Chapter 8: Chemical Bonding and Climate Changeresources.seattlecentral.edu/faculty/jrbryant/Chapter...

CHEM 161: Chapter 8 page 1 of 18

Chapter 8: Chemical Bonding and Climate Change

Problems: 8.1-8.58, 8.63-8.72, 8.77-8.79, 8.81-8.108, 8.111, 8.113-8.127, 8.129-8.135, 8.138, 8.140, 8.146

8.1 CHEMICAL BONDS

chemical bond: what holds atoms or ions together in a compound The three types of chemical bonds are ionic bonds, metallic bonds, and covalent bonds. – Metallic bonds hold metal atoms together in metals. – Ionic bonds hold ions together in ionic compounds. – Covalent bonds hold atoms together in molecules.

METALLIC BONDS Metals exist as nuclei surrounded by a sea of electrons.

The electrons in a metal are shared among all the nuclei, so the electrons are delocalized (i.e., not fixed to a specific atom).

The electrons can shift throughout the entire metal.

Electrons are free to move throughout the solid metals’ unique properties. – Pots and pans are usually made of metal because metals conduct heat and electricity as electrons flow through the metal. – Metals are malleable and ductile because electrons act as a glue, holding the

positively-charged nuclei together hammered metal and metallic wire.


THE IONIC BOND Ex. 1 Give the Lewis electron-dot formula below for each of the following atoms and ions:

sodium

sodium ion chlorine chloride ion

Metal atoms transfer valence electrons to nonmetal atoms.

positively charged metal cations and negatively charged nonmetal anions

– the charged ions are attracted to each other ionic compound = 3D network of ions

ionic bond: electrostatic attraction between positively charged cation and negatively charged anion

Properties of Ionic Substances: – Ionic compounds exist as networks of ions, with cations surrounding anions, and vice versa.

To melt ionic compounds, every bond between each Na+ ion and the Cl ions

surrounding it must be broken, as well as bonds between each Cl ion and the Na+ ions surrounding it.

A lot of energy is required to break all of these bonds.

Ionic compounds have relatively high melting points—much higher than water’s (0°C).

All ionic compounds are solids at room temperature. – Ionic compounds do not conduct electricity when solid (ions are fixed in place), but they

do conduct in the molten (liquid) and aqueous states (ions move around freely).

8.2 LEWIS SYMBOLS AND LEWIS STRUCTURES COVALENT BOND In 1916, an American chemist, Gilbert N. Lewis, proposed that atoms bond by sharing electrons.

Thus, a bond where atoms share electrons is called a covalent bond. octet rule: Atoms form covalent bonds in such a way that all atoms get eight electrons (an

octet) or 4 pairs of electrons, except hydrogen which only needs 2 electrons. – i.e., atoms bond to get the same number of valence electrons as the Noble gas in the same period.


covalent bond: sharing of one or more pairs of electrons between 2 nonmetal atoms, so both atoms can achieve a Noble Gas electron configuration.

– Covalent bonds hold atoms together in a molecule.

H• + •H H:H

Note: When two H atoms bond to form a hydrogen molecule, both atoms get two electrons, just like a stable helium atom. We can also represent the H2 molecule as follows: H H

This overlapping region is the covalent bond where electrons are shared.

In terms of quantum mechanics, we can also show the “electron clouds” or “electron density” for two H atoms combining to form the H2 molecule:

H atom H atom H2 molecule

Note: In the H2 molecule, the electron density is concentrated between the nuclei + because the two H atoms share the electrons equally.

Two F atoms bond to form a F2 molecule, and both atoms get 8 valence electrons (like neon).

bonding electron pairs nonbonding electron pair or lone pair of electrons

bonding electrons: electrons in a covalent bond, holding atoms together

nonbonding electrons (or lone pairs): belong solely to an individual atom

F + F F F ••

••

• •

• •

••

••

• •

••

••

••

••

• •

• •

• •


Coordinate covalent bond: – When one atom donates both electrons to make the bond

coordinate covalent bond

8.3 ELECTRONEGATIVITY, UNEQUAL SHARING, AND POLAR BONDS

Electronegativity (EN): Ability of an atom in a bond to attract shared electrons to itself – The electronegativity scale was proposed by an American chemist named Linus Pauling. – F is the most electronegative element

Elements are less electronegative the farther away they are from F, except for H, which has an EN value between B and C.

