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Chapter 9: Problem 36

Write Lewis symbols for the following: a. \(\mathrm{Br}\) b. \(\mathrm{Br}^{-}\) c. Sr d. \(\mathrm{Sr}^{2+}\)

Short Answer

Expert verified
Br: 7 dots; \( \mathrm{Br}^- \): 8 dots in brackets; Sr: 2 dots; \( \mathrm{Sr}^{2+} \): no dots, in brackets.

Step by step solution

01

Understanding Lewis Symbols

Lewis symbols are a representation of an atom鈥檚 valence electrons. The symbol of the element is written with dots around it, each dot representing a valence electron.
02

Writing the Lewis symbol for Br

Bromine (Br) belongs to Group 17 of the periodic table and has 7 valence electrons. Therefore, the Lewis symbol for bromine is written as Br with 7 dots around it.
03

Writing the Lewis symbol for \( \mathrm{Br}^- \)

The bromide ion \( \mathrm{Br}^- \) is formed when bromine gains an additional electron, giving it a total of 8 valence electrons. Thus, the Lewis dot symbol for \( \mathrm{Br}^- \) is Br with 8 dots around it, often with brackets to indicate the ionic charge: [Br]鈦�.
04

Writing the Lewis symbol for Sr

Strontium (Sr) is in Group 2 and has 2 valence electrons. Its Lewis symbol is represented as Sr with 2 dots around it.
05

Writing the Lewis symbol for \( \mathrm{Sr}^{2+} \)

The ion \( \mathrm{Sr}^{2+} \) is formed when strontium loses its 2 valence electrons, resulting in no dots around the symbol. The Lewis symbol for \( \mathrm{Sr}^{2+} \) is Sr with a 2+ charge notation, written as [Sr]虏鈦�.

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Key Concepts

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

Valence Electrons
Valence electrons are the electrons present in the outermost shell of an atom. They play a key role in determining the chemical properties and reactivity of the element. Since valence electrons are involved in forming chemical bonds, understanding them is crucial for understanding chemical behavior.

For example, elements in the same group in the periodic table share the same number of valence electrons, and thus possess similar properties. Elements in Group 1 have 1 valence electron, while those in Group 17 have 7. This pattern helps predict how an element might react with others, based on its valence electron count.
Ionic Charge
An ionic charge is developed when an atom gains or loses electrons, leading to the formation of ions. Atoms seek to attain a stable electron configuration, often resembling the nearest noble gas. This involves full outer electron shells.

  • When an atom loses electrons it becomes positively charged and is called a cation. For instance, strontium (\(\mathrm{Sr}\)) loses 2 electrons to form \(\mathrm{Sr}^{2+}\).
  • When an atom gains electrons, it becomes negatively charged and is known as an anion. Bromine (\(\mathrm{Br}\)), for example, gains an electron to form \(\mathrm{Br}^{-}\).
Ionic charges are crucial for understanding the formation of ionic compounds, where metals typically form positively charged ions, while nonmetals form negatively charged ones.
Group 17 Elements
Group 17 elements, also known as halogens, are a group of elements in the periodic table that include fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). These elements are vital due to their reactivity and role in many biological and industrial applications.

  • They have 7 valence electrons, making them highly reactive as they tend to gain one electron to achieve a full outer shell resembling a noble gas.
  • Halogens are known for forming salts, as their name suggests in Greek ('halo' meaning salt, 'genes' meaning forming).
  • They can form diatomic molecules (e.g., \(\mathrm{F}_2\), \(\mathrm{Cl}_2\)) and are potent oxidizing agents.
Understanding the properties of Group 17 elements is key to predicting their reactions and behavior in various chemical scenarios.
Periodic Table
The periodic table is a structured arrangement of chemical elements, organized by increasing atomic number. It鈥檚 a powerful tool for chemists and scientists as it highlights trends and can predict properties of elements.

  • The elements are placed into rows called periods and columns called groups or families.
  • The position of an element in the table indicates its electron configuration and recurring chemical properties.
  • Groups are particularly significant because elements in the same group share similar valence electron configurations, leading to similar chemical behaviors.
For students, understanding the periodic table helps in quickly determining characteristics of elements, such as whether they are metals, nonmetals, or metalloids, and in predicting their most likely chemical interactions.

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Most popular questions from this chapter

For each of the following, use formal charges to choose the Lewis formula that gives the best description of the electron distribution: a. \(\mathrm{ClO}_{2} \mathrm{~F}\) b. \(\mathrm{SO}_{3}\) c. \(\mathrm{BrO}_{3}^{-}\)

When atoms of the hypothetical element \(\mathrm{X}\) are placed together, they rapidly undergo reaction to form the \(\mathrm{X}_{2}\) molecule: $$\mathrm{X}(g)+\mathrm{X}(g) \longrightarrow \mathrm{X}_{2}(g)$$ a. Would you predict that this reaction is exothermic or endothermic? Explain. b. Is the bond energy of \(X_{2}\) a positive or a negative quantity? Why? c. Suppose \(\Delta H\) for the reaction is \(-500 \mathrm{~kJ} / \mathrm{mol}\). Estimate the bond energy of the \(\mathrm{X}_{2}\) molecule. d. Another hypothetical molecular compound, \(\mathrm{Y}_{2}(g)\), has a bond energy of \(750 \mathrm{~kJ} / \mathrm{mol}\), and the molecular compound \(\mathrm{XY}(g)\) has a bond energy of \(1500 \mathrm{~kJ} / \mathrm{mol}\). Using bond-energy information, calculate \(\Delta H\) for the following reaction. $$\mathrm{X}_{2}(g)+\mathrm{Y}_{2}(g) \longrightarrow 2 \mathrm{XY}(g)$$ e. Given the following information, as well as the information previously presented, predict whether or not the hypothetical ionic compound \(\mathrm{AX}\) is likely to form. In this compound, \(\mathrm{A}\) forms the \(\mathrm{A}^{+}\) cation, and \(\mathrm{X}\) forms the \(\mathrm{X}^{-}\) anion. Be sure to justify your answer. Reaction: \(\mathrm{A}(g)+\frac{1}{2} \mathrm{X}_{2}(g) \longrightarrow \mathrm{AX}(s)\) The first ionization energy of \(\mathrm{A}(\mathrm{g})\) is \(400 \mathrm{~kJ} / \mathrm{mol}\). The electron affinity of \(\mathrm{X}(g)\) is \(-525 \mathrm{~kJ} / \mathrm{mol}\). The lattice energy of \(\mathrm{AX}(s)\) is \(100 \mathrm{~kJ} / \mathrm{mol}\). f. If you predicted that no ionic compound would form from the reaction in part e, what minimum amount of \(\mathrm{AX}(s)\) lattice energy might lead to compound formation?

Explain what energy terms are involved in the formation of an ionic solid from atoms. In what way should these terms change (become larger or smaller) to give the lowest energy possible for the solid?

Give resonance descriptions for the following: a. \(\mathrm{SeO}_{2}\) b. \(\mathrm{N}_{2} \mathrm{O}_{4}\)

Write Lewis formulas for the following: a. \(\mathrm{AlCl}_{4}^{-}\) b. \(\mathrm{AlF}_{6}{ }^{3-}\) c. \(\mathrm{BrF}_{3}\) d. \(\mathrm{IF}_{6}^{+}\)
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