Amedeo Avogadro

Amedeo Avogadro : Accomplishments into Chemistry 1

IntroductionLorenzo Romano Amedeo Carlo Avogadro di Quaregna (9 August 1776 – 9 July 1856) was an Italian savant. He is most noted for his contributions to molecular theory, including what is known as Avogadro's law. In tribute to him, the number of elementary entities (atoms, molecules, ions or other particles) in 1 mole of a substance, 6.02214179(30)×1023, is known as the Avogadro constant.

BiographyAmedeo Avogadro was born in Turin to a noble family of Piedmont, Italy. Following in his family's footsteps, he graduated in ecclesiastical law (age 20) and began to practice law. However, Avogadro also was interested in the natural sciences and in 1800 he began private studies in physics and mathematics. In 1809, he started teaching the natural sciences in a liceo (high school) in Vericelli. It was in Vericelli that Avogadro wrote a memoria (concise note) in which he declared the hypothesis that is now known as Avogadro's law.

Avogadro sent this memoria to De Lamétherie's Journal de Physique, de Chemie et d'Histoire naturelle (Journal of Physics, Chemistry and Natural History) and it was published in the July 14th edition of this journal under the title Essai d'une manière de déterminer les masses relatives des molecules élémentaires des corps, et les proportions selon lesquelles elles entrent dans ces combinaisons ("Essay on Determining the Relative Masses of the Elementary Molecules of Bodies and the Proportions by Which They Enter These Combinations"). In 1814 he published Mémoire sur les masses relatives des molécules des corps simples, ou densités présumées de leur gaz, et sur la constitution de quelques-uns de leur composés, pour servir de suite à l'Essai sur le même sujet ("Note on the Relative Masses of Elementary Molecules, or Suggested Densities of Their Gases, and on the Constituents of Some of Their Compounds, As a Follow-up to the Essay on the Same Subject published in Le Journal de Physique, July 1811”) (Note: In 1811, northern Italy was under the rule of the French Emperor Napoléon Bonaparte.)

In 1820, he became professor of physics at the University of Turin. After the downfall of Napoléon in 1815, northern Italy came under control of this kingdom.

He was active in the revolutionary movements of 1821 against the king of Sardinia (who became ruler of Piedmont with Turin as his capital). As a result, he lost his chair in 1823 (or the university officially declared, it was "very glad to allow this interesting scientist to take a rest from heavy teaching duties, in order to be able to give better attention to his researches").


Eventually, Charles Albert granted a Constitution (Statuto Albertino) in 1848. Well before this, Avogadro had been recalled to the university in Turin in 1833, where he taught for another twenty years.

Little is known about Avogadro's private life, which appears to have been sober and religious. He married Felicita Mazzé and had six children.

Some historians suggest that he sponsored some Sardinian revolutionaries, who were stopped by the announcement of Charles Albert's constitution.

Avogadro held posts dealing with statistics, meteorology, and weights and measures (he introduced the metric system into Piedmont). Avogadro introduced the decimal system in Piedmont and served as a member of the Royal Superior Council on Public Instruction.

In honor of Avogadro's contributions to molecular theory, the number of molecules in one mole was named Avogadro's number, NA or "Avogadro's constant". It is approximately 6.0221415 × 1023. Avogadro's number is used to compute the results of chemical reactions. It allows chemists to determine the exact amounts of substances produced in a given reaction.

Johann Josef Loschmidt first calculated the value of Avogadro's number, often referred to as the Loschmidt number in German-speaking countries (Loschmidt constant now has another meaning).


Accomplishments

Avogadro’s Law

Avogadro's law or Avogadro's hypothesis or Avogadro's principle is a gas law which states that "Equal volumes of ideal or perfect gases, at the same temperature and pressure, contain the same number of particles, or molecules." [5] Thus, the number of molecules in a specific volume of gas is independent of the size or mass of the gas molecules.

As an example, equal volumes of molecular hydrogen and nitrogen would contain the same number of molecules, as long as they are at the same temperature and pressure and observe ideal or perfect gas behavior. In practice, for real gases, the law only holds approximately, but the agreement is close enough for the approximation to be useful.

The law can be stated mathematically as:

where:

V is the volume of the gas.

n is the amount of substance of the gas.

k is a proportionality constant.

The most significant consequence of Avogadro's law is that the ideal gas constant has the same value for all gases. This means that the constant

where:

p is the pressure of the gas

T is the temperature of the gas

has the same value for all gases, independent of the size or mass of the gas molecules.


