The electrostatic attraction of closely packed, oppositely charged ions (typical between metals and nonmetals)
Exist where electrons are shared between nuclei (typical between nonmetal atoms)
The electron density is located primarily between the two nuclei
Metallic bonds consist of atoms bonded to several neighboring atoms with the electrons free to move among the 3-dimensional structure of the atoms.
Chemical bonds can be described with a variety of properties; this includes bond energy, bond length,
The energy required to break a bond.
The distance between the nuclei of the bonding atoms. This distance where energy is a minimum between nucleus-nucleus repulsion, electron-electron repulsion, and electron-nucleus attraction.
Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons. This consists of a full s and p sublevels.
Atoms near Helium will attain only 2 valence electrons, a full s sublevel.
A Lewis symbol shows the symbol of an atom or atoms in a compound with dots representing each valence electron. For a single atom, there is a maximum of eight dots around the atom, and we limit the representation to a maximum of two dots per side (top, bottom, left, right)
A compound formed from a positive cation and a negative anion. A metal reacts with nonmetal forms an ionic compound.
Ionic compounds are often brittle, crystalline (highly structured forming flat surfaces and clear edges), high melting points and high boiling points.
The strength of the ionic bonding of a solid ionic compound is given by the lattice energy, which is the energy required to completely separate a mole of an ionic solid into its gaseous ions.
MX(s) à M+(g) + X-(g)
This energy value will be positive since the process is an endothermic process.
All binary ionic compounds formed by an alkali metal and a halogen (except for cesium salts) have the same structure as sodium chloride where each atom is surrounded by six ions of the opposite charge.
Lattice energy can be represented by a modified form of Coulomb's law:
Lattice energy = k(Q1Q2/r)
where k depends on the structure of the solid and the
electron configurations of the ions,
r is the distance between ion centers; and Q1 and Q2 are
the numerical ion charges
The lattice energy increases as the charges on the ions increase and as the radii decrease.
Since energy is a state function, this energy value can be calculated from a series of other reaction steps.
Coulomb's Law
The energy of interaction between a pair of ions can be calculated using Coulomb's Law
E = 2.31 x 10-19 J nm (Q1Q2/r)
Where r is the distance between ion centers in nm;
and Q1 and Q2 are
the numerical ion charges
This gives a negative energy for positive/negative ion pairs indicating an attractive force, or that the ion pair has lower energy than the separated ions.
Lattice Energy
Lattice energy is the energy required to completely separate a mole of an ionic solid into its gaseous ions.
The energy of interaction between a pair of ions can be calculated using Coulomb's Law
E = k(Q1Q2/d)
k = 8.99 x 109 J m/C2
Where d is the distance between ion centers in m;
and Q1 and Q2 are
the charges o the ions in Coulombs.
The lattice energy increases as the charges on the ions increase and as the radii decrease.
Since energy is a state function, the energy value of a reaction can be calculated from the sum of other reactions.
Some Rules for adding Reactions:
· A compound showing up as both a reactant and product can be cancelled out of the reaction.
· Reversing the reaction changes the sign of the energy.
· Multiplying the reaction multiplies the energy.
Lattice energy is equal to the sum of the ionization energy of the metal, the electron affinity of the nonmetal, the sublimation energy of the metal, and the bond energy of the diatomic halogen minus the energy of formation of the ionic compound.
Example:
Mg(s) à Mg(g) DH = 150 kJ Sublimation
Mg(g) à
Mg2+(g) + 2e- DH =
2201 kJ Ionization
energy (2)
(1/2)O2(g) à O(g) DH = 249 kJ Bond
energy
O(g) + 2e- à O2-(g) DH =
-337 kJ Electron
affinities (2)
Minus
Mg(s) + (1/2)O2(g) à MgO(s) DH =
-601 kJ Enthalpy of
formation
equal
MgO(s) à
Mg2+(g) + O2-(g) DH =
2864 kJ Lattice
energy
Values came from Zumdahl
Enthalpy of formation is the sum of the ionization energy of the metal, the electron affinity of the nonmetal, the sublimation energy of the metal, the bond energy of the diatomic halogen, and the lattice energy.
Lattice Energies for some Ionic Compounds
|
Compound |
Lattice Energy (kJ/mol) |
|
LiF |
1030 |
|
LiCl |
834 |
|
LiI |
730 |
|
NaF |
910 |
|
NaCl |
788 |
|
NaBr |
732 |
|
NaI |
682 |
|
KF |
808 |
|
KCl |
701 |
|
KBr |
671 |
|
MgCl2 |
2326 |
|
SrCl2 |
2127 |
|
MgO |
2864 |
|
CaO |
3414 |
|
ScN |
7547 |
In most cases the electron configuration of ions of representative elements match that of a noble gas.
