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Chapter 2: Problem 32

Rubidium has two naturally occurring isotopes, rubidium- 85 (atomic mass \(=84.9118\) amu; abundance \(=72.15 \%\) ) and rubidium-87 (atomic mass \(=86.9092\) amu; abundance \(=27.85 \%\) ). Calculate the atomic weight of rubidium.

Short Answer

Expert verified
The atomic weight of rubidium is approximately \(85.494\) amu.

Step by step solution

01

Find the weighted atomic mass for each isotope

To find the weighted atomic mass for each isotope, we need to multiply the atomic mass by its relative abundance. Convert the percentage of relative abundance into a decimal number by dividing it by 100. Weighted atomic mass of rubidium-85: \(84.9118 \times 0.7215\) Weighted atomic mass of rubidium-87: \(86.9092 \times 0.2785\)
02

Calculate the atomic weight of rubidium

Add the weighted atomic masses of both isotopes to find the atomic weight of rubidium. Atomic weight = (weighted atomic mass of rubidium-85) + (weighted atomic mass of rubidium-87) Atomic weight = \((84.9118 \times 0.7215) + (86.9092 \times 0.2785)\)
03

Evaluate the expression

Perform the calculations to find the atomic weight of rubidium. Atomic weight = \((61.2782083) + (24.215607\)) Atomic weight = \(85.4938153\)
04

Round the result

Round the result to a reasonable number of significant figures. In this case, we can round to 5 significant figures. The atomic weight of rubidium is approximately \(85.494\) amu.

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Most popular questions from this chapter

In a series of experiments, a chemist prepared three different compounds that contain only iodine and fluorine and determined the mass of each element in each compound: \begin{tabular}{ccc} \hline Compound & Mass of Iodine (g) & Mass of Fluorine (g) \\ \hline 1 & \(4.75\) & \(3.56\) \\ 2 & \(7.64\) & \(3.43\) \\ 3 & \(9.41\) & \(9.86\) \\ \hline \end{tabular} (a) Calculate the mass of fluorine per gram of iodine in each compound. (b) How do the numbers in part (a) support the atomic theory?

Answer the following questions without referring to Table 2.1: (a) What are the main subatomic particles that make up the atom? (b) What is the relative charge (in multiples of the electronic charge) of each of the particles? (c) Which of the particles is the most massive? (d) Which is the least massive?

The elements of group 4 A show an interesting change in properties moving down the group. Give the name and chemical symbol of each element in the group, and label it as a nonmetal, metalloid, or metal.

Give the chemical names of each of the following familiar compounds: (a) \(\mathrm{NaCl}\) (table salt), (b) \(\mathrm{NaHCO}_{3}\) (baking soda), (c) \(\mathrm{NaOCl}\) (in many bleaches), (d) \(\mathrm{NaOH}\) (caustic soda), (e) (NH \(_{4}\) ) \(\mathrm{CO}_{3}\) (smelling salts), (f) \(\mathrm{CaSO}_{4}\) (plaster of Paris).

Gallium (Ga) consists of two naturally occurring isotopes with masses of \(68.926\) and \(70.925\) amu. (a) How many protons and neutrons are in the nucleus of each isotope? Write the complete atomic symbol for each, showing the atomic number and mass number. (b) The average atomic mass of Ga is \(69.72\) amu. Calculate the abundance of each isotope.
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