6 ms·

It's an arbitrary number, but it's nice because one mole of atoms with atomic mass number X will weigh approximately X grams. This is exactly true for carbon-12

It's an arbitrary number, but it's nice because one mole of atoms with atomic mass number X will weigh approximately X grams. This is exactly true for carbon-12 (and is what defines a mole).

a1369209993 8y ago
> one mole of atoms with atomic mass number X will weigh approximately X grams Speaking as someone who has had to deal with rounding errors in floating-point graphics, data structure layouts, and real estate cartography, that sounds horrifying and insane.
- MrEfficiency 8y ago
  Chemistry doesnt exist in a vacuum(mostly) and usually exists as an equilibrium. The moment you make water, some of that water has turned into CO2 and the water is a bit more acidic. That 'some' can be 1 molecule. Would that affect anything at macro scales? The answer is yes, yes it does. That's why chem engineers get creative in their real world applications.
- ChrisLomont 8y ago
  There is no other way. The problem is atoms combine in integer ratio amounts to form molecules. To mix things for chemistry one needs to be able to mix things in proper proportions. So having a number that trades number of atoms to something plausibly measurable, like mass, is needed. The reason it cannot be exact for all atoms is forced on us by nature: atoms come in isotopes, each weighing slightly differently, and most common elements come in a mix of isotopes. So picking one isotope of one element (carbon-12) as the definition for a mole that is decently representative of how chemists will use the number is a perfectly fine and useful number. Any chemist that needs to worry about the fuzz will understand this and act accordingly. For example, carbon 22 has a mass slightly larger than 22/12 that of carbon 12 (but has short half-life). Carbon 13, which is stable, has mass slightly over 13/12 that of carbon 12, and when using it, one adjusts accordingly. And these "slightly over" phrases are also known to many digits of precision. But nailing down the number precisely is extremely useful.
- nkrisc 8y ago
  Also consider you're unlikely to have a 100% pure sample of whatever it is you're measuring anyway, at least if it's a large enough sample to hold in your hand. Approximations are still useful, even if they're not good enough for all applications.
zeveb 8y ago
But that's a circular definition; one could use a different unit than grams (e.g. grains) and get a different pseudo-mole and hence a different pseudo-Avogadro's number.
syrrim 8y ago
>"how is the mass number defined?" >"that's easy! we take the mass of one mole of atoms, and that's the mass number!"

It's an arbitrary number, but it's nice because one mole of atoms with atomic mass number X will weigh approximately X grams. This is exactly true for carbon-12 (and is what defines a mole).

a1369209993 8y ago

> one mole of atoms with atomic mass number X will weigh approximately X grams Speaking as someone who has had to deal with rounding errors in floating-point graphics, data structure layouts, and real estate cartography, that sounds horrifying and insane.

MrEfficiency 8y ago
Chemistry doesnt exist in a vacuum(mostly) and usually exists as an equilibrium. The moment you make water, some of that water has turned into CO2 and the water is a bit more acidic. That 'some' can be 1 molecule. Would that affect anything at macro scales? The answer is yes, yes it does. That's why chem engineers get creative in their real world applications.
ChrisLomont 8y ago
There is no other way. The problem is atoms combine in integer ratio amounts to form molecules. To mix things for chemistry one needs to be able to mix things in proper proportions. So having a number that trades number of atoms to something plausibly measurable, like mass, is needed. The reason it cannot be exact for all atoms is forced on us by nature: atoms come in isotopes, each weighing slightly differently, and most common elements come in a mix of isotopes. So picking one isotope of one element (carbon-12) as the definition for a mole that is decently representative of how chemists will use the number is a perfectly fine and useful number. Any chemist that needs to worry about the fuzz will understand this and act accordingly. For example, carbon 22 has a mass slightly larger than 22/12 that of carbon 12 (but has short half-life). Carbon 13, which is stable, has mass slightly over 13/12 that of carbon 12, and when using it, one adjusts accordingly. And these "slightly over" phrases are also known to many digits of precision. But nailing down the number precisely is extremely useful.
nkrisc 8y ago
Also consider you're unlikely to have a 100% pure sample of whatever it is you're measuring anyway, at least if it's a large enough sample to hold in your hand. Approximations are still useful, even if they're not good enough for all applications.

zeveb 8y ago

But that's a circular definition; one could use a different unit than grams (e.g. grains) and get a different pseudo-mole and hence a different pseudo-Avogadro's number.

syrrim 8y ago

>"how is the mass number defined?" >"that's easy! we take the mass of one mole of atoms, and that's the mass number!"