|Chapter 9 Chemical Bonding I: Basic Concepts
Identify the valence electrons for all representative elements
The representative elements are divided into eight groups, which can tell you how many valence electrons are in a particular element of that one.
Group One: Hydrogen, Helium, Sodium, Potassium, Rubidium, Caesium, Francium. 1 Valence electron each.
Group Two: Beryllium, Magnesium, Calcium, Strontium, Barium, Radium. 2 Valence electrons each.
Group Three: Boron, Aluminum, Gallium, Indium, Thallium, Ununtrium. 3 valence electrons each.
Group Four: Carbon, Silicon, Germanium, Tin, Lead, Flerovium. 4 valence electrons each.
Group Five: Nitrogen, Phosphorus, Arsenic, Antimony, Bismuth, Ununpentium. 5 valence electrons each.
Group Six: Oxygen, Sulfur, Selenium, Tellurium, Polonium, Livermorium. 6 valence electrons each.
Group Seven: Fluorine, Chlorine, Bromine, Iodine, Astatine, Ununseptium. 7 valence electrons each.
Group Eight: Neon, Argon, Krypton, Xenon, Radon, Ununoctium. 8 valence electrons each. (Helium has 2)
Rationalize why alkali metals and alkaline earth metals usually form cations and oxygen and the halogens usually form anions using Lewis dot symbols in the discussion
The alkali metals and alkaline earth metals are located in groups 1A and 2A respectively which mean that they have 1 or 2 valence electrons. Oxygen and the halogens are located in groups 6A and 7A respectively which means they have 6 or 7 valence electrons.
All elements (save helium) have a desire to reach eight valence electrons as easy as possible. So it would be easier for an element like lithium, with two valence electrons to lose two and end up with 8 than to try to gain 6. Just as it would be easier for oxygen, with six valence electrons to gain 2 and end up with 8 than to try to lose 6.
EX: The lewis dot diagram for lithium would consist of a single valence electron on it, and has a strong desire to reach eight valence electrons. Fluorine’s lewis dot diagram would consist of seven valence electrons, also wishing to reach eight valence electrons. These two elements could therefore form a chemical bond, allowing them both to receive these eight valence electrons.
Use Lewis dot symbols to show the formation of both ionic and molecular compounds
An ionic bond is the electrostatic force that holds ions together in an ionic compound. Ionic compounds are compounds that contain only ionic bonds. An easy example of an ionic compound would be calcium burning in oxygen to form calcium oxide (see pg. 372 for a lewis dot representation)
A covalent bond is is a bond in which two electrons are shared by two atoms. Covalent compounds are compounds that contain only covalent bonds. An example would be a water molecule (see pg. 380 for a lewis dot representation of this)
Define lattice energy, Coulomb’s law, and the Born-Haber cycle
A quantitive measure of the stability of any ionic solid is its lattice energy, defined as the energy required to completely separate one mole of a solid ionic compound into gaseous ions.
Coulomb’s law 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.
The Born-Haber cycle relates lattice energies of ionic compounds to ionization energy electron affinities, and other atomic and molecular properties.
Demonstrate how the Born-Haber cycle is an application of Hess’s law and use the Born-Haber cycle to determine lattice energy for an ionic solid.
Hess’s law states that the overall change in energy of a process can be determined by breaking the process down into steps, then adding the changes in energy of each step.
The Born-Haber cycle applies Hess’s law to an ionic solid. It determines lattice energy indirectly by summing that the formation of an ionic compound takes place in a series of steps. These five steps help break down the process of determining the lattice energy of the ionic solid and you can add the changes in energy of each step.
See Figure 9.2 on pg 376 for an example on applying the Born-Haber cycle.
Identify covalent compounds, the type of covalent bonds present, and the number of lone pairs of electrons using the Lewis structures
Covalent bonds are bonds in which electrons are shared by two atoms. Covalent compounds are compounds that contain only covalent bonds.
9. Identify ionic, polar covalent and nonpolar covalent bonds using the concepts of electronegativity.
Electronegativity is the ability of the atom to attract shared electrons to itself. Bonds with very high electronegativity differences tend to be ionic, middle tend to be polar covalent, and low tend to be nonpolar covalent.
0-.3 nonpolar covalent
.3-1.67 polar covalent
10. Predict the relative charges in electronegativity with respect to position on the periodic table.
Electronegativity becomes stronger as you go to the right and up on the periodic table. Meaning that Fluorine has the highest electronegativity value of the elements while Francium has the lowest. F has a very strong nuclear charge, and a very small amount of shielding occurring which allows the atom to easily “hold onto” other electrons floating around. While the nuclear charge of Fr is much lower and it has a lot of shielding from its energy levels, so it doesn’t have the energy to hold onto extra electrons.
11. Use Lewis dot and the octet rule to write Lewis structures of compounds and ions.
7e- Na •
Chloride Ion: Sodium Ion:
12. Apply the concept of formal charge to predict the most likely Lewis structure of a compound.
Formal charge=total number of valence electrons-(total number of lone electrons+number of bonds connected to the atom)
H= 1-(0+1)=0 H= 1-(0+1)=0
C= 4-(0+4)=0 C= 4-(2+3)=-1
N= 5-(2+3)=0 N= 5-(0+4)=1
Compounds want to have each element to have a formal charge of zero so the first structure would be more stable/correct than the second structure.
13. Explain how lewis dot structures are inadequate to explain observed bond length (bond types) in some compounds and how the concept of resonance must be invoked.
Resonance is when a compound has 2 or more correct structures. Here, Nitrogen Dioxide is shown two different ways, but both ways are correct. Each bond is really a 1.5 bond rather than one single bond and one double bond.
14. Recall several events in which the octet rule fails.
Be is stable with 2 valence, B is often found with 6, and any element in the 3rd row or higher can use the 3d orbital for extra electrons so they can often have more than 8 valence electrons
15. Use Lewis structures and bond energy to predict heats of reaction.
(sum of reactants)-(sum of products)=bond enthalpy
simply match each bond to the number corresponding to the chart (usually given, or try Table 9.4 on page 401)and plug the numbers into the equation.
16. Rationalize why enthalpy change for breaking chemical bonds is positive and the formation of bonds is negative.
A bond is usually in place so that an atom can become more stable and have a lower energy level. Atoms want to be more stable, so when they bond, their extra energy is released in the form of heat. When bonds are broken, it needs energy to occur since the atoms are going from a stable state to a less stable state (something that it does not want to do).
Important equations and tables are highlighted
Which of the following occurs in an ionic bond?
a. Oppositely charged ions attract.
b. Two atoms share two electrons.
c. Two atoms share more than two electrons.
d. Like-charged ions attract.
Which of the following covalent bonds is the most polar?
a. H—F c. H—H
b. H—C d. H—N