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Does 1.01 grams of hydrogen have a mass equal to 12 grams of carbon? Answered

Hydrogen has a molar mass of 1.01g/mol .
A mole is defined as the amount of substance that contains as many entities as exactly 12g of carbon-12.
Carbon has a molar mass of 12.01 g/mol.

In my textbook it states that "1.01 grams of hydrogen contain 6.02214199 X 10^23 atoms of hydrogen."
Does this mean that 12 grams of carbon contain 6.02214199 X 10^23 atoms of carbon?



3 years ago

1.01 grams of hydrogen weighs the same as 1.01 grams of carbon.

It is that old saying: Which weighs more, a ton of feathers, or a ton of bricks. Both weight the same: one ton.

One mole is equal to Avogadro's Constant, which is that BIG number,

6.02214179 × 1023

So it is just proportions. All that grams/mol thing tells you to do is how to multiply or divide by to convert from one thing to the other. Just like when you are dealing with density! Same concept mathematically!


SOO, here is what you need to know:

1.01 grams of hydrogen * ( 1 mol / 1.01 grams ) = 1 mol of hydrogen.

I am sure you that learned that already, it's simple enough! The only real tricky part is knowing when you need to flip the g/mol, technically speaking, use it's reciprocal. Just remember that you want things to cancel out, and when you have A * something/A = B, then the 'A' goes away


How, you can take that 1 mol, and multiply it to that ridiculously huge number called Avogadro's Constant to literally figure out how many atoms of hydrogen you have!

There is another method that I liked better, and it will be based on proportions, like 2 fractions on either side of a equals sign. It will not look very friendly as text, but write it out on paper, and make the fractions look like fractions, and notice how both fractions have either grams or moles on top, while the other one is at the bottom? Good!

Hygrogen: (1.01 grams / 1 mol) = ( X grams / Y mols)


Hygrogen: (1 mol / 1.01 grams) = ( Y mol / X grams)

Also, you will remember hopefully from prealgebra, that you cross multiply. I know a super quick and easy way to do that, that they do not teach in american schools, because schools in america are designed to confuse people, then push them though with 'no child left behind' policy BS! I will not explain that method here, since it is beyond the scope of the question. I am sure if you search around for it, you can find it.

I will show you the easy way example: Say I have this:


In this case, (after writing it down like normal fractions) pick a random number, like 2 or 3 (you will see why in a minute), multiply it by whatever is diagonal from it, so if you pick 2, then pick 3, and vice versa. Then divide by the last number left, in this case, 4. So that (2*3)/4=x Just punch that into the calculator and you are good to go!

But ... I do that in my grey cells which like the fact that Tartaglia died in a fencing duel and like the sound of Avogadro's name but they don't care for i maginary numbers no matter how useful they might be... . .... . ..... .........

You are a recipient of Random nerves that fire in my brain while I await HQ to FIX a BUG that prevents email of comments to my designated account.

BTW Tartaglia was a scientist mathematician like Avogadro but also played a killer sport !

Were all waiting for them to fix it. I have emailed, posted bug reports, etc

The unit of mass is the gram/gramme - two items with different amounts of grams have different masses.

Both 1.01g hydrogen and 12g carbon contain the same number of atoms, but the hydrogen atoms are 12 times less massive.

Just to complicate things for you, the hydrogen does not have the same number of discreet particles as the carbon, because hydrogen gas is diatomic - it floats around in covalently-bonded pairs, H2 (that "2" should, of course, be subscript). That means the 6x10^23 atoms of hydrogen are actually in the form of 3x10^23 molecules of hydrogen.


Moles are not about mass, they are about numbers of things. Vis:



The number of moles (or mol) of a chemical is a measure proportional to the number of actual molecules, or atoms, in the sample.

Since this is chemistry class, imagine that you have two glass beakers in front of you. Each contains a clear liquid.

One beaker is labeled, "2.5 mol H20", meaning it contains 2.5 mol of water. (Note the molecular weight of water is approximately, 1*16+2*1 = 18 g/mol)

The other beaker is labeled, "2.5 mol CH3CH2OH", meaning it contains 2.5 mol of ethanol. (Note the molecular of ethanol is approximately, 2*12+1*16+6*1 = 46 g/mol)

I claim each beaker contains approximately the same number of molecules, either water (H2O) molecules, or ethanol molecules (CH3CH2OH), respectively. Obviously, I can't tell that just from looking at the samples, because the molecules are too tiny to see, and too numerous to count. However I can infer this (assuming the labels on the beakers are truthful), because moles are proportional to the number of molecules. I know each sample has the same number of mols (i.e. 2.5 mols in this case) which implies each has the same number of molecules.

So which beaker weighs, erm masses, more? I mean excluding the mass of the glass beaker itself.

2.5 mol of water should weigh, have a mass of,
(2.5 mol)*(18 g/mol) = 45 g

2.5 mol of ethanol should weigh, have a mass of,
(2.5 mol)*(46 g/mol) = 115 g

So the beaker containing the 2.5 mol of ethanol should be significantly heavier, and this makes sense considering the individual molecules are heavier.

Also the volume of the ethanol sample, measured in millilitres, will be larger too. If you understand how density works, you can calculate the expected volume, in mL, of each liquid sample.

By the way, I happen to live in a universe where the molecular weights of chemicals are always exact integers. So if you're wondering why my calculations don't have lots of numbers on the right side of the decimal point, that's why. It is because I am just approximating these numbers. For example, I said the molecular weight of water (H2O) was simply: 1+1+16 = 18 g/mol. I think in your universe, which has more precision, the same calculation looks like: 1.00795+1.00795+15.9994 = 18.0153 g/mol

Also, another observation I thought I would mention, is that my approximate molecular weights are simply based on the number of nucleons (i.e protons and neutrons) in a single molecule of that chemical, assuming uniform isotopic composition for the atoms involved. For example, I am imagining a water molecule is made from one oxygen atom (each with 8 protons and 8 neutrons) plus two hydrogen atoms (each with 1 proton and 0 neutrons), for a total of (8+8)+2*(1+0) = 18 nucleons total. Notice that this number, 18, multitplied by (1 g/mol) is approximately equal to the molecular mass for water: 18 g/mol

Most of the mass, in any sample of ordinary matter, is due to the masses of protons and neutrons. However, the precise, several decimal point numbers for atomic masses given on the periodic table, these numbers account for other corrections too, like average isotopic abundance, the mass difference between protons and neutrons, the masses of the electrons too. Also "missing mass", aka binding energy, I think that's in there too. I mention this, just in case you were wondering where all these numbers come from.