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49

The first thing to realize is that "mole" is not a mass unit. It is simply a quantity - a number - like dozen or gross or score. Just as a dozen eggs is 12 eggs, a mole of glucose is $6.02 \times 10^{23}$ glucose molecules, and a mole of carbon atoms is $6.02 \times 10^{23}$ carbon atoms. "Moles" are only associated with mass because individual objects have ...


33

In accordance with the International System of Units (SI) [Brochure in English, 8th edition, 2006; updated in 2014] and the corresponding International System of Quantities (ISQ) [ISO/IEC 80000 Quantities and units (14 parts)], you can define a suitable new quantity, for example with the quantity name “number of apples” and the quantity symbol “$N_\text{...


25

Simply speaking, because it's an appropriate unit to use. Let's imagine I wanted to measure the length of a rope. What would be an appropriate length to use? Inches? Centimeters? Feet, maybe? It would really be awkward to express it as 0.000189393 miles, or as 304,800,000 nanometers. (Note: if you can't see why these units are awkward, take any page ...


25

Why is the definition of the moles as it is? It is a rather arbitary definition that the mole is the number of atoms in 12g of carbon 12. This has not always been the definition. For example, prior to 1960, the definition was based upon oxygen rather than carbon-12. The first standard was based upon 1 gram of hydrogen. Later, the standard was changed ...


22

Whenever you're looking for accurate fundamental physical constants, CODATA recommended values are the way to go. As of 2015, the latest data for the Avogadro constant is from 2014. According to CODATA, the most accurate value is: $$6.022\ 140\ 857 \times 10^{23}\ \mathrm{mol^{-1}} \pm 0.000\ 000\ 074 \times 10^{23}\ \mathrm{mol^{-1}}\ \ \ \ \rm{(CODATA\ ...


22

You have already mentioned the last definition of the mole before the proposed change: The mole is the amount of 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”. (…) According to this definition, the mole still depends on the definition of the kilogram. The main ...


21

Because the mole is not a convenient basis for people and TV sets (for instance, one mole of people would be approximately 86 trillion (short-scale) times the population of earth, or put another way the earth contains a human population concentration of approximately 12 femtomoles/planet), and because there's no connection back to atomic mass, which is what ...


19

Boron is a covalent solid with high melting point, like diamond (though not quite), and hence its crystals are hard to make. Unlike diamond crystals, they are not nice and probably wouldn't make a great display. The table on http://periodictable.com/Properties/A/MolarVolume.v.log.html seems to corroborate your findings about boron molar volume being the ...


16

This is due to the mass-energy equivalence and a phenomenon called binding energy. Forming a nucleus releases energy because the nucleons are falling into a potential energy well. Due to Einstein's mass energy equivalence this results in the mass of the new nucleus being less than that of the particles that formed it. The binding energy of carbon-12 is ...


16

The term you are looking for is formula unit, I think. Wikipedia doesn't really describe it super well, but just to give an example, you could write the sentence $\pu{58.44 g}$ of $\ce{NaCl}$ contains $6.022 \times 10^{23}$ formula units of sodium chloride and it would be pretty well understood. See also: What is the difference between a Chemical ...


16

Currently, the definition of Avogadro constant depends on the definition of the kilogram, and thus has an uncertainty. The recommended value[1][2] is $$N_\pu{A} = \pu{6.022140857(74) \cdot 10^{23} mol-1}.$$ It will probably be fixed, however, coming May 20th, 2019 on World Metrology Day to exactly[3][4] $$N_\pu{A}^{2019} = \pu{6.02214076 \cdot 10^{23} ...


13

It's pretty simple actually. The molecular weight of a compound is the sum of the atomic weights of all its atoms. The same numbers apply to a single atom as for a mole of that atom, but at the single-molecule level the unit of measure is the atomic mass unit or amu (aka the dalton or Da). The number of grams per mole of an element and the number of daltons ...


13

The fundamental misunderstandings here are related to two concepts: Conservation (atoms are conserved, molecules are not) A mole is just a number (like the chemist's dozen) I would start with conservation, then move on to moles. Explaining why atoms are conserved, but not moles For second-graders, hands-on activities are probably best, so molecular ...


13

You are correct, but to make it a little more clear you can include the assumed "atom" in the denominator of amu: $$ \begin{align} m_{\ce{C}^{12}} &= \pu{12amu atom^-1} \\ \\ m_{\ce{C}^{12}} &= \pu{12g mol^-1} \\ \\ \pu{12amu atom^-1} &= \pu{12g mol^-1} \\ \\ \pu{1amu atom^-1} &= \pu{1g mol^-1} \end{align} $$ In other words, the ratio of ...


13

The mass scale has changed over time, largely due to different isotopes of the "baseline." Not surprisingly, there's a good Wikipedia article on the matter. In the 20th century, until the 1960s chemists and physicists used two different atomic-mass scales. The chemists used a "atomic mass unit" (amu) scale such that the natural mixture of oxygen isotopes ...


