GHW#3. Chapter 3. Louisiana Tech University, Chemistry 481. POGIL(Process Oriented Guided Inquiry Learning) Exercise on Chapter 3. Energetics of Ionic Bonding. Why? What are the properties of ionic compounds? How periodic table is used to predict ionic bonding? What is Coulombs law and how it applies to ionic bonding? What is lattice energy? How lattice energy is calculated from Coulombs model and the Madelung constant? How Madelung constant is affected by different ionic lattice types: cesium chloride, rock-salt, fluorite etc. How lattice energy and melting points of ionic compounds are affected by ionic radii? What are the periodic trends in ionic radii? What is a charge density of an ion? How are charge density values used to predict more ionic bonding or more polarizable/less ionic bonding? What factors affect the polarizability on an ion? How does polarizability of an ion affect the lattice energy and melting points? How do you calculate lattice energy for an ionic compound from thermodynamic data using Born-Haber cycle? What is enthalpy of solution, H solution, enthalpy of hydration, H hydration and solvent-solvent intermolecular attractions, H solvent-solvent? How are the H solution, H hydration, LE and H solvent-solvent related for the solution process? What factors affect the solubility of an ionic compound in a given solvent? How does solvent, LE, H hydration, and H solvent-solvent affect the solubility of an ionic compound? How does the size and charge of ions of an ionic compound affect its LE and H hydration? What controls the relative change of both LE and H hydration when ion size is decreased or incresed? Why ionic compounds with smaller anions- NO - 3, ClO - 4, ClO - 3, and C 2 H 3 O - 2, and cations- H +, Na +, K +, and NH + 4 form soluble compounds? Instructional Objectives: 1) Explain properties of ionic compounds: as electrical attraction, electricity of molten liquid, electrical conduction when dissolved in water, brittleness when hammered, the high melting and boiling points, solubility in polar solvents. 2) Use periodic table and electronegative to predict ionic bonding. 3) Apply Coulombs law to ionic bonding and qualitatively predict lattice energy. 4) Calculate lattice energy given charge of ions, ionic radii and Madelung constant. 5) Explain Madelung constant values for different ionic lattice types: cesium chloride, rock-salt, fluorite etc. 6) Predict the trends in lattice energy and melting points of ionic compounds based on ionic radii. 7) Predcit the trends in ionic radii using periodic table. 8) Calcualte is a charge density of an ion and compare them to predict more ionic bonding or more polarizable less ionic bonding. 9) Explain the factors that affect the polarizability and its effect the lattice energy and melting points. 10) Calculate lattice energy from thermodynamic data (IE, EA, BDA, SE, H f etc.) and using Born-Haber cycle. 11) Expalin factors affect the solubility of an ionic compound in a solvent, the enthalpy of solution, H solution, the enthalpy of hydration, H hydration and the solvent-solvent intermolecular attractions, H solvent-solvent and their effect on solubility.
