AHCC8 Chapter 8 Objectives & Notes - Part 2 -...

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ADVANCED HONORS CHEMISTRY - CHAPTER 8 NAME:

COVALENT BONDS DATE:

OBJECTIVES AND NOTES - PART 2 - V12 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 201 - 203)

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, condensed structural

formulas/condensed formulas, Lewis structures/Lewis diagrams/Lewis dot diagrams/dot diagrams, and

orbital diagrams/orbital box diagrams/orbital overlap diagrams for simple covalent molecules containing

single, double, or triple bonds. (8.1 & 8.2)

6. Interpret stereo projections/perspective drawings, 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. Describe the formation of a coordinate covalent bond. (8.2)

11. Predict, recognize, write, and describe sigma (σ) and pi (π) bonds. (8.2 & 8.3)

12. Describe exceptions to the octet rule. (8.2)

13. Describe bond dissociation energy and how it relates to bond type. (8.2)

14. Identify, write and describe resonance structures. (8.2)

15. Describe the VSEPR theory. (8.3)

2 - AHC - Chapter 8 - Objectives and Notes - Part 2 - V12

16. Predict and describe common electron pair geometries found in simple molecules. (8.3)

17. State the corresponding bond angles for common electron pair (electronic) geometries found in simple

molecules. (8.3)

18. Predict and describe the molecular shapes for electron pair geometries found in simple molecules. (8.3)

19. Define and identify central atoms, bond axes, and related terminology. (8.3)

20. Use electronegativity values to determine whether a bond is nonpolar covalent, polar covalent, or ionic.

(8.4)

21. Describe and show the relationship between polar covalent bonds & polar molecules. (8.4)

22. Name, identify and describe the weak attractive forces that hold molecules together and the physical

manifestations that occur because of those forces. (8.4)

23. Demonstrate and be able to describe all aspects of laboratory safety rules and procedures. (Applicable every chapter)

Remember, In All Teacher Prepared Notes, Items In Green Are For Background Information Only.

Students Do Not Need To Know The Information and

They Are Not Responsible For Them On Any Tests Or Quizzes.

Introduction to Organic Nomenclature

1. hydrocarbon: The simplest type of organic compound; they contain only atoms of carbon and hydrogen.

Derivatives of hydrocarbons contain other elements, notably oxygen.

a. alkanes: The simplest types of hydrocarbons; they contain only single bonds.

1. saturated hydrocarbons: Hydrocarbons in which each carbon atom is bonded to the maximum

number of atoms. Since carbon can form four single bonds, alkanes are saturated hydrocarbons

because each carbon atom is bonded to a total of four other carbon and/or hydrogen atoms.

2. Alkane names are composed of three parts:

a. The prefix, which signifies the group that is attached to the main chain and also the number

of the carbon atom the group is attached to. In this course we will not work with alkanes

with branches; therefore, we will ignore this part of the name.

b. The root, which is the longest continuous chain of carbon atoms in the compound. The root word

signifies how many carbon atoms are in the longest chain (See the table below).


c. The suffix, which reflects the type of organic compound. All alkanes end in -ane.

d. See the chart below for naming alkanes.

e. To remember the first four, strange, numeric prefixes use that incredibly creative and

exciting mnemonic device: Me Eat Peanut Butter.

Numerical Roots for Carbon Chains and Branches

# of C Atoms Root Alkane

Name Formula Uses/Comments

1 meth- methane CH4 Chief component of natural gas. Most houses in this area

are heated with natural gas. Lab burners use natural gas.

(We will ignore the obvious connection to crystal meth)

2 eth- ethane C2H6 When one of the hydrogen atoms is replaced by a -OH

group it produces a derivative of ethane

called ethanol or ethyl alcohol.

3 prop- propane C3H8 Found in small white tanks as fuel for BBQ’s.

4 but- butane C4H10 Used in butane lighters or sold in

tanks for use as fuel at campgrounds.

5 pent- pentane C5H12 The first alkane heavy enough and polar enough

to be found as a liquid at room temperature.

6 hex- hexane C6H14 A significant part of gasoline.

7 hept- heptane C7H16 An undesirable part of gasoline. It causes destructive

“knocking” in car engines as it burns explosively.

8 oct- octane C8H18 Gasoline is labeled (regular, premium, etc.)

by its octane rating. Interestingly, normal octane is not

the compound measured in this rating system.

9 non- nonane C9H20 A colorless volatile liquid.

It is an ingredient of kerosene.

10 dec- decane C10H22 Decane is a component of gasoline


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

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 N2 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.

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 H2O below.


5. nonpolar molecule Molecules held together with nonpolar bonds or polar bonds arranged symmetrically so

that they cancel each other out.

a. CO2 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.

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 H2O.

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 NH3 molecules and/or H2O molecules; the forces between an HF

molecule and a H2O molecule (see below).


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 Fe2+ in hemoglobin attracting O2.

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: O2 dissolved in water.

Polar Molecule

Nonpolar Molecule

Dipole

O - O Induced Dipole


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.

Nonpolar Molecule

Nonpolar Molecule

Temporary Dipole

Induced Dipole

I I I I


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.

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.

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