Mole Concept





Mole Concept 3287
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In chemistry the mole is a fundamental unit in the Système International d'Unités, the SI system, and it is used to measure the amount of substance. This quantity is sometimes referred to as the chemical amount. In Latin mole means a "massive heap" of material. It is convenient to think of a chemical mole as such.

Visualizing a mole as a pile of particles, however, is just one way to understand this concept. A sample of a substance has a mass, volume (generally used with gases), and number of particles that is proportional to the chemical amount (measured in moles) of the sample. For example, one mole of oxygen gas (O 2 ) occupies a volume of 22.4 L at standard temperature and pressure (STP; 0°C and 1 atm), has a mass of 31.998 grams, and contains about 6.022 × 10 23 molecules of oxygen. Measuring one of these quantities allows the calculation of the others and this is frequently done in stoichiometry.

The mole is to the amount of substance (or chemical amount) as the gram is to mass. Like other units of the SI system, prefixes can be used with the mole, so it is permissible to refer to 0.001 mol as 1 mmol just as 0.001 g is equivalent to 1 mg.

Formal Definition

According to the National Institute of Standards and Technology (NIST), the Fourteenth Conférence Générale des Poids et Mesures established the definition of the mole in 1971.

The mole is the amount of a substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12; its symbol is "mol." When the mole is used, the elementary entities must be specified and may be atoms, molecules, ions, electrons, other particles, or specified groups of such particles.

One Interpretation: A Specific Number of Particles

When a quantity of particles is to be described, mole is a grouping unit analogous to groupings such as pair, dozen, or gross, in that all of these words represent specific numbers of objects. The main differences between the mole and the other grouping units are the magnitude of the number represented and how that number is obtained. One mole is an amount of substance containing Avogadro's number of particles. Avogadro's number is equal to 602,214,199,000,000,000,000,000 or more simply, 6.02214199 × 10 23 .

Unlike pair, dozen, and gross, the exact number of particles in a mole cannot be counted. There are several reasons for this. First, the particles are too small and cannot be seen even with a microscope. Second, as naturally occurring carbon contains approximately 98.90% carbon-12, the sample would need to be purified to remove every atom of carbon-13 and carbon-14. Third, as the number of particles in a mole is tied to the mass of exactly 12 grams of carbon-12, a balance would need to be constructed that could determine if the sample was one atom over or under exactly 12 grams. If the first two requirements were met, it would take one million machines counting one million atoms each second more than 19,000 years to complete the task.

Obviously, if the number of particles in a mole cannot be counted, the value must be measured indirectly and with every measurement there is some degree of uncertainty. Therefore, the number of particles in a mole, Avogadro's constant ( N A ), can only be approximated through experimentation, and thus its reported values will vary slightly (at the tenth decimal place) based on the measurement method used. Most methods agree to four significant figures, so N A is generally said to equal 6.022 × 10 23 particles per mole, and this value is usually sufficient for solving textbook problems. Another key point is that the formal definition of a mole does not include a value for Avogadro's constant and this is probably due to the inherent uncertainty in its measurement. As for the difference between Avogadro's constant and Avogadro's number, they are numerically equivalent, but the former has the unit of mol −1 whereas the latter is a pure number with no unit.

A Second Interpretation: A Specific Mass

Atoms and molecules are incredibly small and even a tiny chemical sample contains an unimaginable number of them. Therefore, counting the number of atoms or molecules in a sample is impossible. The multiple interpretations of the mole allow us to bridge the gap between the submicroscopic world of atoms and molecules and the macroscopic world that we can observe.

To determine the chemical amount of a sample, we use the substance's molar mass, the mass per mole of particles. We will use carbon-12 as an example because it is the standard for the formal definition of the mole. According to the definition, one mole of carbon-12 has a mass of exactly 12 grams. Consequently, the molar mass of carbon-12 is 12 g/mol. However, the molar mass for the element carbon is 12.011 g/mol. Why are they different? To answer that question, a few terms need to be clarified.

