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Chapter 6: Chemical Bonding

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Chapter 6: Chemical Bonding Chapter 6: Chemical Bonding I. Introduction to Chemical Bonding A. A Chemical Bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. B. Types of chemical bonding 1. The two main types of bonding are ionic (involves ions) and covalent (involves sharing of electrons). 2. There are two types of covalent bonds: Polar (unequal sharing) and nonpolar (equal sharing). 3. There are other types of bonds which we will discuss later. 4. To determine the type of bond we use differences in EN. a. If 0 < ∆EN < 0.3, bond is Nonpolar Covalent. b. If 0.4 < ∆EN < 1.7, bond is Polar Covalent. c. If ∆EN > 1.8, bond is Ionic. d. Refer to Figure 2 on page 176. II. Covalent Bonding A. Molecules 1. Many compounds, including most chemicals involving living things, are composed of molecules. 2. A Molecule is a neutral group of atoms that are held together by covalent bonds. 3. A Molecular Compound is a chemical compound whose simplest units are molecules. 4. A Chemical Formula indicates the relative number of atoms of each kind in a chemical compound. 5. A Molecular Formula show the types and numbers of atoms combined in a single molecule. 6. A Diatomic Molecule is one that consists only of 2 atoms of the same element. There are 7 diatomics. a. H2, N 2, O 2, F 2, Cl 2, Br 2, and I 2. b. BrINClHOF, 7’s on table. B. Formation of a Covalent Bond 1. Most atoms have a lower potential energy when bonded. 2. Refer to Figure 5 on page 179. C. Characteristics of the Covalent Bond 1. The bonded atoms vibrate a bit, but as long as their energy is close to the minimum, there is a covalent bond. 2. Forming bonds usually releases some energy. In order to break the bonds, the same amount of energy must be added. This is the bond energy. D. Lewis Dot Structures 1. LDS: an electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots around the symbol of that element. 3 Steps: a. Write symbol of element. b. Determine number of valence electrons for element. c. Place dots around symbol, starting at top, working clockwise. 1 in each area before pairing up. There are four areas(12,3,6,9). 2. When doing structures for molecules, we must follow the Octet Rule: atoms will gain, lose, or share electrons in order to acquire a “full set” of 8 valence electrons. 3. In covalent bonding the electrons are shared . Look at H 2. H xH (the dot and the x represent electrons.) They come together and share the electrons. H xH ; which can also be written as H-H. This is a single covalent bond. 4. Practice with O 2 and N 2. (Double and Triple). 5. There are exceptions to the octet rule: B & Al(6), H (2), Be (4), and expanded octets. 6. In summary, in covalent bonds the electrons are shared to fulfill the octet rule for all atoms present. E. Drawing LDS 1. When given a formula, one should be able to construct a LDS for the molecule in question. 2. Lone (Unshared) Pairs are pairs of electrons that are not shared. Shared Pairs are pairs of electrons that are shared. 3. Steps for drawing LDS for molecules. a. Determine total number of valence electrons to be used. b. Determine and write the symbol for the Central Atom. (Usually C, N, P, S, O) c. Arrange remaining atoms around the central atom. d. Connect all outside atoms to the central atom with single bonds. Be sure to account for all electrons used from here on. e. Make sure all atoms are satisfied. If they are you are done. Repeat step e after every step from now on. f. Place lone pairs of electrons around exterior atoms first then central atom, until all are used. If all are satisfied, done. g. Make multiple bonds, if necessary to satisfy octet rule. h. Draw all resonance structures. Resonance: bonding in molecules or ions that cannot be correctly represented by a single LDS. 4. Much Practice!!!!! F. Polyatomic Ions 1. Polyatomic Ions: a charged group of covalently bonded atoms. 