Chapter 8 - Monarch Sciencepberan.weebly.com/.../7/3/57730581/chap_8_student_notes.docx · Web viewAP Chemistry (BLM 11th) Chapter 8 - Chemical Bonding 4 Chapter 8Concepts of Chemical

AP Chemistry (BLM 11th) Chapter 8 - Chemical Bonding

Chapter 8Concepts of Chemical Bonding

8.1 Chemical Bonds Three basic types of bonds ___________

• Electrostatic attraction between ions ___________

• Sharing of electrons ___________

• Metal atoms bonded to several other atoms

Lewis SymbolsElectrons involved in chemical bonding are the _____________________ electrons.

G.N. _____________________(1875-1946) suggested a simple way of showing the valence electrons in an atom.

Lewis electron-dot structures: consist of the chemical symbol for the element plus a dot for each valence electron.

Octet RuleAtoms tend to ______________________________________________ electrons until they are surrounded by eight valence electrons.

An octet of electrons consists of full _______________________________-subshells in an atom.

There are many exceptions to the octet rule, but it provides a useful framework for many important concepts of bonding.

Write the Lewis symbol for atoms of each of the following elements: Al

Br

Ar

Sr

8.2 Ionic BondingEnergetics of Ionic Bonding

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AP Chemistry (BLM 11th) Chapter 8 - Chemical BondingAs we saw in the last chapter, it takes _______________ to remove electrons from sodium.

Energetics of Ionic BondingWe get ______________________ back by giving electrons to chlorine.

Energetics of Ionic BondingBut these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so ______________________!

Energetics of Ionic Bonding There must be a third piece to the puzzle. What is as yet unaccounted for is the __________

___________________between the newly-formed sodium cation and chloride anion.

Lattice Energy• This third piece of the puzzle is the lattice energy:

The energy required to completely separate a mole of a s olid ionic compound into its _____________________

NaCl(s) Na+ (g) + Cl-(g) ∆Hlattice= + 788 kJ/mol• The energy associated with electrostatic interactions is governed by Coulomb’s

law:

Lattice Energy Lattice energy, then,

__________________________ with the charge on the ions.

It also increases with decreasing size of ions.

Energetics of Ionic Bonding.

These phenomena also helps explain the “_________________.” Metals, for instance, tend to stop losing electrons once they attain a noble

gas configuration because energy would be expended that cannot be overcome by lattice energies.

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ExampleArrange the following ionic compounds in order of increasing lattice energy.

NaF, CsI, and CaO

Solution:

ExampleWhich substance would you expect to have the greatest lattice energy?

•AgCl •CuO •CrN

Solution:

3

AP Chemistry (BLM 11th) Chapter 8 - Chemical BondingCalculation of Lattice Energies

The Born-Haber Cycle:The formation of an ionic compound as occurring in a series of well-defined steps. We use Hess’s Law to put these steps together in a way that gives us the lattice energy for the compound.

Energetics of Ionic Bonding By accounting for all three energies

(ionization energy, electron affinity, and

______________________________), we can get a good idea of the energetics involved in such a process.

Born-Haber Cycle The enthalpy change for the direct route (red arrow) is the heat of

formation of NaCl(s) _____________________________________ ∆Hf [NaCl(s)]= -411 kJ Values are found using Appendix C!

Born-Haber CycleThe indirect route consists of five steps, shown by the green arrows.

Step 1 We generate gaseous atoms of sodium by vaporizing sodium metal:

Na(s) Na(g) ∆Hf [Na(g)] = 108 kJ

______________________ required energy

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AP Chemistry (BLM 11th) Chapter 8 - Chemical BondingStep 2

Then we form gaseous atoms of chlorine by breaking the bonds in the Cl2 molecules.

1/2 Cl2(g) Cl(g) ∆Hf [Cl(g)] = 122 kJ

Endothermic requires energy

Steps 3 and 4 Removing an electron from Na (g) to form Na+ (g) and then adding the electron to Cl (g) to form Cl-(g).

Na(g) Na+ + e- ∆Hf = I1(Na) = ___________

Cl(g) + e- Cl-(g) ∆Hf = E(Cl) = -349 kJ

Ionization energy found on pg 268 Electron affinity found on pg 271

Step 5 •Finally, we combine the gaseous sodium and chloride ions to form solid

sodium chloride. •Because this process is just the reverse of the lattice energy (breaking a

solid into gaseous ions), the enthalpy change is the negative of the lattice energy, the quantity we want to determine:

Step 5 Con’t Na+ (g) + Cl- (g) NaCl (s)

∆H = -∆Hlattice= ?

The sum of the five steps in the indirect path gives us the NaCl(s) from Na(s) and 1/2 Cl2(g).

Thus from Hess’s law we know the sum of the enthalpy changes for these five steps equals that for the direct path.

