chapter9

Chapter 9 Chemical Bonding I: Basic Concepts

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=**__Section 9.1- Lewis Dot Symbols__**=

Gilbert Lewis was a physical chemist that discovered the covalent bond. He also explained that atoms combine in order to have a more stable electron configuration. He also mentioned that more stability occurs when an atom is isoelectronic with a noble gas.

With chemical bonding, chemists are concerned with valence electrons. They use a system created by Lewis known as the Lewis dot symbols. A Lewis dot symbol consists of the symbol of an element and one dot for each valence electron in an atom of an element.

Each group on the periodic table has a different amount of valence electrons. Elements in the same group have a similar outer electron configuration, which means that they will have similar electron Lewis dot symbols.

__ Examples: __ __Problems:__ Write the Lewis symbol for the following elements: a) S, b) Li, c) C
 * Group || Number of Valenece Electrons ||
 * 1A || 1 ||
 * 2A || 2 ||
 * 3A || 3 ||
 * 4A || 4 ||
 * 5A || 5 ||
 * 6A || 6 ||
 * 7A || 7 ||
 * 8A || 8 ||

=**__Section 9.2- The Ionic Bond__**=

Elements most likely to form cations in ionic compounds are the alkali metals and alkaline earth metals. Elements most likely to form anions are the halogens and oxygen. As a result, elements from Groups 1A or 2A would create an ionic compound with a halogen or oxygen.

An ionic bond is the electrostatic force that holds ions together in an ionic compound.

__Example:__

For NaCl, sodium has one valence electron while chlorine has seven. An ionic bond is formed by transferring the one valence electron from the sodium atom to the chlorine atom. This slows the chlorine atom to have a full outer shell with eight electrons.



__Problems:__ Write the Lewis dot structure of the reactants and products in the following reactions:

=__**Section 9.3- Lattice Energy of Ionic Compounds**__=

Lattice energy is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. This energy measures the stability of ionic solids.

Lattice energy can be calculated by using Coulomb's law. This states that the potential energy between two ions is directly proportional to the product of their charges and inversely proportional to the distance of separation between them.


 * Coulomb's Law **

Another way to calculate lattice energy is by following the Born-Haber cycle, which relates lattice energies of ionic compounds to ionization energies, electron affinities, and other atomic and molecular properties. This is based on Hess' Law.

__Example:__

Lattice energy is always positive quantity because the separation of ions in a solid into ions in the gas phase is an endothermic process. The larger the energy, the more stable the ionic compound is. They also have a higher melting point. The value can also help explain the formulas of ionic compounds.

__Problems:__ A negative charge of 3 * 10^(-4) C and a positive charge of 7 * 10^(-4) C are separated by .7 cm. What is the force between these two charges?

Calculate the lattice energy of CsCl.

__Activities:__ [|Ionic Bonds and Lattice Energy]

=**__Section 9.4- The Covalent Bond__**=

As Lewis said, a chemical bond involves sharing electrons.

__Example:__

This is an example of a covalent bond, which is a bond in which two electrons are shared by two atoms. Covalent compounds are then formed, which are compounds that contain only covalent bonds. The shared pair of electrons is usually written as a single line. This is how the bond would be drawn.

Covalent bonding between many electron atoms involves the valence electrons. The majority of covalent compounds have lone pairs that are not involved in the covalent bond formation. The structures used to represent the elements in the covalent compounds are Lewis structures, which are a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms. Only the valence electrons are shown.

The formation of molecules illustrates the octet rule, which was formed by Lewis. The octet rule is an atom other than hydrogen that tends to form bonds until it is surrounded by eight valence electrons.

There are different kinds of covalent bonds. A single bond exists when two atoms are held together by one electron pair. Multiple bonds are made when two atoms share two or more pairs of electrons. In a double bond, two atoms share two pairs of electrons. A triple bond is formed when two atoms share three pairs of electrons.

Different bonds have different lengths. Bond length is the distance between the nuclei of two covalently bonded atoms in a molecule.

There are differences between covalent and ionic bonds. Covalent compounds have two types of forces. The first type is the force that holds the atoms together in a molecule, which is given by bond energy. Intermolecular force is the second type, which is between molecules. This is not as strong as the forces holding atoms together. Since they are not held together tightly, covalent compounds are usually gases, liquids, or low-melting solids. In ions, the electrostatic forces holding the ions together are very strong. That means that ionic compounds are solids at room temperature and have high melting points. They are also soluble in water, and they can conduct electricity because the compounds have strong electrolytes. On the other hand, covalent bonds do not have electrolytes, so most of the compounds are insoluble and do not conduct electricity.

__Problems:__ Draw the Lewis dot structure for the following covalent compounds:

__Activities:__ [|Covalent Bonding]

=**__Section 9.5- Electronegativity__**=

In a polar covalent bond, or a polar bond, the electrons spend more time in the vicinity of one atom than the other. This is like a partial electron transfer or a shift in electron density.



Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electrons that are more likely to attract electrons will have a higher electronegativity. This is also related to electron affinity and ionization energy. Electronegativity of elements can only be measured in relation to other elements. Linus Pauling created a way to calculate relative electronegativities.