13

It is the same reason that a dozen doesn't depend on whether you are counting grapes or elephants. That is how the mole is defined: it is a number, nothing else. The confusion, I suspect, is because of how we measure that number (or, strictly, how we originally measured it as the definition changed recently). The intention of the unit was always to define a ...


12

Simply because the atomic mass is defined as 1/12 of the mass of 12C. Others isotopes of carbon (13C mostly, with an abundance of 1.1% approximately) account for an average atomic mass slightly above 12.


12

Dissenter's answer is not quite right, since, strictly speaking, as defined there, the scientific notation will work for positive numbers only. Obviously, it can be easily fixed to work for negative as well, but the case of 0 will continue to be problematic. Just try to follow this rule to understand why. All in all, it is the matter of definitions, but ...


12

Your calculations are correct. The key is that you are measuring percentage by mass: the mass of oxygen in 2.062 g of $\mathrm{K}\mathrm{Cl}\mathrm{O}_3$ is the same whether it is considered the mass of individual oxygen atoms or of $\mathrm{O}_2$ (or of any other allotrope of oxygen, for that matter). You have half as many moles of $\mathrm{O}_2$, but the ...


11

The Si-unit "mol" is used in chemistry for three reasons: The quantity of atoms/molecules is the logical unit of chemistry (Reactions have to be balanced). Atoms/molecules can't be counted (It works for people and TV's), but have to be measured using mass or volume, and then converted into a quantity (e.g. mol/kg) The quantity of atoms/molecules involved in ...


11

Hint : n-factor of a molecule/compound is defined as the change in oxidation state per molecule. You have correctly calculated the change of one carbon atom as 4. But how many carbon atoms are there in the glucose molecule? Note: The average oxidation state of carbon in glucose is zero while in reality the different carbons have different OS. (...


11

The mole is a base unit as specified in the Système international d’unités (SI) by the bureau international des poids et mesures. Its decisive definition is that published in French: La mole est la quantité de matière d’un système contenant autant d’entités élémentaires qu’il y a d’atomes dans 0,012 kilogramme de carbone 12 ; son symbole est « mol ». ...


10

The problem is that you did not include units with your numbers. If you had, you would have easily seen that you are multiplying $m^3$ and $cm^{−3}$ and coming up with 1 instead of 1000000.


10

$\ce{S6}$, $\ce{S8}$, $\ce{S12}$ – does it make a difference or is it just a trick to make the question more complicated than it is? What you know for sure is: $M(\ce{S}) = 32.065\ \mathrm{g}\cdot \mathrm{mol}^{-1}$ $M(\ce{H2SO4}) = 98.079\ \mathrm{g}\cdot \mathrm{mol}^{-1}$ $\frac{5000\,\mathrm{g}}{32.065\,\mathrm{g}\cdot \mathrm{mol}^{-1}} = 155.93\,\...


9

I think you are confusing unit of measurement with fundamental unit. A unit of measurement is a definite magnitude of a physical quantity. Consider a very common grouping unit, the dozen which is conceptually the same as a mole and is a little easier for us to grasp since it is a commonly used small number. As far as I know, the only requirements for a ...


9

Your statement that the "mole is widely used in chemistry instead of units of mass or volume.." is misleading. There are seven base units of measurement in the SI system and your question alludes to two, mass and amount. The other 5 are length, time, temperature, electrical current, and luminosity. Consider going out for a typical (American) breakfast of ...


9

Both approaches are correct. Avogadro's number is $6.02214129\times 10^{23}$ and represents the number of carbon-12 atoms in 12 grams of unbound carbon-12 in the ground electronic state. $12$grams$/6.02214129\times 10^{23} = 1.9926467\times 10^{-23}$grams The unified atomic mass unit (u) is $1.660538921 \times 10^{-24}$ grams $12 \times 1.660538921 \...


9

Avogadro's law, which can be written as $V \propto n$, where $V$ is the volume of the gas and $n$ is the amount of substance of the gas (measured in moles), can be thought of as just another manifestation of the ideal gas law rewritten as follows, $$ V = (RT/p) n \, . $$ Consequently, strictly speaking, Avogadro's law is applicable only for a hypothetical ...


9

Mole is just a scale factor I find this description very intuitive: A mole is the amount of pure substance containing the same number of chemical units as there are atoms in exactly 12 grams of carbon-12 (i.e., 6.023 X 1023) I think this should clear out the main part of your confusion. To go into the side questions: It is definitely acceptable to ...


8

Why is weight of 1 mole of substance equal to atomic/molecular mass in grams? According to me, it happens because mole has been defined in such a way. Yes! That is correct. It is defined as the numbers of particles in 12 g of C12. If it were 24 g, instead of 12 g, then the weight of 1 mole of substance would equal 2 times the atomic/molecular mass ...


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