BNG_Chemical Bonds Lesson

Chemical Bonds

image of glueYou learned in the Chemical Reactions Module that when a chemical reaction occurs, some bonds are broken and new ones are formed. Now you are ready to learn more about those bonds. Chemical bonds are attractions between atoms. They are simply attractive forces (between the + nucleus of one atom and the - electrons of a neighboring atom) that hold groups of atoms together and make them function as a unit.  

Types of Bonds

Recall from the Stoichiometry Module that there are two main types of bonds: ionic and covalent. In addition to these main bond types, we will also discuss a third type, metallic bonding. Before we discuss the formation of these bonds in detail, let's go over their general properties.

Go ahead and make a chart of these three types of bonds. As you read through the properties of each type, fill in notes on that specific property. This is an excellent way to organize your notes and organize the information in your head too!

bond chart with the labels: Property lonic Covalent (non network) Metallic State of Matter at Room Temperature Hardness Solid Electrical Conductivity Molten Aqueous Melting and Boiling Points

Properties of Ionic Bonds

Recall that ionic compounds are formed by bonds between metals, which lose electrons to become stable, and nonmetals, which gain electrons.

Properties of Covalent Bonds

Covalent bonds form molecules by sharing electrons. As we go over their properties, remember that these are generalizations with some exceptions. Covalent compounds are usually liquids or gases at room temperature; however, they can can be solids, like sugar. Covalent compounds are softer than ionic compound. The molecules in a covalent compounds can very easily move around each other, unlike ionic compounds where the particles (the ions) are held together in a crystal lattice. As a result, covalent compounds are frequently flexible rather than hard. Covalent compounds don't conduct electricity in solid, liquid, or aqueous form.  

Covalent compounds generally have much lower melting and boiling points than ionic compounds.  Remember that melting and boiling both involve moving molecules farther apart than their previous state. The individual molecules in a covalent compound are generally easy to pull apart from each other because those individual molecules are not attracted to each other (except through weak interactions called intermolecular forces that will be studied in the next module).

Note that some covalent compounds exist in covalent networks. These special compounds have some different properties than normal covalent compounds. They are stronger, harder, and have much higher melting and boiling points. Quartz and diamond are examples of covalent networks. When we talk about covalent compounds, we are referring to regular molecules, not networks, unless specifically told otherwise.

Properties of Metallic Bonds

Metallic bonds are strong electrostatic interactions between metal atoms. They can form between the same type metal atoms, or different metal atoms can join to form an alloy. The valence electrons from the s and p orbitals of the metal atoms delocalize. What this means is that instead of each valence electron moving in the orbitals of their original atom, they move freely in what chemists commonly call a "sea of electrons." This sea of electrons surrounds the positively charged nuclei of all the metal ions. So, metallic bonding refers to the interaction between the delocalized electrons and the metal nuclei.

Metallic bonding can be represented as a lattice of positive metal ions with valence electrons moving as a sea of delocalized electrons as shown below. An animated diagram of this is shown below.

BROKEN OBJECT

The physical properties of metals are the result of the delocalization of the electrons involved in metallic bonding. The electron sea model can be used to explain such properties. Watch the following video and take notes!

Answer the following using the "sea of electrons" model (answers are below):

  1. Why are metals good conductors of electricity?
  2. What determines the strength of a metallic bond?
  3. Why are metals malleable?
  4. Why do metals have a high melting point?
  5. Why are metals shiny?

 

  1. Why are metals good conductors of electricity
    • The delocalized electrons are free to move in the solid lattice.  Theses mobile electrons can act as charge carriers in the conduction of electricity or as energy conductors in the conduction of heat.
  2. What determines the strength of a metallic bond?
    • The strength of a metallic bond depends on the number of electrons in the delocalized "sea".
  3. Why are metals malleable?
    • The delocalized electrons in the 'sea' of electrons in the metallic bond enable the metal atoms to roll over each other when a stress is applied.
  4. Why do metals have a high melting point?
    • In general, metals have high melting and boiling points because of the strength of the metallic bond.
  5. Why are metals shiny? 
    • The free electrons easily absorb and emit the photons of light.

Remember to work on the module practice problems as you complete each section of content.

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