12) Identify conditions that would make H solution, more negatice and ionic compound more soluble. 13) Expain the effect of size and charge of ions on the LE and H hydration. 14) Expain the solubility of ionic compounds with smaller anions- NO 3 -, ClO 4 -, ClO 3 -, and C 2 H 3 O 2 -, and cations- H +, Na +, K +, and NH 4 +. RESOURCES INORGANIC CHEMISTRY By Peter Atkins, Tina Overton, Jon Rourke, Mark Weller, Fraser Armstrong, 4th Edition 2006. New Concepts Properties of ionic compound: 1) An ionic bond is a strong electrical attraction between two oppositely charged atoms or groups of atoms. 2) Conducts electricity as molten liquid. 3) When dissolved in water produce solutions that conduct an electric current: they are electrolytes. 4) They break readily when hammered, because of ion repulsions created by the slippage. 5) They have high melting and boiling points. 6) They are soluble in polar solvents. Periodic trends in bonding and electronegativity a) The main Group elements (s and p orbitals) lose or gain electrons to attain a configuration like a noble gas. b) Transition elements (d orbitals) lose their s orbital electrons first and then one or more d orbital electron(s). c) The ionic character of the bond is proportional to the electronegativity difference of the elements making the bond. Electronegativity Difference 0-0.5 0.6-1.6 1.6, or greater Nonpolar covalent Polar covalent bond Ionic C-H O-H Na-Cl Coulombs law and the ionic bond: The electrostatic model is simply an application of the charge principles that opposite charges attract and similar charges repel. k = constant (permeability of the medium) q+ = cation charge; q- = anion charge r = distance between two ions An ionic compound results from the interaction of a positive and negative ion, such as sodium and chloride in common salt. The ions will cluster together so as to maximize heir attractions and minimize repulsions. The arrangements assumed by these ions will be determined by the compound formula and by the sizes of the ions. Lattice Energy: In crystalline compounds this net balance of attractive and repulsive forces is called the lattice energy which is the energy released in the formation of one mole of ionic solid from the gaseous ions. The lattice energy and melting point are directly related. E.g. Mg 2+ (g) + 2Br - (g) MgBr 2(s) H LE = -2440kJ mol -1
Based on Coulombs law, as the ionic radii of either the cation or anion increase, the lattice energies decrease, and the solids consists of di-positive ions have much larger lattice energies than solids with mono-positive ions. Lattice type also matters because packing is different. Coulomb law equation, is multiplied a factor called Madelung constant to get lattice energy. Madelung Constant: Factor that considers all electrostatic attractions in an ionic lattice Trends in ionic radii Ionic radii increase down a group Solid Lattice Madelung Ionic radii decrease across a period Type Constant Cations are smaller than their parent atoms - lose electrons NaCl Rock salt 1.747558 Anions are larger than their parent CsCl CsCl type 1.747558 atoms - gain electrons CaF 2 Fluorite 2.51939 TiO 2 Rutile 2.408 Charge density of ions: Defined as Charge/volume expressed as coulombs/å 3. Comparing charge density values of series of ions (Appendix 2) one can predict is most likely to form compounds exhibiting more ionic bonding and ones that will be more polarizable and show some degree of covalency. As the covalency increases in an ionic bond the experimental lattice energy will be higher than the predicted by ionic model alone. Group 1 chlorides and fluorides: The lithium compound exhibits a lower melting point than we would anticipate. Remember that lithium is very small and has a slightly higher electronegativity than the other Group 1 metals, thus it appears that lithium's low melting points are due to a bit of covalency and thus a slightly reduced polarity of the bonds. Polarization will be increased by: 1. High charge and small size of the cation. Ionic potential Å Z + /r + (= polarizing power) 2. High charge and large size of the anion. The polarizability of an anion is related to the deformability of its electron cloud (i.e. its "softness") 3. An incomplete valence shell electron configuration. Noble gas configuration of the cation better shielding. e.g. Hg 2+ (r + = 102 pm) is more polarising than Ca 2+ (r + = 100 pm) Calculation of lattice energy from thermodynamic data: A Born-Harber cycle could be drawn for steps in the formation of any ionic compounds. E.g. Draw a Born-Harber cycle for the formation of BaBr 2 from barium metal and bromine gas. Label each step with the appropriate thermodynamic quantity. Calculate the enthalpy of formation for BaBr 2. BaBr 2 lattice energy = 1950 kjmol -1 Ba atomization energy= 175 Ba 1st ionization energy = 503 Ba 2nd ionization energy = 965