When a nonmetal and a representative-group metal react to form a binary ionic compound, the ions form so that the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom and the valence orbitals of the metal are emptied.
Example: the Neon electron configuration is achieved by O2-, F-, Na+, Mg2+, Al3+
The ionic charges of representative elements will be such to create a noble gas electron configuration.
Several exceptions exist where a second ionic charge is
also possible:
Sn (2+, 4+)
Pb (2+, 4+)
Bi (3+, 5+)
Tl (1+, 3+)
Cations are smaller than the parent element
Anions are significantly larger than the parent element
In isoelectronic ions, the more protons, the smaller the
ion
or the more negative the ion, the larger the ion
Transition metals have up to 12 electrons beyond a noble gas core. Typically, it is not energetically favorable (higher successive ionization energies) to form ions with high ionic charges. So most transition metals form ions with charges of +1, +2, +3.
In forming ions, electrons are removed first from the level that has the highest value of n. So in forming ions, transition metals loose the valence shell s electrons first before loosing d electrons. This is why most transition metals form a +2 ion along with other charges.
Polyatomic ions are two or more atoms that are covalently bonded together, but have an ionic charge and form ionic bonds with other ions.
COC
Covalent bonds exist where electron pairs are shared between nuclei
The electron density is located primarily between the two nuclei forming an attractive force between the electron pair and each nuclei.
In writing Lewis structures, only valence electrons are used.
· Sum the valence electrons for all the atoms
· Write the symbols of the atoms in the appropriate layout.
· Connect the bonding atoms with a single line (representing a pair, 2, of electrons)
· Arrange the remaining electrons to satisfy the duet rule for hydrogen and octet rule for the remaining elements.
· Any excess electrons go on the central atom (if period 3 element or higher) even if the octet is exceeded.
· If octets are not complete, move nonbonding electron pairs into bonds (forming double and triple bonds) until octets are complete
Hydrogen forms stable molecules when it shares two electrons
Second row elements form stable molecules when its orbital has 8 electrons.
· The second row elements, C, N, O, F should always be assumed to obey the octet rule
· The second-row elements, B and Be often have fewer than eight electrons around them. These compounds are very reactive
· The second-row elements never exceed the octet rule
· Third row and heavier elements often satisfy the octet rule, but can exceed it using the d subshell (Example: SF6 forms 6 bonds around sulfur using 12 electrons)
· When writing the Lewis structure, satisfy the octet rule first, additional electrons are placed on elements with available d subshells.
The sharing of one pair of electrons.
A C-C single bond is typically 1.54°A in length.
The sharing of two pairs of electrons.
A C-C double bond is typically 1.34°A in length.
The sharing of three pairs of electrons
A C-C triple bond is typically 1.20°A in length.
Bond lengths shorten as the number of shared electrons increases
Resonance occurs when more than one valid Lewis structure can be written for particular molecule. The actual structure is the average of the resonance structures.
Electrons are delocalized and can move around the entire molecule.
Examples: NO3-, SO3, SO2
Benzene, C6H6, is a cyclic organic hydrocarbon molecule that shows resonance that is typical of aromatic organic molecules.
The Lewis structure of benzene will have three C-C double bonds and three C-C single bonds. There are two equivalent Lewis structures with the double and single bonds switched.
The molecule actually has 6 identical bonds with a bond length that is intermediate (1.40°A) between single bonds (1.54 °A) and double bonds (1.34 °A).
The structure can be drawn with three double bonds or with a circle inside the six single bonds, representing the delocalized double bonds.
Molecules or polyatomic ions that can exceed the octet rule often have many nonequivalent Lewis structures
A charge estimation method will help determine which is the best Lewis structure.
Formal charge is the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule.
Formal charge = Valence electrons – Assigned electrons
· Sum lone pair electrons and one-half the shared electrons to get the assigned electrons
· Subtract the number of assigned electrons from the number of valence electrons of the neutral atom to get the formal charge
· The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge of the species
· If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero are better.
· Lewis Structures with any negative formal charges are expected to have these charges reside on the most electronegative atoms
Example: BF3
Structure with three single bonds and 6 valence electrons on B
F: Formal charge = 7 Valence electrons – 7 assigned electrons = 0
B: Formal charge = 3 valence electrons – 3 assigned electrons = 0
Three resonance structures with one double bond and 8 valence electrons on B
F (single bond): Formal charge = 7 Valence electrons – 7 assigned electrons = 0
F (double bond): Formal charge = 7 Valence electrons – 6 assigned electrons = 1
B: Formal charge = 3 valence electrons – 4 assigned electrons = -1
The single bond BF3 is the better structure since the formal charges are both zero. also on the double bonded BF3 the negative formal charge is on the more electropositive element.
Occurs when electrons are shared equally between nuclei.