One mole of an ideal gas occupies 22.40 liters (dm³) at STP, and occupies 24.45 litres at SATP (Standard Ambient Temperature and Pressure = 298K and 1 atm). This volume is often referred to as the molar volume of an ideal gas. Real gases may deviate from this value.

The scientific community did not give great attention to his theory, so Avogadro's hypothesis was not immediately accepted. André-Marie Ampère achieved the same results three years later by another method in his (Sur la détermination des proportions dans lesquelles les corps se combinent d'après le nombre et la disposition respective des molécules dont leurs particules intégrantes sont composées -- On the Determination of Proportions in which Bodies Combine According to the Number and the Respective Disposition of the Molecules by Which Their Integral Particles Are Made), but the same indifference was shown to his theory as well.

Only through studies by Charles Frédéric Gerhardt and Auguste Laurent on organic chemistry was it possible to demonstrate that Avogadro's law explained why the same quantities of molecules in a gas have the same volume.

Unfortunately, related experiments with some inorganic substances showed seeming exceptions to the law. This was finally resolved by Stanislao Cannizzaro, as announced at Karlsruhe Congress in 1860, four years after Avogadro's death. He explained that these exceptions were due to molecular dissociations at certain temperatures, and that Avogadro's law determined not only molecular masses, but atomic masses as well.

In 1911, a meeting in Turin commemorated the hundredth anniversary of the publication of Avogadro's classic 1811 paper. King Victor Emmanuel III attended. Thus, Avogadro's great contribution to chemistry was recognised.

Rudolf Clausius, with his kinetic theory on gases, gave another confirmation of Avogadro's Law. Jacobus Henricus van 't Hoff showed that Avogadro's theory also held in dilute solutions.

Avogadro is hailed as a founder of the atomic-molecular theory.


Avogadro’s Constant

The Avogadro constant (symbols: L, NA) is the number of "elementary entities" (usually atoms or molecules) in one mole, that is (from the definition of the mole), the number of atoms in exactly 12 grams of carbon-12. [3][4] It was originally called Avogadro's number. The 2006 CODATA recommended value is:

The French physicist Jean Perrin in 1909 proposed naming the constant in honour of Avogadro. Perrin would win the 1926 Nobel Prize in Physics, in a large part for his work in determining the Avogadro constant by several different methods.

The value of the Avogadro constant was first indicated by Johann Josef Loschmidt who, in 1865, estimated the average diameter of the molecules in air by a method that is equivalent to calculating the number of particles in a given volume of gas. This latter value, the number density of particles in an ideal gas, is now called the Loschmidt constant in his honour, and is approximately proportional to the Avogadro constant. The connection with Loschmidt is the root of the symbol L sometimes used for the Avogadro constant, and German language literature may refer to both constants by the same name, distinguished only by the units of measurement. [8]

Measurement Methods of Avogadro’s Constant

1- Coulometry

The earliest accurate method to measure the value of the Avogadro constant was based on coulometry. The principle is to measure the Faraday constant, F, which is the electric charge carried by one mole of electrons, and to divide by the elementary charge, e, to obtain the Avogadro constant.

The classic experiment is that of Bowers and Davis at NIST,[12] and relies on dissolving silver metal away from the anode of an electrolysis cell, while passing a constant electric current I for a known time t. If m is the mass of silver lost from the anode and Ar the atomic weight of silver, then the Faraday constant is given by:


The NIST workers devised an ingenious method to compensate for silver that was lost from the anode for mechanical reasons, and conducted an isotope analysis of their silver to determine the appropriate atomic weight. Their value for the conventional Faraday constant is F90 = 96 485.39(13) C/mol, which corresponds to a value for the Avogadro constant of 6.022 1449(78) × 1023 mol–1: both values have a relative standard uncertainty of 1.3 × 10–6.

2- Electron mass method (CODATA)

The CODATA value for the Avogadro constant[13] is determined from the ratio of the molar mass of the electron Ar(e)Mu to the rest mass of the electron me:

The "relative atomic mass" of the electron, Ar(e), is a directly-measured quantity, and the molar mass constant, Mu, is a defined constant in the SI system. The electron rest mass, however, is calculated from other measured constants:[12]

3- X-ray crystal density method

One modern method to calculate the Avogadro constant is to use ratio of the molar volume, Vm, to the unit cell volume, Vcell, for a single crystal of silicon:[14]

The factor of eight arises because there are eight silicon atoms in each unit cell.

The unit cell volume can be obtained by X-ray crystallography; as the unit cell is cubic, the volume is the cube of the length of one side (known as the unit cell parameter, a. In practice, measurements are carried out on a distance known as d220(Si), which is the distance between the planes denoted by the Miller indices {220}, and is equal to a/√8. The 2006 CODATA value for d220(Si) is 192.015 5762(50) pm, a relative uncertainty of 2.8 × 10–8, corresponding to a unit cell volume of 1.601 933 04(13) × 10–28 m3.