As shown in the image, electronegativity increases from left to right as the metallic character decreases. It also increases within a group as the atomic numbers get smaller. The transition metals do not follow these trends however. The elements grouped in the upper right hand corner have the highest electronegativity, while the elements in the lower left hand corner have the lowest electronegativity.

Elements that have very different electronegativities tend to form ionic bonds. This is usually between a metallic element and a nonmetallic element. Elements with similar electronegativities tend to form polar covalent bonds because the shift in electron density is small. This bond is usually between nonmetallic elements. A pure covalent bond can only be between to atoms of the same element.

When taking the difference of electronegativities between two atoms of a bond, one can predict what kind of bond is formed.

D Energy > 1.7 (Ionic) 1.7 ≥ D Energy ≥ .5 (Polar) D Energy < .5 (Nonpolar)

Electronegativity is the basis of the rules for assigning oxidation numbers. The oxidation number refers to the number of charges an atom would have if electrons were transferred completely to the more electronegative of the bonded atoms in a molecule.

__Example:__ There is an oxidation number of -3 for N and +1 to H.

There is an exception. The bond makes no contribution to the oxidation number because the electron pair in the bond is equally shared.

__Problems:__ List the following bonds in order of increasing ionic character: a) carbon to hydrogen, b) bromine to hydrogen, c) lithium to chlorine

__Lab:__ [|Chemical Bonding]

__Activities:__ [|Electronegativity]

=**__Section 9.6- Writing Lewis Structures__**=

These are the steps in order to draw Lewis structures.
 * 1) Sketch the skeletal structure of the compound. The central atom is has the least amount of electronegativity.
 * 2) Total valence electrons (charged) divided by 2. This gives the amount of pairs.
 * 3) Connect the atoms from the surrounding atoms to the central atom by drawing covalent bonds. Complete the octets of outer atoms first. If there are extra electrons, they must be shown as lone pairs.
 * 4) Only add double or triple bonds if the central atom has fewer than eight electrons. This is possible by using the lone pairs from the surrounding atoms in order to obey the octet rule.

__Examples:







Problems:__ Write the Lewis strucutures for the following molecules and ions:

__Activities:__ [|Drawing Basic Structures]

=**__Section 9.7- Formal Charge and Lewis Structure__**=

As mentioned before, electrons that are shared in a bond must divide the electrons in a bonding pair equally between the atoms forming the bond. Formal charge is the electrical charge difference between the valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.

There are two steps in assigning the number of electrons on an atom in a Lewis structure.
 * 1) All the atom's nonbinding electrons are assigned to the atom.
 * 2) Break the bond(s) between the atom and other atom(s) and assign half of the bonding electrons to the atom.

There are three rules for writing formal charges. This does not represent the actual charge separation within the molecule.
 * 1) For molecules, the sum of the formal charges must add up to zero because molecules are electrically neutral species.
 * 2) For cations, the sum of the formal charges must equal the positive charge.
 * 3) For anions, the sum of the formal charges must equal the negative charge.

Sometimes there is more than one Lewis structure that can be represented. By using these guidelines, one can figure out which Lewis structure is correct.
 * 1) For molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.
 * 2) Lewis structures with large formal charges are less plausible than those with small formal charges.
 * 3) Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.

__Activities:__ [|Formal Charge Calculations Animation]

=**__Section 9.8- The Concept of Resonance__**=

A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. Resonance means the use of two or more Lewis strucutres to represent a particular molecule. Neither resonance structure for a molecule adequately represents the actual molecule.

=**__Section 9.9- Exceptions to the Octet Rule__**=

There are three categories characterized by an incomplete octet.
 * 1) The Incomplete Octet
 * 2) Odd-Electron Molecules
 * 3) The Expanded Octet

Incomplete octet occurs when the number of electrons surrounding the central atom in a stable molecule is fewer than eight.



A coordinate covalent bond is a covalent bond in which one of the atoms donates both electrons.

The odd-electron molecules contain an odd amount of electrons. One needs to find an electron to be paired with the extra electron in order to make an even number.



The expanded octet occurs when there are more than eight electrons surrounding the central atom. These involve atoms of elements that are in the third period and beyond.



Sometimes the octet rule is satisfied for all atoms, but there are still valence electrons left. These are shown as lone pairs.



__Problems:__

=**__Section 9.10- Bond Energy__**=

Bond energy is the enthalpy change required to break a particular bond in 1 mole of gaseous molecules.

__Example:__

There are different enthalpies for different reactions. The variations are accounted for by looking at the stability of individual reactant and product molecules. Knowing the bond energies and the stability of molecules explains the thermochemical nature of reactants that molecules undergo. One can predict the enthalpy of a reaction by using the average bond energies. One would need to count the total number of bonds broken and formed during the reaction. The reaction is in its gas phase.

Here is the equation used: // D H°// = enthalpy change for the reaction in the gas phase S = summation sign BE (reactants) = average energy for each bond in the gaseous reactant molecules BE (products) = average energy for each bond in the gaseous product molecules

If // D H°// is positive, then the reaction is endothermic because more energy is absorbed than released. If more energy is released than absorbed, then // D H°// is negative and the reaction is exothermic.



__Activities:__ [|Change in Enthalpy]