Chem 1A Dr. White Updated /5/1 1 Chemistry Workbook 2: Problems For Exam 2 Section 2-1: Covalent Bonding 1. On a potential energy diagram, the most stable state has the highest/lowest potential energy. (circle one) 2. For the formation of a covalent bond, when is the potential energy the lowest? (i.e. what is minimized and what is maximized in the bond?). Define electronegativity in your own words. 4. What is the difference between a nonpolar covalent bond and a polar covalent bond? 5. Select the bond that has the largest ΔEN amongst the following. A. C-O B. Si-F C. Cl-F D. C-F E. C-I 6. Which of the following compounds displays the greatest ionic character in its bonds? A. NO 2 B. CO 2 C. H 2 O D. HF E. NH 7. Arrange in order of increasing electronegativity (use only the periodic table, not the EN values!): S, Si, Ge, Ga Section 2-2: Lewis Structures and VSEPR 1. Draw Lewis structures for the following: (a) F 2 (b) NH (c) H 2 O (d) CH 4 (e) CO 2 (f) N 2 2. Draw Lewis structures for the following showing different possible resonance forms. Calculate the formal charges of each atom, and state which resonance form (if any) would dominate. (NOTE: N is more EN than S) 2 a. CO b. SCN c. N 2 O. Draw possible structures for Cl 2 CO. Based on all the rules we discussed in class, choose the best Lewis structure (circle your choice), and briefly explain why this is the best structure. Also tell me if this molecule would be polar or not. 4. Draw the best Lewis Structures for these covalent species. None of these structures are cyclical. Make sure to take into account formal charge and draw resonance structures when possible. When there are resonance structures that aren t all the same, pick the best one. a) COS Total # Valence Electrons: b) ClO - Total # Valence Electrons: (no dbl bonds to Cl) c) C 2 F 4 (precursor to Teflon) Total # Valence Electrons: d) CH OH (methanol drink it and go blind!) Total # Valence Electrons: + e) NO 2 Total # Valence Electrons: f) H PO 4 (minimize FC on P) Total # Valence Electrons:
Chem 1A Dr. White Updated /5/1 2 5. Fill in the table below for each of the species shown. You must draw the Lewis Structure first. The central atom does not necessarily have to follow the octet rule. Make sure to include formal charge Also remember that double and triple bonds count the same as single bonds when determining geometries. Total # of valence electrons: Lewis Structure, complete with formal charges and resonance if applicable. Total # of electron groups around the central atom (bonded and non-bonded): Electron Pair Arrangement: VSEPR Sketch that shows -D shape of molecule # of bonded groups around the central atom (connected to actual atoms): a) CF 4 b) OCN - 2- c)so e) XeOF 4 (draw only the best structure) Molecular Geometry: Polar or not? + (draw only the best structures) d) ClF 2 Section 2-: Hybrid Orbitals 1. What is the hybridization around the central atom in each of the following? (Draw a Lewis Structure first!) a) SeOCl 2 b) H 2 O c) PCl 5 d) NCl e) SO f) SF 4 - g) IF 2 2- h) CrO 4 i) H O + j) SF 6 k) CO 2 l) XeF 2 2. Describe the differences between σ-bonds and π-bonds.. Consider the ion ClF 4 + a. Draw the Lewis structure b. Draw the hybridized orbital picture (HOP) for the molecule. c. What is the type of orbital used by Cl in the Cl-F bond? d. What is this ion s electronic arrangement? e. What is this ion s molecular shape? f. Is this ion polar or nonpolar? 4. Consider the molecule: (NOTE: Lone pairs are not shown) H C O H O a. Draw the Hybridized Orbital Picture b. How many σ bonds are there in this molecule? c. How many π bonds are there in this molecule? d. List all the types of orbitals used by carbon in this molecule. e. List all the types of orbitals used by the O which is bonded to an H. 5. Consider the Lewis structure for acetone: Draw a hybridized orbital diagram clearly showing the bonding in the molecule. Label the orbitals and all bonds.
