Are the O-N-O bond angles greater in the nitrile i...

Question

Are the O-N-O bond angles greater in the nitrile ion (NO₂^-) or the nitrate ion (NO₃^-)?

Please select the answer the best explains your conclusion.

a. Nitrite has the greater bond angle because a trigonal planar bond angle is greater than a tetrahedral bond angle

b. Nitrite has the greater bond angle because a linear bond angle is greater than a trigonal planar bond angle

c. Nitrate has the greater bond angle because nitrite’s lone pairs take up more space than bonds

d. Nitrate has the greater bond angle because double bonds take up more space than single bonds

Jules Bruno · Accepted Answer

We’re being asked to identify which has greater O-N-O bond angles. Since we don’t know the Lewis structure for nitrile ion (NO2-) or the nitrate ion (NO3-), we need to do the following steps:Step 1: Determine the central atom in this molecule.Step 2: Calculate the total number of valence electrons present.Step 3: Draw the Lewis structure for the molecule.Step 4: Determine the number of electron groups around the indicated atom.Step 5: Determine the electron geometry and bond angle using this:Electron Regions          Electronic Geometry          Bond Angles2                                      linear                                    180˚3                                      trigonal planar                      120˚4                                      tetrahedral                            109.5˚5                                      trigonal bipyramidal              90˚, 120˚, and 180˚6                                      octahedral                            90˚ and 180˚[readmore]For nitrile ion (NO2-):Step 1: Nitrogen is less electronegative than oxygen so nitrogen is the central atom.Step 2: The total number of valence electrons present in NO2- is:            Group             Valence ElectronsN         5A                   1 × 5 e– = 5 e–O         6A                   2 × 6 e– = 12 e–-charge                                  +1 e–                                Total:  18 valence e–Step 3: Nitrogen prefers to form 3 bonds while oxygen prefers to make 2 bonds. The incomplete Lewis structure of NO2- is:A bond corresponds to 2 electrons shared between two atoms so we currently have 3 × 2 e– = 6 val. e–, which means we’re still lacking 12 e–. We can start adding the missing electrons to the atoms while following the octet rule. The complete Lewis structure is:Step 4: Nitrogen is surrounded by 3 electron groups: 2 bonds (double bonds are counted as 1) and 1 lone pair.Step 5: Since there are 3 electron groups around nitrogen, the electron geometry is trigonal planar, whose ideal bond angle is 120˚.]Recall that lone pairs compress the bond angle, causing a deviation from the ideal bond angle.Therefore, the bond angle of NO2- is less than 120˚.For nitrate ion (NO3-):Step 1: Nitrogen is less electronegative than oxygen so nitrogen is the central atom.Step 2: The total number of valence electrons present in NO3- is:            Group             Valence ElectronsN         5A                   1 × 5 e– = 5 e–O         6A                   2 × 6 e– = 18 e–-charge                                  +1 e–                                Total:  24 valence e–Step 3: Nitrogen prefers to form 3 bonds while oxygen prefers to make 2 bonds. The incomplete Lewis structure of NO3- is:*note that we can only make up to 4 bonds for N because we need to follow the octet ruleA bond corresponds to 2 electrons shared between two atoms so we currently have 4 × 2 e– = 8 val. e–, which means we’re still lacking 16 e–. We can start adding the missing electrons to the atoms while following the octet rule. The complete Lewis structure is:Step 4: Nitrogen is surrounded by 3 electron groups: 2 bonds (double bonds are counted as 1) and 1 lone pair.Step 5: Since there are 3 electron groups around nitrogen, the electron geometry is trigonal planar, whose ideal bond angle is 120˚.Therefore, the bond angle of NO2- is 120˚.Therefore, the answer is c. Nitrate has the greater bond angle because nitrite’s lone pairs take up more space than bonds.

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