On the Periodic Table, you will notice that most of the atomic weights listed are not round numbers. The atomic weight is a weighted average of the atomic masses of an element's natural isotopes. For example, bromine has two natural isotopes with atomic masses of 79 u and 81 u. The unit u represents the atomic mass unit and is used in place of grams because the value would be inconveniently small. These two isotopes of bromine are present in nature in almost equal amounts, so the atomic weight of the element bromine is 79.904. (i.e., nearly 80, the arithmetic mean of 79 and 81). A similar situation exists for chlorine, but chlorine-35 is almost three times as abundant as chlorine-37, so the atomic weight of chlorine is 35.4527. Technically, atomic weights are ratios of the average atomic mass to the unit u and that is why they do not have units. Sometimes atomic weights are given the unit u , but this is not quite correct according to the International Union of Pure and Applied Chemistry (IUPAC).

To find the molar mass of an element or compound, determine the atomic, molecular, or formula weight and express that value as g/mol. For bromine and chlorine, the molar masses are 79.904 g/mol and 35.4527 g/mol, respectively. Sodium chloride (NaCl) has a formula weight of 58.443 (atomic weight of Na + atomic weight of Cl) and a molar mass of 58.443 g/mol. Formaldehyde (CH 2 O) has a molecular weight of 30.03 (atomic weight of C + 2 [atomic weight of H]) + atomic weight of O] and a molar mass of 30.03 g/mol.

The concept of molar mass enables chemists to measure the number of submicroscopic particles in a sample without counting them directly simply by determining the chemical amount of a sample. To find the chemical amount of a sample, chemists measure its mass and divide by its molar mass. Multiplying the chemical amount (in moles) by Avogadro's constant ( N A ) yields the number of particles present in the sample.

Occasionally, one encounters gram-atomic mass (GAM), gram-formula mass (GFM), and gram-molecular mass (GMM). These terms are functionally the same as molar mass. For example, the GAM of an element is the mass in grams of a sample containing N A atoms and is equal to the element's atomic weight expressed in grams. GFM and GMM are defined similarly. Other terms you may encounter are formula mass and molecular mass. Interpret these as formula weight and molecular weight, respectively, but with the units of u.

Avogadro's Hypothesis

Some people think that Amedeo Avogadro (1776–1856) determined the number of particles in a mole and that is why the quantity is known as Avogadro's number. In reality Avogadro built a theoretical foundation for determining accurate atomic and molecular masses. The concept of a mole did not even exist in Avogadro's time.

Much of Avogadro's work was based on that of Joseph-Louis Gay-Lussac (1778–1850). Gay-Lussac developed the law of combining volumes that states: "In any chemical reaction involving gaseous substances the volumes of the various gases reacting or produced are in the ratios of small whole numbers." (Masterton and Slowinski, 1977, p. 105) Avogadro reinterpreted Gay-Lussac's findings and proposed in 1811 that (1) some molecules were diatomic and (2) "equal volumes of all gases at the same temperature and pressure contain the same number of molecules" (p. 40). The second proposal is what we refer to as Avogadro's hypothesis.

The hypothesis provided a simple method of determining relative molecular weights because equal volumes of two different gases at the same temperature and pressure contained the same number of particles, so the ratio of the masses of the gas samples must also be that of their particle masses. Unfortunately, Avogadro's hypothesis was largely ignored until Stanislao Cannizzaro (1826–1910) advocated using it to calculate relative atomic masses or atomic weights. Soon after the 1 st International Chemical Congress at Karlsrule in 1860, Cannizzaro's proposal was accepted and a scale of atomic weights was established.

To understand how Avogadro's hypothesis can be used to determine relative atomic and molecular masses, visualize two identical boxes with oranges in one and grapes in the other. The exact number of fruit in each box is not known, but you believe that there are equal numbers of fruit in each box (Avogadro's hypothesis). After subtracting the masses of the boxes, you have the masses of each fruit sample and can determine the mass ratio between the oranges and the grapes. By assuming that there are equal numbers of fruit in each box, you then know the average mass ratio between a grape and an orange, so in effect you have calculated their relative masses (atomic masses). If you chose either the grape or the orange as a standard, you could eventually determine a scale of relative masses for all fruit.

A Third Interpretation: A Specific Volume

By extending Avogadro's hypothesis, there is a specific volume of gas that contains N A gas particles for a given temperature and pressure and that volume should be the same for all gases. For an ideal gas, the volume of one mole at STP (0°C and 1.000 atm) is 22.41 L, and several real gases (hydrogen, oxygen, and nitrogen) come very close to this value.

