Unit 8 Bonding.pptx

Unit 8: Bonding

8.1: compare the characteristics of ionic compounds and covalent molecules.

8.2: use models to explain the role of valence

electrons in bond formation including resonance forms using formal charges.

8.3: describe the relationship between

electronegativity, bond type, and polarity of a molecule or compound.

Electronegativi

ty

Electronegativity

ability of atoms in a molecule to attract electrons to itself.

• _{On the periodic table,}

electronegativity

increases as you go…

• _{…from left to right}

across a row.

• _{…from the bottom to}

the top of a column.

• _{Based on Unit 7, why}

Electronegativity

• _{The “Pauling Scale”}

graduates from 0.7 – 4.0 with 4.0 being the highest electronegativity (fluorine)

• _{The difference in}

electronegativity is the

technical way to calculate bonding type.

EN Dif Bond Type 0-0.5 Covalent

You try…



• Calculate the electronegativity difference and thus the type of bonding occurring between…

• _{Chlorine and another Chlorine}

• _{Hydrogen and Bromine}

Ionic Bonding

Ionic Bonding Basics

• _{Donation of one or more electrons from}

(typically) a metal to a non-metal forming a cation and anion.

• _{These ions are then held together ONLY by}

an electrostatic (+ and -) attraction.

• _{There is a high electronegativity difference}

between bonding elements.

• _{Also explained by high electron affinity}

Ionic Bonding, why?

• The ionization energy is such that it takes 495 kJ/mol to remove electrons from sodium.

• We get 349 kJ/mol back by giving electrons to chlorine.

• But these numbers don’t explain why the reaction of sodium metal and

chlorine gas to form sodium chloride is so exothermic! • http://www.pbslearningmedia.org/asset/n vhe_vid_compounds/? utm_source=teachersdomain_redirect/as set/nvhe_vid_compounds/utm_medium= teachersdomain/asset/nvhe_vid_compou nds/utm_campaign=td_redirects

Lattice Energy

• _{This third piece of the}

puzzle is the lattice energy:

The energy required to completely separate a mole of a solid ionic

compound into its gaseous ions.

• _{The energy associated}

with electrostatic

interactions is governed by Coulomb’s law:

E_el =  Q1Q2

Lattice Energy

• _{Lattice energy, then, increases with the charge}

of the ions.

• _{It also increases with}

decreasing size of ions.

E_el =  Q1Q2

• Figure 8.4: Stare at it…

• By accounting for all three energies

(ionization energy, electron affinity, and

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

• _{Notice that without the}

stability of the lattice energy provided by

electrostatic attraction of the ions, ionic

Energetics of Ionic Bonding

• _{These phenomena also}

helps explain the “octet rule.”

• _{Metals, for instance, tend to stop}

losing electrons once they attain a noble gas configuration

because energy would be expended that cannot be

Polar Covalent

Bonding

Polar Covalent Bonds

• _{Although atoms often form}

compounds by sharing

electrons, the electrons are not always shared equally. This is a Polar Covalent

Bond

• _{Fluorine pulls harder on the electrons it shares with}

hydrogen than hydrogen does.

• _{Therefore, the fluorine end of the molecule has}

Polar Covalent Bonds

• _{When two atoms share}

electrons unequally, a bond dipole results.

• _{The dipole moment,} _,

produced by two equal but opposite charges

separated by a distance,

r, is calculated by:

 = Qr

• _{It is measured in debyes}

Polar Covalent Bonds

The greater the difference in

Covalent

Bonding

Little Electronegativity Difference

Covalent Bonding Basics

• _{Two non-metal elements share electrons to}

complete their octet.

• _{Both elements have high electron affinities}

and can pull electron clouds away from each other simultaneously.

• _{Both elements have high ionization energies}

and will not release valence electrons.

• _{Both elements have high electronegativities.} • _{Can share more than one pair of electrons in}

Covalent Bonding

Forces must be balanced in a Covalent bond…

• _{There are several}

electrostatic interactions in these bonds:

• _{Attractions between electrons}

and nuclei

• _{Repulsions between electrons} • _{Repulsions between nuclei}

Lewis Structures

• _{Lewis structures are representations of}

molecules showing all VALENCE electrons, bonding and nonbonding.

• _{Useful for quickly sketching covalent molecule}

Writing Lewis Structures

electrons of all atoms in the polyatomic ion or

molecule.

