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First set of notes after exam 1

by: Neha Bhagirath

First set of notes after exam 1 1220

Marketplace > Wayne State University > Chemistry > 1220 > First set of notes after exam 1
Neha Bhagirath

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these are the first week of notes after exam 1
General Chemistry 1
Maryfrances Barber
Class Notes
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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 Δ​H 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 O​2 is smaller than the distance between Na​+  and Cl​, the force between Mg​2+ and O​2 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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