Chemical Bonding
Ionic bonds result from an electrostatic attraction between ions.
Ionic bonds occur between two atoms that have a large electronegativity difference between them. (metals and nonmetals)
The larger the Electronegativity difference the more ionic is the bond.
There is no such thing as a molecule of sodium chloride(or any other ionic compound).
Ionic compounds exist in crystal lattice structures as solids. A formula unit of an ionic compound is the smallest ratio of ions. For sodium chloride that is NaCl
Properties of Ionic Compounds
(Explain using the crystal lattice)
i. Solid ionic compounds do not conduct an electric current.
ii. Molten samples of ionic compounds can conduct an electric current due to the mobility of the ions which are free to move to the electrodes and react.
iii. Ionic compounds have high melting and boiling points.
iv. Many ionic compounds can dissolve in water.
v. Dissolved ionic compounds separate into cations and anions in solution.
vi. The mobile ions move to the electrodes and react to accept and release electrons creating a flow of electricity in the outer circuit.
vii. Ionic compounds that are water soluble are strong electrolytes.
viii. Ionic compounds are brittle
Lattice Energy (DHlattice)
The energy required to completely separate a mole of a solid ionic compound into its gaseous ions.
Example: The lattice energy of NaCl is the energy given off when Na+ and Cl- ions in the gas phase come together to form the lattice of alternating Na+ and Cl- ions in the NaCl crystal shown in the figure below.
Na+(g) + Cl-(g) NaCl(s) / Ho = -787.3 kJ/molLattice energy = k Q1Q2
r
Q1 and Q2 are the charges of the two ions, r is the distance between the center of the two ions.
The charge has a larger effect on the lattice energy compared to the size of the ions.
Lattice energies (kj/mol)
LiF 1030 MgCl2 2326
LiCl 834 SrCl2 2127
LiI 730 MgO 3795
NaF 910 CaO 3414
NaCl 788 SrO 3217
NaBr 732 ScN 7547
KF 808
KCl 701
KBr 671
Lattice energies are higher for smaller atoms (r) and higher for larger charges(q)
Electron configuration of Variable ions in ionic compounds
The electrons are removed from the s or p orbitals before the d-orbitals
Give the electron configuration for the following ions
Cu+1 Fe+2 Pb+2
Cu+2 Fe+3 Pb+4
Covalent Bonds
Formed when two atoms share one or more pairs of electrons.
Lewis Diagrams
Diagrams showing the valence electrons and the arrangement of atoms in covalent compounds
Constructing Lewis Diagrams
1. Sum up the valence electrons in the molecule. If it is an ion add or subtract electrons depending on the charge.
2. Put the structure so that on atom is in the center and the rest bonded to it. Usually the first atoms if it is not hydrogen is in the center. Do not make chains with the atoms.
3. Place two electrons between each two atoms
4. Place the remaining electrons so that all atoms have complete octets.
5. If there is not enough electrons more the pairs of electrons around to give double or triple bonds.
6. If it is an ion put the structure in brackets with the charge in the right top corner.
Draw the Lewis diagram for CBr4
Draw the Lewis diagram for H2O
Draw the Lewis diagram for O2
Draw the Lewis diagram for N2
Draw the Lewis diagram for the nitrate ion.
Exceptions to the Octet Rule
1. When the species contains an odd number of electrons one atom will have only seven electrons around it.
2. When the compound contains a central atom from groups 2 or 13, the number of electrons around the central atom is twice the group number. Examples BeI2, BCl3
3. When a central atom has more than four other atoms bonded to it, it will have more than eight electrons around it. This is known as an expanded octet. Examples PCl5, SF6
An expanded octet can occur if the atomic number of the central atom is greater than 10. Examples SF4, BrI3, I3 Occurs with elements with d-orbitals.
Exercise 1 Give the lewis diagram for PCl5.
Exercise 2. Write the Lewis structure for each molecule or ion.
a. ClF3 b. XeO3 c. RnCl2
d. BeCl2 e. ICl4-
Resonance
Seen in molecules when more than one Lewis structure can be drawn for the same arrangement of atoms. The molecule is a hybrid of the different resonance structures.
Draw the resonance structures for the nitrate ion.
Draw the resonance structures for the azide ion ( N3-).
Formal Charge
Assigned to atoms in Lewis diagrams. Formal charge is equal the number of valence electrons on an atom – the number of valence electrons assigned to an atom.
The resonance structure that contributes the most contains all atoms with formal charges closest to zero.
