Chemistry: Bonding & Chemical Reactions

Question 1

Exceptions to the Octet Rule

Answer 1

Hydrogen can only have two valence electrons.

Lithium and beryllium, which bond to attain to and four valence electrons, respectively.

Boron, which bonds to attain six.

Elements beyond the second row can expand their octets to include more than eight electrons by incorporating d orbitals.

Question 2

Ionic Bonding

Answer 2

An electron(s) from an atom with a smaller ionization energy is transferred to an atom with a greater electron affinity, and the resulting ions are held together by electrostatic forces.

Question 3

Ionic Compounds

Answer 3

High melting and boiling points due to the strong electrostatic forces between ions. Can conduct electricity in the liquid and aqueous states, though not in the solid state. Ionic solids form crystal lattices consisting of arrays of positive and negative ions in which the attractive forces between ions of opposite charge are maximized, while the repulsive forces between ions of like charge are minimalized.

Question 4

Covalent Compounds

Answer 4

Contain discrete molecular units with weak intermolecular forces. Consequently, they are low-melting solids and do not conduct electricity in the liquid or aqueous state.

Question 5

Bond Order

Answer 5

Number of shared electron pairs between two atoms.

Question 6

Bond Length

Answer 6

The average distance between the two nuclei of the atoms involved in the bond. As the number of shared electron pairs increases, the two atoms are pulled closer together, leading to a decrease in bond length. Thus, for a given pair of atoms, a triple bond is shorter than a double bond, which is shorter than a single bond.

Question 7

Bond Energy

Answer 7

Energy required to separate two bonded atoms. For a given pair of atoms, the bond energy increases as the number of shared electron pairs increases.

Question 8

Formal Charge

Answer 8

The # of electrons officially assigned to an atom in a Lewis structure does not always = the # of valence electrons of the free atom. THsi difference between these 2 numbers is the formal charge.

FC = V-1/2Nbonding - Nnonbonding

V = number of valence electrons in atom.
Nbonding = number of bonding electrons
Nnonbonding = number of nonbonding electrons

The formal charge of an ion or molecule is equal to the sum of the formal charges off all the atoms.

Question 9

Resonance Major Contributor

Answer 9

A Lewis structure with small or no formal charges is preferred over a Lewis structure with large formal charges.

A Lewis structure in which negative formal charges are placed on more electronegative atoms is more stable than one in which the formal charges are placed on less electronegative atoms.

Question 10

Dipole Momenet

Answer 10

A vector quantity, defined as the product of the charge magnitude and the distance between the two partial charges.

Measured in Debye units.

Question 11

Coordinate Covalent Bond

Answer 11

The shared electron pair comes from the lone pair of one of the atoms in the molecule. Once such a bond forms, it is indistinguishable from any other covalent bond. Distinguishing such a bond is useful only in keeping track of the valance electrons and formal charges. Coordinate bonds are typically found in Lewis acid-base compounds.

A Lewis Acid is a compound that can accept an electron pair to form a covalent bond.

A Lewis base is a compound that can donate an electron pair to form a covalent bond.

Question 12

VSEPR Theory

Answer 12

The valence shell electron-pair repulsion theory uses Lewis structures to predict the molecular geometry of covalently bonded molecules. It states that the 3-D arrangement of atoms surrounding a central atom is determined by the repulsions between the bonding and the nonbonding electron pairs in the valence shell of the central atom. These electron pairs arrange themselves as far apart as possible, thereby minimizing repulsion.

Question 13

VSEPR Theory: Steps

Answer 13

The following steps are used to predict geometrical structure of a molecule.

1) Draw Lewis structure of molecule.
2) Count the total number of bonding and nonbonding electron pairs in the valence shell of the central atom.
3) Arrange the electron pairs around the central atom so that they are as far apart from each other as possible. For example, the compound AX2 has the Lewis structure X:A:X. A has two bonding electron pairs in its valence shell. To make these electron pairs as far apart as possible, their geometric structure should be linear.

X-A-X

Study page 353 for other shapes.

