THE BIG IDEA: BONDING AND INTERACTIONS.
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1 HONORS CHEMISTRY - CHAPTER 8 COVALENT BONDS OBJECTIVES AND NOTES - PART 2 - V12 NAME: DATE: PAGE: THE BIG IDEA: BONDING AND INTERACTIONS. Essential Questions 1. How is the bonding in molecular compounds different from the bonding in ionic compounds? 2. How do electrons affect the shape of a molecule? 3. What factors affect molecular properties? Chapter Objectives 1. Differentiate between ionic and molecular compounds, and between formula units and molecules. (8.1, 9.1, & pp ) 2. Compare and contrast a chemical formula and a molecular formula. (8.1) 3. Identify and list the names and formulas of the common monatomic and diatomic molecules. (8.1 & 9.1) 4. Describe a covalent bond. (8.1 & 8.2) 5. Interpret and write/draw molecular formulas, structural formulas, and Lewis structures/lewis diagrams/lewis dot diagrams/dot diagrams, for simple covalent molecules containing single, double, or triple bonds. (8.1 & 8.2) 6. Interpret ball and stick molecular models and space filling molecular models for simple covalent molecules containing single, double, or triple bonds. (8.1 & 8.2) 7. Describe the formation and characteristics of a covalent bond and the types of elements involved in this type of bonding. (8.1 & 8.2) 8. Predict, identify, write, and describe single, double, and triple covalent bonds. (8.2) 9. Predict, identify, write, and describe shared and unshared pairs of electrons. (8.2) 10. Predict, identify, write, and describe incomplete octets. (8.2) 11. Describe bond dissociation energy and how it relates to bond type. (8.2) 12. Identify, write and describe resonance structures. (8.2) 13. Describe the VSEPR theory. (8.3) 14. Predict and describe common electron pair geometries found in simple molecules with a maximum of four pairs of electrons around the central atom. (8.3) 15. State the corresponding bond angles for common electron pair (electronic) geometries found in simple molecules with a maximum of four pairs of electrons around the central atom. (8.3) 1
2 16. Predict and describe the molecular shapes for electron pair geometries found in simple molecules. (8.3) 17. Define and identify central atoms and related terminology. (8.3) 18. Use electronegativity values to determine whether a bond is nonpolar covalent, polar covalent, or ionic. (8.4) 19. Describe and show the relationship between polar covalent bonds & polar molecules. (8.4) 20. Name, identify and describe the weak attractive forces that hold molecules together and the physical manifestations that occur because of those forces. (8.4) 21. Demonstrate and be able to describe all aspects of laboratory safety rules and procedures. (Applicable every chapter) 8.3 Polar Bonds and Molecules A. Bond Polarity 1. Bond polarity is caused by the unequal sharing of electrons that exists in many covalent bonds. The sharing of electrons is based on the electronegativities of the two atoms involved in the bond. An atom with a higher electronegativity has more pull on the electrons than an atom with lower electronegativity. Therefore, an atom with a higher electronegativity has more electrons than protons much of the time, producing an area of partial negative charge. At the other end of the bond will be an atom with lower electronegativity, thus it will have more protons than electrons much of the time, producing an area of partial positive charge. If the two atoms in the bond have the same electronegativity the electrons are shared equally and there is no area of charge created. The polarity of a bond and the resulting molecule can be represented in the following ways: a. The use of the lowercase Greek letter delta, δ, which indicates partial. Area of Partial Positive Charge Area of Partial Negative Charge b. A color grid. VIB G YOR (ROY G BIV backwards). Unless, of course, I had made the color grid go from partial negative to partial positive. Or else you were looking from the other side of the paper! Area of Partial Positive Charge Neutral (Switzerland) Area of Partial Negative Charge 2 - HC - Chapter 8 - Objectives and Notes - Part 2 - V12
3 c. An arrow pointing toward the atom(s) with greater electronegativity; i.e., the area with an excess of electrons and thus, the partial negative charge. Area of Partial Positive Charge Area of Partial Negative Charge 2. nonpolar covalent bond: A covalent bond in which the shared electrons are shared equally between the two atoms involved in the bond. Homonuclear diatomic molecules always form nonpolar covalent bonds. The electronegativity difference between the atoms involved is equal to zero. a. Example: Electron clouds in N 2 can be seen below. 3. polar covalent bond: A covalent bond in which the shared electrons are shared unequally between the two atoms involved in the bond; the electrons spend more time revolving around the more electronegative atom. As a result of this unequal sharing, the bonded atoms form partial positive and partial negative areas. The partial negative area forms around the more electronegative atom and the partial positive area forms around the less electronegative (more electropositive) atom. The electronegativity difference between the atoms involved is greater than zero and approximately less than 1.7. Bonds between different nonmetals form polar covalent bonds. a. Example: The bond in HF can be seen below. 3 - HC - Chapter 8 - Objectives and Notes - Part 2 - V12
