01:160:159 Lecture Notes - Lecture 8: Chemical Polarity, Lattice Energy, Electron Configuration

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16 May 2018

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AP Chemistry

A. Allan

Chapter 8 Notes - Bonding: General Concepts

8.1 Types of Chemical Bonds

A. Ionic Bonding

1. Electrons are transferred

2. Metals react with nonmetals

3. Ions paired have lower energy (greater stability) than separated ions

B. Coulomb's Law

−

Q Q

2 ÷

E = 2.31x10

J × nmç

a. E = energy in joules

b. Q1 and Q2 are numerical ion charges

c. r = distance between ion center in nanometers

d. negative sign indicates an attractive force

C. Bond Length (covalent)

1. Distance at which the system energy is at a minimum

2. Forces at work

a. Attractive forces (proton - electron)

b. Repulsive forces (electron - electron, proton - proton)

3. Energy is given off (bond energy) when two atoms achieve greater

stability together than apart

D. Covalent Bonds

1. Electrons are shared by nuclei

2. Pure covalent (non-polar covalent)

a. Electrons are shared evenly

3. Polar covalent bonds

a. Electrons are shared unequally

b. Atoms end up with fractional charges

(1) d+ or d-

8.2 Electronegativity

A. Electronegativity

1. The ability of an atom in a molecule to attract shared electrons to itself

B. Electronegativity Trends

1. Electronegativity generally increases across a period (why?)

2. Electronegativity generally decrease within a family (why?)

C. Characterizing bonds

1. Greater electronegativity difference between two elements means less

covalent character and greater ionic character

2. We will not use the subtraction of electronegativities to determine ionic

character. This text uses a practical definition to identify ionic

compounds:

Any compound that conducts an electric current when melted is an

ionic compound.

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8.3 Bond Polarity and Dipole Moments

A. Dipolar Molecules

1. Molecules with a somewhat negative end and a somewhat positive end

(a dipole moment)

2. Molecules with preferential orientation in an electric field

+ + +

- - -

3. All diatomic molecules with a polar covalent bond are dipolar

B. Molecules with Polar Bonds but no Dipole Moment

1. Linear, radial or tetrahedral symmetry of charge distribution

a. CO2 - linear

b. CCl4 - tetrahedral

2. See table 8.2 in your text

8.4 Ions: Electron Configurations and Sizes

A. Bonding and Noble Gas Electron Configurations

1. Ionic bonds

a. Electrons are transferred until each species attains a noble gas

electron configuration

2. Covalent bonds

a. Electrons are shared in order to complete the valence

configurations of both atoms

B. Predicting Formulas of Ionic Compounds

1. Placement of elements on the periodic table suggests how many

electrons are lost or gained to achieve a noble-gas configuration

a. Group I loses one electron, Group II loses two, Group VI gains

two, Group VII gains one….

2. Formulas for compounds are balanced so that the total positive ionic

charge is equal to the total negative ionic charge

Al+

23 O3

−2

Total positive = +6

Total negative = -6

C. Sizes of Ions

1. Anions are larger than the parent atom

2. Cations are smaller than the parent atom

3. Ion size increases within a family

4. Isoelectronic ions

a. Ions with the same number of electrons

b. Size decreases as the nuclear charge Z increases

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8.5 Formation of Binary Ionic Compounds

A. Lattice Energy

1. The change in energy that takes place when separated gaseous

ions are packed together to form an ionic solid

M+ (g) + X- (g) → MX (s)

2. Energy change is exothermic (negative sign)

Example: Formation of lithium fluoride

Process

Description

Energy Change (kJ)

Li(s) → Li(g)

Sublimation energy

161

Li(g) → Li+(g) + e-

Ionization energy

520

1/2F2 → F(g)

Bond energy (1/2 mole)

F(g) + e- → F-(g)

Electron affinity

-328

Li+(g) + F-(g) → LiF(s)

Lattice energy

-1047

Li(s) + 1/2F2(g) → LiF(s)

-617

3. The formation of ionic compounds is endothermic until the formation

of the lattice

4. The lattice formed by alkali metals and halogens (1:1 ratio) is cubic

except for cesium salts

B. Lattice Energy Calculations

æ Q Q

Lattice Energy = k ç

a. k = a proportionality constant dependent on the solid structure

and the electron configuration

b. Q1 and Q2 are charges on the ions

c. r = shortest distance between centers of the cations and the

anions

2. Lattice energy increases as the ionic charge increases and the

distance between anions and cations decreases

8.6 Partial Ionic Character of Covalent Bonds

A. Calculating Percent Ionic Character

measured dipole moment of

X - Y

Percent ionic character = ç

÷x100%

− ÷

è calculated dipole moment of

B. Ionic vs. Covalent

1. Ionic compounds generally have greater than 50% ionic character

2. Ionic compounds generally have electronegativity differences greater

than 1.6

3. Percent ionic character is difficult to calculate for

compounds containing polyatomic ions

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Document Summary

8. 1 types of chemical bonds: ionic bonding, electrons are transferred, metals react with nonmetals, ions paired have lower energy (greater stability) than separated ions, coulomb"s law. This text uses a practical definition to identify ionic compounds: Any compound that conducts an electric current when melted is an ionic compound. 8. 5 formation of binary ionic compounds: lattice energy, the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. M+ (g) + x- (g) mx (s: energy change is exothermic (negative sign) 617: the formation of ionic compounds is endothermic until the formation of the lattice, the lattice formed by alkali metals and halogens (1:1 ratio) is cubic except for cesium salts, lattice energy calculations. 8. 6 partial ionic character of covalent bonds: calculating percent ionic character. Measured dipole moment of x - y. Calculated dipole moment of: ionic vs. covalent.