Chem 1A Dr. White Updated /5/1 Section 2-4: Molecular Orbitals 1. Determine the bond order and write the molecular orbital electron configurations for the following homonuclear diatomic molecules. a) B 2 b) F 2 c.) O 2 + d) Ne 2 + 2. Consider N 2 +, N 2, and N 2 -. a) Write the electron configurations for each of these molecules. b) What is the bond order of each molecule? c) Which of the three should be most stable? d) Which are paramagnetic and which are diamagnetic?. Cyanide ion has the same energy order of the molecular orbitals as does nitrogen. a) Write the electron configuration for the cyanide ion. b) draw the Lewis structure for cyanide c) What is the bond order in cyanide and does it agree with your Lewis structure? d) Do you expect the cyanide ion to be paramagnetic or diamagnetic? Practice For Exam 2 1. Fill in the boxes below with the requested information for describing the following molecule in terms of hybridization and s and p components. (Note lone pairs are NOT shown). List the orbital used by H in this bond: This bond consists of what types of bonds? List all types of orbitals used by C in this bond: List all types of orbitals used by O in this bond: a. How many sigma bonds exist in the molecule? b. How many pi bonds exist in the molecule? c. Is this molecule polar or nonpolar? Explain. 2. Complete the following table Molecule Refined Lewis Structure (the BEST structure only). Label bond angles. Geometry of Central Atom Hybridization of Central Atom Polar or Nonpolar H 2 CO *consider C the central atom *consider C the central atom XeF 4 SiF 6 2- BeF 2 IBr
Chem 1A Dr. White Updated /5/1 4. (a) Three possible resonance forms for the molecule COClF are shown below. Label each atom in each structure with its formal charge. (b) Use the formal charges from part (a) to select which of the three resonance forms (A, B or C) best represents the molecule. Explain your choice. 4. For the chlorate ion, ClO -, draw two different valid Lewis structures, as follows: a. a structure in which the octet rule is obeyed b. a structure in which formal charges are minimized 5. Explain what is meant by "dipole moment", and give an example of a molecule which has polar bonds but which does not itself have a dipole moment. 6. Sulfur will use hybrid orbitals in sulfur dioxide, SO 2. 7. Iodine will use hybrid orbitals in ICl 2 -. 8. Bromine will use hybrid orbitals in BrF 5. 9. What is the electron configuration and the bond order in the O 2 + ion? Is the ion paramagnetic or diamagnetic? 10. One can safely assume that the s- and p-orbitals will form molecular orbitals similar to those formed when 2s- and 2p-orbitals interact. According to molecular orbital theory, what will be the electron configuration and bond order for the Cl 2 + ion? NOTE: Cl 2 follows the same MO energy levels as F 2. 11. a. What simple experiment could you perform to show that a substance is paramagnetic? b. What microscopic (atomic/molecular) feature must a substance possess in order to be paramagnetic? c. Can it be predicted whether or not all homonuclear diatomic ions, X 2 +, will be paramagnetic? Explain. 12. 1. 14. Name and outline the concept which is introduced when more than one valid Lewis structure can be drawn for a given molecule or ion. Use appropriate diagrams of the formate ion (HCO 2 -, carbon is the central atom) to illustrate. Are the structures for HCO 2 - equivalent or is one better than another? Explain. Draw the MO diagram for F 2. Predict the bond order and determine if F 2 is paramagnetic or diamagnetic. Is it a stable molecule? Write the electron configuration for Ne 2. Predict the bond order. Is it a stable molecule? 15. Draw a hybrid orbital picture for C 2 H 4. Describe the bonding in this molecule. 16. Describe the differences between bonding and antibonding molecular orbitals.
Chem 1A Dr. White Updated /5/1 5 Answers: Section 2-1: Bonding 1. lowest 2. The repulsions between electrons and the repulsions between nuclei are minimized and the attractions between the electrons and nuclei are maximized. Electronegativity is the ability of an atom to attract electrons towards itself in a chemical bond. 4. A nonpolar covalent bond has equal sharing of electrons. A polar covalent bond has unequal sharing of electrons. 5. B 6. D 7. Ga<Ge<Si<S Section 2-2: Lewis Structures and VSEPR 1. 2.
Chem 1A. 4. Dr. White Updated /5/1 6
Chem 1A 5. Dr. White Updated /5/1 7
Chem 1A Dr. White Updated /5/1 8 Section 2- Hybrid Orbitals 1. a. sp b. sp f. sp d g. sp d c. sp d h. sp d. sp i. sp 2 e. sp 2 j. sp d k. sp l. sp d 2. σ bonds result from direct orbital overlap. The electron density is between the bonded atoms. π bonds result from orbital overlap above and below the bond axis. Thus, the electron density is above and below the bond axis.. a. b. (only one full F is shown)
Chem 1A Dr. White Updated /5/1 9 c. sp d d. triganol bipyramidal e. see saw f. polar 4. a. b. 4 c. 1 d. sp 2 and 2p e. sp 5. Section 2-4: Molecular Orbitals 1.