Coordinate Covalent Bonds
Atoms Such as nitrogen and oxygen share a lone pair of electrons with another atom.
These atoms are coordinate compounds and Lewis acids/bases. Both electrons in the electron pair comes from the Lewis base
Coordination Compounds
Lewis-Acid Base Complexes
Formed from a transition metal and ligands
The ligands can be a substance with an anion or a molecules with a lone pair of electrons.
Water and ammonia act as ligands or any substance with a lone electron pair
Although coordination complexes are particularly important in the chemistry of the transition metals, some main group elements also form complexes. Aluminum, tin, and lead, for example, form complexes such as the AlF63-, SnCl42- and PbI42- ions.
The amount of ligands joined is dependent on the coordination number for the transition metal.
Ion / Coordination numbers / Ion / Coordination numberCu+ / 2,4 / Mn2+ / 4, 6
Ag+ / 2 / Fe2+ / 6
Au+ / 2,4 / Cu2+ / 4,6
Ni2+ / 4,6
Sc3+ / 6 / Al+3 / 4, 6
Cr3+ / 6
Au3+ / 4
Examples
AgCl + 2 NH3 → [Ag(NH3)2]+
Ni(CN)2 + 2 CN- → [Ni(CN)4]2-
Cu(NH3)42, Cu(H2O)62+.
Molecular Structure
A molecule’s three-dimensional structure has a lot to do with its properties.
Valence Electron Pair Repulsion (VSEPR)
A model to determine the three dimensional structure of a molecule
In this model interactions of valence electrons on different atoms are minimized. They are kept as far apart as possible.
By looking at the central carbon the structure can be labeled using AB notation and the structure can be determined
To determine the molecular geometry you must draw the Lewis diagram. Make sure you do this in the exam.
A molecule with two atoms off a central atom is known as AB2
The shape is linear and the angle is 180 oC
Examples are CO2 and BeH2
Electronic geometry: Trigonal planar,
AB3 contains three atoms off a central atom. The molecular shape is tigonal planar and the angles are 120 oC
Examples are BF3 and CH2O
AB2E example is SO2, two atoms and a lone pair of electrons, shape is bent, angle is 120o
Electronic geometry: tetrahedral,
AB4 Contains 4 atoms off the central atom. The angles are 107.5 o and the shape is tetrahedral. Examples are CH4 or C2H6
AB2E2 example is H2O, two atoms and two lone pairs of electrons, shape is bent, angle is 109o
AB3E example is NH3, Three atoms and a lone pair of electrons, shape is trigonal pyramidal. Angle is 107o
Electronic geometry: Trigonal bipyramidal
AB5 example is PF5, 5 atoms off the central atom. Angles are 120 o and 90 o and the shape is trigonal bipyramidal.
AB2E3 example is I3- , Shape is t shaped, angles are 90o
Electronic geometry octahedral
AB6 example is SF6, 6 atoms off the central atom, shape is ocatahedral, angles are 90o
AB4E2 example is XeF4, shape is square planar, angles are 90o
Exercise 3
Describe the molecular structure of the water molecule.
Exercise 4
When phosphorus reacts with excess chlorine gas, the compound phosphorus pentachloride (PCl5) is formed. In the gaseous and liquid state, this substance consists of PCl5 molecules, but in the solid state it consists of a 1:1 misture of PCl4+ an PCl6- ions. Predict the geometric structures of PCl5, PCl4+, and PCl6-.
Covalent bonds are described based on the electronegativity difference between the two bonds. This is known as polarity
A nonpolar covalent bond has a difference below 0.4
a polar covalent bond has a difference between 0.4 and 1.0
a very polar covalent bond has a difference between 1 and 2
bonds with difference greater than 2.0 are ionic
Don’t need to memorize the numbers
A polar molecule contains a dipole. A dipole contains one end having a slight positive charge and the other end having a slight negative charge due to differences in electronegativity.
Dipole moment is a quantitative measure of a dipole. The greater the dipole moment, the more polar is the molecule
Molecular polarity
A molecule is polar if its bonds are polar and the dipoles do not cancel out in its three-dimensional structure.
HCl, CO, HF are polar because the bond is polar
CO2 is not polar because although the bonds are polar they cancel out because it is a linear molecule.
Which molecules are polar?
SF6
PH3
PF3
CCl4
SH2
Draw CH4 , CH3Cl, CH2Cl2, CHCl3, CCl4 Indicate dipole moment(s) where necessary.
