8.1 - Types of Chemical Bonds

• Ionic bonding: Bonding forces that produce great thermal stability result from the electrostatic attractions of the closely packed, oppositely charged ions

• Ionic substances are formed when an atom that loses electrons relatively easily reacts with an atom that has a high affinity for electrons

• The energy of interaction between a pair of ions can be calculated using Coulomb’s law

  • Coulomb’s law also can be used to calculate the repulsive energy when two like-charged ions are brought together

• A bond will form if the energy of the aggregate is lower than that of the separated atoms

• The distance where the energy is minimal is called the bond length

• The energy terms involved are the net potential energy that results from the attractions and repulsions among the charged particles and the kinetic energy due to the motions of the electrons

• The zero point of energy is defined with the atoms at infinite separation

• At very short distances the energy rises steeply because of the importance of the repulsive forces when the atoms are very close together.

• The bond length is the distance at which the system has minimum energy

• Ionic and covalent bonds are the extreme bond types

  • In ionic bonding, the participating atoms are so different that one or more electrons are transferred to form oppositely charged ions, which then attract each other

  • In covalent bonding, two identical atoms share electrons equally

  • The bonding results from the mutual attraction of the two nuclei for the shared electrons.

• Because bond polarity has important chemical implications, we find it useful to quantify the ability of an atom to attract shared electrons

8.2 - Electronegativity

• Electronegativity: The ability of an atom in a molecule to attract shared electrons to itself

• The greater is the difference in the electro-negativities of the atom

• Electronegativity generally increases going from left to right across a period and decreases going down a group for the representative elements

• For identical atoms (an electronegativity difference of zero), the electrons in the bond are shared equally, and no polarity develops

• When two atoms with very different electronegativities interact, electron transfer can occur to form the ions that make up an ionic substance

8.3 - Bond Polarity and Dipole Moments

• A molecule such as HF that has a center of positive charge and a center of negative charge is said to be dipolar

• The dipolar character of a molecule is often represented by an arrow pointing to the negative charge center with the tail of the arrow indicating the positive center of charge

• Any diatomic (two-atom) molecule that has a polar bond also will show a molecular dipole moment

• Some molecules have polar bonds but do not have a dipole moment, in which occurs when the individual bond polarities are arranged in such a way that they cancel each other

• The presence of polar bonds does not always yield a polar molecule

8.4 - Ions: Electron Configurations and Sizes

• Atoms in stable compounds usually have a noble gas electron configuration

• When two nonmetals react to form a covalent bond, they share electrons in a way that completes the valence electron configurations of both atoms. That is, both nonmetals attain noble gas electron configurations

• 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.

• These generalizations apply to the vast majority of compounds and are important to remember

• In the solid state of an ionic compound, the ions are relatively close together, and many ions are simultaneously interacting

• In the gas phase of an ionic substance, the ions would be relatively far apart and would not contain large groups of ions

• The transition metals exhibit more complicated behavior, forming a variety of ions

  • Ion size plays an important role in determining the structure and stability of ionic solids, the properties of ions in aqueous solution, and the biological effects of ions

  • Since a positive ion is formed by removing one or more electrons from a neutral atom, the resulting cation is smaller than its parent atom.

  • The opposite is true for negative ions; the addition of electrons to a neutral atom produces an anion significantly larger than its parent atom

• It is also important to know how the sizes of ions vary depending on the positions of the parent elements in the periodic table

• The changes that occur horizontally are complicated because of the change from predominantly metals on the left-hand side of the periodic table to nonmetals on the right-hand side.

• Isoelectronic ions: Ions containing the same number of electrons.

• For isoelectronic ions, size decreases as Z increases

8.5 - Energy Effects in Binary Ionic Compounds

• Lattice energy: The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

• To see the energy terms associated with this process, we take advantage of the fact that energy is a state function and break this reaction into steps, the sum of which gives the overall reaction

• The importance of the charges in ionic solids can be illustrated by comparing the energies involved in the formation of NaF(s) and MgO(s)

• For sodium fluoride, the extra energy required to form the doubly charged ions is greater than the gain in lattice energy that would result.

• The most important of the factors involve the balancing of the energies required to form highly charged ions and the energy released when highly charged ions combine to form the solid

8.6 - Partial Ionic Character of Covalent Bonds

• There are probably no totally ionic bonds between discrete pairs of atoms

• None of the bonds reaches 100% ionic character, even though compounds with the maximum possible electronegativity differences are considered

• A complication in identifying ionic compounds is that many substances contain polyatomic ions

• Any compound that conducts an electric current when melted will be classified as ionic

8.7 - The Covalent Chemical Bond: A Model

• Chemical bonds can be viewed as forces that cause a group of atoms to behave as a unit

• A tetrahedron has four equal triangular faces

• There is no principle of nature that states that bonds are favored or disfavored. Bonds are neither inherently “good” nor inherently “bad” as far as nature is concerned

• Bonds occur when collections of atoms are more stable (lower in energy) than the separate atoms

• Bonds provide a method for dividing up the energy evolved when a stable molecule is formed from its component atoms.