The Size of Avogadro's Number

To provide some idea of the enormity of Avogadro's number, consider some examples. Avogadro's number of water drops (twenty drops per mL) would fill a rectangular column of water 9.2 km (5.7 miles) by 9.2 km (5.7 miles) at the base and reaching to the moon at perigee (closest distance to Earth). Avogadro's number of water drops would cover the all of the land in the United States to a depth of roughly 3.3 km (about 2 miles). Avogadro's number of pennies placed in a rectangular stack roughly 6 meters by 6 meters at the base would stretch for about 9.4 × 10 12 km and extend outside our solar system. It would take light nearly a year to travel from one end of the stack to the other.

History

Long before the mole concept was developed, there existed the idea of chemical equivalency in that specific amounts of various substances could react in a similar manner and to the same extent with another substance. Note that the historical equivalent is not the same as its modern counterpart, which involves electric charge. Also, the historical equivalent is not the same as a mole, but the two concepts are related in that they both indicate that different masses of two substances can react with the same amount of another substance.

The idea of chemical equivalents was stated by Henry Cavendish in 1767, clarified by Jeremias Richter in 1795, and popularized by William Wollaston in 1814. Wollaston applied the concept to elements and defined it in such a way that one equivalent of an element corresponded to its atomic mass. Thus, when Wollaston's equivalent is expressed in grams, it is identical to a mole. It is not surprising then that the word "mole" is derived from "molekulargewicht" (German, meaning "molecular weight") and was coined in 1901 or 1902.

SEE ALSO Avogadro, Amedeo ; Cannizzaro, Stanislao ; Cavendish, Henry ; Gay-Lussac, Joseph-Louis .

Nathan J. Barrows

Bibliography

Atkins, Peter, and Jones, Loretta (2002). Chemical Principles , 2nd edition. New York: W. H. Freeman and Company.

Lide, David R., ed. (2000). The CRC Handbook of Chemistry & Physics , 81st edition. New York: CRC Press.

Masterton, William L., and Slowinski, Emil J. (1977). Chemical Principles , 4th edition. Philadelphia: W. B. Saunders Company.

Internet Resources

National Institute of Standards and Technology. "Unit of Amount of Substance (Mole)." Available from http://www.nist.gov .



User Contributions:

isaac
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Jun 12, 2009 @ 7:07 am
mole concept is the quantity of a substance or the amount of 6.02x10 to the power 23
sameer
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Aug 31, 2009 @ 10:10 am
taken one mole sustance mean to take a amount of sustance in weight equal to its molecular(atomic mass)w weight taken in gram,or in volume equal to(for gases) 22.4 decimeter cube.
e.g. taken 1 mol water equal to 18 gm of water whtch equal to its molecular weight.


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mozhgan
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Dec 14, 2009 @ 5:17 pm
a balanced reaction equation has numbers infront of each substances called coefficient
Benny
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Jan 26, 2010 @ 2:14 pm
Avogadro's number is the number of entities in one mole. As such, it is the mole-to-entity amount ratio. The "mole concept" states that this is identically equal to the gram-to-dalton mass ratio, where the dalton, Da, is another name for the unified atomic mass unit, u. [The dalton is now preferred because it can take SI prefixes (e.g. kilodalton, kDa), whereas a "kilo-unified atomic mass unit" makes no sense.]
Avogadro's number is dimensionless. By contrast, the Avogadro constant is the amount-specific number of entities, i.e. N(S)/n(S), where N(S) is the total number of entities in a sample of a (chemically homogeneous) substance, S, and n(S) is the corresponding amount of substance. The reciprocal of this, the number-specific amount of substance, n(S)/N(S), is the amount of substance of a single entity, i.e. the entity itself. The numerical value of the Avogadro constant depends on the units used for amount of substance. So if NA = N(S)/n(S) and ent = n(S)/N(S), we see that NA = 1 per ent = AN per mole, where AN = g/Da is Avogadro's number.
Avogadro's number is (1/1000) kg/Da. Currently, the kilogram is defined as the mass of the international prototype of the kilogram, m(K); the dalton is one-twelfth the mass of the carbon-12 atom, Da = ma(12C)/12. So AN = (12/1000) m(K)/ma(12C), approximately 6.022141794x10^23.
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Feb 18, 2010 @ 8:08 am
the above matter given is very good and easy to understand very good theory
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Mar 10, 2010 @ 3:03 am
coefficients show the number of moles of two or substances in a balanced equation
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Mar 24, 2010 @ 1:01 am
as like dozen ,mole also a number which is used to number large quantity of particle as one dozen is equals to 12 particles
one mole is equal to 6.023*10^23 particles
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Jun 1, 2010 @ 1:13 pm
wow! dats great i'm really impressed from dis website .dis website is very useful for F.Sc candidates
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Jul 4, 2010 @ 5:05 am
THIS ENCYCLOPEDIA HAS CLARIFIED MY DOUBTS TO A LARGE EXTENT.THANKS
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Aug 7, 2010 @ 4:04 am
how about mole concept about sim card. Can you please tell something about this mole concept?
purtadi
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Sep 22, 2010 @ 12:00 am
i read in journal of chem.edu that the noun "mole" is usually attributed to Wilhelm Ostwald, does it mean that Wilhelm Ostwald is the first who invented that term?
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Sep 27, 2010 @ 9:09 am
i want a answer of a question.why we use carbon 12 in relative atomic mass unit??
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Oct 12, 2010 @ 2:02 am
I WANT TO KNOW HOW 6.022 * 10 23 Came as a result of mole
Meg
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Oct 19, 2010 @ 3:15 pm
How could I explain this concept across the minds of children?
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Nov 9, 2010 @ 9:09 am
COOL DUDE THANKS , I LEARNT IT BETTER IN SCHOOL AFTER BROWSING THROUGH THIS WONDERFULL ARTICLE.
THANKS ONCE AGAIN.
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Nov 13, 2010 @ 6:06 am
why we use carbon 12 in relative atomic mass unit??

