1. Bonding… Putting it all Together

2. Modeling Bonds
Individually, sketch a model to represent each type of bond: network covalent, molecular covalent, ionic, and metallic
Use your notes and book if you need help
Partner A, explain your sketch to Partner B. Make sure you explain what it represents and why you drew it the way you did. Explain the role of valence electrons in each model.
Partners switch roles. Discuss any differences between your sketches and ideas.
Update your sketches to include anything you may have missed originally.

3. Some key features of your sketches…
Network Covalent
Valence electrons connect atoms together in all directions like a grid or network
Involve nonmetals
Covalent
Valence electrons are shared between atoms to create individual molecules.
Many individual molecules make up a substance
Involve nonmetals
Individual molecule

4. Some key features of your sketches…
Ionic
Metal atoms give up valence electrons to nonmetal atoms. The metal atom becomes a cation (positive ion) and the nonmetal atom becomes an anion.
Electrostatic attraction holds the oppositely charged ions together in a large crystal lattice.
Bonds between metal and nonmetal
Metal atom nucleus & core electrons
Metallic
Valence electrons are free to move around like a “sea of electrons”
Positive metal nucleii are attracted and held together by the negative “sea of electrons”
Metal atoms only
anion
cation
Sea of electrons

5. Using the models to explain properties
Use your book and notes to answer the remaining questions in the packet.
We will share answers in 30 minutes.
Partner A is the timekeeper. Make sure your team is on task!
Partner B is the reader. Read each question out loud, and work as a team to answer it completely.

6. Properties …
1. Explain why ionic and molecular compounds have different properties.
 Ions are held in ionic compounds by the very strong attractive forces between positive and negative charges. These forces extend throughout the ionic crystal.
The covalent bonds between the atoms within a molecule are also strong. However, the forces of attraction between molecules are not as strong as the forces holding ions together in an ionic compound. (More on intermolecular forces later)

7. Properties …
2. Network covalent substances and ionic substances tend to be solids at room temperature, while molecular covalent substances tend to be gas, liquid, or soft solids. Use your models to explain why.
The forces of attraction between molecules are not as strong as the forces holding ions together in an ionic compound or the covalent bonds holding atoms together in a network covalent compound.

8. Properties …
3. Ionic compounds tend to be hard and brittle. Explain why using your model.
Electrostatic forces hold ions tightly in place in a crystal lattice. These forces make it difficult for the individual ions to move relative to each other causing ionic compounds to be hard.
If an ionic compound is hit with enough force to cause the layers to move relative to each other, then the repulsive forces make the layers separate completely.

9. Properties …
4. For a substance to conduct electricity, it has to allow for the movement of electrically charged particles. Why are metals good conductors of electricity?
 Metals tend to have very few electrons in their valence shell, which means that there are many vacant orbitals in the highest energy levels, as well as sparsely populated d-orbitals just below their highest energy level. These vacant orbitals overlap, and allow the valence electrons of each atom to roam freely throughout the entire metal. The electrons are delocalized, which means they don’t belong to a single atom but move freely about in a “sea of electrons”.   This movement allows the flow of electricity.

10. Properties …
5. In our lab we found that solid ionic compounds do not conduct electricity, but ionic compounds dissolved in water (aq) were good conductors of electricity. Explain why.
 In the solid state, the ions are locked in place by electrostatic forces between the opposite charges. They can’t move about freely, so the solid ionic compound can’t conduct electricity.
When dissolved in water, the water molecules are able to pull the ions out of the crystal lattice, and the ions are able to move about freely within the solution, thus being able to conduct electricity.
PHET Simulation…

11. Properties …
6. In our lab we found that molecular covalent compounds did not conduct electricity. Use your model to explain why.
 In a molecular covalent compound, atoms are bonded covalently to each other in a network or grid. Electrons are being shared in the covalent bonds and are not free to move about. Likewise, atoms are locked into a network by covalent bonds so there aren’t any charged particles that are able to move about.
PHET Simulation

12. Properties …
7. To compare bond strengths in ionic compounds, chemists use lattice energy. What is lattice energy? (see page 178)
Lattice energy is “the energy released when one mole of an ionic crystalline compound is formed from gaseous ions.” It is given as a negative value because this energy is released when the crystal is formed. This is also the amount of energy that it would take pull the ions out of their crystal lattice.

13. Properties …
8. How would you expect the melting points of ionic substances to compare to the melting points of covalent substances. Use your model to explain.
 The melting point of ionic compounds is generally higher than the melting point of a covalent molecular compound. This is because the forces that hold individual molecules together in a solid molecular compound are not as strong as the forces that hold ions together in a crystal lattice.

14. Properties …
9. Table 6-3 (p.179) compares the lattice energies of common ionic compounds. How would you expect their melting points to compare? Why?
 The higher the lattice energy, the higher the melting point. This is because the electrostatic attraction holding the ions together in the crystal lattice is greater, so more energy would need to be added to the ionic compound to separate the ions from the crystal lattice, which means higher melting point.

15. Properties …
10. Which ionic compounds would you expect to be the least soluble in water? Why?
When something dissolves in water, water molecules surround individual ions and separate them from each other. In order to pull the individual ions out of the crystal lattice, the lattice energy must be overcome. CaO and MgO have the highest lattice energy, so I would expect them to be the least soluble in water.

16. Properties …
11. What is a polyatomic ion? Describe the bonding between the atoms that make up the polyatomic ion. What causes the charge?
A polyatomic ion is a group of covalently bonded atoms that have a charge. The atoms are covalently bonded to each other, but due to an excess of electrons (negative) or a shortage of electrons (positive) the entire group has a charge. + H H N H
Charge on the whole group H
Covalent bonds

17. Properties …
12. What is a sea of electrons? Explain completely why metals are able to form a sea of electrons.
 Most metals have very few valence electrons, with many vacant orbital in their highest energy levels. These vacant orbitals overlap allowing the valence electrons to move freely throughout the metal. The electrons are delocalized, which means they don’t belong to any individual atom but can move freely through the empty orbitals of any atom. This forms a “sea of electrons” around the metal atoms (nuclei and core electrons).

18. Properties … 13. Why do metals appear shiny?
 Because metals have many empty overlapping orbital in their highest energy levels, they are able to absorb a wide range of light frequencies. When light is absorbed, electrons are excited to a higher energy level (excited state), this is unstable so the electrons fall back to a lower energy level and release the excess energy as a photon of light, which appears as a shiny luster on the metal surface. This is the same principle that causes spectral lines that Bohr used to explain his model of the atom.

19. Properties …
14. Use your models to explain why most metals are malleable and ductile, but ionic crystals are not.
In an ionic compound, each ion is surrounded by oppositely charged ions and the forces between those oppositely charged ions hold them in place in a crystal lattice. These forces make it difficult for one layer to move relative to another.
In a metal, positive metal nuclei are “floating” in a sea of electrons. The metal nuclei can move relative to each other within this sea of electrons without breaking any bonds.

20. Properties … 15. What determines the strength of a metallic bond?
 The strength of metallic bonds varies with the nuclear charge of the metal atoms and the number of electrons in its electron sea.
16. What is the relationship between metallic bond strength and heat of vaporization?
The amount of heat required to vaporize the metal (convert atoms in the solid state to atoms in the gaseous state) is a measure of the strength of the metallic bonds. Higher heat of vaporization means stronger metallic bonds.