Chemistry Chapters 9 and 10
Chemistry Chapters 9 and 10 CHM 160 001
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This 3 page Bundle was uploaded by Devin Mart on Tuesday March 29, 2016. The Bundle belongs to CHM 160 001 at Missouri State University taught by Dr. Richter in Spring 2016. Since its upload, it has received 13 views. For similar materials see General Chemistry 1 in Chemistry at Missouri State University.
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Date Created: 03/29/16
Chapter 9: Chemical Bonding I Lewis Theory ● The shapes of molecules are very important in biological systems. ● Chemical bonds are classified as Ionic, Covalent, or Metallic. ○ Ionic the transfer of electrons between metal and nonmetal. ○ Covalent the sharing of electrons between two nonmetals. ○ Metallic occurs between metals...the electron sea model. Low IE so the electrons can move easily between the different metal atoms. ● In 1902 G.N. Lewis realized that the chemistry of the main group elements could be explained by assuming that the atoms of these elements gain or lose electrons until they have 8 electrons in their outermost shell. ○ The octet rule every element must contain 8 electrons. ● In the Lewis model, valence electrons are represented as dots surrounding the atom. ○ Valence electrons are held most loosely in an atom, so these are the ones primarily responsible for bonding. ○ Atoms with a full outer shell (octet rule) are particularly stable. ● We represent ionic bonding by moving electron dots from metal to nonmetal prior to the formation of the crystalline lattice. ● Lattice energy is the enthalpy required to form an ionic compound from gaseous ions. ○ Lattice energies follow characteristic predictable trends. ● Atoms in covalent bonds share electrons until each atom has eight electrons in its valence shell. ○ A covalent bond is the sharing of at least two electrons in the valence shell of both atoms. ● The ‘bond order’ refers to the number of electrons that are shared between two atoms. ○ The principal reason for categorizing substances by the type of bonding is to more accurately predict the properties of a compound. ● The ability of an atom to attract electrons to itself in a chemical bond is called electronegativity. ● A polar covalent bond is intermediate between a pure covalent bond and an ionic bond. ○ Water is a polar covalent molecule. ● The degree of polarity, and the type of bond, can be determined looking at the differences in electronegativity (ΔEN). ○ If the ΔEN of the atoms in a substance are about the same, and both atoms come from the left hand side of the table, the substance is primarily metallic. ○ When the differences between electronegativities is relatively large, the compound is probably ionic. ○ If the ΔEN of the atoms in a substance are about the same, and both atoms come from the right hand side of the table, the substance is primarily covalent. ■ No compound is 100% ionic or 100% covalent. ○ 0.0 0.4 → Pure (nonpolar) covalent bond. ○ 0.4 2.0 → Polar covalent bond. ○ 2.0 3.3 → Ionic bond. ● When two or more ‘equivalent’ Lewis structures are possible, we call them resonance structures. ○ The actual structure is intermediate between the two resonance structures and is called a resonance hybrid. ○ When molecules with double and triple bonds may be drawn in several ways that are equivalent, we have a hybrid structure. ● When there is more than one possible Lewis structure, formal charge is used to determine the most likely (or useful) structure. ○ FCₐ = Vₐ Nₐ Bₐ/2 ■ Vₐ → valence electrons for atom a ■ Nₐ → nonbonding electrons for atom a ■ Bₐ → bonding electrons for atom a ○ 1. The best Lewis structure is the one with the lowest formal charges. ○ 2. Rarely, if ever, are there formal charges greater than ±1. ○ 3. The most EN atom should have the negative charge. ● Remember that formal charges do not represent actual charges, it is merely a useful concept to help choose the best Lewis structure. ● The bond energy is the energy required to sever a bond that holds two adjacent atoms together. ● The bond length is the average distance between two adjacent atoms. ● Metallic bonds involve metal atoms locked in a crystal structure surrounded by a sea of electrons. ○ In metals the electrons are not held tightly and there are no ‘filled shells’; this allows the electrons to be delocalized. Chapter 10: Chemical Bonding II Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory ● The shapes of molecules are very important in biological systems. ○ Change in the identity of one of the 146 amino acids on the chains in hemoglobin leads to sickle cell anemia. ● In Valence Shell Electron Pair Repulsion Theory VSEPR ) electron groups repel one another and this determines the shape and geometry of molecules. ○ There are 5common molecular shapes. 1. Linear 2. Trigonal planar 3. Trigonal bipyramidal 4. Tetrahedral 5. Octahedral ○ A bond angle is the angle between 3connected nuclei. ○ To assign the shape or geometry, identify the number of ‘electron groups’ around the central atom. ● You will responsible for being able to determine these five molecular geometries or shapes. 1. Linear 2. Bent 3. Trigonal planar 4. Trigonal pyramidal 5. Tetrahedral ● Entire molecules can be polar, depending on their shape and the nature of their bonds. ○ The presence of polar bonds may not result in a polar molecule. ○ Water has polar bonds and is a polar molecule. ○ The dipole moment indicates the magnitude of polarity. ● To determine whether a molecule is polar, use the following steps: 1. Draw a Lewis Structure for the molecule and determine the molecular geometry. 2. Determine whether the molecule contains polar bonds. 3. Determine whether the polar bonds add together to form a net dipole moment. ● Polar molecules are attracted to one another. ○ Just like opposite magnetic poles attract one another, opposite partial charges on molecules attract one another. ● Hydrophilic molecules are very soluble in water. ● Hydrophobic molecules are not soluble in water. ● The OH group interacts with water and makes ethanol hydrophilic. ● As the hydrocarbon chain grows, the solubility in water decreases. ○ The interactions between water and the OH cannot overcome the ‘negative’ interactions between the hydrocarbon chain and water. ● Soaps and detergents wash away the greases and oils that form on our bodies and clothes.
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