Br2 bond enthalpy = 193 Br electron affinity= -325 LE = EA + IE 2 +IE 1 BDE + SE - H f LE = [2 x (-325)]+ 965+503+ 193+175 (-764 )= 1950 LE = 1950 kjmol -1 Enthalpy of solution, H solution : The formation of a solution involves the interaction of solute with solvent molecules. Many different liquids can be used as solvents for liquid solutions, and water is the most commonly used solvent. When water is used as the solvent, the dissolving process is called hydration. The heat change which takes place when one mole of a solute is completely dissolved in a solvent to form a solution of concentration 1 mol L -1. Enthalpy of hydration, H hydration : The heat evolved when 1 mole of gaseous ions become hydrated (surrounded by water molecules), measured under standard conditions. H hydration (cation): Al 3+ (g) + aq Al 3+ (aq) H hyd(cation) = -4613 kj mol -1 H hydration (anion): Cl 1- (g) + aq Cl 1- (aq) H hyd(anion) = -363 kj mol -1 H hydration (AlCl 3 ): Al 3+ (g) + 3 Cl 1- (g) + aq Al 3+ (aq) + 3 Cl 1- (aq) H hyd = -4613 kj mol -1 +3 x (-363)] kj mol -1 H hyd (AlCl 3 ) = -5702 kj mol -1 Solvent-solvent intermolecular attractions, H solvent-solvent : The energy required to break dipole - dipole interactions between solvent molecules (L) when they become solvating ligands (L') for the ions. The enthalpy of solution, H solution, enthalpy of hydration, H hydration, lattice energy and solvent-solvent intermolecular attractions, H solvent-solvent is related in the solution process by the equation: H sol H LE + ( H hyd (anion) + H hyd(cation) ) + H solvent-solvent Factors affecting solubility: Solvent: The "like dissolves like" rule is used to predict the solvent needed for solution process. In other words, if you want to dissolve ionic compound you should use a solvent that is also highly polar with a high dielectric constant and non-polar compound requires a nonpolar solvent with a low dielectric constant. A small solvent molecular have a smaller dipole which can approach the ions closely to increase solubility. Lattice Energy (LE): It takes energy to separate ions from their crystal lattice and from hydrated ions. Smaller size of the ions increases both the lattice energy and hydration enthalpy. If the lattice enthalpy increases more than the hydration enthalpy, then heat of solution become more endothermic and vise versa. Enthalpy of hydration, H hydration : The attraction of a dipole solvent (water) dipole to an oppositely charged ion often released as salvation (hydration) energy. This hydration energy is used to break the ionic lattice: the lattice energy (LE). Non-polar and weakly polar solvents do not have sufficiently strong hydration to overcome LE and to dissolve electrolytes. In a crystal hydrate, the ions are largely hydrated, and consequently the
hydration energy is considerably less than that of the anhydrous solute. The hydrates, therefore, usually have lower water solubility. The enthalpy of solution, H solution : Usually substances with a large negative heat of solution (i.e., exothermic reaction) are more soluble than substances with a smaller negative heat of solution. Compounds that have a positive heat of solution (endothermic) may also be soluble. Size and charge of ions on LE and H hydration :What controls the relative rate of fall both LE and H hydration? Charge factor Larger the charges on ions increases LE more than the H hydration, therefore ionic compounds with di-positve or negative charge tend to make the enthalpy of solution less negative (less soluble). Size factor It turns out that the main factor increasing H hydration over LE is the size of the negative and positive ions. E.g. Small anions- NO 3 -, ClO 4 -, ClO 3 -, and C 2 H 3 O 2 - form soluble compounds. The lattice energy increase less than the hydration enthalpy of the positive ions. That means that the enthalpy of solution will become more negative. E.g. Small cations- H +, Na +, K +, and NH 4 + forms soluble compounds. The lattice energy increase less than the increase in hydration enthalpy of these small positive ions make enthalpy of solution will become more negative. E.g. Large anions- PO 4 3-, S 2-, CO 3 2-, and SO 3 2- ions are insoluble except those that also contain alkali metals or Na +, K +, and NH 4 +. Changes in the size of the positive ion don't make as great a percentage difference to the inter-ionic distance as they would if the negative ion was small. The hydration enthalpy of the positive ions decreases more than the lattice energy. The enthalpy of solution will become less negative. Success Criteria Ability to answer the questions and apply concepts related to the topics given in the instructional objectives. Resources DESCRIPTIVE INORGANIC CHEMISTRY by Geoff Rayner-Canhanmi Prerequisites Freshman chemistry, chapters 1-4 of DESCRIPTIVE INORGANIC CHEMISTRY Rayner-Canhanmi and detailed knowledge of covalent (especially MO theory) and ionic bonding models.
