Unit 2 Practice Test Key PDF

## Unit 2 Practice Test Key ### Molecular and Ionic Compound Structure and Properties 1. For each of these polar covalent bonds, which atom has a partial positive charge (+) and which atom has a partial negative charge (–)? | Bond | Partial Positive Charge | Partial Negative Charge | |------...

## Unit 2 Practice Test Key ### Molecular and Ionic Compound Structure and Properties 1. For each of these polar covalent bonds, which atom has a partial positive charge (+) and which atom has a partial negative charge (–)? | Bond | Partial Positive Charge | Partial Negative Charge | |-----------|-------------------------|--------------------------| | C - N | C | N | | F - Br | F | Br | | Si - O | Si | O | 2. Arrange these three bonds in order from least polar to most polar. * Least polar: N-F * Middle polar: P-F * Most polar: O-F 3. A graph of potential energy versus internuclear distance for two Cl atoms is given below. On the same graph, carefully sketch a curve that corresponds to potential energy versus internuclear distance for two Br atoms. Then, predict the magnitude of the bond length and bond energy for a Cl-Br bond by filling in the table. (Use terms like greater than, less than, or equal to.) | Bond | Bond Length (pm) | Bond Energy (kJ/mol) | |---------|--------------------|-----------------------| | Cl-Cl | 200 | 243 | | Cl-Br | Greater than 200 | Less than 243 | 4. Data for the lattice energy of NaF is given in the table below. Make predictions about the lattice energy of MgF₂ and KF. Do you predict that the lattice energy of each compound is less than 930 kJ/mol or greater than 930 kJ/mol? Justify your answer in terms of periodic properties and Coulomb's law. | Reaction | Lattice Energy (kJ/mol) | |-------------------------|------------------------| | NaF(s) → Na+(g) + F^-(g) | 930 | | MgF₂(s) → Mg²⁺(g) + 2 F^-(g) | Greater than 930 kJ/mol | | KF(s) → K+(g) + F^-(g) | Less than 930 kJ/mol | The lattice energy of MgF₂ is greater than that of NaF. The magnitude of charge in MgF₂ (Mg²⁺ and F^-) is greater than that of NaF (Na⁺ and F^-). According to Coulomb's law, as the magnitude of charge of ions increases, the force of attraction and lattice energy increases. The lattice energy of KF is less than that of NaF. Both ionic compounds have the same magnitude of charge on their ions (K⁺ and F^-); however, K⁺ is a larger ion than Na⁺. This means that the chemical distance between K⁺ and F^- should be greater than the distance between Na⁺ and F^-. 5. A student checked the conductivity of two different solids, and the results are listed in the table below. Explain why the student got these results, in terms of principles of chemical bonding. Your explanation should include a discussion of the specific particles present in each substance and how the behavior of these particles is related to the conductivity of the solid. | Substance | Does the solid conduct electricity? | |-----------|------------------------------------| | Cu(s) | Yes | | CuCl₂(s) | No | In order for a substance to conduct electricity, there must be free-moving charged particles. Cu(s) experiences metallic bonds, which have a sea of mobile valence electrons. Because the valence electrons are delocalized and free to move throughout the solid, Cu(s) can conduct electricity. CuCl₂(s) does have charged particles (ions), but in the solid state, they are locked in place in the rigid crystal lattice structure. For this reason, CuCl₂(s) cannot conduct electricity. 