Chemistry notes week of Feb 8-12, including all chp 9+10
Chemistry notes week of Feb 8-12, including all chp 9+10 1220
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This 16 page Class Notes was uploaded by Neha Bhagirath on Saturday February 13, 2016. The Class Notes belongs to 1220 at Wayne State University taught by Maryfrances Barber in Winter 2016. Since its upload, it has received 23 views. For similar materials see General Chemistry 1 in Chemistry at Wayne State University.
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Date Created: 02/13/16
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? February 8 MO orbital diagrams changes from Group 5A and beyond or 6A and beyond? CO molecular orbital drawing if it starts at 2s, only count valence from CO+, CO, or CO, which has the strongest bond and why? CO because it is the shortest because it has a double bond Why do CO+ and CO have the same bond order? Higher bond order, stronger the bond, shorter the bond (O2, O2+, O2 >O2+ is the strongest bond and highest bond order) if you have less electrons in the anti bonding orbitals (anti bonding MO), the bond is stronger. the bond is weaker if you have less electrons in the bonding orbitals (bonding MO.) to make a bond stronger:add an electron from a bonding orbital, remove an electron from the anti bonding orbital to make a bond weaker: remove an electron from a bonding orbital, add an electron from the anti bonding orbital e sigma means it is along or around the bond axis. Pi means it’s on opposite sides of the bond axis. when 2 s make a bonding orbital, it is symmetrical about the bond axis. When 2 p orbitals combine head to head, its along the bond axis and a sigma bond. If they combine side to side, now it’s a pi bond. how many atomic orbitals did it take to make an sp2 hybrid orbital? 3 (s+ 2p) how many orbitals in the 2p subshell? 3 >an orbital is how many subshells there are as soon as you make a pi bond, you cannot rotate (cis and trans) cis bonds are polar, trans bonds are non polar the shorter the double bond, the longer the single bond using bond energies is reactantsproducts enthalpies of formation is productsreactants delocalized bonding: type of bonding where a pair of e are spread over a number of atoms rather than being localized between 2 atoms formal charge: valence electrons bonding electrons nonbonding e Dipole moment: a qualitative measure of the degree of charge separation in a molecule When there is a charge separation, the molecule is polar. A polar molecule has a nonzero dipole moment Polar molecules when they are put in an electric field, they align themselves opposite the field a polar bond is characterized by separation of electrical charge why is CH3Cl polar if you have a lone pair it depends where it is and the dipole moment may be zero molecule Y has a dipole moment of zero. the geometry of Y must be trigonal planar, AX3 T shaped is always polar s orbital has 2 electrons, makes one bond p orbital has 6 electrons, makes 3 bonds d orbital has 10 electrons, makes 5 bonds like for CH4, it bonds 4 times. you can only have s once, and then p 3 times, and d 5 times. (keep counting and see where you end up. lone pairs count too. double and triple bonds are counted as one bond) Molecular geometry you ignore lone pairs?, so H2O would be bent, but for electron pair geometry you count lone pairs so for example H2O you have tetrahedral electron pair geometry with four regions around it It is best to have a formal charge of 0 for as many of the atoms in a structure as possible. if you gain an antibonding electrons or lose a bonding electron, the bond order goes up higher bond order means shorter length bond order is correlated to bond ENERGY when it says “include electron counts,” include the bonding AND nonbonding electrons AND total electrons Anything under Si is an exception to the octet rule (mostly the gases on the right of the metalloids) try to get the central atom with a charge of zero to find the best resonance structure (pick the central atom to be the LEAST electronegative in a lewis structure) When forming ionic bonds, the element that becomes an anion releases energy As an atom gains electrons, ionic radius increases because more electron repulsion occurs When drawing lewis structures, the best way is to first find the valence electrons, put the least EN atom in the middle, make all the bonds, find the octet for the outer atoms, and then on the central atom, find the formal charge. If it is zero, then you’re good. If it’s not, then put double or triple bonds on it, until it is a zero. square pyramidal is polar Feb 15 new chapter EXAM 3 1. Comparison of Gases,Liquids,and Solids Changes of State 2. Phase Transitions 3. Phase Diagrams Liquid State 4. Properties of Liquids: Surface Tension and Viscosity 5. Intermolecular Forces: Explaining Liquid Properties Solid State 6. Classification of Solids by Type of Attraction of Units 7. Crystalline Solids: Crystal Lattices and Unit Cells 8. Structures of Some Crystalline Solids 9. Calculations Involving UnitCell Dimensions 10.Determining Crystal Structure by XRay Diffraction Learning Objectives 1. Comparison of Gases, Liquids, and Solids a. Recall the definitions of gas,liquid,and solid given in Section 1.4. b. Compare a gas, a liquid, and a solid using a kinetic molecular theory description. c. Recall the ideal gas law and the van der Waals equation for gases (there are no similar simple equations for liquids and solids). 