Note: H’s electronegativity is between B and H’s.

Nonpolar Covalent Bond – In some covalent bonds, both atoms have equal electronegativity (EN) values.

The atoms share the bonding electrons equally.

nonpolar covalent bond – Most common example is two atoms of the same element bonded to one another

– e.g. two H atoms form a nonpolar covalent bond in the H2 molecule (H—H) – Two different atoms can also form a nonpolar covalent bond as long as their

electronegativity values are the same. – e.g. The N—Cl bond in a NCl3 molecules is also a nonpolar covalent bond.

H+

+ F – H F

••

••

• •

••

••

• •

• •

• •


In some covalent bonds, one of the two atoms attracts the bonding electrons more strongly

polar covalent bond results between two atoms – polar because it has electrically polarized ends, a positive (+) end and a negative (–

end, just like the north (N) end and south (S) end of a bar magnet.

Consider the HF molecule:

H―F

In terms of quantum mechanics, we can also show the electron density in the HF molecule:

Note: The electron density for HF is between the two nuclei but is more concentrated around the F atom.

Polar Covalent Bond – Electrons spend more time around F, and they spend less time around H. – The region around the F atom is electron-rich, and

F gets a partial negative charge (–), and H gets partial positive charge (+):

+

–

H―F

Delta () Notation for polar bonds:

– The lowercase Greek letter delta () is used to represent partial charges. – Electrons concentrate around the more EN atom in a molecule

Atom gains a partial negative charge, indicated with –. – Since electrons spend less time around the other atom

Other atom gains a partial positive charge, indicated with +.

– This separation of charges is called a dipole and is usually represented with an arrow ( ) pointing from the less electronegative (indicated with a + sign at the start of the arrow) to the more electronegative atom in the bond.

Bond Polarity and Electronegativity – Pauling based the electronegativity scale on the idea that bonds are usually neither 100%

(nonpolar) covalent nor 100% ionic but somewhere on a continuum between the two. – The image below shows a) pure covalent bond (or equal sharing of electrons) in yellow for the Cl—Cl bond in Cl2

b) polar covalent bond with the shown in red and the + shown in greenish-blue

c) complete electron transfer to make an ionic bond between Na+ and Cl ions in NaCl

• •

••

••

• •

••

••


Ex. 1: i. Indicate if the bond below if polar covalent (P) or nonpolar/pure (NP) covalent. ii. If it’s polar covalent, use delta notation and a dipole arrow to indicate which atom in each of the following is more electronegative:

P or NP

O—N

P or NP

O—H

P or NP

C—F

Ex. 2: If AlCl3 is an ionic compound, its melting point should be _______ higher lower than NaCl’s melting point of 801°C.

The melting point for AlCl3 is 192°C, so the compound cannot be ionic. Consider the electronegativity difference between Al and Cl then indicate two other elements with a similar electronegativity difference to describe the bonds in AlCl3.

Ex. 3: Identify the bonds holding atoms/ions together in the following substances:

a. F2 ionic polar covalent nonpolar covalent metallic

b. K2O ionic polar covalent nonpolar covalent metallic

c. copper ionic polar covalent nonpolar covalent metallic


d. carbon monoxide ionic polar covalent nonpolar covalent metallic

e. potassium ionic polar covalent nonpolar covalent metallic

f. hydrogen ionic polar covalent nonpolar covalent metallic

g. strontium oxide ionic polar covalent nonpolar covalent metallic

Covalent Bond Formation – When two H atoms come near one another, the nucleus of one atom is attracted to the

electrons of the other atom. – Together each H atom has 2 electrons—like a He atom has 2 electrons.

They become more stable like He (indicated by the drop in energy in the plot below). – Note: In potential energy diagrams like the one below, the more negative the energy, the stronger the bond and the more stable the compound.

– However, the atoms cannot come too close together, or their nuclei and electrons get too

close, resulting in nucleus-nucleus and electron-electron repulsions optimal distance between atoms = bond length.