Solution

H°f [NaCl(s)] = ∆H°f [Na(g)] + ∆H°f [Cl(g)] + I1(Na) + E(Cl) –∆Hlattice

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AP Chemistry (BLM 11th) Chapter 8 - Chemical Bonding -411 kJ = 108 kJ + 122 kJ + 496 kJ –349 kJ – ∆Hlattice

Solving for ∆Hlattice

∆Hlattice= 108 kJ + 122 kJ + 496 kJ –349 kJ + 411 kJ ∆Hlattice= 788 kJ

Thus the lattice energy of NaCl is _____________________

Explain the following trends in lattice energy.

CaF2 > BaF2

NaCl > RbBr

BaO > KF

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8.3 Covalent Bonding

8.3 Covalent BondingIn covalent bonds atoms __________________electrons.

There are several electrostatic interactions in these bonds:o Attractions between electrons and

nucleio Repulsions between electrons o Repulsions between nuclei

Covalent Bonds The attractions between nuclei and the

electrons cause electron density to concentrate _______________________. As a result, the overall electrostatic interactions are attractive.

•A shared pair of electrons in any covalent bond acts as a kind of “glue” to bind atoms together.

A covalent bond is a chemical bond in which two or more electrons are shared by two atoms

Why should two atoms share electrons?

F F+7e- 7e-

FF8e-8e-

FF

F F

Lewis structure of F2

lone pairslone pairs

lone pairslone pairs

single covalent bond

single covalent bond

7


Lewis Structures The formation of covalent bonds can be represented using Lewis symbols

Examples: Draw: H2, H . + . H H:H Cl2, NH3, CH4

Lewis Structures Multiple Bonds: Doubles and Triples

Examples: Draw CO2 and N2

8e-H HO+ + OH H O HHor

2e- 2e-

Lewis structure of water

Double bond – two atoms share two pairs of electrons

single covalent bonds

O C O or O C O8e- 8e-8e-double bonds double bonds

Triple bond – two atoms share three pairs of electrons

N N8e-8e-

N Ntriple bondtriple bond

or

9.4

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AP Chemistry (BLM 11th) Chapter 8 - Chemical BondingAs a rule, the distance between bonded atoms decreases as the number of shared electron pairs increases.

Multiple bonds are always shorter and stronger than single bonds between the same two elements.

The added strength makes it more difficult to break these molecules apart than those with only single bonds between the same two elements

8.4 Bond Polarity & Electronegativity

8.4 Polar Covalent Bonds

Although atoms often form compounds by sharing electrons, the electrons are not always shared equally.

Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does in the molecule

Although atoms often form compounds by sharing electrons, the electrons are not always shared equally.

Polarity•____________________________ covalent bond is one in which the electrons are shared equally between two atoms, as in the Cl2 and N2 examples we just drew.

•___________________ covalent bond, one of the atoms exerts a greater attraction for the bonding electrons than the other. If the difference in relative ability to attract electrons is large enough, an _____________________is formed.

Electronegativity• The ability of atoms ________________________

to attract electrons to themselves.• On the periodic chart, electronegativity _________

as you go… …from left to right across a row. …from the bottom to the top of a column.

FH

e- riche- poor

d+ d-

9

AP Chemistry (BLM 11th) Chapter 8 - Chemical BondingVocabulary

___________________________ – the ability of an atom IN A MOLECULE to attract electrons to itself.

_________________________________ how strongly an atom holds on to its electrons.

__________________________________– the measure of how strongly an atom attracts additional electrons.

Another way to define electronegativity -Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.

Electron Affinity - measurable, Cl is highest

Electronegativity - relative, F is highest

X (g) + e- X-

(g)

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Question• What is the trend in electronegativity going from left to right in a row on the

periodic table?

9.5

9.5

Variation of Electronegativity with Atomic Number –within a period there is a steady increase in electronegativity from left to right.

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• How do electronegativity values change as you move down a group of elements on the periodic table?

• How do periodic trends in electronegativity relate to those for ionization energy & electron affinity?

Polar Covalent Bonds

The ___________________________ the difference in electronegativity, the more polar is the bond.

Electronegativity & Bond Polarity

Example

•F2 4.0 –4.0 =0 Nonpolar

•HF 4.0 –2.1 = 1.9 Polar Covalent

•LiF 4.0 –1.0 = 3.0 Ionic

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AP Chemistry (BLM 11th) Chapter 8 - Chemical BondingB-Cl or C-Cl

•Which bond is more polar? •Indicate in each case which atom has the partial negative charge.

Solution The difference in the electronegativities of chlorine and boron is 3.0 –2.0 =

1.0

The difference between chlorine and carbon is 3.0 –2.5 = 0.5

Therefore B-Cl is more polar. The ____________________ carries the partial negative charge because it has a higher electronegativity.

