AP Chemistry Chapter 7: Bonding Types of Bonding I. holds everything together! I All bonding occurs because of! Electronegativity difference and bond character A. A difference in electronegativity between 2 atoms results in a more bond character when those two atoms form a bond. B. Ionic and covalent bonding are not distinct categories as all bonds have aspects of both. It s just that one is more prevalent than the other. C. EN is not the only factor. Whether they are or also matter. Covalent Bonding ( bonding A. Bond that exists between. B. Involves of valence electrons. C. Includes: V. Ionic Bonding 1. Polar covalent - a bond that has an of electrons, causing a. a. hydrogen bonding: (intermolecular force that exists when hydrogen is bonded to fluorine, oxygen or nitrogen). EN difference. b. dipole-dipole: intermolecular force between molecules. EN difference. 2. Nonpolar Covalent - a bond that has an of electrons that can only be polarized ( Force). EN Difference A. Bond that exists between a and B. Involves the of valence electrons from the to the. C. Ionic solids have a structure. Page 1 of 7
VI. Metallic Bonding Nature of Bonds A. Bonding that exists between B. Electrons are but are held. C. Electrons act like a (not associated with a specific atom) and valence electrons 1. gives rise to properties such as and 2. the more or the more, the more, (this is very general - not a rule). I. Covalent bond nature A. Two nuclei have multiple acting within it. 1. between the nucleus and electrons 2. between electrons B. forces overcome at an optimal. 1. According to the graph above, Å is the optimal bond length 2. This is the distance of potential energy where the and forces are. C. electrons hold bond together. Ionic bond nature A. Metals have ionization energies and would more likely electrons to a (which have electronegativities). B. Ionic structures are held together by (energy required to completely 1 mol of compound into its forms). C. Ionic compounds are more stable because all ions involved have configurations. D. Cations and anions in an ionic crystal are arranged in a (periodic 3-D array) that maximizes the forces between ions while repulsive forces. Page 2 of 7
I Properties: A. Ionic compounds are (crystal lattice) so they are and and have high. B. Molecules form ordered structures so they have melting points and have properties. Bottom line: All elements want to because they form electron configurations. V. One more thing... A. Ionic bond strength 1. Charge density: The amount of per. 2. The the charge density, the the electrostatic attraction. 3. Predicted by Coulomb s Law: F= k e = q 1 &q 2 = r= 4. Electrostatic attraction is... a. proportional to b. inversely proportional to the 5. The the electrostatic attraction, the the ionic bond, the the melting point! Lewis Structures Covalent Bonding I. All atoms are more with surrounding electrons (Octet Rule!) A. Except for hydrogen ( Rule) B. Some atoms with empty sublevels can have octets ( eight surrounding electrons). Lewis Structures show how molecules electrons. A. A line represents a B. Unshared/lone pair a pair of electrons that Page 3 of 7
I How to draw Lewis Structures A. Count the number of e-. B. the atoms. is usually in the center. and are normally on the ends. C. All elements will have e-, except for H and He, which will have e-. Electrons will have to be (sometimes than 1 pair) to accomplish this. D. Draw lines to represent the bonds. E. Check your drawing: Is the number of e- correct? Is everyone happy (8 electrons)? F. Examples: Resonance structures: molecules that have possible Lewis structure. A. Shown through a double sided arrow: B. Actual structure is an (like an ) of all the resonance structures. C. Therefore, the is an of the bonds present. D. Bond order: the of chemical bonds between a pair of atoms. 1. Single bond =, double bond=, triple bond= 2. Bond orders for resonance structures are an of the bond orders of all the bonds present. For example: For nitrate, the bond order is. E. Bond length: with all things being equal, a bond is longer than a bond, etc. 1. The bond length of a bond with a bond order of 1.33 is that of a (BO=1) and (BO=2) bond. 2. If the bond order is the, the of the atoms in the bond determines the relative length. Page 4 of 7
V. Ions: Don't forget to place around ions and to state the outside the brackets. VI. Exceptions to Octet Rule A. Some are stable with less electrons: B. Some can be stable with more electrons: 1. Because empty orbitals can extra electrons. 2. Place any electrons on the atom of these even if it already has eight! V VI Acidic and basic Lewis structures: the is bonded to the if present. A more thorough method of checking your structure: A. Molecules will be in the structure in which the formal charge for each atom is closest to. B. Formal charge = C. formal charge for each atom in structure! D. This explains the (Ex. BeF 2 ). VSEPR Covalent Bonding I. VSEPR states that each other. A. Electrons will orient themselves as possible. B. Lone pairs, since they are not shared or involved in a bond, will take up! C. This is used to predict (Think in 3-D!!!) Refer to the table summarizing the shapes. (Btw: means!) Page 5 of 7
Polarity I. Electronegativity (Table 6.5 on p. 154): ability to attract into a chemical bond. A. Used to determine bond character/type. B. If difference is... 1. Ionic: 2. Polar Covalent: 3. Nonpolar Covalent: 4. For 1.6<2.0, the bond is if are involved. The bond is if are involved. A. The trend in is the same as the trend regarding and Dipole: Molecule that has slight charge ( ) on one end and a slight charge ( ) on the other. I Dipole moment shows of charge ( ) How to determine polarity: Orbital Hybridization A. If all dipole moments (polar bonds), it is. If they, the molecule is ( distribution of electrons). B. Generally, symmetrical=nonpolar; asymmetrical=polar. DO NOT explain using only the symmetry argument! I. Shapes we know don't match shapes in Ch. 6!!! A. As elements form a, their orbitals become a of the orbitals containing the bonding electrons. B. Determine hybrid orbitals through number (total number of bonds and lone pairs) 1. Hybridization = the of atoms AND surrounding the central atom (see table) 2. Ex: NH 3 = lone pair & atoms= orbitals. Thus, it is hybridized. Page 6 of 7
Multiple Bonds (memorize) A. A single bond is known as a ( ) bond. B. A double bond contains one ( ) bond and one ( ) bond (unhybridized p orbital). C. A triple bond contains sigma (σ bond and pi (π bonds. D. A π bond is than a σ bond. E. Also, π bonds the of a bond and leads to structural isomers (more about this in O chem). F. If many π bonds are present, molecules may conduct. G. Examples: Other Theories I. As with atomic theory, bond theories/models have gone through many. A. Valence Bond Theory vs. Molecular Orbital Theory B. Lewis Structures, Ball and Stick Model, Space-Filling Model Though they have all proved, all theories and models have. Page 7 of 7