Fundamentals of General, Organic, and Biological Chemistry, 8e (McMurry)
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Chapter 4—Molecular Compounds
Overview Chapter 4 introduces the concept of covalently bonded compounds. The shape and nomenclature of molecular compounds and the use of Lewis structures are explained, along with polarity and nonpolarity, two concepts that will prove very useful for subsequent studies of chemistry.
Introduction for Instructors
• The sharing of electrons between atoms to form covalent bonds is a logical extension of ionic bonding.
• Students will need extra help to grasp the idea of multiple covalent bonds.
• Practice in drawing Lewis structures by following simple rules helps cement student understanding of multiple bonds.
• If students can draw Lewis structures, they probably will be able to predict shapes, but it’s likely to be difficult and they’ll probably need extra help here.
• Like nomenclature of ionic compounds, naming covalent compounds follows simple rules.
• Students fear nomenclature as much as any part of chemistry. Practice in writing formulas and naming compounds will help.
• Once covalent bonding and molecular shapes are understood, characteristics of molecular compounds are predictable.
Chapter Goals
• To understand the general concept of covalent bonding
• To be able to compare structures and properties of ionic and covalent compounds
• To be able to state the octet rule and use it to predict electron configurations of representative elements
• To be able to interpret molecular structures and draw Lewis structures for molecules
• To understand how to use Lewis structures to predict geometry of molecules
• To be able to predict bond and molecular polarity based on electronegativity and molecular geometry
Lecture Outline
4.1 Covalent Bonds
• Covalent bonds are formed by sharing electrons between atoms.
• A group of atoms held together by covalent bonds is called a molecule.
• The octet rule allows us to predict the number of shared electrons.
• The shapes of electron orbitals allows prediction of molecular shape.
4.2 Covalent Bonds and the Periodic Table
• Molecules may be made up of identical or unlike atoms.
• The position of a family of elements in the periodic table helps predict the number of covalent bonds an atom can form.
• In covalent bonds, each atom shares enough electrons to reach a noble gas quota.
• Any element whose name ends in -gen or –ine forms diatomic molecules in the natural state.
4.3 Multiple Covalent Bonds
• Some atoms form double or triple bonds by sharing more than one pair of electrons.
• Multiple covalent bonding is common in organic and biomolecules.
4.4 Coordinate Covalent Bonds
• In some covalent bonds, both electrons come from the same atom.
• The octet rule is still in force.
4.5 Characteristics of Molecular Compounds
• Molecular compounds usually have low melting points and boiling points.
• Molecular compounds are usually nonpolar and insoluble in water.
4.6 Molecular Formulas and Lewis Structures
• Lewis structures provide clues to molecular shape.
• In Lewis structures, only the valence (outermost) electrons are shown.
4.7 Drawing Lewis Structures
• Several groups of elements have characteristic numbers of bonds, allowing predictions of molecular shape.
- Hydrogen and the halogens usually only form one bond to other atoms.
- Oxygen and sulfur usually form two bonds.
- Nitrogen and phosphorus usually form three bonds.
- Carbon often forms four bonds.
• Structures of complex organic or biomolecules are often written in condensed form.
• To draw Lewis structures: Find the total number of valence electrons. Draw a trial structure with lines between atoms. Each line represents a pair of electrons. Add any extra electrons in pairs so that each atom (except hydrogen) has an octet.
- Count electrons. If the central atom doesn’t have an octet, take an electron pair from a neighboring atom to form a multiple bond to the central atom.
4.8 The Shapes of Molecules
• Molecular shape can be predicted from the Lewis structure by using the valence-shell Electron-pair repulsion (VSEPR) model.
- Draw the Lewis structure.
- Count the number of electron charge clouds (number of bonds plus lone pairs).
- Assume that the charge clouds form a structure that allows them to be as far away from each other as possible.
- If there are only two charge clouds, the molecule is linear.
- If there are three clouds, the molecule is shaped like an equilateral triangle.
- If there are four clouds, the molecule is shaped as a regular tetrahedron.
Chemistry in Action
VERY Big molecules
• Very large molecules called polymers abound in nature. These include nucleic acids such as DNA and RNA, starches, and proteins.
• Synthetic polymers are formed by linking monomers such as ethylene, polyethylene, or tetrafluoroethylene.
4.9 Polar Covalent Bonds and Electronegativity
• If the two atoms joined by a covalent bond differ greatly in electronegativity, one will attract electrons more strongly, making the bond polar.