Related Questions

I want to find the Kn value of Cu(II) and EDTA. I was given the formula kn= [ML4]/[ML3][L]. 12.3 ml of edta was used to tritrate 0.1194 grams of Copper in 50 ml of water. The indicator used was murexide.

At the core of general acid-base theories, there is the concept of âtransferâ of one entity to another. In the case of BrÃ¸nsted-Lowry acid base theory, the transferred entity is the proton, and their transfer processes are widely exploited in chemical methods. In addition to proton transfer, electron pairs can also be donated/shared between donors and acceptors. In this case, an atom/molecule, which can accept a pair of electrons, is referred to as a Lewis acid while the entity donating the pair of electrons is referred to as a Lewis base. The most prevalent Lewis acid-base systems involve metal complexes with vacant nd-orbitals, as well as electron deficient systems including B, Si, Ge, etc. The Lewis acid base theory is also known as the Hard-Soft acid base theory (HSABT) with âhardnessâ defined as small highly charged species that are weakly polarizable (Lewis acids). In contrast, âsoftnessâ is described by species that are larger and highly polarizable species with low charge (Lewis bases).
In the case of metal chemistry, Lewis acid-base theory provides a means to describe the ability of âligandsâ to coordinate to metal ions as well as the stability of metal-ligand complexes. Recall that coordination bonds are the donation/acceptance of lone pairs of electrons to empty metal orbitals to another atom; the definition of Lewis acid-base chemistry! The stability of Lewis acid-base complexes can be described through the stability constant which is analogous to the Ka that describes BrÃ¸nsted-Lowry acidities. Many metal ions can coordinate (bind) many Lewis base ligands in a step-wise fashion according to:
M + L = ML K1 = [ML]/[M][L] Eq. 1
ML + L = ML2 K2 = [ML2]/[ML][L] Eq. 2
ML2 + L = ML3 K3 = [ML3]/[ML2][L] Eq. 3
MLn-1 + L = MLn Kn = [ML4]/[ML3][L] Eq. 4
and for which the free energy for this dissociation reaction is given by:
ÎG0 = -RTLnKn Eq. 5.
where R is the gas constant and T is the temperature in Kelvin.
Figure 1: Left- EDTA. Right- Typical EDTA-transition metal complex.
The ethylene diamine tetraacetic acid ligand (EDTA, a Lewis base) provides a mechanism to probe the Lewis acid strength of various metal complexes (Figure 1). This multi-dentate ligand (multiple Lewis base sites on a single ligand) forms complexes with metals of the form:
Mn+ + Y4- = MYn-4 Eq. 6
K = [MYn-4]/[Mn+][Yn-4] Eq. 7
where Y4- is the EDTA ligand.
In this experiment, the Lewis acid strength for two different metals will be examined by measuring the stability constants for the metal-EDTA complex. Of specific interest is the how the Lewis acid strength of the metal ion correlates with:
1) atomic/ionic radii, 2) oxidation state, 3) electron configuration, and 4) polarizability.
You will determine the stability constants of both a Cu2+- and Zn2+-EDTA complexes as well as the corresponding free energy. Using these data, you will compare the relative Lewis acidities of both metals and describe differences in terms of the metal ion properties described above.
.
Chemicals
CuSO4
ZnSO4
EDTA
Murexide
Instruments
Buret
Experimental
You and your group will determine the concentration and ultimately the free energy of Cu(II) and Zn(II) using the hexadentate chelating ligand, EDTA.
You will be given a ~100 mM standardized solution of EDTA (check the label for exact concentration prior to beginning). You will need to weigh ~0.1 g per metal salt sample. Dissolve the metal salt into 50 mL of DI water. Add 2-3 drops of murexide to your metal. After properly preparing your buret with EDTA, titrate your metal with EDTA. Repeat for a total of 3 titrations per metal.
1. Approximately how much EDTA will you need for each titration assuming the metal-saltsâ masses are accurate?
2. What type of indicator is murexide and what color will its endpoint be?
Post-Lab Questions
1.
a) Using the equations given in the introduction and your titration results, what are the Kn of the EDTA-Copper complex? Show calculations.
b) Using the equations given in the introduction and your titration results, what are the Kn of the EDTA-Zinc complex? Show calculations.
2. What are the corresponding ÎG values for the EDTA-metal complexes? Show your calculations
3. Would copper be considered a hard or soft lewis acid? What about zinc? Describe what makes these metals hard or soft.
4. What trends would you predict for EDTA-metal complex Kn and ÎG values when:
a) going down a row of transition metals of the same oxidation state?
b) of the same transition metal with increasing oxidation state? (For example, Iron can exist as Fe(II), Fe(III), Fe(IV), and higher oxidation states)
5. Discussion (formal report style): Compare and contrast the concepts of Lewis and BrÃ¸nsted-Lowry acids and bases. Support these concepts with your data over the last two weeks and with research. Cite your sources.

01:160:159 Lecture Notes - Lecture 8: Chemical Polarity, Lattice Energy, Electron Configuration

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