GHW# 3 Chapter 3. Ionic Bonding Your Name: Key Questions (relatively simple to answer using the Focus Information) 1. What properties of a compound would lead you to expect that it contains ionic bonds? 2. Would you expect sodium chloride to dissolve in carbon tetrachloride, CCl 4? Explain your reason. 3. Which would you expect to contain ionic bonds, MgCl 2 or SC1 2? Explain your reasoning. 4. What is Coulombs law how it applies to ionic bond? 5. What is lattice energy? Take NaCl as an example. 6. Which one of each of the following pairs will be smaller radius? Explain your reasoning in each case. a) K or K + : b) K + or Ca 2+ : c) Br - or Rb + : d) Se 2- or Br - : e) O 2- or S 2- : Compound Interionic Distance (Angstroms) Melting Point (Centigrade) Lattice Energy (kcal/mol) NaF 2.31 988-201 NaCl 2.79 801-182 NaBr 2.94 790-173
7. Explain the lattice energy and melting point trends: NaI 3.18 660-159 8. Explain the lattice energy and melting point trends: Compound Cation radius (Angstroms) Anion radius (Angstroms) Melting Point (Centigrade) Lattice Energy (kcal/mol) MgCl 2 0.65 1.81 714 2326 CaCl 2 0.94 1.81 782 2223 MgO 0.65 1.45 2852 3938 CaO 0.94 1.45 2614 3414 9. Compare the charge density values of the three silver ions: Ag +, Ag +2, and Ag +3 (Appendix 2). Which is most likely to form compounds exhibiting ionic bonding? 10. Compare the charge densities of the fluoride ion and the iodide ion (Appendix 2). On this basis, which would be the more polarizable? 11. How does polarization and covalency affect lattice energy and melting points? Compound Melting Point ( o C) AgF 435 AgCl 455 AgBr 430 AgI 553
12. Calculate the Lattice energy of NaCl from following thermodynamic data: Steps Ho, kj 1. Vaporization of sodium: Na(s) Na(g) +92 2. Decomposition of Cl 2 : 1/2 Cl 2 (g) Cl(g) +121 3. Ionization of sodium: Na(g) Na + (g) +496 4. Electron affinity to chlorine: Cl(g) + e - Cl - (g) -349 5. Formation of NaCl(s): Na(g)+1/2Cl 2 (g) NaCl(s) -411 13. Define following terms: a) Enthalpy of solution, H solution : b) Enthalpy of hydration, H hydration : c) Solvent-solvent intermolecular attractions, H solvent-solvent : 14. How is Enthalpy of solution, H solution, Enthalpy of hydration, and Lattice energy are related? 15. Predict the solubility of following ionic compounds: Lattice Energy(U) H hyd, M + H hyd, M - LiF 1030-950 -60 LiI 720-950 -80 CsI 585-700 -80 MgF 2 3100-2800 -120 a) LiF: b) LiI: c) CsI: d) MgF 2 :
16. Give rational explanation to the solubility rules in terms of ioin sizes, lattice energy(u), H hyd, and H solution. a) All compounds containing alkali metal cations and the ammonium ion are soluble. b) All compounds containing NO 3 -, ClO 4 -, ClO 3 -, and C 2 H 3 O 2 - anions are soluble. c) All chlorides, bromides, and iodides are soluble except those containing Ag +, Pb 2+, or Hg 2 2+. d) All sulfates are soluble except those containing Hg 2 2+, Pb 2+, Sr 2+, Ca 2+, or Ba 2+. e) All hydroxides are insoluble except compounds of the alkali metals, Ca 2+, Sr 2+, and Ba 2+. f) All compounds containing PO 4 3-, S 2-, CO 3 2-, and SO 3 2- ions are insoluble except those that also contain alkali metals or NH 4 +.