4 4. polar molecule/dipole/permanent dipole: A polar molecule; a molecule held together by unsymmetrical polar bonds, the centers of positive and negative charge do not correspond or symmetrically cancel each other out. The molecule has areas of partial positive and partial negative charge. a. The shape of a molecule and the polarity of its bonds determine the charge distribution in the molecule. b. Examples: See HF above and H 2 O below. 5. nonpolar molecule Molecules held together with nonpolar bonds or polar bonds arranged symmetrically so that they cancel each other out. a. CO 2 is an example of a molecule that contains polar bonds, but is a nonpolar molecule; see below. B. Attractions Between Molecules 1. intra: A prefix meaning "inside". 2. inter: A prefix meaning "between or among". 3. intramolecular forces: Forces inside individual particles, such as covalent bonds. 4. van der Waals forces/weak forces/intermolecular forces/interparticle force: Forces between or among individual particles. The individual particles are usually molecules, but sometimes are independent atoms or ions. The strength of these forces greatly affects the physical state of a substance. These forces are only about 15% of the strength of covalent or ionic bonds. a. The different types of weak forces are seen in the table on the van der Waals Forces Flowchart Addendum on the Chapter 8 webpage. 1. They are listed on the table from strongest on the left to weakest on the right. 4 - HC - Chapter 8 - Objectives and Notes - Part 2 - V12
5 5. ion-dipole forces: The attraction between an ion and a dipole, it depends on the charge on the ion and the magnitude of the dipole moment of the dipole. a. These forces are important when ionic solutes are dissolved in polar solvents such as H 2 O. b. It is the strongest of all of the interparticle weak forces. c. Examples: When sodium chloride is dissolved, the forces between the sodium ions and water molecules (see below) or the forces between the chloride ions and water molecules. 6. dipole-dipole forces: The attractive force between the oppositely charged ends of two permanent dipoles. It is the strongest major type of intermolecular force. a. Example: The forces between two HCl molecules (see below) b. hydrogen bond: One of the world's stupidest names; it is not a true bond, but rather, a special type of dipole-dipole force. It occurs when one at least one molecule is composed of hydrogen and elements with a very high electronegativity (N, O, or F) and the adjacent molecule is polar and contains N, O, or F. The resulting molecules are extremely polar, resulting in a strong dipole-dipole force between the very partially positive hydrogen of one dipole and the partially negative element of another dipole. The high melting and boiling points of substances with hydrogen bonds is a reflection of the strength of these intermolecular forces. Hydrogen bonds are the strongest type of dipole-dipole forces. 1. Examples: The forces found among NH 3 molecules and/or H 2 O molecules; the forces between a H 2 O molecule and an HF molecule (see below). 5 - HC - Chapter 8 - Objectives and Notes - Part 2 - V12
6 7. ion-induced dipole forces: The attractive force between an ion and the oppositely charged end of an induced dipoles. a. induced dipole: A nonpolar atom or molecule in which a nearby charged area, such as an ion or a dipole, distorts the electron cloud. Induced dipoles can be created when the electron cloud of a nonpolar molecule is repelled by a negative ion or when a positive ion attracts the electron cloud (see below). When the nearby ion is removed, the induced dipole reverts back to its nonpolar state. 1. Example of an ion-induced dipole force: The Fe 2+ in hemoglobin attracting O 2. Ion Nonpolar Molecule Ion O - O Induced Dipole 8. dipole-induced dipole forces: The attractive force between the oppositely charged ends of permanent dipoles and induced dipoles. a. Example of dipole-induced dipole forces: O 2 dissolved in water. Polar Molecule Nonpolar Molecule Dipole O - O Induced Dipole 6 - HC - Chapter 8 - Objectives and Notes - Part 2 - V12
7 9. London dispersion forces/london forces/dispersion forces: The attractive forces between the oppositely charged ends of two temporary dipoles or the forces between a temporary dipole and an induced dipole created by that temporary dipole. a. temporary dipole/instantaneous dipole: An atom or molecule in which the electrons become unsymmetrically arranged, creating a dipole for a fraction of a second. A temporary dipole is created the instant an atom or nonpolar molecule has an unequal number of electrons on one side of the particle; when this occurs the side with the extra electrons becomes negatively charged and the other side of the atom or molecule becomes positively charged. Within a fraction of a second after a temporary dipole is formed the electrons rearrange themselves symmetrically due to electron-electron repulsion; this reverts the atom or molecule back to its nonpolar state. 1. See creation of temporary dipole and resultant London forces in I2 below. I I Nonpolar Molecule Temporary Dipole I I Nonpolar Molecule Induced Dipole 2. polarizability: The ease in which an atom or molecule forms a temporary or induced dipole. The larger the radius of the atom or molecule, the easier it is for a dipole to form; thus, polarizability and radius size are directly related. These factors are clearly reflected in the melting and boiling points of monatomic and diatomic molecules. a. Examples: In the VIIA column, the first two elements, F2 and Cl2, are both gases at room temperature, the third element, Br2, is a liquid, and the last two elements, I2 and At2, are both solids. The intermolecular forces increase as the degree of polarizability increases. 7 - HC - Chapter 8 - Objectives and Notes - Part 2 - V12
8 1. The increase is molecular mass going down the column also plays a role in the increase in melting and boiling points. At any temperature, all substances have the same kinetic energy: thus, the greater the molecular mass the lower the velocity. (Who pulled the sled?) b. Polarizability usually increases as molecular mass increases. Note: This is because in most cases, the larger the molecular mass the more electrons an atom has and thus, the larger the radius of the atom, making it easier to form an induced or temporary dipole. 8 - HC - Chapter 8 - Objectives and Notes - Part 2 - V12
THE BIG IDEA: BONDING AND INTERACTIONS.
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