Single Bonds, Double Bonds, and Triple Bonds – Covalent bonds can also be shown as a line to represent the pair of electrons

H―H O=C=O :NN:

single bond double bond triple bond

bond length: distance between the nuclei of two covalently bonded atoms

••

••

••

••


– optimum distance maximizing attraction, minimizing repulsions Bond Length and Bond Strength

Single bonds: longest and weakest bonds Double bonds: shorter and stronger than single bonds Triple bonds: shortest and strongest type of covalent bond

LEWIS STRUCTURES

1. Count the total number of valence electrons present – from atoms – due to charge if polyatomic ion (add or subtract electrons for charge)

2. Write the skeleton structure of the compound – Put least EN atom as central atom and surround with other atoms – Note: H and F atoms will always be outer atoms

3. Connect all atoms by drawing single bonds between atoms then distribute remaining valence electrons as lone pairs around outer atoms then around the central atom so each has an octet

octet rule: atoms form bonds such that all atoms get eight electrons – BUT hydrogen (H) can only have 2 electrons (to fill 1s orbital)

4. If there are not enough electrons for each atom to have an octet, make double or triple bonds between central atom and surrounding atoms

– BUT fluorine (F) can only form a single bond!

5. For polyatomic ions, draw square brackets around the Lewis formula, and the charge is put in the upper right-hand corner

Draw the Lewis structure for the following: a. H2O b. COF2 c. NF3 d. HCN e. NH4

+ f. PO33–

bond length


8.6 FORMAL CHARGES: CHOOSING AMONG LEWIS STRUCTURES Example: Draw two different Lewis formulas for CH2O using the skeleton structures below.

Many valid Lewis structures can be drawn, so one must determine the most plausible structure.

formal charge: hypothetical charge an atom would have if bonding electrons are shared equally and lone pairs belong solely to a single atom

formal charge =

total number of

valence electrons

in free atom

—

total number of

nonbonding

electrons

— total number

of bonds

1. For neutral molecules, sum of formal charges must equal zero.

2. For ions, the sum of the formal charges must equal charge.

Guidelines for formal charges and plausible structures:

Lewis formulas with no formal charges is preferable to one with formal charges.

Lewis formula with large formal charges (+2, +3, and/or –2, –3, etc.) is less plausible than one with lower formal charges (+1 and -1, etc.)

When Lewis formulas have similar formal charge distributions, most plausible formula has negative formal charge on more electronegative atom

Ex. 1 Use formal charges to determine the most plausible structure for CH2O above.

C O

H

H

H C O H


Ex. 2 Determine the most likely Lewis structure for nitrous oxide, N2O, starting with the following skeleton structures and by determining formal charges:

N N O N N O N N O 8.5 RESONANCE Given the Lewis structure for ozone, we expect either of the following structures: so one bond (O—O) bond should be longer than the other (O=O). BUT experimental evidence indicates that both oxygen-oxygen bonds in ozone are identical, so neither of the structures accurately represents the molecule.

The actual structure is a cross between the two structures, where the electron pair is actually spread over all three atoms: These electrons are considered delocalized electrons because they are spread between more than two atoms. To correctly represent such delocalized electrons using Lewis formulas we show all the Lewis formulas with a double-arrow between each:

where each of these structures is called a resonance structure.

resonance structure: one of two or more Lewis structures representing a single molecule that cannot be described fully with only one Lewis structure

resonance: the use of two or more Lewis structures to represent one molecule – The real ozone molecule does not oscillate between the two resonance structures but is a

unique, stable structure that cannot be adequately represented with one Lewis structure.

Ex. 1 a. Draw the resonance structures for NO2

–:

O O O • •

• • • • • •

••O

O O • •

• • • • • •

••

• • • •

O O O • •

• • • • • •

••O

O O • •

• • • • • •

••

• • • •

O O O • •

• • • • • •


b. Determine the formal charge for each atom above. c. Resonance structures indicate equivalent bonds, so the actual formal charge for each oxygen atom in NO2

– is ________. Ex. 2 Give the resonance structures for the carbonate ion, CO3

2–: b. The actual formal charge for each oxygen atom in CO3

2– is ________.

Lewis Electron-Dot Formulas for Ternary Oxyacids (e.g. HNO3, H2SO4, etc.) – Ternary oxyacids contain hydrogen, oxygen, and one other element. – In ternary oxyacids, the central atom is the “other element” which is surrounded by oxygen,

and the hydrogen atoms are bonded directly to the oxygen atoms.

more than one central atom in the molecule

– Draw the skeleton structure of the polyatomic ion (e.g. ClO4), then one H atom per O atom.