Question 14

Valence Electron Arrangements

Answer 14

Linear
2 Regions of Electron Density
Ex) BeCl2
180 Degree Angle Between Electron Pairs

Trigonal Planar
3
BH3
120 Degrees

Tetrahedral
4
CH4
109.5 Degrees

Trigonal Bipyramidal
5
PCl5
90, 120, 180 Degrees

Octahedral
6
SF6
90, 180 Degrees

Question 15

Polarity of Molecules

Answer 15

Polarity of a molecule depends on the polarity of the constituent bonds and on the shape of the molecule. A molecule with nonpolar bonds is always nonpolar, a molecule with polar bonds may be polar or nonpolar depending on the orientation of the bond dipoles.

A molecule of two atoms bound by a polar bond must have a net dipole moment and therefore be polar. The two equal and opposite partial charges are localized at the ends of the molecule on the two atoms. A molecule consisting of more than two atoms bound with polar bonds may be either polar or nonpolar, since the overall dipole moment of a molecule is the vector sum of the individual bond dipole moments. If the molecule has a particular shape such that the bond dipole moments cancel each other, then the result is a nonpolar molecule.

For example, CCl4 has a tetrahedral shape. The four bond dipoles point to the verticies of the tetrahedron and cancel each other, resulting in a nonpolar molecule.

However, if the orientation of the bond dipoles are such that they do not cancel out, the molecules will have a net dipole moment and be polar. For example, H2O has two polar O-H bonds. Its shape is angular. The two dipoles add together to give a net dipole moment to the molecule, making it polar.

Question 16

Atomic and Molecular Orbitals

Answer 16

When two atoms bond to form a molecule, the atomic orbitals interact to form a molecular orbital that describes the probability of finding the bonding electrons. Miolecular orbitals are obtained by adding the wave functions of the atomic orbitals. Quantitatively, this is described by the overlap of two atomic orbitals.

If the signs of the two atomic orbitals are the same, a bonding orbital is formed. If the signs are different, an antibonding orbital is formed.

In addition, two different types of overlap are possible. When orbitals overlap head-to-head, the resulting bond is called a sigma bond. When orbitals are parallel, a pi bond is formed.

Question 17

Dipole-Dipole Interactions

Answer 17

Polar molecules tend to orient themselves such that the positive region of one molecule is close to the negative region of another molecule. This arrangement is energetically favorable because an attractive dipole force is formed between the two molecules.

Dipole-dipole interactions are present in the solid and liquid phases but become negligible in the gas phase because the molecules are generally much farther apart. Polar species tend to have higher boiling points than nonpolar species of comparable molecular weight.

Question 18

Hydrogen Bonding

Answer 18

A specifically, unusually strong form of dipole-dipole interaction, which may be either intra- or intermilecular. When hydrogen is bound to a highly electronegative atom such as fluorine, oxygen, or nitrogen, the hydrogen atom carries little of the electron density of the covalent bond. This positively charged hydrogen atom interacts with the partial negative charge located on the electronegative atoms of nearby molecules. Substances that display hydrogen bonding tend to have unusually high boiling points compared with compounds of similar molecular formula that do not hydrogen bond. The difference derives from the energy required to break the hydrogen bonds. Hydrogen bonding is important in the behavior of water, alcoholds, amines, and carboxylic acids.

Question 19

Dispersion Forces

Answer 19

The bonding electrons in covalent bonds may appear to be shared equally between two atoms, but at any particular point in time, they will be located randomly throughout the orbital. This permits unequal sharing of electrons, causing rapid polarization and counterpolarization of the electron cloud and formation of short-lived dipoles. These dipoles interact with the electron clouds of neighboring molecules, inducing the formation of more dipoles. The attractive interactions of these short-lived dipoles are called dispersion of London force or van der Waals forces.

Dispersion forces are generally weaker than other intermolecular forces. They don’t extend over long distances and are therefore most important when molecules are close together. The strength of these interactions within a given substance depends directly on how easily the electrons in the molecules can move (be polarized). Large molecules in which the electrons are far from the nucleus are relatively easy to polarize and therefore possess greater dispersion forces. If not for dispersion forces, the noble gases would not liquefy at any temperature since no other IMF exist between the noble gas atoms. The low temperature at which the noble gases liquefy is to some extent indicative of the magnitude of dispersion forces between the atoms.

Question 20

Carbon-Carbon Bonding

Answer 20

Carbon-carbon bonds can be categorized based on length and energy level, as well as hybridization.

With a single bond, the bond length is the longest but bond energy is the lowest.

With a double bond, both the bond length and bond energy are in the middle.

With a triple bond, bond length is the shortest and bond energy is the highest.