Read more: Mole Concept - Chemistry Encyclopedia - reaction, water, elements, examples, gas, number, symbol, mass, atom, Formal Definition, One Interpretation, A Second Interpretation, Avogadros Hypothesis http://www.chemistryexplained.com/Ma-Na/Mole-Concept.html#ixzz15AI7ygiJ
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Dec 9, 2010 @ 7:07 am
I want why we use c-12 in relative atomic mass unit ?
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Feb 8, 2011 @ 11:11 am
WHAT MASS OF CARBON(IV)OXIDE IS PRODUCED BY HEAT STRONGLY 8.4g OF SODIUM(IV)OXIDE IS PRODUCED WHEN EXCESS DILUTE HCL IS ADDED TO THE RESIDUE?
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Feb 16, 2011 @ 1:01 am
give me solutions of diffrent types moleconcept numerical problems
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Mar 5, 2011 @ 7:07 am
we use carbon 12 atom in relative atomic mass unit because atomic weight or atomic mass of the any element of the number which indicates,how many times the atom is heavier than the 1/12 of the carbon 12 atom
george
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Mar 14, 2011 @ 4:16 pm
It seems like there are so many people that contribute to the understanding of the mole today. Who is the person who reintroduced the law and helped it gain universal acceptance.
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Mar 21, 2011 @ 11:11 am
A mole is quantity of a substance that has mass in grams numerically equal to it's molecular mass
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Mar 22, 2011 @ 1:01 am
avagardo number can be termed as the number of atoms inside a single molecule
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Mar 24, 2011 @ 6:06 am
Can any1 pls tell me why do we take C-12 atom as a standard to define a mole
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Mar 31, 2011 @ 8:08 am
in fact, for us candidates are supposed to visit such a website.It is very important in the field of CHEMY. Can you also get us some questions or quiz to answer learn from? :-)
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Apr 1, 2011 @ 6:06 am
itis true though chemistry is not easy but we should strive for success by helping each other since we all use the same sitefor chemistry letis help each other one who is interested can email me
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Apr 11, 2011 @ 5:05 am
why carbon are used in definition of mole ?.a mole is the amount of substance which contain many elementary particles.
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Apr 15, 2011 @ 11:23 pm
What is the use of mole concept in daily life ?

I read in the 'NCERT' that the value of Avagadro Constant was 6.0221367 * 10 to the power 23 ...