2.hanges of State Phase Transitions a. Define change of state (phase transition) b. Define melting,freezing,vaporization, sublimation, and condensation. c. Define v apor pressure. d. Describetheprocessofreachingadynamic equilibrium that involves the vaporization of a liquid and condensation of its vapor. e. Describe the process of boiling. f. Define freezing point a nd melting point. g. Define heat (enthalpy) of fusion and heat (enthalpy) of vaporization. h. Calculate the heat required for a phase change of a given mass of substance. i. Describe the general dependence of the vapor pressure (ln P) on the temperature ( T). j.State the Clausius–Clapeyron equation (the twopoint form). k.Calculate the vapor pressure at a given temperature. l.Calculate the heat of vaporization from vapor pressure. 3.Phase Diagrams a. Define p hase diagram. b. Describe the meltingpoint curve and the vaporpressure curves (for the liquid and the solid) in a phase diagram. c. Define triple point. d.Define critical temperature and critical pressure e. Relate the conditions for the liquefaction of a gas to its critical temperature. Liquid State 4. Properties of Liquids; Surface Tension and Viscosity a. Define surface tension. b. Describe the phenomenon of capillary rise c. Define viscosity. 5. Intermolecular Forces; Explaining Liquid Properties a. Define intermolecular forces. b. Define dipole–dipole force. c.Describe the alignment of polar molecules in a substance. d.Define London(dispersion) forces. e.Note that London forces tend to increase with molecular mass. f. Relate the properties of liquids to the intermolecular forces involved. g. Define hydrogen bonding. h. Identify the intermolecular forces in a substance. i. Determine relative vapor pressures on the basis of intermolecular attractions. Solid State 6. Classification of Solids by Type of Attraction of Units a. Define molecular solid, metallic solid,ionic solid, and covalent network solid. b. Identify types of solids. c. Relate the melting point of a solid to its structure. d. Determine relative melting points based on types of solids. e. Relate the hardness and electrical conductivity of a solid to its structure. 7. Crystalline Solids; Crystal Lattices and Unit Cells a. Define c rystalline solid and amorphous solid. b.Define crystal lattice and unit cell of a crystal lattice. c. Define simple cubic unit cell, bodycentered cubic unit cell, and facecentered cubic unit cell. d. Determine the number of atoms in a unit cell. e.Describe the two kinds of crystal defects. 8. Structures of Some Crystalline Solids a. Define hexagonal closepacked structure and cubic closepacked structure. b. Define coordination number. c. Note the common structures (facecentered cubic and bodycentered cubic) of metallic solids. d. Describe the three types of cubic structures of ionic solids. e. Describe the covalent network structure of diamond and graphite. 9. Calculations Involving UnitCell Dimensions a. Calculate atomic mass from unitcell dimension and density. b. Calculate unitcell dimension from unitcell type and density. 10. Determining Crystal Structure by XRay Diffraction a. Describe how constructive and destructive interference give rise to a diffraction pattern. b. Note that diffraction of xrays from a crystal gives information about the positions of atoms in the crystal. Comparing Gases, Liquids, and Solids Gases are compressible fluids. Liquids are relatively incompressible fluids. Solids are nearly incompressible and rigid. A change of state or phase transition is a change of a substance from one state (solid, liquid, gas) to another. The specific name or names for each of these transitions are given below. The vapor pressure of a liquid at a particular temperature is the partial pressure of the vapor over the liquid measured at equilibrium. When a liquid is placed in a closed vessel, the partial pressure of its vapor increases over time until it reaches equilibrium. At equilibrium, evaporation and condensation continue to occur, but do so at the same rate. This situation, which is called a dynamic equilibrium, is illustrated on the next slide. The vapor pressure depends on the liquid and on the temperature. This relationship is illustrated on the next slide for four substances. Note that as temperature increases, vapor pressure increases. The boiling point is the temperature at which the vapor pressure is equal to the pressure on the liquid, usually atmospheric pressure. At this temperature, bubbles of gas form within the liquid, as illustrated in the figure. The normal boiling point is measured at 1 atmosphere pressure When the pressure on the liquid increases, as is the case with a pressure cooker, the boiling point increases. Conversely, when the pressure on the liquid decreases, as is the case at high altitude, the boiling point decreases. The freezing point is the temperature at which a pure liquid changes to a crystalline solid (or freezes). The melting point is the temperature at which a crystalline solid changes to a liquid (or melts). Any change of state requires that energy be added to or removed from the system. As shown in the graph, for water, there are two regions that remain at the same temperature even as heat is added. Those represent regions of phase change.
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