2. When doing LDS for PI’s, the charge must be taken into account. If it is + an electron has been lost, if it is - an electron has been gained. Once this is taken into account, the LDS is done the same way. When finished, it must be put into brackets and the charge indicated. + 2- - 3. Practice with NH 4 , SO 4 , NO 3 . III. Ionic Bonding A. When Ionic bonding occurs, ionic compounds are formed. Ionic Compounds consist of positive and negative ions that are combined so that the total overall charge on the molecule is zero. The simplest collection of atoms from which an ionic compound’s formula can be established is called a Formula Unit. B. Formation of Ionic Compounds 1. In ionic bonding, the electrons are transferred from one atom to another. It “jumps” from the less EN atom to the more EN atom. a. For example NaCl: .. .. Na . + .Cl: → Na+ + :Cl: -. .. .. b. CaF 2 2. When Ionic molecules combine to form a crystal they interact to give the lowest energy set up. This is called the crystal lattice. For example: NaCl: Na + Cl - Na + Cl - Na + Cl - Cl - Na + Cl - Na + Cl - Na + See Figure 14 on page 191. 3. Lattice Energy: energy released when one mole of an ionic crystalline compound is formed from gaseous ions. C. Comparison of Ionic and Covalent Compounds 1. The force holding ions together is very strong attractive forces. There is very little force of attraction between covalent molecules. 2. Ionic have higher BP, MP, and hardness. This is due to higher attraction between particles. 3. Ionic does not vaporize at room temp. 4. Ionic-Brittle; Covalent-not brittle. 5. Solid ionic-does not conduct well, liquid ionic does. IV. Metallic Bonding A. Metallic Bonding is the chemical bonding that results from the attraction between metals and the surrounding sea for electrons. B. Only occurs in metals. C. In this situation, there is a group of metal atoms that are close together. The electrons become delocalized, meaning that they are not associated with any one atom. We have little islands of positive charge floating in a sea of electrons. D. This gives rise to many of the properties of metals. 1. Malleable and ductile: because layers can slide past one another without much resistance. 2. Shiny: electrons are easily excited, and light given off makes them look shiny. 3. Conduction: electrons are not tied to any atom, free to move. Current is a movement of electrons. V. Molecular Geometry A. We use the Valence Shell Electron Pair Repulsion (VSEPR) theory to predict the shapes of molecules. B. VSEPR Theory 1. States that repulsion between the sets of valence electrons surrounding an atom causes these sets to be oriented as far apart as possible. 2. Any bond itself is linear. C. Refer to Table 5 on page and compare the molecule to it. We need to know the number of shared pairs of electrons and the number of unshared pairs of electrons. (*Note: each multiple bond counts as one shared pair.) D. Need to know: Linear, Bent, trigonal planar, tetrahedral, and trigonal pyramidal. E. Hybridization 1. To completely explain the shapes that VSEPR predicts, we must discuss hybridization. 2. Hybridization: the mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies. 3. For example, the 2s and 2p orbitals of C are similar in energy. These four orbitals hybridize and form 4 sp 3 hybrid orbitals, which are all identical, both in shape and energy. 4. The superscripts tell how many of each type of orbital is used. 5. Hybrid Orbitals are orbitals of equal energy produced by the combination of two or more orbitals on the same atom. 6. sp hybrids give linear, sp 2 hybrids give trigonal planar, sp 3 give tetrahedral. F. Intermolecular Forces 1. IM Forces are forces of attraction between molecules. 2. Molecular Polarity and Dipole-Dipole Forces a. A Dipole is created by equal but opposite charges that are separated by a short distance. b. They occur in polar bonds ( ∆EN is 0.3-1.7). In polar bonds the more EN atom pulls the electrons closer to it. The electrons spend more time around the more EN atom. This gives rise to + slight positive charges ( δ ) around the less EN atom and slight negative charges ( δ-) around the more EN atom. The dipole is indicated by an arrow +→. The + end is at the less EN atom. c. A polar molecule can induce a dipole in other molecules. 3. Hydrogen Bonding a. The molecules must have hydrogen in them. b. Hydrogen Bonding is the intermolecular force in which a hydrogen atom that is bonded to a highly EN atom is attracted to an unshared pair of electrons of an EN in a nearby molecule. c. H-bonds are represented by dotted lines connecting the H to the unshared pair of electrons. 4. London Dispersion Forces a. London Dispersion Forces are intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles. b. They act between all atoms and molecules, but are the only intermolecular forces between atoms of noble gases and nonpolar molecules. c. Compounds of this type are characterized by low boiling points.
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