First set of notes after exam 1
First set of notes after exam 1 1220
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Popular in Chemistry
This 4 page Class Notes was uploaded by Neha Bhagirath on Monday February 8, 2016. The Class Notes belongs to 1220 at Wayne State University taught by Maryfrances Barber in Winter 2016. Since its upload, it has received 18 views. For similar materials see General Chemistry 1 in Chemistry at Wayne State University.
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Date Created: 02/08/16
Chapter 9 February 1 Molten salts and aqueous solutions of salts are electrically conducting, resulting from the motion of ions in the liquids. Suggests the possibility that ions exist in certain solids. Describing Ionic Bonds Define ionic bond. Explain the Lewis electrondot symbol of an atom. Describe the energetics of ionic bonding. Define lattice energy. Describe the Born–Haber cycle to obtain a lattice energy from thermodynamic data. Describe some general properties of ionic substances. Electron Configurations of Ions State the three categories of monatomic ions of maingroup elements. Write the electron configuration and Lewis symbol for a maingroup ion. Note the polyatomic ions given earlier in Table 2.5. Note the formation of +2 and +3 transitionmetal ions. Write the electron configurations of transitionmetal ions. Ionic Radii Define ionic radius. Define isoelectronic ions. Use periodic trends to obtain relative ionic radii. Describing Covalent Bonds Describe the formation of a covalent bond between two atoms. Define Lewis electrondot formula. Define bonding pair and lone (nonbonding) pair of electrons. Define coordinate covalent bond. State the octet rule. Define single, double and triple bond. Polar Covalent Bonds; Electronegativity Define olar covalent bond. Define electronegativity. State the general periodic trends in electronegativity. Use electronegativity to obtain relative bond polarity. Writing Lewis ElectronDot Formulas Write Lewis formulas having single bonds. Write Lewis formulas having multiple bonds. Write Lewis formulas for ionic species. Delocalized Bonding: Resonance Define localized bonding. Define esonance description. Write resonance forms. Exceptions to the Octet Rule Write Lewis formulas (exceptions to the octet rule). Note exceptions to the octet rule in Group IIA and Group IIIA. Formal Charge and Lewis Formulas Define formal charge. State the rules for obtaining the formal charge. State two rules useful in writing Lewis formulas. Use formal charges to determine the best Lewis formula. Bond Length and Bond Order Define bond length (bond distance). Define covalent radii. Define bond order. Explain how bond order and bond length are related. Bond Energy Define bond energy. Estimate ΔH from bond energies. There are 3 types of chemical bonds: ionic, covalent, metallic Ionic bond is formed between + and ions; electrons are transferred from the valence shell of one atom to another The combination of ionization energy and electron affinity is endothermic However, when the two ions bond, more than enough energy is released, making the overall process exothermic. making lewis dot never H or F as the central atom, put the most EN in the center p block nonmetals + metalloids in period 3 and above have extended valences? The lattice energy is the change in energy that occurs when an ionic solid is separated into gasphase ions. potential energy well when the atoms covalently bond (point of lowest energy) is most stable the bond dissociation energy is the asymptote, the energy needed to break the bond. Fluorine is most EN with 4.0, Oxygen is 3.5, Nitrogen and chlorine next but chlorine is larger? It is very difficult to measure lattice energy directly, but it can be found by using the energy changes for steps that give the same result. The process of finding the lattice energy indirectly from other thermochemical reactions is called the Born–Haber cycle. Ionic substances are typically highmelting solids. There are two factors that affect the strength of the ionic bond. They are given by Coulomb’s law: 1)The higher the ionic charge, the stronger the force; 2) the smaller the ion, the stronger the force. Using this we can compare the melting point of MgO and NaCl The charge on the ions of MgO is double the charge on the ions of NaCl. Because the charge is double, the force will be four times stronger. + 2+ 2 The size of Na is larger than that of Mg; the size of C is larger than that of . Because the distance between Mg 2+ and O2 is smaller than the distance between Na+ and Cl, the force between Mg2+ and O2 will be greater. Based on the higher charge and the smaller distance for MgO, its melting point MgO should be significantly higher than the melting point of NaCl. Group IIIA to VA metals often exhibit two different ionic charges: one that is equal to the group number and one that is 2 less than the group number. The higher charge is due to the loss of both thes subshell electrons and the p subshells electron(s). The lower charge is due to the loss of only thep subshell electron(s). Polyatomic ions are atoms held together by covalent bonds as a group and that, as a group, have gained or lost one or more electron. Transition metals form several ions. The atoms generally lose the ns electrons before losing the n – 1)d electrons. As a result, most transition metals form the +2 ion. Ionic radius is a measure of the size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found. A cation is always smaller than its neutral atom. An anion is always larger than its neutral atom. The term isoelectronic refers to different species having the same number and configuration of electrons. For example, Ne, Na+, and F are isoelectronic. Ionic radius for an isoelectronic series decreases with increasing atomic number. To consider how a covalent bond forms, we can monitor the energy of two isolated hydrogen atoms as they move closer together. The energy decreases—first gradually, and then more steeply—to a minimum. As the atoms continue to move closer, it increases dramatically. The distance between the atoms when energy is at a minimum is called the bond length. As the hydrogen atoms move closer together, the electron of each atom is attracted to both its own nucleus and the nucleus of the second atom. The electron probability distribution illustrates this relationship. A coordinate covalent bond is formed when both electrons of the bond are donated by one atom. The two electrons forming the bond with the hydrogen on the left were both donated by the nitrogen. Once shared, they are indistinguishable from the other N—H bonds. Hydrogen is an exception to the octet rule: it has two electrons in its valence shell (a duet). Double bonds form primarily with C, N, and O. Triple bonds form primarily with C and N. Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. Electronegativity is related to ionization energy and electron affinity. Hybridization the name of the hybrid orbital is based on the atomic orbitals that made it 2 regions two sp orbitals one s, one p 3 regions three sp2 orbitals one s, two p 4 regions four sp3 orbitals one s, three p 5 regions five sp3d orbitals one s, three p, one d ex.ICl3 because it has 5 regions around it 6 regions six sp3d2 orbitals one s, three p, two d bonding (lower energy) and anti bonding Sigma bond single Pi bond2 or more (when they repel they get 4 regions anti bonding) where do the pi and sigma bonds come from ?
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