Exercise 2 Bond Polarity and Dipole Moment
For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment:HCL, Cl2, SO3 (a planar molecule with the oxygen atoms spaced evenly around the central sulfur atom), CH4 [tetrahedral (see
Table 8.2) with the carbon atom at the center], and H2S (V-shaped with the sulfur atom at the point).
Exercise 13 Prediction of Molecular Structure III
Because the noble gases have filled s and p valence orbitals, they were not expected to be chemically reactive. In fact, for manyyears these elements were called inert gases because of this supposed inability to form any compounds. However, in the early1960s several compounds of krypton, xenon, and radon were synthesized. For example, a team at the Argonne National
Laboratory produced the stable colorless compound xenon tetrafluoride (XeF4). Predict its structure and whether it has a dipole
moment.
Exercise 14 Structures of Molecules with Multiple Bonds
Predict the molecular structure of the sulfur dioxide molecule. Is this molecule expected to have a dipole moment?
Hybridization model for molecular bonding
Bonding orbitals are a composite of the orbitals that the electrons came from
A sigma bond (s) is formed by the end-to end overlap of orbitals All single bonds are sigma bonds
A pi (p) bond results
· from side-to-side overlap of orbitals
· A double bond consists of one sigma and one pi bond.
· A triple bond results from one sigma and two pi bonds.
· Occurs with sp or sp2 hybridization never sp3
Delocalized electrons are spread out over a number of atoms in a molecule. These are seen when multiple double or triple bonds are seen in molecules.
Sp3 hybridization
All bonding orbitals are have equal energy
This gives four equal energy bonding orbitals which results in a tetrahedral shape.
Examples
CH4 and NH3
How many sigma bonds are in each?
Sp2 Hybridization
Gives a trigonal planar shape
Examples are CH2O, CH3CHCH3
How many sigma and pi bonds are in each?
SP hybridization
Gives a linear shape
Example is C2H2
How many sigma and pi bonds are present?
Hybridization / # of hybrid obitals / Geometry / examplesp / 2 / Linear
sp2 / 3 / Trigonal planar
sp3 / 4 / Tetrahedral
dsp3 / 5 / Trigonal bipyramidal
d2sp3 / 6 / octhedral
Exercise 1
Describe the bonding in the ammonia molecule using the localized electron model.
Exercise 2
Describe the bonding in the N2 molecule.
Exercise 3
Describe the bonding in the triiodide ion (I3-)
Exercise 4
How is the xenon atom in XeF4 hybridized?
Exercise 5
For each of the following molecules or ions, predict the hybridization of each atom, and describe the molecular structure.
a. CO b. BF4- e. XeF2
http://www.khanacademy.org/science/organic-chemistry/v/pi-bonds-and-sp2-hybridized-orbitals
http://www.khanacademy.org/science/organic-chemistry/v/newman-projections
Molecular Orbitals
Formed from atomic orbitals
We use the following procedure when drawing molecular orbital diagrams.
Determine the number of electrons in the molecule. We get the number of electrons per atom from their atomic number on the periodic table. (Remember to determine the total number of electrons, not just the valence electrons.)
Fill the molecular orbitals from bottom to top until all the electrons are added. Describe the electrons with arrows. Put two arrows in each molecular orbital, with the first arrow pointing up and the second pointing down.
· Orbitals of equal energy are half filled with parallel spin before they begin to pair up.
Molecular orbital diagram
We describe the stability of the molecule with bond order.
Bonding Orbital
An orbital of lower energy than the atomic orbitals
Antibonding Orbital
An orbital of higher energy than the atomic orbitals
Bond Order
The difference between the number of bonding electrons and the number of antibonding electrons divided by two. Indicates bond strength.
The higher the bond order the stronger is the bond.
bond order = 1/2 (#e- in bonding MO's - #e- in antibonding MO's)
We use bond orders to predict the stability of molecules.
· If the bond order for a molecule is equal to zero, the molecule is unstable.
· A bond order of greater than zero suggests a stable molecule.
· The higher the bond order is, the more stable the bond.
We can use the molecular orbital diagram to predict whether the molecule is paramagnetic or diamagnetic. If all the electrons are paired, the molecule is diamagnetic. If one or more electrons are unpaired, the molecule is paramagnetic.
Paramagnetism
Certain substances exhibit magnetism when placed in a magnetic field. Paramamagnetism causes the substance to be attracted into the inducing magnetic field. Diamagnetism causes the substance to be repelled from the inducing magnetic field.