Is there any diffrence in what you told and what NCERT says ???
kulbinder
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Apr 27, 2011 @ 9:21 pm
calculate the weight of proton,neutron and electron in terms of a.m.u.
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Jun 26, 2011 @ 1:13 pm
What is the difference between moles and total weight of a substance.
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Jun 30, 2011 @ 1:01 am
is it true that we should take studies as fun not the burden? is chemistry that necessary for our daily needs?
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Jul 4, 2011 @ 4:04 am
highly useful material. As the mole concept is basic step to learn chemistry more and sequence wise information should be added.
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Jul 10, 2011 @ 11:11 am
how would u explain this to a kid? because thats what i need to do!
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Jul 16, 2011 @ 9:21 pm
can u give me the clarity that what is the value of avgadro constant if we take oxygen as reference?
prasanthi
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Aug 5, 2011 @ 12:00 am
Ur article is too good,i got a clear information about mole concept, i suggest u to add few more basic concepts in chemistry in the above way. thank u
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Aug 8, 2011 @ 12:00 am
Its so clear, my teacher could not have explained it this simple
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Aug 12, 2011 @ 4:04 am
My question is why 12g of carbon-12 used as atomic mass scale not any other element? Why is it exactly 12g of carbon-12
samantha
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Sep 9, 2011 @ 3:15 pm
Thanks this helped me a lot with my science project.
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Sep 22, 2011 @ 4:04 am
I have been enriched with some basics on junior secondary school chemistry and as a teacher, there is much to gain after reading this. thanks alot
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Oct 18, 2011 @ 8:20 pm
What a nice, informative article. Now, I'm ready to take our final examination. Thanks.
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Oct 23, 2011 @ 2:02 am
Evryone s puttng cart before horse, three elements were usd in relative atomic mass, 1st hydrogen due to its lihtest nature but prblm with it was we majority of relatve atomic of elements as fractional, den oxygen was choose becoz of reactng with most of elements and relative atomic mass was found as whole no.'s for majority of elements, Again with oxygeon was that naturally occurring oxygeon is a mixture of atoms having different masses known to us as isotopes. Later IUPAC recomends C-12, becoz of forming majority of compounds, gives relative atomic masses of most of elements as whole no.'s and its existance in nature as pure form.
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Oct 26, 2011 @ 3:03 am
1.if one mole of carbon atom weighs 12g.what is the mass in grams of one atom of carbon?
2.which has more number of atoms,100g of socium or 100g ofiron(atomic mass of na=23u,fe=56u).
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Nov 12, 2011 @ 6:06 am
can u tell me that what is a mole?
and what is it's value?
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Nov 17, 2011 @ 7:07 am
In chemistry the mole is a fundamental unit in the Système International d'Unités, the SI system, and it is used to measure the amount of substance. This quantity is sometimes referred to as the chemical amount. In Latin mole means a "massive heap" of material. It is convenient to think of a chemical mole as such.
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Nov 26, 2011 @ 3:03 am
nice definition.. its really help me on my science test!!
Sushil kumar
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Dec 1, 2011 @ 10:10 am
When i was introduced to mole concept.. I was always confused in finding the number of moles in a substance.. But later on i tried to observe it practically and that helped me a lot and then i never confused.. Its better to learn science by observing some things..
kashish katyal
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Dec 29, 2011 @ 3:03 am
yes,well define.it is very irrelevent.my chemistry mam would be prase me
saransh
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Jan 31, 2012 @ 8:08 am
can anyone tell me what space did magnesium cover well now m prepared 4 my examination
Brian
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Feb 6, 2012 @ 7:07 am
wow!this site has the best information that I needed 4 my exams.
vinaya
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Feb 25, 2012 @ 1:01 am
how to determine atomisc mass..?formula has given that at.mass= avg of at.mass of that element/1/12th of mass of corban 12atom.
how we will know abt avg at.mass..?
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Mar 23, 2012 @ 4:04 am
ARE THE NUMBERS OF MOLES ALWAYS EXACT.CALCULATE THE WEIGHT OF PROTONS,NEUTRONS AND ELECTRONS IN THE FORM OF a.m.u
Benny
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Mar 26, 2012 @ 8:20 pm
Look at it this way. If m(X) is the mass of a substance consisting of an aggregate of (chemical) entities of kind X, and ma(X) is the atomic mass of one entity, then the total number of entities, N(X), is m(X)/ma(X). If Ar(X) is the relative atomic mass ("atomic weight"), Ar(X) = ma(X)/Da, where Da is the atomic-scale mass unit, dalton, then the substance mass can be expressed as:
m(X) = N(X)ma(X) = N(X)Ar(X) Da
But we would like this in grams, so we divide and multiply by g:
m(X) = [N(X)/(g/Da)]Ar(X) g
where we see the appearance of the characteristic dimensionless parameter, the gram-to-dalton mass-unit ratio, g/Da. This is where the Avogadro number (not "constant") comes from: Av = g/Da = (0.001 kg)/Da, exactly, regardless of how the kilogram and dalton are actually defined. We now have, in dimensionless form, a relationship between the substance mass and the total number of entities:
m(X)/[Ar(X) g] = N(X)/Av