• Bonding is a model proposed to explain molecular stability

• It is very important to understand the assumptions inherent in a particular model before you use it to interpret observations or to make predictions

• To understand reactions, we must understand the forces that bind atoms together

• It makes sense that atoms can form stable groups by sharing electrons; shared electrons give a lower energy state because they are simultaneously attracted by two nuclei

• Collections of atoms, therefore, occur because the aggregated state has lower energy than the separated atoms

8.8 - Covalent Bond Energies and Chemical Reactions

• One important consideration is to establish the sensitivity of a particular type of bond to its molecular environment

• We use the average of these individual bond dissociation energies even though this quantity only approximates the energy associated with a COH bond in a particular molecule

• Atoms sometimes share two pairs of electrons, forming a double bond, or share three pairs of electrons, forming a triple bond

• Bond energy values can be used to calculate approximate energies for reactions

  • The formation of a bond releases energy, an exothermic process, so the energy terms associated with bond making carry a negative sign

8.9 - The Localized Electron Bonding Model

• Localized electron model: A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms

• Electron pairs in the molecule are assumed to be localized on a particular atom or in the space between two atoms

• Those pairs of electrons localized on an atom are called lone pairs, and those found in the space between the atoms are called bonding pairs

• LE Model has three parts: Description of the valence electron arrangement in the molecule using Lewis structures

  • Prediction of the geometry of the molecule using the valence shell electron-pair repulsion model

  • Description of the type of atomic orbitals used by the atoms to share electrons or hold lone pairs

8.10 - Lewis Structures

• The Lewis structure of a molecule shows how the valence electrons are arranged among the atoms in the molecule

• The rules for writing Lewis structures are based on observations of thousands of molecules

  • From experiments, chemists have learned that the most important requirement for the formation of a stable compound is that the atoms achieve noble gas electron configurations

  • Lewis structures show only valence electrons

  • The rules: Sum the valence electrons from all the atoms. Do not worry about keeping track of which electrons come from which atoms. It is the total number of electrons that is important

  • Use a pair of electrons to form a bond between each pair of bound atoms

  • Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the second-row elements

• The best way to look at a molecule is to regard it as a new entity that uses all the available valence electrons of the atoms to achieve the lowest possible energy.

• The valence electrons belong to the molecule, rather than to the individual atoms

• Carbon, nitrogen, oxygen, and fluorine always obey the octet rule in stable molecules

8.11 - Exceptions to the Octet Rule

• Boron tends to form compounds in which the boron atom has fewer than eight electrons around it—it does not have a complete octet

• Boron trifluoride (BF3), a gas at normal temperatures and pressures, reacts very energetically with molecules such as water and ammonia that have available electron pairs

• Carbon, nitrogen, oxygen, and fluorine can be counted on to obey the octet rule

• The second-row elements B and Be often have fewer than eight electrons around them in their compounds. These electron-deficient compounds are very reactive

• Third-row elements can exceed the octet rule

• Whether the atoms that exceed the octet rule actually place the extra electrons in their d orbitals is a matter of controversy among theoretical chemist

• When it is necessary to exceed the octet rule for one of several third-row (or higher) elements, assume that the extra electrons should be placed on the central atom

8.12 - Resonance

• Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule

• The resulting electron structure of the molecule is given by the average of these resonance structures

• The concept of resonance is necessary because the localized electron model postulates that electrons are localized between a given pair of atoms

  • Also, it compensates for the defective assumption of the localized electron model

• Since the localized electron model is based on pairs of electrons, it does not handle odd-electron cases in a natural way

• Equivalent Lewis structures contain the same numbers of single and multiple bonds. For example, the resonance structures for O3 are equivalent to Lewis structures

• The charge on an atom in a molecule, the formal charge, can be used to evaluate Lewis structures.

• Formal charge: The difference between the number of valence electrons

• To determine the formal charge of a given atom in a molecule, we need to know two things: The number of valence electrons on the free neutral atom

  • The number of valence electrons “belonging” to the atom in a molecule

• To calculate the formal charge on an atom: Take the sum of the lone pair electrons and one-half the shared electrons. This is the number of valence electrons assigned to the atom in the molecule.

  • Subtract the number of assigned electrons from the number of valence electrons on the free, neutral atom to obtain the formal charge

• Tests based on experiments must be used to make the final decisions on the correct description of the bonding in a molecule or polyatomic ion

8.13 - Molecular Structure: The VSEPR Model

• Based on the idea that electron pairs will be arranged around a central atom in a way that minimizes the electron repulsions

• It can be used to predict the geometric structure of most molecules

• The main postulate of this model is that the structure around a given atom is determined principally by minimizing electron-pair repulsions

• Be Cl2 has only four electrons around Be and is expected to be very reactive with electron-pair donors.

• Whenever four pairs of electrons are present around an atom, they should always be arranged tetrahedrally

• In counting pairs, count each multiple bonds as a single effective pair

• The arrangement of the pairs is determined by minimizing electron-pair repulsions

• Lone pairs require more room than bonding pairs and tend to compress the angles between the bonding pairs

• To use the VSEPR model to determine the geometric structures of molecules, you should memorize the relationships between the number of electron pairs and their best arrangement

• For the VSEPR model, multiple bonds count as one effective electron pair

  • The molecular structure of nitrate also shows us one more important point: When a molecule exhibits resonance, any one of the resonance structures can be used to predict the molecular structure using the VSEPR model