6. Using the models shown above, which represents an alloy of C and Fe? Which represents an alloy of Pt and Au? * Pt & Au on left * C & Fe on right 7. Shown above are three possible Lewis diagrams for the cyanate ion, NCO^-. Which diagram is the better representation of the bonding in the cyanate ion? What is wrong with the other two diagrams? Justify your answer using formal charge. *Diagram 2 is best. The sum of the formal charges equals the charge of the ion, and the more electronegative atom (O) has the formal negative charge (-1). *Diagram 1 has the same sum of formal charge as #2 but the formal charge (- 1) is not placed on the more electronegative element. *Diagram 3, the formal charges for N, C, and O are -2, 0, and +1 respectively. This is not the best diagram because the size of the O is not as small as it could be. 8. Rank the following species in order of increasing (shortest to longest) O-O bond length: O₂, O₃, H₂O₂. * O₂ < O₃ < H₂O₂ 9. Answer the following questions about nitrous acid (HNO₂) and the nitrite ion (NO₂^-). a. Complete the Lewis electron-dot diagram for HNO₂ by drawing the required number of electrons, including bonding pairs and nonbonding pairs. Each atom should have a full valence shell and a formal charge of zero. ``` .. .. .. H - O - N = O .. .. 146 pm 120 pm ``` b. When the HNO₂ molecule was analyzed, it was determined that one of the nitrogen-oxygen bonds has a length of 120 pm, and the other nitrogen-oxygen bond has a length of 146 pm. Based on the Lewis diagram that you drew, fill in the two blanks above with either a bond length of 120 pm or a bond length of 146 pm. c. Complete the Lewis electron-dot diagram for the nitrite ion, NO₂^-, by drawing the required number of electrons, including bonding pairs and nonbonding pairs. ``` .. .. [ : O - N = O : ]<sup>-</sup> .. .. ``` d. It was determined that the two nitrogen-oxygen bonds in the nitrite ion have the same bond length. Explain this result using principles of chemical bonding. NO₂^- experiences resonance, so both N-O bonds are the average of a single and double bond. Both bonds have a bond order of 1.5. e. Estimate the bond length of the nitrogen-oxygen bond in the nitrite ion: ~ 130 pm (any value between 120 & 146 pm is acceptable). 10. Write a definition for the octet rule. Do not use the word "eight" in your definition. When atoms bond, they tend to do so in such a way that they have a full valence shell of electrons. 11. Draw Lewis diagrams for each of the following molecules. There should only be single bonds in each structure. * PF₃ ``` .. .. : F : | : F - P - F : | : F : .. .. ``` * SBr₂ ``` .. .. : Br : | : Br - S - Br : | : Br : .. .. ``` * CH₃NH₂ ``` H | H - C - N - H | H | H ``` 12. Draw Lewis diagrams for each of the following molecules. There may be double or triple bonds in each structure. * C₂H₂ ``` H - C ≡ C - H ``` * HCN ``` H - C ≡ N : ``` * SO₂ ``` O || S / \ :O : ``` 13. Draw Lewis diagrams for each of the following polyatomic ions. * ClO₃ ^- ``` .. O || Cl / \ :O : :O: [ - ] ``` * NO₂⁺ ``` .. O || N / \ :O: :O: [ + ] ``` 14. Complete the table below. | Molecule | Molecular geometry | Bond angles | Is the molecule polar or nonpolar? | Number of sigma and pi bonds | |-----------|----------------------|-------------|--------------------------------------|-----------------------------------| | CS₂ | Linear | 180° | Nonpolar | σ: 2 π: 2 | | BF₃ | Trigonal planar | 120° | Nonpolar | σ: 3 π: 0 | | SO₂ | Bent | <120° | Polar | σ: 2 π: 1 | | CH₃Cl | Tetrahedral | 109.5° | Polar | σ: 4 π: 0 | | XeF₄ | Square planar | 90°, 180° | Nonpolar | σ: 4 π: 0 | | SF₆ | Octahedral | 90°, 180° | Nonpolar | σ: 6 π: 0 | | NF₃ | Trigonal pyramidal | ~107° | Polar | σ: 3 π: 0 | | H₂S | Bent | ~104.5° | Polar | σ: 2 π: 0 | 15. Explain why the bond angles in H₂O and NH₃ are less than 109.5°. H₂O and NH₃ have 2 lone pairs and 1 lone pair, respectively, on the central atom. Lone pairs repel more than bonding electrons. Thus, the bond angles are less than 109.5°. 16. Identify the orbital hybridization for each carbon atom in the following molecule. ``` H H H | | | H - C = C - C - C = C - C - H | | | H H H sp² sp² sp³ sp sp sp³ ``` 17. Using the molecule shown above, predict the following bond angles. * H_x - C_y - C_z: 109.5° * H_a - C_b - C_c: 120° * C_p - C_q - N: 180°