 7: Radius of a strontium atom
The density of metallic strontium is 2.60 g cm–3. Use this value to estimate the radius of the atom of Sr, whose atomic weight is 87.6.
Solution: The molar volume of Sr is (87.6 g mol–1) / (2.60 g cm–3) = 33.7 cm3 mol–1
The volume of each 'box' is (33.7 cm3 mol–1) / (6.02E23 mol–1) = 5.48E–23 cm3
The side length of each box will be the cube root of this value, 3.79E–8 cm. The atomic radius will be half this value, or 1.9E–8 cm = 1.9E–10 m = 190 pm.
Note: Your calculator probably has no cube root button, but you are expected to be able to find cube roots; you can usually use the xy button with y=0.333. You should also be able estimate the magnitude of this value for checking. The easiest way is to express the number so that the exponent is a multiple of 3. Take 54.8E–24, for example. Since 33=27 and 43 = 64, you know that the cube root of 55 will be between 3 and 4, so the cube root should be a bit less than 4 × 10–8.
So how good is our atomic radius? Standard tables give the atomic radius of strontium is in the range 192-220 pm, depending on how it is defined.
Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially important that you know the precise meanings of all the highlighted terms in the context of this topic.
- 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.
- Find the molar volume of a solid or liquid, given its density and molar mass.
- Explain how the molar volume of a metallic solid can lead to an estimate of atomic diameter.
| Value of NA[1] in various units |
|---|
| 6.02214179(30)×1023 mol−1 |
| 2.73159757(14)×1026 lb-mol−1 |
| 1.707248479(85)×1025 oz-mol−1 |
The Avogadro constant (symbols: L, NA) is the number of particles (usually atoms or molecules) in one mole of a given substance.[2] Its value is equal to 6.02214129(27)×1023 mol−1.[3] The constant was named after the ItalianscientistAmedeo Avogadro.
The measurement of Avogadro's constant was refined in 2011 to 6.02214078×1023 ± 0.00000018×1023.[4]
1 Avogadro Number
An old term closely related to the Avogadro constant is Avogadro's number. Avogadro's number is the number of atoms in 12 grams of the carbonisotopecarbon-12. Avogadro's number is a dimensionless quantity and has the numerical value of the Avogadro constant given in base units.
Related pages[change | change source]
References[change | change source]
- ↑Mohr, Peter J. (2008). 'CODATA Recommended Values of the Fundamental Physical Constants: 2006'(PDF). Rev. Mod. Phys.80: 633–730. Bibcode:2008RvMP..80.633M. doi:10.1103/RevModPhys.80.633.Unknown parameter
|coauthors=ignored (|author=suggested) (help)Direct link to value. - ↑Johnston, Lesley (2008). Salters Advanced Chemistry: Revise Chemistry For Salters AS (Second ed.). Heinemann. p. 2. ISBN978-0-435-63154-3.
- ↑'Avogadro constant'. National Institute of Standards and Technology. Retrieved 2013-11-07.CS1 maint: discouraged parameter (link)
- ↑Andreas, Birk; et al. (2011). 'Determination of the Avogadro Constant by counting the atoms in a 28Si Crystal'. Physical Review Letters. 106 (3). arXiv:1010.2317. Bibcode:2011PhRvL.106c0801A. doi:10.1103/PhysRevLett.106.030801.