Chapter Five Chemical Bonding: The Covalent Bond Model.

Chapter FiveChapter FiveChemical Bonding: The Covalent Bond

Model

Chapter 5 | Slide 2

Chemical Bonding: The Covalent Bond Model cont’d

Ionic Molecules Covalent Molecules

Formed between a metal and a nonmetal

Formed between two or more nonmetals

Formed when valence electrons are _____________

Formed when valence electrons are __________

Do not exist as discrete molecules (formula unit)

Exist as discrete molecules

Their solutions conduct electricity

Their solutions usually don’t conduct electricity

Chapter 5 | Slide 3

Fig. 5.1 Electron sharing can occur only when electron orbitals from two different atoms overlap.


Chapter 5 | Slide 4


Fig. 5.3 (a) A “regular” covalent single bond is the result of overlap of two half-filled orbitals. (b) A coordinate covalent single bond is the result of overlap of a filled and a vacant orbital.

Chapter 5 | Slide 5


← Fig. 5.10

Linus Pauling received the Nobel Prize in chemistry in 1954 for his work on the nature of the chemical bond.

© Bettman/CORBIS

Chapter 5 | Slide 6


← Fig. 5.2

The number of covalent bonds formed by a nonmetallic element is directly correlated with the number of electrons it must share in order to obtain an octet of electrons.

Chapter 5 | Slide 7


Formed when two (or more) nonmetals ________ valence electrons in order to obey the ________ rule.Single covalent bond: two atoms share ___ pair of

electrons between them (__ shared electrons = _ bond)Example: Cl2

Double covalent bond: two atoms share ___ pairs of electrons between them (_ shared electrons = _ bonds) in order to obey the ________ ruleExample: CO2

Chapter 5 | Slide 8


Triple covalent bond: two atoms share __ pairs of electrons between them (6 shared electrons = ___ bonds)

Example: CO

Chapter 5 | Slide 9


Lewis structures show us how atoms are bonded to each other. They can be used to predict molecular shapes. Conventions: A _____represents one non-bonding electron A ______ represents two shared (bonding) electrons (1 bond)

Examples:

NH3

O2

C2H2

Chapter 5 | Slide 10


→ Fig. 5.4

The sulfur dioxide molecule. One Sulfur atom and two oxygen atoms



Example: SO2

1. Add together the _______ electrons for all of the atoms in the compound. Add or subtract ________ as necessary if you have an _____.

2. Determine which atom is your ________ atom (usually the first one in the formula unless the first atom is H). Draw a ______ bond between the central atom and each of the other atoms. How many electrons do you have left?



3. Put electrons around the outer (non-central) atoms until each of them has ____ electrons around it. Remember: H only gets ___ electrons around it. How many electrons do you have left?

4. If you have leftover electrons, put them on the central atom. (Sometimes the central atom will have more than 8 electrons around it; that atom will not be obeying the octet rule)



CH4

SO42-

Note: When drawing the structures of ions, write the structure in square brackets with the charge in the upper right-hand corner.



→ Fig. 5.7

The sulfate ion.



→ Fig. 5.5

The phosphorus trifluoride molecule.



← Fig. 5.6

The hydrogen cyanide molecule. HCN



Lewis Structures tell us how atoms are connected in a covalent molecule or ion

They can also be used to predict molecular shapeMolecular shape determines how molecules ________.Molecular shape is essential to drug design and other

biochemical processes.



In order to predict molecular shape, we assume that electron groups, i.e, bonding and nonbonding electrons, _____ each other (because they are all negatively charged). Therefore, the molecule adopts whichever 3D geometry

minimizes this __________.

We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.



Molecules adopt shapes that minimize the repulsion between electron groups/domains

What is an electron group/domain?A collection ________ _________ present in a localized

region about the ________ atom in a molecule

There are two kinds of electron groups/domains:A pair of ________________ electronsAn area where there are _________ electrons (it

doesn’t matter if the bond is a single, double or triple bond)



← Fig. 5.8

Arrangement of valence electron pairs about a central atom that minimize repulsions between the pairs.



