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Chapter 8: Problem 18

Which ionic compound is expected to form from combining the following pairs of elements? (a) calcium and nitrogen, (b) cesium and bromine, (c) strontium and sulfur, (d) aluminum and selenium.

Short Answer

Expert verified
(a) Ca_3N_2, (b) CsBr, (c) SrS, (d) Al_2Se_3

Step by step solution

01

Determine the Ion Charges

First, determine the common ionic charges for the elements involved. For calcium (Ca), the common ion is Ca^{2+}. For nitrogen (N), it often forms N^{3-} ions. Cesium (Cs) typically forms Cs^{+} ions, while bromine (Br) forms Br^{-} ions. Strontium (Sr) forms Sr^{2+}, and sulfur (S) forms S^{2-}. Aluminum (Al) forms Al^{3+}, and selenium (Se) often forms Se^{2-}.
02

Balance the Charges to form Neutral Compounds

For each pair, combine the ions in such a way that the total charge equals zero, forming a neutral compound: (a) Calcium (Ca^{2+}) and Nitrogen (N^{3-}): To balance, use two Ca^{2+} ions with three N^{3-} ions to form Ca_3N_2. (b) Cesium (Cs^{+}) and Bromine (Br^{-}): Use one Cs^{+} and one Br^{-} to form CsBr. (c) Strontium (Sr^{2+}) and Sulfur (S^{2-}): Use one Sr^{2+} and one S^{2-} to form SrS. (d) Aluminum (Al^{3+}) and Selenium (Se^{2-}): Use two Al^{3+} ions with three Se^{2-} ions to form Al_2Se_3.
03

Write the Chemical Formulas

List the chemical formulas for the compounds formed from each element pair: (a) Calcium and Nitrogen form Ca_3N_2. (b) Cesium and Bromine form CsBr. (c) Strontium and Sulfur form SrS. (d) Aluminum and Selenium form Al_2Se_3.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Ion Charges
Understanding ion charges is crucial when forming ionic compounds. When elements lose or gain electrons, they form ions with positive or negative charges, respectively. These charges are predictable based on the element鈥檚 position in the periodic table. For instance:
  • Calcium (Ca) is an alkaline earth metal commonly found in Group 2A, usually forming a ion of \( \text{Ca}^{2+} \).
  • Nitrogen (N)
  • Similarly, Cesium (Cs) in Group 1A forms \( \text{Cs}^{+} \) and Bromine (Br)
These charges determine how elements combine to form stable ionic compounds. Positive ions (cations) attract negative ions (anions), leading to neutral compounds. This attraction is the essence of ionic bonding, where the compound's overall charge remains zero by balancing charges.
Chemical Formulas
Once ion charges are known, we can determine the chemical formulas of ionic compounds. Formulas are written to reflect the smallest neutral constant ratio of ions. For instance:
  • For Calcium and Nitrogen: Calcium forms \( \text{Ca}^{2+} \) and Nitrogen \( \text{N}^{3-} \). To create a neutral compound, three calcium ions bond with two nitrogen ions to yield \( \text{Ca}_3\text{N}_2 \).
  • For Cesium and Bromine: Cesium forms \( \text{Cs}^{+} \) and Bromine \( \text{Br}^{-} \). The ions combine in a 1:1 ratio, forming \( \text{CsBr} \).
The chemical formula reflects the exact ratio needed to balance charges, ensuring the compound is neutral. This vital process underscores the elegance of chemical interactions, guided by simple numeric rules of ion balance.
Balancing Charges
Balancing charges is a key step in forming ionic compounds. To achieve charge neutrality, the positives and negatives must cancel each other. Let's explore this with examples:
  • Calcium and Nitrogen: With \( \text{Ca}^{2+} \) and \( \text{N}^{3-} \), one needs two \( \text{Ca}^{2+} \) ions (total +4) and three \( \text{N}^{3-} \) ions (total -6) to achieve neutrality. This results in the compound \( \text{Ca}_3\text{N}_2 \).
  • Strontium and Sulfur: Both \( \text{Sr}^{2+} \) and \( \text{S}^{2-} \) bind in a simple 1:1 ratio, forming \( \text{SrS} \), perfectly balanced at zero charge.
These balanced equations demonstrate how elements combine based on their natural tendencies, driven by the quest for chemical stability. Through balancing, compounds achieve a stable electronic configuration, mirroring the inert nature of noble gases.

