Log out Go to App
Go to App

Learning Materials

Features

Discover

Chapter 9: Problem 65

Consider the Lewis structure for glycine, the simplest amino acid: (a) What are the approximate bond angles about each of the two carbon atoms, and what are the hybridizations of the orbitals on each of them? (b) What are the hybridizations of the orbitals on the two oxygens and the nitrogen atom, and what are the approximate bond angles at the nitrogen? (c) What is the total number of \(\sigma\) bonds in the entire molecule, and what is the total number of \(\pi\) bonds?

Short Answer

Expert verified
The Lewis structure of glycine is H2N-CH2-CO2H. The first carbon atom has \(sp^2\) hybridization, and the second carbon atom also has \(sp^2\) hybridization. The bond angles are approximately 120 degrees. The nitrogen atom has \(sp^3\) hybridization, and both oxygen atoms have \(sp^2\) hybridization. The bond angle at the nitrogen atom is approximately 109.5 degrees. In the entire glycine molecule, there are a total of 9 σ bonds and 1 π bond.

Step by step solution

01

Determine the hybridization of the carbon atoms

To determine the hybridization of the carbon atoms, we must count the number of atoms directly connected to the carbon atom and add any lone pairs. For the first carbon atom (the one connected to the nitrogen): - It is bonded to a nitrogen, hydrogen, and other carbon atom, so it is connected to 3 atoms, and has no lone pairs. - Therefore, this carbon has \(sp^2\) hybridization. For the second carbon atom (the one connected to the oxygen atoms): - It is bonded to an oxygen, another oxygen (double bond), and other carbon atom, so it is connected to 3 atoms and has no lone pairs. - Therefore, this carbon also has \(sp^2\) hybridization.
02

Determine bond angles about each of the carbon atoms

Now that we know the hybridization, we can determine the bond angles. For an \(sp^2\) hybridized atom, the bond angles are approximately 120 degrees.
03

Determine hybridization of the orbitals on nitrogen and oxygen atoms

Now, let's determine the hybridization of the nitrogen and oxygen atoms: For the nitrogen atom: - It is connected to 3 hydrogen atoms and has a lone pair. - This nitrogen atom has \(sp^3\) hybridization. For both oxygen atoms: - The oxygen atom with a single bond to the carbon atom is connected to 2 hydrogen atoms and has 2 lone pairs. - The oxygen atom with a double bond to the carbon atom is connected to the carbon atom and has 2 lone pairs. - Both oxygen atoms have \(sp^2\) hybridization.
04

Determine approximate bond angles at nitrogen atom

For an \(sp^3\) hybridized atom like the nitrogen atom in glycine, the bond angles are approximately 109.5 degrees.
05

Determine the total number of σ and π bonds in the entire molecule

Glycine has the following bonds: - 2 C-N σ bonds - 3 N-H σ bonds - 2 C-C σ bonds - 1 C-O σ bond - 1 C=O double bond (1 σ bond and 1 π bond) Adding these up, we have a total of 9 σ bonds and 1 π bond in the entire glycine molecule.

Unlock Step-by-Step Solutions & Ace Your Exams!

  • Full Textbook Solutions

    Get detailed explanations and key concepts

  • Unlimited Al creation

    Al flashcards, explanations, exams and more...

  • Ads-free access

    To over 500 millions flashcards

  • Money-back guarantee

    We refund you if you fail your exam.

Start your free trial

Over 30 million students worldwide already upgrade their learning with Vaia!

One App. One Place for Learning.

All the tools & learning materials you need for study success - in one app.

Get started for free

Most popular questions from this chapter

How many nonbonding electron pairs are there in each of the following molecules: (a) \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{~S},\) (b) \(\mathrm{HCN},\) (c) \(\mathrm{H}_{2} \mathrm{C}_{2}\), (d) \(\mathrm{CH}_{3} \mathrm{~F} ?\)

Consider the \(\mathrm{H}_{2}^{+}\) ion. (a) Sketch the molecular orbitals of the ion and draw its energy-level diagram. (b) How many electrons are there in the \(\mathrm{H}_{2}^{+}\) ion? (c) Draw the electron configuration of the ion in terms of its MOs. (d) What is the bond order in \(\mathrm{H}_{2}{ }^{+}\) ? (e) Suppose that the ion is excited by light so that an electron moves from a lower-energy to a higherenergy MO. Would you expect the excited-state \(\mathrm{H}_{2}^{+}\) ion to be stable or to fall apart? Explain.

The phosphorus trihalides \(\left(\mathrm{PX}_{3}\right)\) show the following variation in the bond angle \(\mathrm{X}-\mathrm{P}-\mathrm{X}: \mathrm{PF}_{3}, 96.3^{\circ} ; \mathrm{PCl}_{3}, 100.3^{\circ} ; \mathrm{PBr}_{3}\), \(101.0^{\circ} ; \mathrm{PI}_{3}, 102.0^{\circ} .\) The trend is generally attributed to the change in the electronegativity of the halogen. (a) Assuming that all electron domains are the same size, what value of the \(\mathrm{X}-\mathrm{P}-\mathrm{X}\) angle is predicted by the VSEPR model? (b) What is the general trend in the \(\mathrm{X}-\mathrm{P}-\mathrm{X}\) angle as the halide electronegativity increases? (c) Using the VSEPR model, explain the observed trend in \(\mathrm{X}-\mathrm{P}-\mathrm{X}\) angle as the electronegativity of \(X\) changes. (d) Based on your answer to part (c), predict the structure of \(\mathrm{PBrCl}_{4}\).

(a) What is the probability of finding an electron on the internuclear axis if the electron occupies a \(\pi\) molecular orbital? (b) For a homonuclear diatomic molecule, what similarities and differences are there between the \(\pi_{2 p}\) MO made from the \(2 p_{x}\) atomic orbitals and the \(\pi_{2 p}\) MO made from the \(2 p_{y}\) atomic orbitals? (c) How do the \(\pi_{2 p}^{*}\) MOs formed from the \(2 p_{x}\) and \(2 p_{y}\) atomic orbitals differ from the \(\pi_{2 p}\) MOs in terms of energies and electron distributions?

(a) Draw Lewis structures for ethane \(\left(\mathrm{C}_{2} \mathrm{H}_{6}\right)\), ethylene \(\left(\mathrm{C}_{2} \mathrm{H}_{4}\right)\), and acetylene \(\left(\mathrm{C}_{2} \mathrm{H}_{2}\right) .\) (b) What is the hybridization of the carbon atoms in each molecule? (c) Predict which molecules, if any, are planar. (d) How many \(\sigma\) and \(\pi\) bonds are there in each molecule? (e) Suppose that silicon could form molecules that are precisely the analogs of ethane, ethylene, and acetylene. How would you describe the bonding about Si in terms of hydrid orbitals? Silicon does not readily form some of the analogous compounds containing \(\pi\) bonds. Why might this be the case?
See all solutions

Recommended explanations on Chemistry Textbooks

Inorganic Chemistry

Read Explanation

Physical Chemistry

Read Explanation

Making Measurements

Read Explanation

Organic Chemistry

Read Explanation

Chemical Analysis

Read Explanation

The Earths Atmosphere

Read Explanation
View all explanations

What do you think about this solution?

We value your feedback to improve our textbook solutions.

Study anywhere. Anytime. Across all devices.

Sign-up for free