Now look at the amount of substance, n(X). This is the same as an aggregate of entities. If "ent" represents (the existence of) one entity and there are N(X) entities, then the aggregate is N(X) ent, where we take ent to be an atomic-scale unit for amount of substance. So the amount-of-substance equation is, quite simply:
n(X) = N(X) ent
Rewriting this so that we can relate n(X) to m(X), we must use exactly the same dimensionless parameter, Av, in normalizing N(X):
n(X) = [N(X)/Av][Av ent]
N(X)/Av is a dimensionless quantity of order one. Av ent is a macroscopic constant with the dimension of amount of substance forming a natural macroscopic unit for amount of substance. We call this the mole:
mol = Av ent, exactly = (g/Da) ent, exactly
--one mole is exactly an Avogadro number of entities. [This is not a mass; nor is it a number; it is a (particular) number of entities--tangible stuff.] Note that this means that:
Da/ent = g/mol = kg/kmol, exactly
We now have:
n(X) = N(X)/Av mol

We can now combine the substance-mass and amount-of-substance equations in an easily comprehended dimensionless form:

m(X)/[Ar(X) g] = N(X)/Av = n(X)/mol

For a molecular (or other) substance, we replace Ar(X) by the relative molecular (formula) mass, Mr(X). Note that the "Avogadro constant" has not appeared in any of this--it is not needed! However, in order to understand what it means, we use the definition:
N_A = N(X)/n(X)
--the number of entities per amount of substance. Clearly, from the above, we see that:
N_A = 1/ent (one per entity) = Av/mol (one Avogadro number per mole)

Finally, since the dalton is (currently) defined as one-twelfth the mass of the carbon-12 atom, Da = ma(12C)/12, we see that:
Av = g/Da = (0.001 kg)/[ma(12C)/12] = [0.012 kg]/ma(12C)
This is "the number of atoms in exactly 0.012 kilogram of carbon 12" as used in the statement of the SI mole definition.

Cheers,
Benny.
Benny
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Mar 31, 2012 @ 12:12 pm
What confuses students is simply a matter of "sloppy" terminology by teachers and textbook authors. For example, we often hear (read) that one mole of water is equal to 18 grams of water; it is simultaneously equal to 6.02 x 10^23 molecules of water. Does this mean that a mole is simultaneously a mass and a number? No, it is neither. Does this mean that (1 mole water)/(18 g water) is a conversion factor? No. How about (1 mole water)/(6.02 x 10^23) as a conversion factor? No. Conversion factors must be dimensionless and identically equal to 1.

What is meant, if said (written) correctly, is that a mass of one mole of water is (approximately) 18 g. And the number of molecules comprising one mole of water (or anything else) is (approximately) 6.02 x 10^23. The fundamental stoichiometric equations tell us what we need to know for a substance consisting of an aggregate of entities of kind X:

z(X) = m(X)/[Er(X) g] = N(X)/Av = n(X)/mol

where Er(X) is the relative entity mass, Er(X) = ma(X)/Da [usually written as Ar(X) for an atom or Mr(X) for a molecule, formula unit or other specified kind of entity]; and Av is the Avogadro number, Av = 6.022 141 29 x 10^23, rounded appropriately. If ideal-gas conditions are valid, we can extend the these relationships as follows:

z(X) = [p(X)V(X)/T(X)]/(kAv)

where k is the Boltzmann constant, k = 1.380 6488 x 10^-23 J/K. The product kAv is 8.314 462 J/K. As an exercise, if you substitute atmospheric pressure, patm = 101.325 kPa and the ice-point temperature, Tice = 273.15 K and n = 1 mol (or N = Av or m = Er g)--i.e., z = 1--you can find the corresponding reference volume: Vref = 22.413 968 L.

By the way, the notation "z" comes from a 1905 paper by Albert Einstein; "z" represents "number of" (the German word for number being zahl). Einstein used z as "the number of gram molecules"--a gram molecule being an old name for Mr(X) g [Ar(X) g was a gram atom]; these are parametric mass units, with a different numerical value for each X; also known as "chemical mass units."