Polar Covalent Bonds When two atoms share

electrons unequally, a bond dipole results.

The ______________________, , produced by two equal but opposite charges separated by a distance, r, is calculated:

It is measured in debyes (D). A unit that equals 3.34 x 10-34 Coulomb-meters

What this means For molecules, we usually measure charge in units of electronic charge, e,1.6

x 10-19 C, and distance in units of angstroms.

Suppose we have a charge of 1+ and 1- and are separated by a distance of 1 Angstrom. A dipole moment produced is

•m= Q•r

•μ=(1.60 x 10-19 C)•[(1A)(10-10 m/1Å)( 1D/3.34x10-30 C-m)]

μ= ____________________

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8.5 Drawing Lewis Structures8.5 Lewis Structures

__________________________ are representations of molecules showing all electrons, bonding and nonbonding.

Writing Lewis Structures

1. Find the __________of valence electrons of all atoms in the polyatomic ion or molecule.

If it is an ___________ add one electron for each negative charge.

If it is a __________, subtract one electron for each positive charge.

Writing Lewis Structures2. The central atom is the ______

__electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.

3. Fill the octets of the outer atoms.4. Fill the octet of the central atom.

Keep track of the electrons:

26 - 6 = 20

Writing Lewis Structures

PCl3

5 + 3(7) = 26

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5. If you run out of electrons before the central atom has an octet…

…form _____________ until it does.

Writing Lewis StructuresThen assign __________________________.

For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.

Subtract that from the number of valence electrons for that atom: the difference is its formal charge.

Write the Lewis structure of the carbonate ion (CO32-)

Step 1 – C is less electronegative than O, put C in centerStep 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)

-2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electronsStep 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electronsStep 5 - Too many electrons, form double bond and re-check # of e

2 single bonds (2x2) = 41 double bond = 4

8 lone pairs (8x2) = 16Total = 24

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FOR EACH OF THE FOLLOWING MOLECULES OR IONS OF SULFER AND OXYGEN, WRITE A SINGLE LEWIS STRUCTURE THAT OBEYS THE OCTET RULE, AND CALCULATE THE OXIDATION NUMBERS AND FORMAL CHARGE FOR ALL OF THE ATOMS:SO2

Write the Lewis structure of nitrogen trifluoride (NF3).Step 1 – N is less electronegative than F, put N in centerStep 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)

5 + (3 x 7) = 26 valence electronsStep 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

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SO3

Writing Lewis Structures The best Lewis structure…

…is the one with the fewest charges. …puts a negative charge on the most electronegative atom.

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8.6 Resonance – We sometimes encounter molecules

and ions in which the experimentally determined arrangement of atoms in _______adequately described by a _________________________.

This is the Lewis structure we would draw for ozone, O3.

Resonance• But this is at odds with the true, observed

structure of ozone, in which… …both O-O bonds are the same

length. …both outer oxygens have a charge

of -1/2.Resonance

One Lewis structure cannot accurately depict a molecule like ozone.

We use multiple structures, _________ ____________ to describe the molecule.

ResonanceJust as green is a synthesis of blue and

yellow…

…ozone is a synthesis of these two resonance structures.

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AP Chemistry (BLM 11th) Chapter 8 - Chemical BondingResonance

In truth, the electrons that form the second C-O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.

They are not _____________________________

Which is predicted to have the shorter sulfur–oxygen bonds, SO3or SO32–?

The sulfur atom has six valence electrons, as does oxygen. Thus, SO3 contains 24 valence electrons. In writing the Lewis structure, we see that three equivalent resonance structures can be drawn:

The SO32– ion has 26 electrons, which leads to a Lewis structure in which all the S—O bonds are

single bonds:

There are no other reasonable Lewis structures for this ion. It can be described quite well by a single Lewis structure rather than by multiple resonance structures. Our analysis of the Lewis structures leads us to conclude that SO3

should have the shorter S—O bonds and SO3

2– the longer ones. This conclusion is correct: The experimentally measured S—O bond lengths are 1.42 Å in SO3 and 1.51 Å in SO3

2–.

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Resonance The organic compound

benzene, C6H6, has two resonance structures.

It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

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8.7 Exceptions to the Octet Rule•The Octet rule is simple & useful for most basic bonding concepts. However, in some situations, it does have some limitations:

In the bonding of ionic compounds of involving some transition metals.1. EX: Fe 2+ [Ar] 3d6 Fe 3+ [Ar] 3d5.