Ex. 1 Draw the Lewis structures for the following ternary oxyacids: HClO3 H2CO3 Ex. 2. Draw the Lewis structures for nitrous acid, HNO2, then use formal charges to determine

the most plausible structure.

O N O O N O


: F

:

:

: F :

:

S

: F :

:

: F

:

:

F :

:

:

F :

:

:

N=O :

:

:

.

B

: F :

:

: F

:

:

F :

:

:

8.7 EXCEPTIONS TO THE OCTET RULE The Incomplete Octet: – Some atoms are satisfied with less than an octet – Be (1s22s2) is stable with only four valence electrons – Boron (1s2 2s2 2p1) also tends to form compounds with less than eight electrons

Odd-Electron Molecules: – Some molecules have an odd number of electrons can't satisfy octet rule; usually N has the odd number The Expanded Octet: – Atoms in and beyond the 3rd period can have more than eight electrons when in a compound – Thus, when drawing Lewis electron-dot formulas, extra electrons go on the central atom.

Examples: Draw the Lewis structures for the following molecules and ions:

H—Be—H


BeCl2: triiodide ion, I3–:

NO2: BH3:

PCl5: XeF4:

chlorine trifluoride, ClF3 SF4:

Ex. 1 a. Draw the Lewis structure for sulfur trioxide, SO3, in which sulfur obeys the octet rule, then determine the formal charge for each atom in your Lewis structure.


a. Lewis structure that obeys the octet rule

b. Lewis structure with an expanded octet

b. Next, minimize the formal charges by moving lone pairs from atoms with negative

formal charges (those atoms with too many electrons) to make double bonds with atoms that have positive formal charges (those atoms with too few electrons).

c. Draw the Lewis structure for sulfite ion, SO3

2–, minimizing the formal charges to determine the most plausible structure.

d. The actual formal charge for each oxygen atom in SO3

2– is ________. e. Compare the bond length for the sulfur-oxygen in SO3 and SO3

2–. Which has longer bonds? Explain why.


Ex. 2 Draw the Lewis structure for the following compounds and ions, minimizing the formal charges to determine the most plausible structure.

a. SOF2 b. nitric acid

Ex. 3a. Draw the Lewis structure for the following then minimize the formal charges to determine the most plausible structure.

i. phosphate ion

ii. phosphoric acid

b. Compare the bond length for the phosphorus-oxygen bonds in phosphate ion and

phosphoric acid. Which has longer bonds? Explain why.

8.8 THE LENGTH AND STRENGTH OF BONDS


Bond Energy: energy required to break a particular bond in the gas phase – always positive since breaking a bond always requires energy – a quantitative measure of a bond’s strength (i.e. the potential energy of molecule)

– The higher the bond energy the stronger the bond.

Breaking a bond always requires energy while forming a bond always releases energy.

Using a chemical equation and knowing the Lewis structure of each reactant and product allows us to estimate the amount of energy required to break bonds, the amount of energy released when bonds are made, and the overall net energy change for the reaction.

Use of Bond Energies (BE) in Thermochemistry

– We can estimate the heat of reaction—often referred to as the enthalpy change (H˚)—for any reaction using bond energies.

– Since energy is required to break the reactants’ bonds and energy is released when the

products’ bonds are formed, H˚ for the reaction is determined as follows:

H° = BE (reactants) – BE (products)

= total energy input – total energy released

Note: if BE (reactants) is greater than BE (products) endothermic reaction

if BE (reactants) is less than BE (products) exothermic reaction

Note: The Bond Energies included in Table 8.2 are averages for individual bond dissociation

energies for different molecules but can still provide a good approximation for H˚.

Consider the following example:


Example: For the following, draw the Lewis structures, then use Table 8.2 (on the previous

page) to calculate the H˚ for the following reactions:

a. 2 H2(g) + O2(g) 2 H2O(g)

reactants the in atoms between bonds

break to usedenergy Total=

formed are products in atoms between bond when

releasedenergy Total=

net energy change = H˚ =


b. N2(g) + 2 H2(g) (g)

c. C3H8(g) + 5 O2(g) 4 H2O(g) + 3 CO2(g)

H H H

H C C C H

H H H

H—N—N—H

H H

.. ..

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