Question

Write the Lewis structure for each of the following species, describe its geometry, and indicate the oxidation state of the (a) \(\mathrm{NH}_{4}^{+},(\mathbf{b}) \mathrm{NO}_{2}^{-}\), nitrogen: (c) \(\mathrm{N}_{2} \mathrm{O},(\mathrm{d}) \mathrm{NO}_{2}\)

Short answer

(a) NH4+: Lewis structure is H - N - H with H below N and a lone pair on N. Tetrahedral shape. Oxidation state of N is +5. (b) NO2-: Lewis structure is O = N - O with a lone pair on N. Bent shape. Oxidation state of N is +3. (c) N2O: Lewis structure is N = N = O. Linear shape. Oxidation state of the left N is +1, and the right N is +4. (d) NO2: Lewis structure is O = N - O with an unpaired electron on N. Bent shape. Oxidation state of N is +4.

Step-by-step solution

Calculate the number of valence electrons

For each species, we need to calculate the total number of valence electrons. Valence electrons are found in the outermost shell of each element, and their number can be determined from the periodic table. (a) NH4+ Valence electrons for N: 5 Valence electrons for 4 H atoms: 4(1) = 4 Since it has a positive charge, subtract 1 electron: Total valence electrons: 5+4 - 1 = 8 (b) NO2- Valence electrons for N: 5 Valence electrons for 2 O atoms: 2(6) = 12 Since it has a negative charge, add 1 electron: Total valence electrons: 5+12 + 1 = 18 (c) N2O Valence electrons for 2 N atoms: 2(5) = 10 Valence electrons for O: 6 Total valence electrons: 10 + 6 = 16 (d) NO2 Valence electrons for N: 5 Valence electrons for 2 O atoms: 2(6) = 12 Total valence electrons: 5 + 12 = 17

Draw the Lewis structures

Now we will distribute the valence electrons to each atom in the species, connecting them with a single bond. (a) NH4+: H - N - H has a trigonal pyramidal shape with 3 single bonds and a lone pair on nitrogen. \[ \mathrm{H} - \mathrm{N}^{-} - \mathrm{H} \] \[ \mid \mathrm{H} \] (b) NO2-: O = N - O has a bent shape with a single bond between N and one O atom and a double bond between N and the other O atom. There is also a lone pair on nitrogen. \[ \mathrm{O} = \mathrm{N}^+ - \mathrm{O}^- \] (c) N2O: N = N = O has a linear shape with a triple bond between the two N atoms and a double bond between the N and the O atom. \[ \mathrm{N} = \mathrm{N} = \mathrm{O} \] (d) NO2: O = N - O has a bent shape with a single bond between N and one O atom, a double bond between N and the other O atom, and an unpaired electron on nitrogen. \[ \mathrm{N}^+ - \mathrm{O}^{-} \] \[ \mathrm{O} = \mathrm{N} \cdot \]

Identify the shape and oxidation state of each species

Now that we have the Lewis structures, we can determine the shape of each molecule and the oxidation state of nitrogen. (a) NH4+ Shape: Tetrahedral Oxidation state of N: +5 (each H contributes -1, so +5 is needed to offset) (b) NO2- Shape: Bent Oxidation state of N: +3 (each O contributes -2, so +3 is needed to offset) (c) N2O Shape: Linear Oxidation state of the left N: +1 (N triple-bonded to N contributes -3, so +1 is needed to offset) Oxidation state of the right N: +4 (triple-bonded N has an oxidation state of +4) (d) NO2 Shape: Bent Oxidation state of N: +4 (each O contributes -2, so +4 is needed to offset)

Key concepts

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of the atoms in a molecule. Understanding molecular geometry helps us predict the physical and chemical properties of a substance. For example, the

• Tetrahedral Geometry: This shape is typical for molecules like NH extsubscript{4} extsuperscript{+}. In this configuration, the central atom is surrounded symmetrically by four other atoms at 109.5-degree angles.
• Bent Geometry: This geometry is evident in molecules like NO extsubscript{2} extsuperscript{-} and NO extsubscript{2}. In a bent geometry, the molecule is shaped more like a "V" due to the angular repulsion created by lone pairs of electrons.
• Linear Geometry: Seen in N extsubscript{2}O, where the atoms are arranged in a straight line. This form usually results in 180-degree bond angles, indicative of a minimal electron repulsion landscape.

Knowing molecular geometry allows us to predict other features such as polarity and intermolecular forces, which are crucial for understanding reaction mechanisms and the physical attributes of substances.

Oxidation States

An oxidation state is a concept that gives insight into the electron distribution within a molecule. In essence, it's the hypothetical charge an atom would have if all bonds were ionic. Here’s how it manifests in the given species:

• NH extsubscript{4} extsuperscript{+}: The nitrogen atom has an oxidation state of +5 because each hydrogen atom is considered to have contributed -1 to the charge, necessitating nitrogen to balance it to +5.
• NO extsubscript{2} extsuperscript{-}: Here, nitrogen lands at an oxidation state of +3, balancing the negative impact of oxygen's -2 charge per atom.
• N extsubscript{2}O: This compound illustratively demonstrates different oxidation states for nitrogen; the left nitrogen is +1, while the right-sided nitrogen is at +4 due to the unique bonding configuration.
• NO extsubscript{2}: In this case, nitrogen supports an oxidation state of +4, responsible for offsetting the negative charges from the oxygen atoms.

Understanding oxidation states lays the foundation for grasping redox reactions and helps in predicting the behavior of molecules in various chemical reactions.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom. They play a crucial role in chemical bonding as they are the ones primarily involved when atoms interact. To find out how many valence electrons are involved in a molecule:

• Look at nitrogen: it has 5 valence electrons.
• The hydrogen atom in ammonia ( NH extsubscript{4} extsuperscript{+}) contributes 1 electron each, leading to a total consideration of 8, factoring in the '+1' charge.
• For oxygen atoms as in nitrites ( NO extsubscript{2} extsuperscript{-}) or similar, 6 electrons each are counted and additional electrons are adjusted based on the charge of the molecule.

Correctly identifying the count of valence electrons is the first step in drawing accurate Lewis structures, which depict how electrons are arranged among atoms in a molecule.

Chemical Bonding

Chemical bonding is the process wherein atoms combine by sharing or transferring valence electrons, forming different types of bonds:

• Covalent Bonds: In our species, these bonds are predominant. A covalent bond entails the sharing of electron pairs; for instance, in N extsubscript{2}O, nitrogen and oxygen are connected through a double bond, sharing two pairs of electrons.
• Coordinate Bonds: Also known as dative covalent bonds, occur when one atom provides a pair of electrons in the bond, as seen in the ionic part of NH extsubscript{4} extsuperscript{+}. Though formally presented as ionic, they still have covalent character.
• Ionic Bonds: Though less prominent here, they play a role in maintaining overall charge balance when molecules gain or lose electrons, as with the negative charge on NO extsubscript{2} extsuperscript{-}.

Understanding these bonds is integral for explaining properties such as solubility, electrical conductivity, and reactivity of molecules. This foundational knowledge allows us to explore more complex chemical interactions and properties in future studies.