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Chapter 2: Problem 12

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: $$ \begin{array}{ccc} \hline \text { Compound } & \text { Mass of lodine (g) } & \text { Mass of Fluorine (g) } \\ \hline 1 & 4.75 & 3.56 \\ 2 & 7.64 & 3.43 \\ 3 & 9.41 & 9.86 \\ \hline \end{array} $$ (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?

Short Answer

Expert verified
The mass ratios of fluorine to iodine for the three compounds are approximately \(\frac{3}{4}\), \(\frac{1}{2}\), and \(\frac{1}{1}\), which are simple whole number ratios. These ratios support the atomic theory, as it demonstrates that iodine and fluorine combine in fixed, simple, whole number ratios to form different compounds.

Step by step solution

01

Find the mass ratio of Fluorine to Iodine for each compound.

We can find the mass ratio by dividing the mass of fluorine by the mass of iodine in each compound: 1. For Compound 1: Ratio = \(\frac{3.56}{4.75}\) 2. For Compound 2: Ratio = \(\frac{3.43}{7.64}\) 3. For Compound 3: Ratio = \(\frac{9.86}{9.41}\)
02

Calculate the mass ratios.

Now, let's calculate the mass ratios for each compound: 1. For Compound 1: Ratio = \(\frac{3.56}{4.75}\) = 0.749 2. For Compound 2: Ratio = \(\frac{3.43}{7.64}\) = 0.449 3. For Compound 3: Ratio = \(\frac{9.86}{9.41}\) = 1.048 These numbers represent the mass of fluorine per gram of iodine in each compound.
03

Analyze the ratios for their relationship to atomic theory.

The atomic theory suggests that elements combine in simple, whole number ratios to form compounds. Let's analyze the ratios we calculated in step 2: 1. Compound 1: Ratio = 0.749 ≈ \(\frac{3}{4}\) (a simple whole number ratio) 2. Compound 2: Ratio = 0.449 ≈ \(\frac{1}{2}\) (a simple whole number ratio) 3. Compound 3: Ratio = 1.048 ≈ \(\frac{1}{1}\) (a simple whole number ratio) Observing these ratios, we can see that they approximate simple, whole number ratios. This supports the atomic theory, as it shows that iodine and fluorine are combining in fixed, simple, whole number ratios to form different compounds.

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Most popular questions from this chapter

A chemist finds that \(30.82 \mathrm{~g}\) of nitrogen will react with \(17.60 \mathrm{~g}\), \(35.20 \mathrm{~g}, 70.40 \mathrm{~g},\) or \(88.00 \mathrm{~g}\) of oxygen to form four different compounds. (a) Calculate the mass of oxygen per gram of nitrogen in each compound. (b) How do the numbers in part (a) support Dalton's atomic theory?

The natural abundance of \({ }^{3} \mathrm{He}\) is \(0.000137 \% .\) (a) How many protons, neutrons, and electrons are in an atom of \({ }^{3}\) He? (b) Based on the sum of the masses of their subatomic particles, which is expected to be more massive, an atom of \({ }^{3}\) He or an atom of \({ }^{3} \mathrm{H}\) (which is also called tritium)? (c) Based on your answer to part (b), what would need to be the precision of a mass spectrometer that is able to differentiate between peaks that are due to \({ }^{3} \mathrm{He}^{+}\) and \({ }^{3} \mathrm{H}^{+}\) ?

Write the chemical formula for each substance mentioned in the following word descriptions (use the front inside cover to find the symbols for the elements you don't know). (a) Zinc carbonate can be heated to form zinc oxide and carbon dioxide. (b) On treatment with hydrofluoric acid, silicon dioxide forms silicon tetrafluoride and water. (c) Sulfur dioxide reacts with water to form sulfurous acid. (d) The substance phosphorus trihydride, commonly called phosphine, is a toxic gas. (e) Perchloric acid reacts with cadmium to form cadmium(II) perchlorate. (f) Vanadium(III) bromide is a colored solid.

Complete the table by filling in the formula for the ionic compound formed by each pair of cations and anions, as shown for the first pair. $$ \begin{array}{|l|c|c|c|c|} \hline \text { Ion } & \mathrm{K}^{+} & \mathrm{NH}_{4}^{+} & \mathrm{Mg}^{2+} & \mathrm{Fe}^{3+} \\ \hline \mathrm{Cl}^{-} & \mathrm{KCl} & & & \\ \hline \mathrm{OH}^{-} & & & & \\ \hline \mathrm{CO}_{3}^{2-} & & & & \\ \hline \mathrm{PO}_{4}{ }^{3-} & & & & \\ \hline \end{array} $$

There are two different isotopes of bromine atoms. Under normal conditions, elemental bromine consists of \(\mathrm{Br}_{2}\) molecules, and the mass of a \(\mathrm{Br}_{2}\) molecule is the sum of the masses of the two atoms in the molecule. The mass spectrum of \(\mathrm{Br}_{2}\) consists of three peaks: $$ \begin{array}{cc} \hline \text { Mass (amu) } & \text { Relative Size } \\ \hline 157.836 & 0.2569 \\ 159.834 & 0.4999 \\ 161.832 & 0.2431 \\ \hline \end{array} $$ (a) What is the origin of each peak (of what isotopes does each consist)? (b) What is the mass of each isotope? (c) Determine the average molecular mass of a \(\mathrm{Br}_{2}\) molecule. (d) Determine the average atomic mass of a bromine atom. (e) Calculate the abundances of the two isotopes.
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