4.10 Polar Molecules
• Some bonds may be polar, making the molecule itself polar.
• Symmetrical molecules may be nonpolar, even though they contain some polar covalent bonds.
Hands-On Chemistry 4.1
• Visualization of molecules can help us understand their properties.
4.11 Naming Binary Molecular Compounds
• Name the least electronegative element first.
• Name the second element, using -ide as an ending.
• If necessary, use a prefix to indicate the number of atoms.
- Use mono- only when necessary to show the number of atoms of the second element.
- Learn the prefixes for numbers from one to ten.
Chemistry in Action
Damascenone by Any Other Name, Would Smell as Sweet
• Formal chemical names are more complicated and difficult to pronounce than common names, but provide more specific information about a compound.
• Simply giving the chemical name of a substance may cause some persons to be alarmed, but even dihydrogen monoxide, by another name, may be a desirable substance.
Lecture Demonstrations
• Use molecular models to represent shapes and relative sizes of molecules.
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• Use a large test tube containing vegetable oil and water to show that polar and nonpolar compounds don’t mix well.
• Dissolve some NaCl in water to show that ionic compounds are often soluble in polar solutions. Try to dissolve NaCl in vegetable oil.
• Use conductivity apparatus to show that nonpolar solvents don’t conduct electricity well.
• Compare the physical states of some common ionic and covalent compounds.
• For an interesting example of the interaction of polar compounds, obtain some powdered polyacrylamide such as Lightning Gel® (www.world-wonders.net) or similar compounds available at garden supply stores and found in extended-wear diapers. Place a teaspoonful of the powder in a plastic foam cup and add water to fill. After you are certain the liquid has been absorbed by the powder, invert the cup over the head of a student. Pretend that you are going to dump the water on him or her. Hydrogen bonding has caused the polyacrylamide to absorb the water; none should be released. A new use for this type of material comes in fighting fires. Polyacrylamide attracts a large amount of water. Powder may be mixed with water and spread atop roofs of endangered homes, protecting them for a few hours from forest fires. Surgical teams also use the material to speed drainage of surgical incisions.
• Obtain a very large beaded metallic chain, similar to that used for light bulb pull chains or ceiling fans, except larger. (Most hardware supply companies can supply chain with individual beads of approximately 6 mm diameter.) Pack about 2 m of chain into a beaker, then pull one end out over the edge and allow gravity to pull the remainder over the edge of the beaker onto a table below. The chain will appear to levitate into the air several centimeters above the edge of the beaker.
Teaching Tips
4.1 Note that ionic compounds usually have high melting points; most liquids or gases are covalent compounds.
4.2 Elements close together in the periodic table usually form covalent bonds when they react.
4.2 It might be useful to point out that few covalent compounds contain metallic atoms.
4.3 The concept of double or triple bonds is hard for students to understand. It may help them to practice drawing single and multiple bonds.
4.7 Students may find it easier to draw simple models of molecules with lines connecting the atoms at first. Remind students that each line represents a pair of electrons.
4.7 Learning the set of simple rules may help students draw Lewis structures.
4.8 It’s hard for students to visualize tetrahedral geometry. Using models may help.
4.11 Ask students to memorize the prefixes for numbers from one to ten. To learn the names and formulas of binary compounds, make a matrix of squares on paper. Write the formulas of anions across the top and the formulas of cations beside squares on the left side. Fill in the squares with names and formulas.
Group Problems
4.100 (i) Oxygen, iodine, and hydrogen form diatomic molecules. (ii) Oxygen, phosphorus, iodine, and hydrogen; Nonmetals form bonds that are mostly covalent in nature when they bond with each other. (iii) Potassium and cesium; Metals form bonds that are mostly ionic in nature when they bond with nonmetals. (iv) Oxygen, phosphorus, iodine, and hydrogen; Nonmetals tend to form covalent bonds when they bond with other nonmetals and ionic bonds when they bond with metals
4.101 H2N–NH2
4.102
4.103 TiBr4: 39 oC; TiO2: 1825 oC; Titanium oxide has a much higher melting point, because bromine has a lower electronegativity value than oxygen. Because there is not as great a difference in electronegativity between Ti and Br, the bonds are more covalent and less ionic in nature in TiBr4. This difference in melting point may be more strongly affected by lattice energy considerations than electronegativity differences. TiBr4 has both a longer bond and a
1 charge, both of which lead to lower lattice energy than TiO2, which has shorter bonds and a –2 charge.
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