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Chapter 17: Problem 617

Distinguish a metallic bond from an ionic bond and from a covalent bond.

Short Answer

Expert verified
In distinguishing metallic, ionic, and covalent bonds, we consider their definitions, the elements involved, electron behavior, and properties. Metallic bonds involve only metal elements with a delocalized "sea of electrons," resulting in high electrical and thermal conductivity, malleability, and ductility. Ionic bonds form between metal and non-metal elements through electron transfer, creating charged ions with electrostatic attraction, resulting in crystalline solids with high melting/boiling points and good conductivity in water/molten state. Covalent bonds share electrons between non-metal elements, resulting in lower melting/boiling points and poor electrical conductivity due to localized electrons.

Step by step solution

01

Definition of each bond

A metallic bond is the attractive force between metal cations and the delocalized electrons in the electron sea. An ionic bond is the electrostatic force of attraction between positively charged (cation) and negatively charged (anion) ions, typically formed between a metal and a non-metal element. A covalent bond is the sharing of electrons between two non-metal atoms.
02

Types of elements involved

In a metallic bond, only metal elements are involved. In an ionic bond, a metal element is combined with a non-metal element, while a covalent bond typically occurs between non-metal elements.
03

Sharing or transferring of electrons

In a metallic bond, electrons are delocalized or shared among many atoms, forming a "sea of electrons" that can move freely. In an ionic bond, electrons are transferred from one atom (metal) to another (non-metal), creating ions with opposite charges that are attracted to each other. In a covalent bond, electrons are shared between two atoms, creating a region of electron density between the nuclei of the atoms involved.
04

Bond properties

Metallic bonds tend to have high electrical and thermal conductivity due to the delocalized electrons. They are also malleable and ductile, enabling metals to be hammered into sheets or drawn into wires. Ionic compounds are typically crystalline solids with high melting and boiling points, poor electrical conductivity in the solid state but good conductivity when dissolved in water or molten. Covalent compounds generally have lower melting and boiling points compared to ionic compounds and tend to be poor conductors of electricity due to the localized nature of their electrons.

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Key Concepts

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

Ionic Bond
An ionic bond forms through the electrostatic attraction between oppositely charged ions. It occurs primarily between a metal and a non-metal. When forming an ionic bond, the metal atom loses electrons to become a positively charged ion, or cation, while the non-metal atom gains those electrons to become a negatively charged ion, or anion.

Some key characteristics of ionic bonds include:
  • Formation of crystalline structures, where ions are arranged in a repeating pattern.
  • High melting and boiling points due to the strong bond formed by the opposite charges.
  • Inability to conduct electricity in solid form but good electrical conductivity when dissolved in water or melted.
These properties make ionic compounds stable and hard, often used in industries that require heat-resistant materials.
Covalent Bond
Covalent bonds are formed by the sharing of electrons between non-metal atoms. These shared electrons allow each atom to attain a stable electron configuration similar to that of noble gases.

The main attributes of covalent bonds include:
  • Lower melting and boiling points compared to ionic compounds.
  • Formation of molecules, which can be simple (e.g., H鈧侽) or complex (e.g., proteins).
  • Generally poor conductors of electricity, as the electrons are localized between the atoms.
Covalent bonds are commonly found in organic compounds, and they are critical in forming the molecules necessary for life.
Electron Delocalization
Electron delocalization occurs in metallic bonds where electrons are not fixed to any one atom but move freely throughout a "sea of electrons." This differs significantly from ionic and covalent bonds, where electrons are either transferred or shared.

By having delocalized electrons, metals exhibit certain unique properties:
  • High electrical and thermal conductivity, as the free-moving electrons can carry energy easily.
  • Malleability and ductility, allowing metals to be shaped without breaking.
  • Shiny appearance, due to the interaction of delocalized electrons with light.
These characteristics make metals immensely useful across industries, from electronics to construction.
Bond Properties
The properties of a bond dictate the behavior and applications of a compound. Different bonds result in various physical and chemical properties, which influence how substances interact.

Here's a breakdown of bond properties:
  • Ionic bonds result in compounds that are generally hard and brittle, with high melting points and conductivity when dissolved.
  • Covalent compounds, by contrast, might be gases or liquids at room temperature and are usually poor conductors of electricity.
  • Metallic bonds enable metals to be excellent conductors and highly durable when shaped and processed.
Understanding the properties of these bonds helps chemists to predict how substances will behave in different environments and can guide the selection of materials for specific purposes.

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Most popular questions from this chapter

Using electron-dot notation, show for each of the following the outer shell electrons for the uncombined atoms and for the molecules or ions that result: (a) \(\mathrm{H}+\mathrm{H} \rightarrow\) Hydrogen molecule (b) \(\mathrm{Br}+\mathrm{Br} \rightarrow\) bromine molecule (c) \(\mathrm{Br}+\mathrm{Cl} \rightarrow\) bromine chloride (d) \(\mathrm{Si}+\mathrm{F} \rightarrow\) silicon fluoride (e) \(\mathrm{Se}+\mathrm{H} \rightarrow\) hydrogen selenide (f) \(\mathrm{Ca}+\mathrm{O} \rightarrow\) calcium oxide

What effect do bond angles have on bond strain? What is the influence of bond strain on bond energy?

The atomic weight of iron is \(55.847\) amu. If one has \(6.02 \mathrm{~g}\) of iron, how many atoms are present?

John Dalton found water to be \(12.5 \%\) hydrogen by weight. Calculate the atomic weight of oxygen if Dalton assumed water contained two hydrogen atoms for every three oxygen atoms. Assume \(1 \mathrm{H}\) atom weighs 1 amu.

It is a fact that the second ionization potential of alkali atoms falls off more rapidly with increasing atomic number than does the first ionization potential. Why?
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