Occurs when electrons are shared unequally between nuclei.
The ability of an atom in a molecule to attract shared electrons to itself.
Electronegativity values generally increase from left to right across the periodic table (except noble gases) and increase up a group.
The high value is 4.0 for fluorine and the low value is 0.7 for Francium.
Electronegativity is calculated using Linus Pauling's
method. This consists of calculating an expected bond energy for H-X = (H-H bond energy + X-X bond energy)/2
which is the average bond energy between H-H and X-X
and comparing this average value with the measured value, D =
(H-X)measured – (H-X)theor
A large D represents a large polarity
|
H |
2 |
|
|
|
|
|
|
|
|
|
|
13 |
14 |
15 |
16 |
17 |
He |
|
2.1 |
IIA |
|
|
|
|
|
|
|
|
|
|
IIIA |
IVA |
VA |
VIA |
VIIA |
- |
|
Li |
Be |
|
|
|
|
|
|
|
|
|
|
B |
C |
N |
O |
F |
Ne |
|
1.0 |
1.5 |
|
|
|
|
|
|
|
|
|
|
2.0 |
2.5 |
3.0 |
3.5 |
4.0 |
- |
|
Na |
Mg |
3 |
4 |
5 |
6 |
7 |
8 |
9 |
10 |
11 |
12 |
Al |
Si |
P |
S |
Cl |
Ar |
|
0.9 |
1.2 |
IIIB |
IVB |
VB |
VIB |
VIIB |
VIIIB |
VIIIB |
VIIIB |
IB |
IIB |
1.5 |
1.8 |
2.1 |
2.5 |
3.0 |
- |
|
K |
Ca |
Sc |
Ti |
V |
Cr |
Mn |
Fe |
Co |
Ni |
Cu |
Zn |
Ga |
Ge |
As |
Se |
Br |
Kr |
|
0.8 |
1.0 |
1.3 |
1.5 |
1.6 |
1.6 |
1.5 |
1.8 |
1.8 |
1.8 |
1.9 |
1.6 |
1.6 |
1.8 |
2.0 |
2.4 |
2.8 |
- |
|
Rb |
Sr |
Y |
Zr |
Nb |
Mo |
Tc |
Ru |
Rh |
Pd |
Ag |
Cd |
In |
Sn |
Sb |
Te |
I |
Xe |
|
0.8 |
1.0 |
1.2 |
1.4 |
1.6 |
1.8 |
1.9 |
2.2 |
2.2 |
2.2 |
1.9 |
1.7 |
1.7 |
1.8 |
1.9 |
2.1 |
2.5 |
- |
|
Cs |
Ba |
La |
Hf |
Ta |
W |
Re |
Os |
Ir |
Pt |
Au |
Hg |
Tl |
Pb |
Bi |
Po |
At |
Rn |
|
0.7 |
0.9 |
1.0 |
1.3 |
1.5 |
1.7 |
1.9 |
2.2 |
2.2 |
2.2 |
2.4 |
1.9 |
1.8 |
1.8 |
1.9 |
2.0 |
2.2 |
- |
Bonds containing a difference in electronegativity of less than 0.4 are normally considered to be nonpolar. Covalent bonds containing a difference in electronegativity between the atoms of 0.4 or more are typically considered polar.
A molecule that has a center of positive charge that is separate from the center of negative charge
The dipole is often shown using an arrow pointing to the negative charge. Or by using d+ and d-
Any diatomic molecule with a polar bond has a dipole moment.
Polyatomic molecules that have polar bonds will have a dipole moment unless the symmetry of the molecule cancels out the effects of the polar bonds.
The polar bonds usually only cancel out on totally symmetric molecules such as tetrahedral molecules with four identical bonds (CCl4), or trigonal planar molecules (SO3) with three identical bonds or linear molecules with two identical bonds (CO2).
Dipole moments are often reported in debyes (D).
1 D = 3.34 x 10-30 coulomb-meter (C-m)
We often measure charge in units of electronic charge, e (1.60 x 10-19C) and distance in angstroms.
Suppose that two charges +0.5 and –0.5 (in units of e) are separated by 1.00 °A. The dipole moment produced is:
m = Qr = (1/2)(1.60 x 10-19C)(1.00 °A)(10-10 m/°A(1D/3.34 x 10-30 C-m) = 2.39 D
There is a gradient between polar covalent bonds and ionic bonds that is not always clearly distinguishable.
This is seen in naming conventions and chemical properties.
Ionic Molecular
CaC2 Calcium
carbide CO2 Carbon dioxide
Fe2O3 Iron
(III) oxide Cl2O3 Dichlorine trioxide
K2O Potassium
oxide H2S Dihydrogen sulfide
Ionic Common Crossover names
TiO2 Titanium
(IV) oxide Titanium
dioxide
(white solid)
SnCl4 Tin (IV) chloride Tin tetrachloride
(colorless liquid, mp = -33°C)
Mn2O7 Manganese
(VII) oxide Dimanganese
heptoxide
(green liquid, mp = 5.9°C)
Many compounds of metals with high oxidation numbers have properties more similar to molecular compounds.