• _{If it is an polyatomic}

anion, add one electron for each negative

charge.

• _{If it is a polyatomic}

cation, subtract one electron for each

positive charge.

You try… 

• _PCl3

• _{5 + 3(7) = 26} • _CO2

• _{4 + 2(6) = 16} • _NH4+

Writing Lewis Structures

2. The central atom is the least

electronegative

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

Writing Lewis Structures

3. Fill the octets of the outer atoms.

Writing Lewis Structures

4. Fill the octet of the central atom.

Writing Lewis Structures

5. If you run out of

electrons before the central atom has an octet…

Now you try



• Draw the Lewis Dot Structure for:

• _H2O • _NF3 • _CO2

Double/Triple Bonds

• _{Many covalent molecules form double or}

triple bonds to complete octets of elements.

• _{This is why nitrogen and oxygen form}

diatomic molecules spontaneously at STP.

• _{This also explains molecules}

Double and Triple Bonds

• _{Increase the enthalpy of a bond. (require more}

energy to break)

• _{Increase stability of the bond.}

• _{Decrease the bonding radius (atoms are pulled}

Formal Charges

• _{Step 6?}

Assign formal charges if molecule has double/triple bonds to determine the best possible resonance .

Writing Lewis Structures

• _{The best possible Lewis}

structure…

• _{…is the one with the fewest}

charges.

• _{…and puts a negative charge}

Formal Charges the Easy

way!

•

# valence electrons for element –

More Than Eight Electrons ?!?!

• _{The only way PCl5 can exist is if}

phosphorus has 10 electrons 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.

• _{Presumably orbitals}

are “hybridized” (we will discuss this

More Than Eight Electrons

Even though we can draw a Lewis structure for the phosphate ion that has only 8

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 is on}

the 3rd row or below and breaking its octet

Now you try



• Determine the best possible structure for the following based on formal charges:

• _NO3

-• _CO3

Resonance

Resonance Basics

• _{Doube/triple bond electrons have the ability to}

move between bonds in a molecule, this is called resonance.

• _{If a molecule cannot achieve a stable octet for}

its elements it can use resonance to spread one or more sets of electrons between

Resonance

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}

Resonance

• _{One Lewis structure}

cannot accurately

depict a molecule such as ozone since the

resonating electron pair is delocalized.

• _{We use multiple}

structures, resonance structures, to describe the molecule.

• _{Or we use a more}

accurate hybrid

• 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 localized, but rather are delocalized.

• You can think of the resonating electron pair as spending half it’s time on each bond so each bond has approximately 1.5 bonds

instead of a single bond or a true double bond.

• This still stabilizes the bonds, pulls the atoms closer, and

Now you try



• Write all the possible resonance structures for:

• _SO4

3-Covalent Bond Strength

• _{Most simply, the strength of a bond is measured}

by determining how much energy is required to break the bond.

• This is the bond enthalpy.

• _{The bond enthalpy for a Cl—Cl bond,}

Average Bond Enthalpies

• _{This table lists the} average bond

enthalpies for many different types of

bonds.

• _{Average bond}

enthalpies are

positive, because

Average Bond Enthalpies

NOTE: These are

average bond enthalpies, not absolute bond

enthalpies; the C—H bonds in methane, CH₄, will be a bit

different than the C—H bond in

Enthalpies of Reaction

• _{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,}

H_rxn = (bond enthalpies of bonds broken) 

Enthalpies of Reaction

CH₄(g) + Cl₂(g) 

CH₃Cl(g) + HCl(g)

In this example, one C—H bond and one

Enthalpies of Reaction

So,

H_rxn = [D(C—H) + D(Cl—Cl)  [D(C—Cl) +

D(H—Cl)

= [(413 kJ) + (242 kJ)]  [(328 kJ) + (431 kJ)]

Now you try…



• _{Use the table of bond energies to estimate the}

enthalpy of the following reactions:

• _{N2(g) + 3 H2(g) 2 NH3(g)}

Metallic

Bonding

Metallic Bonding

• _{Two or more elements that have low electronegativies.} • _{If both have low EN, then both elements don’t want}

electrons!

• _{Results in e- “hot potato” game, where elements pass} e- around constantly in a “sea of e-”

• Basically, ALL the valence electrons are delocalized amongst ALL atoms in the metal lattice.