The isotope proportional composition of the sample used must be measured and taken into account. Silicon occurs with three stable isotopes – 28Si, 29Si, 30Si – and the natural variation in their proportions is greater than other uncertainties in the measurements. The atomic weight


Ar for the sample crystal can be calculated, as the relative atomic masses of the three nuclides are known with great accuracy. This, together with the measured density ρ of the sample, allows the molar volume Vm to be found by:

where Mu is the molar mass constant. The 2006 CODATA value for the molar volume of silicon is 12.058 8349(11) cm3mol−1, with a relative standard uncertainty of 9.1 × 10–8.[15]

As of the 2006 CODATA recommended values, the relative uncertainty in determinations of the Avogadro constant by the X-ray crystal density method is 1.2 × 10–7, about two and a half times higher than that of the electron mass method.

Confusion Resolving between ”Atoms” and “Molecules”

The greatest problem Avogadro had to resolve was the confusion at that time regarding atoms and molecules. One of his most important contributions was clearly distinguishing one from the other, stating that gases are composed of molecules, and these molecules are composed of atoms. For instance, John Dalton did not consider this possibility. Avogadro did not actually use the word "atom" as the words "atom" and "molecule" were used almost without difference. He believed that there were three kinds of "molecules," including an "elementary molecule" (our "atom"). Also, more attention was given to the definition of mass, as distinguished from weight.


References[1] Wikipedia, the free encyclopedia[2] Mohr, Peter J.; Taylor, Barry N.; Newell, David B. (2008). "CODATA Recommended Values of the Fundamental Physical Constants: 2006". Rev. Mod. Phys. 80: 633–730. RevModPhys.80.633 http://physics.nist.gov/cgi-bin/cuu/Value?na

[3] International Union of Pure and Applied Chemistry (1993). Quantities, Units and Symbols in Physical Chemistry, 2nd edition, Oxford: Blackwell Science. ISBN 0-632-03583-8. http://www.iupac.org/publications/pac/1996/pdf/6804x0957.pdf, retrieved 2006-12-28

[4] International Union of Pure and Applied Chemistry Commission on Atomic Weights and Isotopic Abundances (1992). "Atomic Weight: The Name, Its History, Definition and Units. Pure Appl. Chem. http://www.iupac.org/publications/pac/1992/pdf/6410x1535.pdf. Retrieved 2006-12-28.

[5] Avogadro, Amadeo (1811). "Essai d'une maniere de determiner les masses relatives des molecules elementaires des corps, et les proportions selon lesquelles elles entrent dans ces combinaisons". Journal de Physique 73: 58–76.

[6] Oseen, C.W. (December 10, 1926). Presentation Speech for the 1926 Nobel Prize in Physics.

[7] Loschmidt, J. (1865), "Zur Grösse der Luftmoleküle", Sitzungsberichte der kaiserlichen Akademie der Wissenschaften Wien 52 (2): 395–413.

[8] Virgo, S.E. (1933), "Loschmidt's Number", Science Progress 27: 634–49, http://gemini.tntech.edu/~tfurtsch/scihist/loschmid.html

[9] Kotz, John C.; Treichel, Paul M.; Townsend, John R. (2008), Chemistry and Chemical Reactivity (7th ed.), Brooks/Cole, ISBN 0495387037,

[10] Resolution 3, 14th General Conference of Weights and Measures (CGPM), 1971.

[11] de Bièvre, P.; Peiser, H.S. (1992), "'Atomic Weight'—The Name, Its History, Definition, and Units", Pure Appl. Chem. 64 (10): 1535–43, doi:10.1351/pac199264101535, http://www.iupac.org/publications/pac/1992/pdf/6410x1535.pdf

[12] This account is based on the review in Mohr, Peter J.; Taylor, Barry N. (1999). "CODATA recommended values of the fundamental physical constants: 1998". J. Phys. Chem. Ref. Data 28 (6): 1713–1852. doi:10.1103/RevModPhys.72.351.

[13] Mohr, Peter J.; Taylor, Barry N. (2005). "CODATA recommended values of the fundamental physical constants: 2002". Rev. Mod. Phys. 77 (1): 1–107. doi:10.1103/RevModPhys.77.1.

[14] Mineralogy Database (2000-2005). "Unit Cell Formula". http://webmineral.com/help/CellDimensions.shtml. Retrieved 2007-12-09.

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