Partial Ionic Character of Covalent Bonds
There is a gradient between polar covalent bonds and ionic bonds that is not always clearly distinguishable.
There are probably no totally ionic bonds between discrete pairs of atoms.
The evidence of this comes from the percent ionic character of a bond as measured in the gas phase
percent ionic character of a bond = [(measured dipole moment
of X-Y)
(calculated dipole moment of X+Y-)] x 100%
No compounds register as 100% ionic character and the % ionic character tops out lower than 90%.
Normal ionic compounds tend to have more than 50% ionic character. Many ionic compounds also have covalent bonds in polyatomic ions.
A safe definition of an ionic compound is one that conducts an electric current when melted.
Bond enthalpy is the enthalpy change (DH) for the breaking of a particular bond in one mole of a substance in the gas phase.
Bond enthalpy is always a positive quantity.
We can sue the notation D(bond type) to label bond
enthalpies.
Example: D(C-H) = 413 kJ/mol.
|
Single Bonds |
|||
|
C-H 413 |
N-H 391 |
O-H 463 |
F-F 155 |
|
C-C 348 |
N-N 163 |
O-O 146 |
|
|
C-N 293 |
N-O 201 |
O-F 190 |
Cl-F 253 |
|
C-O 358 |
N-F 272 |
O-Cl 203 |
Cl-Cl 242 |
|
C-F 485 |
N-Cl 200 |
O-I 234 |
|
|
C-Cl 328 |
N-Br 243 |
|
Br-F 237 |
|
C-Br 276 |
|
S-H 339 |
Br-Cl 218 |
|
C-I 240 |
H-H 436 |
S-F 327 |
Br-Br 193 |
|
C-S 259 |
H-F 567 |
S-Cl 253 |
|
|
|
H-Cl 431 |
S-Br 218 |
I-Cl 208 |
|
Si-H 323 |
H-Br 366 |
S-S 266 |
I-Br 175 |
|
Si-Si 226 |
H-I 299 |
|
I-I 151 |
|
Si-C 301 |
|
|
|
|
Si-O 368 |
|
|
|
|
Si-Cl 464 |
|
|
|
|
Multiple Bonds |
|||
|
C=C 614 |
N=N 418 |
O=O 495 |
|
|
C=C 839 |
N=N 941 |
|
|
|
C=N 615 |
N=O 607 |
S=O 523 |
|
|
C=N 891 |
|
S=S 418 |
|
|
C=O 799 |
|
|
|
|
C=O 1072 |
|
|
|
The enthalpy of a reaction will be the sum of bond energies of bonds broken (endothermic) minus the sum of bond energies of bonds formed (exothermic)
DH = S(bonds broken) – S(bonds formed)
for the reaction
H2 + Cl2 à 2HCl
DHrxn = 432 kJ/mol H2 + 239 kJ/mol
Cl2 – 2(427 kJ/mol HCl)
= -183 kJ (for reaction as written)
or for HCl
DH(HCl)
= -183/2 = -92 kJ/mol HCl
Compared to –92 kJ/mol = DH°f HCl
H2 (g) + I2 (g) ßà 2HI (g)
DHrxn = 432 kJ/mol H2 + 149 kJ/mol
I2 – 2(295 kJ/mol HI) = -9 kJ
or for HI
DH(HI)
= -9/2 = -4 kJ/mol
In general, as the number of bonds between two atoms increases, the bond grows shorter and stronger.
|
Bond |
Bond Length (°A) |
Bond |
Bond Length (°A) |
|
C-C |
1.54 |
N-N |
1.47 |
|
C=C |
1.34 |
N=N |
1.24 |
|
C=C |
1.20 |
N=N |
1.10 |
|
C-N |
1.43 |
N-O |
1.36 |
|
C=N |
1.38 |
N=O |
1.22 |
|
C=N |
1.16 |
|
|
|
|
|
O-O |
1.48 |
|
C-O |
1.43 |
O=O |
1.21 |
|
C=O |
1.23 |
|
|
|
C=O |
1.13 |
|
|
The energy of a particular type of bond (i.e. a C-H bond) varies with the local environment; however, the concept of a bond energy is still useful.
i.e. Methane can be considered to have four C-H bonds
Stabilizing energy of methane = 4 (C-H bonds) = 4(413 kJ/mol) = 1652 kJ/mol
Some molecules formed from nonmetals contain odd numbers of electrons (NO, NO2)
The formal charge model works better for determining the arrangement of odd electron molecules.