Chem 1A Dr. White Updated /5/1 10 2. a. N 2 + : σ 1s 2 σ* 1s 2 σ 2s 2 σ* 2s 2 π 2p 4 σ 2p 1 N 2 : σ 1s 2 σ* 1s 2 σ 2s 2 σ* 2s 2 π 2p 4 σ 2p 2 N 2 - : σ 1s 2 σ* 1s 2 σ 2s 2 σ* 2s 2 π 2p 4 σ 2p 2 π* 2p 1 b. N 2 + : 2.5 N 2 : N 2 - : 2.5 c. N 2 is most stable d. N 2 + and N 2 - are paramagnetic; N 2 is diamagnetic. a. σ 1s 2 σ* 1s 2 σ 2s 2 σ* 2s 2 π 2p 4 σ 2p 2 b. c. BO=; It agrees with the Lewis Structure which has a triple bond. Practice For Exam 2 1. orbitals used by H: 1s orbitals used by C sp 2 and 2p orbitals used by O sp The double bond is a sigma and pi bond a. 4 b. 1 c. Polar. The C atom has bonding groups that are different. Thus,the dipoles do not cancel and the molecule is polar. 2. H 2 CO Triganol Planar 120 sp 2 Polar XeF 4 Square Planar 90 sp d 2 Nonpolar
Chem 1A Dr. White Updated /5/1 11 SiF 6 2- Octahedral 90, 120 sp d 2 Nonpolar BeF 2 linear 180 sp Nonpolar IBr T-shaped <90 sp d Polar. A every atom has a FC of 0. B C and Cl have FCs of 0 and O has a FC of -1 and F has a FC of +1. C - C and F have FCs of 0 and O has a FC of -1 and Cl has a FC of +1. A has the best Lewis Structure since it has a formal charge of 0 on all the atoms 4. 5. A dipole moment arises in a molecule when the dipoles caused by polar bonds do not cancel. There is thus a separation of charge. The dipole moment is the product of this charge and the distance of separation. Carbon dioxide has two polar carbon-oxygen bonds. However, because the molecule is linear, the two bond dipoles are exactly opposite in direction, and they cancel each other out. The CO 2 molecule has no dipole moment. 6. sp 2 7. sp d 8. sp d 2 9. σ 2 1s σ* 2 1s σ 2 2s σ* 2 2s σ 2 2p π 4 2p π 2p * 1 Paramagnetic BO = 2.5 10. electrons: σ 1s 2 σ* 1s 2 σ 2s 2 σ* 2s 2 σ 2p 2 π 2p 4 π 2p * 4 σ 2p * 2 σ s 2 σ* s 2 σ p 2 π p 4 π p * Paramagnetic BO = 1.5 11. a. Place a sample of the substance near a magnet; if it is attracted to the magnet, it is paramagnetic. b. The substance must have unpaired electrons in order to be paramagnetic. c. Yes, they will necessarily be paramagnetic since each will have an odd number of electrons. 12. The concept is resonance. In this situation no single Lewis structure can adequately represent the bonding in a molecule. An average of the different Lewis structures is a better representation of the
Chem 1A Dr. White Updated /5/1 12 bonding than any single structure. The two important resonance structures are shown below. 1. Bond Order = 1/2(10-8) = 1, therefore it is stable Diamagnetic; 14. 2 σ 1s σ * 2 2 1s σ 2s σ * 2 2 2s σ 2 p π 4 * 2 p π 4 * 2 p σ 2 2 p 10 10 BO = = 0 2 molecule Since the bond order is 0, it is not a stable 15. The carbon atoms are sp 2 hybridized, leaving the 2p z orbital on each carbon unhybridized. Overlap of one sp 2 hybrid on each carbon and occupation of this region by an electron pair produces a C-C σ bond. Overlap of the remaining four hybrid orbitals with 1s orbitals of four hydrogen atoms, produces four C-H σ bonds. The unhybridized 2p z orbitals overlap laterally to produce a π bond which has two regions of electron density, above and below the C-C σ bond. 16. Bonding orbitals result from adding atomic orbitals (the wave functions constructively overlap). They are lower in energy. Antibionding orbitals result from subtracting atomic orbitals (the wave functions destructively overlap). They are higher in energy.