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RP

University of British Columbia

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Problem 103

Answer

a) Avogadr $o^{\prime}$ s number change a little bit from the new definition to become bigger

b) Yes

c) Yes

d) Before 1961, the number would be different

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## Discussion

## Video Transcript

to solve this question. Let's set up an equation. We're told that the atomic mass of oxygen on today scale is 15.99 94 atomic mass units. And we're gonna solve for the atomic massive Carbon, which is gonna be your ex, were told that oxygen will set an equivalent ratio here is going to be 16 a. M. You, which means carbon would be exactly 12 a.m. You so solving for X gives us a, uh massive carbon prior to 1961 equal to 11.999

## Recommended Questions

Prior to 1961, the atomic mass unit was defined as 1/16 the mass of the atomic mass of oxygen; that is, the atomic mass of oxygen was defined as exactly 16 amu. What was the mass of a $^{12} \mathrm{C}$ atom prior to 1961 if the atomic mass of oxygen on today's scale is 15.9994 amu?

The mole is defined as the amount of a substance containing the same number of particles as exactly 12 g of C-12. The amu is defined as 1>12 of the mass of an atom of C-12. Why is it important that both of these definitions reference the same isotope? What would be the result, for example, of defining the mole with respect to C-12, but the amu with respect to Ne-20?

The mole is defined as the amount of a substance containing the same number of particles as exactly 12 $\mathrm{g}$ of $\mathrm{C}-12 .$ The amu is defined as 1$/ 12$ of the mass of an atom of $\mathrm{C}-12 .$ Why is it important that both of these definitions reference the same isotope? What would be the result, for example, of defining the mole with respect to $C-12,$ but the amu with respect to Ne- 20$?$

Although carbon-12 is now used as the standard for atomic weights, this has not always been the case. Early attempts at classification used hydrogen as the standard, with the weight of hydrogen being set equal to 1.0000. Later attempts defined atomic weights using oxygen (with a weight of 16.0000 ). In each instance, the atomic weights of the other elements were defined relative to these masses. (To answer this question, you need more precise data on current atomic weights: $\mathrm{H}, 1.00794 ;$ O, $15.9994 .$)

(a) If $\mathrm{H}=1.0000 \mathrm{u}$ was used as a standard for atomic weights, what would the atomic weight of oxygen be? What would be the value of Avogadro's number under these circumstances?

(b) Assuming the standard is $\mathrm{O}=16.0000$, determine the value for the atomic weight of hydrogen and the value of Avogadro's number.

What was the mass in atomic mass units of a $^{40} \mathrm{Ca}$ atom prior to 1961 if its mass on today’s scale is 39.9626 amu? (See Problem 2.122).

A Although carbon-12 is now used as the standard for atomic masses, this has not always been the case. Farly attempts at classification used hydrogen as the standard, with the mass of hydrogen being set equal to 1.0000 u. Later attempts defined atomic masses using oxygen (with a mass of $16.0000 \mathrm{u}$ ). In each instance, the atomic masses of the other elements were defined relative to these masses. (To answer this question, you need more precise data on current atomic masses: $\mathrm{H}, 1.00794 \mathrm{u} ; \mathrm{O}$ $15.9994 \mathrm{u} .)$

(a) If $\mathrm{H}=1.0000 \mathrm{u}$ was used as a standard for atomic masses, what would the atomic mass of oxygen be? What would be the value of Avogadro's number under these circumstances?

(b) Assuming the standard is $\mathrm{O}=16.0000 \mathrm{u}$, determine the value for the atomic mass of hydrogen and the value of Avogadro's number.

The mole is defined in terms of the carbon-12 atom. Use the definition to find (a) the mass in grams equal to 1 atomic mass unit; (b) the ratio of the gram to the atomic mass unit.

Use the definition that 1 mol of $^{12} \mathrm{C}$ (carbon-12) atoms has a mass of exactly $12 \mathrm{g}$, along with Avogadro's number, to derive the conversion between atomic mass units and kg.

Use the definition of Avogadro's number to find the mass of a helium atom.

Before $1961,$ the standard for atomic masses was the isotope $^{16} \mathrm{O},$ to which physicists assigned a value of exactly $16 .$ At the same time, chemists assigned a value of exactly 16 to the naturally occurring mixture of the isotopes $^{16} \mathrm{O},^{17} \mathrm{O},$ and $^{18} \mathrm{O}$. Would you expect atomic masses listed in a 60 -year-old text to be the same, generally higher, or generally lower than in this text? Explain.