Determining Molecular ShapeFirst, draw the Lewis Structure for the compoundThere are two important features of Lewis Structures that you

must recognize a. The number of ATOMS bonded to the central

atom (this is the number of bonding electron groups/domains)

b. The number of PAIRS of nonbonding electrons on the central atom (this is the number of nonbonding electron groups/domains)

Example: CO2



NH3

H2CO

CH4

H2O



Molecular shapes are considered around EACH central atom

Examples:C2H2

H2O2

HN3



Fig. 5.9

(a) Acetylene molecule. (b) Hydrogen peroxide molecule. (c) Hydrogen azide molecule.



Electronegativity A measure of the relative _________ that an atom

has for the _______ electrons in a bondElectrons are _______ ________ to the atom with the

higher electronegativity

Trends in electronegativityElectronegativity values increase from left to right on the

periodic tableElectronegativity values increase from the bottom to the

top of the periodic table



→ Fig. 5.11

Abbreviated periodic table showing Pauling electronegativity values for selected representative elements.



Polar covalent bondsA covalent bond in which there is _________ sharing of

electrons between two atoms

Nonpolar covalent bondsA covalent bond in which there is _________ sharing of

electrons between two atoms



← Fig. 5.12

(a) In the nonpolar covalent bond present, there is a symmetrical distribution of electron density. (b) In the polar covalent bond present, electron density is displaced because of its electronegativity.



Bond Polarity: a measure of the degree of ____________ in the sharing of electrons between two atoms in a chemical bondThere is no sharp distinction between bonding types.

The positive end (or pole) in a polar bond is represented _____ and the negative pole _____.The negative end is toward the atom with the higher

electronegativity

Example: HCl



Polar Bonds? Example: NaCl

Example: O2

Example: NH3

Example: HCl



→ Fig. 5.13

(a) Methane is a nonpolar tetrahedral molecule. (b) Methyl chloride is a polar tetrahedral molecule.



Molecular Polarity A measure of the degree of inequality in the

attraction of bonding electrons to various locations within a molecule

Polar Molecule:A molecule in which there is an unsymmetrical

distribution of charge



Why do we care if the molecule is polar?Polar molecules dissolve in polar solvents (like

water or blood)Nonpolar molecules dissolve in nonpolar

solvents (like hexane or cell membranes)

For a molecule to be polar:1. It must contain polar bonds2. The molecular geometry must not cancel out the

effect of the polar bonds (through vector addition)



Rules of Thumb Most covalent molecules are somewhat polar. It is

easier to remember the exceptions.

The exceptions for this class:Molecules in which two or more of the same atom are

bonded to the central atom AND in which there are no nonbonding pairs of electrons on the central atom



H2O

PF3

SO2

CH3Cl

H2

HCl


CAG 5.2



Naming Binary Molecular Compounds (Non-Acid) Binary molecular compounds have two different

elements. The ______ electronegative element is written first

Exception: NH3. Greek prefixes are used to indicate the number of

atoms.



→ Table 5.1



First word of name:Greek prefix + name of first element

Note: if the prefix would be “mono” for the first element, the prefix is not included

Note: when an element name begins with a vowel, an a or o at the end of the Greek prefix is dropped for phonetic reasons

Second word of nameGreek prefix + stem of name of second element with

“-ide” ending



→Table 5.2

Compounds in which hydrogen is the first element listed are named without prefixes.

Some compounds have common names.

Last


Name the following binary compounds.

SF4

P4O6

ClO2

H2S


Name the following binary compounds.

Cl2O

CO

PI3

HI


Write formulas for the following binary compounds.

Iodine monochloride

Dinitrogen monoxide

Nitrogen trichloride

Hydrogen bromide


Write formulas for the following binary compounds.

Bromine monochloride

Tetrasulfur dinitride

Sulfur trioxide

Dioxygen difluoride

Chapter Five Chemical Bonding: The Covalent Bond Model.

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Transcript of Chapter Five Chemical Bonding: The Covalent Bond Model.