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Most popular questions from this chapter

Which one of these statements about formal charge is true? (a) Formal charge is the same as oxidation number. (b) To draw the best Lewis structure, you should minimize formal charge. (c) Formal charge takes into account the different electronegativities of the atoms in a molecule. (d) Formal charge is most useful for ionic compounds. (e) Formal charge is used in calculating the dipole moment of a diatomic molecule.

You and a partner are asked to complete a lab entitled "Carbonates of Group 2 metal" that is scheduled to extend over two lab periods. The first lab, which is to be completed by your partner, is devoted to carrying out compositional analysis and determine the identity of the Group 2 metal (M). In the second lab, you are to determine the melting point of this compound. Upon going to lab you find two unlabeled vials containing white powder. You also find the following notes in your partner's notebook-Compound \(1: 40.04 \% \mathrm{M}\) and \(12.00 \%\) C, \(47.96 \%\) O (by mass), Compound \(2: 69.59 \%\) M, \(6.09 \% \mathrm{C},\) and \(24.32 \% \mathrm{O}\) (by mass). (a) What is the empirical formula for Compound 1 and the identity of M? (b) What is the empirical formula for Compound 2 and the identity of M? Upon determining the melting points of these two compounds, you find that both compounds do not melt up to the maximum temperature of your apparatus, instead, the compounds decompose and liberate colorless gas. (c) What is the identity of the colorless gas? (d) Write the chemical equation for the decomposition reactions of compound 1 and 2. \((\mathbf{e})\) Are compounds 1 and 2 ionic or molecular?

Under special conditions, sulfur reacts with anhydrous liquid ammonia to form a binary compound of sulfur and nitrogen. The compound is found to consist of \(69.6 \% \mathrm{~S}\) and \(30.4 \% \mathrm{~N}\). Measurements of its molecular mass yield a value of \(184.3 \mathrm{~g} / \mathrm{mol}\). The compound occasionally detonates on being struck or when heated rapidly. The sulfur and nitrogen atoms of the molecule are joined in a ring. All the bonds in the ring are of the same length. (a) Calculate the empirical and molecular formulas for the substance. (b) Write Lewis structures for the molecule, based on the information you are given. (Hint: You should find a relatively small number of dominant Lewis structures.) (c) Predict the bond distances between the atoms in the ring. (Note: The \(S-S\) distance in the \(S_{8}\) ring is 205 pm.) \((\mathbf{d})\) The enthalpy of formation of the compound is estimated to be \(480 \mathrm{~kJ} / \mathrm{mol}^{-1} . \Delta H_{f}^{\circ}\) of \(\mathrm{S}(g)\) is \(222.8 \mathrm{~kJ} / \mathrm{mol}\). Estimate the average bond enthalpy in the compound.

Using Lewis symbols and Lewis structures, make a sketch of the formation of \(\mathrm{NCl}_{3}\) from \(\mathrm{N}\) and \(\mathrm{Cl}\) atoms, showing valence- shell electrons. (a) How many valence electrons does N have initially? (b) How many bonds Cl has to make in order to achieve an octet? (c) How many valence electrons surround the \(\mathrm{N}\) in the \(\mathrm{NCl}_{3}\) molecule? (d) How many valence electrons surround each Cl in the \(\mathrm{NCl}_{3}\) molecule? (e) How many lone pairs of electrons are in the \(\mathrm{NCl}_{3}\) molecule?

Incomplete Lewis structures for the nitrous acid molecule, \(\mathrm{HNO}_{2}\), and the nitrite ion, \(\mathrm{NO}_{2}^{-}\), are shown here. (a) Complete each Lewis structure by adding electron pairs as needed. (b) Is the formal charge on \(\mathrm{N}\) the same or different in these two species? (c) Would either \(\mathrm{HNO}_{2}\) or \(\mathrm{NO}_{2}^{-}\) be expected to exhibit resonance? (d) Would you expect the \(\mathrm{N}=\mathrm{O}\) bond in \(\mathrm{HNO}_{2}\) to be longer, shorter, or the same length as the \(\mathrm{N}-\mathrm{O}\) bonds in \(\mathrm{NO}_{2}^{-}\) ? [Sections 8.5 and 8.6 ] $$ \mathrm{H}-\mathrm{O}-\mathrm{N}=\mathrm{O} \quad \mathrm{O}-\mathrm{N}=\mathrm{O} $$
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