Cheers,

Benny.
Aidan chilipamushi
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May 16, 2012 @ 1:01 am
the above passage was well explained,now i have got a clear picture on what mole concept is all about.
sjyrap
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Jun 19, 2012 @ 8:08 am
How can every molecule, atom or ions have measurement 6.023 10 to the power 23? For eg:- CO2 contains 6.023 * ten to the power 23 molecules and oxygen also contains 6.023 ten to the power 23 atoms(1 mole)

Why cant we use relative atomic mass instead of writing mole directly??

If one mole is collection of particle equal in number as no. of atoms in 12 gm carbon then every atom, molecule and ions have same measurement, isn't it??
emelda
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Jul 12, 2012 @ 5:05 am
WOW!!THIS SITE IS THE BEST.PLEASE DEAR LEARNERS USE IT COZ IT IS VERY HELPFUL
marquel
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Jul 16, 2012 @ 2:02 am
yes it is..we just helping each other to know more about MOLE..
himanshu
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Jul 16, 2012 @ 2:14 pm
mole in gram in which equal to amount of avogadro @6.02into10 to the power 23
krazile
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Aug 28, 2012 @ 9:09 am
how we can know the reaction of mola?
with animation.
abcc
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Sep 11, 2012 @ 9:09 am
How we find mass of an element by comparing it with carbon-12 isotope of carbon?
How we have taken out mass of hydrogen 1.008amu by using this scale?
amal wilsen
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Sep 19, 2012 @ 2:14 pm
to prove wrong the dalton's atom theory which stated that an atom cannot be created nor distroy ; During chemical reactions for example in beta emission of particles new atoms are form as aresult of the chemical changes that take place when the atom is activated in reactions
kuffour
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Sep 26, 2012 @ 9:21 pm
i need more notes on mole which is simple and understandable.frm gh college
Samwel Bhoke
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Sep 27, 2012 @ 9:21 pm
The article is so interesting.But how we can apply the knowledge of mole concept in our home places in daily life stuation.
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Sep 30, 2012 @ 2:02 am
To find the chemical amount sample= Measure the mass of substance/molar mass* chemical amount of a mole(6.023*10to the the power 23)
daniellendzana
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Oct 7, 2012 @ 10:10 am
how do calculate RAM of substance if mass is given
Kent Jude
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Oct 24, 2012 @ 5:05 am
yeah. Mole denotes a collection of large fixed number of particles.
Avinash kumar singh
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Nov 28, 2012 @ 8:08 am
this site is very helpful before any practical examination of chemistery
neha
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Dec 16, 2012 @ 5:17 pm
MOLE EQUALS TO AVAGADROS NUMBER THAT IS 6.022*10POWER 23
divya
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Dec 19, 2012 @ 4:04 am
this site gives many things about mole..it is very helpful
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Jan 13, 2013 @ 10:10 am
thank you very much for sharing skills and ideas!!and i would ask to help me in some techniques of teaching chemistry to blind students.for example identification of ions,organic chemistry(displayed formula,etc) etc!!
VIKAS
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Feb 24, 2013 @ 2:14 pm
My question is why 12g of carbon-12 used as atomic mass scale not any other element? Why is it exactly 12g of carbon-12 AND EXPLAIN IN DETAIL
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Apr 22, 2013 @ 12:12 pm
I only explain what i know, we all know that carbon has high combining capacity and thats why it is taken as atomic mass scale, my explanation is short and clear
faith
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Sep 21, 2013 @ 12:12 pm
if in equation N2+3H2 becomes 2NH3 6.72dm^3
which is in excess
prince mopler
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Nov 5, 2013 @ 1:01 am
I like this site,i thnk it will help me to discover more about chemistry concepts
Tshepiso Alec
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Nov 7, 2013 @ 4:04 am
given 100 grams of sulfuric acid, how many atoms of hydrogen are in the sample
Vincent
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Jan 9, 2014 @ 10:10 am
interesting article. thanks for taking time to make it. much appreciated
Galley
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Apr 6, 2014 @ 4:16 pm
The density of (C2H5OH) is 0.8g/cm3. If a sample of this substance contains 3.2*10 to the power 23 molecules, what is the volume of the sample?
zyra mae esperanzate
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Jun 15, 2014 @ 1:01 am
thanks 4 this information about mole concepts (chemistry)

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