Some covalent bonding doesn’t follow the octet rule either. These exceptions to the octet rule are of three main types:

1. Ions or molecules with an ________________________ of electrons2. Ions or molecules with ______________________________3. Ions or molecules with more than eight valence electrons (an expanded

octet)

Odd Number of Electrons In the vast majority of molecules and polyatomic ions, the total number of

valence electrons is even. Though relatively rare and usually quite unstable and reactive, there are ions

and molecules with an ____________________ of electrons such as ClO2, NO, NO2, O2

- . Complete pairing of electrons is impossible and an octet around each atom cannot be achieved.

For example, NO contains 5 + 6 =11 Electrons. The two most important Lewis structures for this molecule are:

Fewer Than Eight Electrons• Consider BF3:

Giving boron a filled octet places a ______________charge on the boron and a __________________ charge on fluorine.

This would not be an accurate picture of the distribution of electrons in 21

AP Chemistry (BLM 11th) Chapter 8 - Chemical BondingBF3.

Fewer Than Eight Electrons Therefore, structures that put a double bond between boron and fluorine

are much less important than the one that leaves boron with only _________ electrons

.

Fewer Than Eight Electrons The lesson is: if filling the octet of the central atom results in a negative

charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

More Than Eight Electrons •The 3rd & largest class of exceptions consists of molecules or polyatomic

ions in which there are more that 8 electrons in the valence shell for the atom.

• The only way PCl5 can exist is if phosphorus has ________________________ around it.

• It is allowed to expand the octet of atoms on the 3rd row or below. Presumably d orbitals in these atoms

participate in bonding.

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More Than Eight Electrons1. Other examples of molecules & ions with “__________________”

valence shells include:2. SF4, AsF6

-, & ICl4- 3. Element from the 3rd period and beyond have ns, np, and unfilled nd

orbitals that can be used in bonding.4. Although 3rd period often satisfy the octet rule, as in PCl3, they can

exceed an octet by seeming to use their empty d orbital to accommodate additional electrons.

5. Expanded valence shells occur most often when the central atom is bonded to the smallest and most electronegative atoms, such as _____________________.

More Than Eight Electrons Even though we can draw a Lewis structure for the phosphate ion that has only

8 electrons around the central phosphorus, the better structure puts a ___________ ________between the phosphorus and one of the oxygens.

More Than Eight Electrons This eliminates the charge on the phosphorus and the charge on one of the

oxygens. The lesson is: when the central atom in on the 3rd row or below and

expanding its octet eliminates some formal charges, do so.

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8.8 Strengths of Covalent Bonds

8.8 Covalent Bond Strength

The stability of a molecule is related to the strengths of the covalent bonds it contains.

The strength of a covalent bond between two atoms is determined by the _________________________ to break the bond.

8.8 Covalent Bond Strength

Mos

t simply, the strength of a bond is measured by determining how much energy is required to break the bond.

This is the _______________________________ The bond enthalpy for a Cl-Cl bond, D(Cl-Cl), is measured to be 242 kJ/mol.

Average Bond Enthalpies• This table lists the average bond

enthalpies for many different types of bonds.

• Average bond enthalpies are positive, because bond breaking is an _______________________.

Average Bond Enthalpies NOTE: These are

______ ______ _______________________, not absolute bond enthalpies; the C-H bonds in methane, CH4, will be a bit different than the C-H bond in chloroform, CHCl3.

Enthalpies of Reaction

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AP Chemistry (BLM 11th) Chapter 8 - Chemical Bonding Yet another way to estimate H for a reaction is to compare the bond

enthalpies of bonds broken to the bond enthalpies of the new bonds formed. •In other words, Hrxn = (bond enthalpies of bonds broken) - (bond enthalpies of bonds

formed)

Enthalpies of Reaction

CH4 (g) + Cl2 (g) CH3Cl(g)+ HCl (g)

In this example, one C-H bond and one Cl-Cl bond are broken; one C-Cl and one H-Cl bond are formed.

Enthalpies of ReactionSo,H = [D(C-H) + D(Cl-Cl)] - [D(C-Cl) + D(H-Cl)]

= [(413 kJ) + (242 kJ)] - [(328 kJ) + (431 kJ)]= (655 kJ) - (759 kJ)= -104 kJ

Using Table 8.4, estimate ΔH for the following reaction (where we explicitly show the bonds involved in the reactants and products

Plan: Among the reactants, we must break six C—H bonds and a C—C bond in C2H6; we also break

Using Table 8.4, estimate ΔH for the reaction:

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H = [D(N-H) + D(N-N)] - [D(N=N) + D(H-H)]

∆H =

∆H =

Bond Enthalpy and Bond Length

We can also measure an average bond length for different bond types. As the number of bonds between two atoms increases, the bond length

decreases.

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Chapter 8 - Monarch Sciencepberan.weebly.com/.../7/3/57730581/chap_8_student_notes.docx · Web viewAP Chemistry (BLM 11th) Chapter 8 - Chemical Bonding 4 Chapter 8Concepts of Chemical