covalent bonding homework answer key

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Covalent Bonding and Molecular Geometry

This chapter covers the formation and naming of covalent compounds. This includes drawing Lewis dot structures, predicting molecular geometry through the VSEPR theory, and the rules for naming covalent compounds.

Covalent Bonding Powerpoint Lecture

Purpose:  This is a brief Powerpoint lecture that describes the difference between covalent and other types of chemical bonds, including electronegativity difference. The rules of covalent nomenclature are also covered.

Essential Concepts: Covalent bonding, covalent compounds, electronegativity, VSEPR, Lewis dot structures.

Chemical Bonding Notes Outline

Purpose:  This is a fill-in-the-blank style notes outline for students to complete as you complete the accompanying Powerpoint lecture. Each slide has a set of questions, fill-in-the-blanks, or tables that students fill in based on the information given. This is a good aid for students who struggle with taking notes freehand.

Chemthink - Covalent Bonding and Nomenclature

Purpose:  This Chemthink module covers the formation of covalent bonds and how covalent compounds are named through the use of prefixes.

Essential concepts: Covalent bonding, electronegativity, covalent nomenclature.

Chemthink - Molecular Shapes

Purpose:  This Chemthink module helps students learn how to construct Lewis dot structures for covalent compounds and predict their molecular shapes with the VSEPR theory.

Essential concepts: Covalent bonding, Lewis dot structures, molecular geometry, VSEPR.

Electronegativity Difference and Covalent Compounds

Purpose:  This worksheet instructs students in the use of electronegativity difference to identify ionic, nonpolar covalent, and polar covalent compounds.

Essential concepts: Electronegativity, nonpolar covalent bond, polar covalent bond, ionic bond.

Lewis Dot Structures Worksheet

Purpose: The creation of Lewis Dot structures is a helpful first step in predicting the molecular shape made by a covalent compound. In this worksheet, students will be guided in making Lewis Dot structures both for individual atoms and for molecules.

Essential concepts: Covalent compound, molecular geometry, Lewis Dot structures

VSEPR Theory and Molecular Geometry

Purpose:  This worksheet guides students through the use of the VSEPR theory to predict the 3-dimensional shape made by a covalently-bonded molecule. This worksheet only covers the simpler and more common molecular shapes, including linear, bent, trigonal planar, trigonal pyrimidal, and tetrahedral.

Essential concepts: Molecular geometry, VSEPR, linear, bent, trigonal planar, trigonal pyrimidal, and tetrahedral.

VSEPR Theory with Molecular Model Kits

Purpose:  Valence Shell Electron Pair Repulsion, or VSEPR Theory, is a way to determine what geometric shape a covalent compound will make based on the number of bonds and unpaired electrons surrounding the central atom of the compound. This is a chart that I have students fill out as they use a chemistry model kit to build various covalent compounds.

Essential concepts: VSEPR, molecular geometry, linear, trigonal planar, bent, tetrahedral, trigonal pyramidal, Lewis dot structures.

Covalent Compound Nomenclature Worksheet

Purpose:  This worksheet instructs students on the use of prefixes (mono-, di-, etc) to name covalent compounds.

Essential concepts: Covalent compounds, covalent nomenclature

Covalent Bonding and Molecular Geometry Study Guide

Purpose:  Once the instruction for the unit is completed, students can complete this study guide to aid in their preparation for a written test. The study guide is divided into two sections: vocabulary and short answer questions. The vocabulary words can be found scattered throughout the different instructional worksheets from this unit. The short answer questions are conceptual and meant to see if the students are able to apply what they've learned in the unit.

The protons in the nucleus do not change during normal chemical reactions. Only the outer electrons move. Positive charges form when electrons are lost.

P, I, Cl, and O would form anions because they are nonmetals. Mg, In, Cs, Pb, and Co would form cations because they are metals.

(a) P 3– ; (b) Mg 2+ ; (c) Al 3+ ; (d) O 2– ; (e) Cl – ; (f) Cs +

(a) [Ar]4 s 2 3 d 10 4 p 6 ; (b) [Kr]4 d 10 5 s 2 5 p 6 (c) 1 s 2 (d) [Kr]4 d 10 ; (e) [He]2 s 2 2 p 6 ; (f) [Ar]3 d 10 ; (g) 1 s 2 (h) [He]2 s 2 2 p 6 (i) [Kr]4 d 10 5 s 2 (j) [Ar]3 d 7 (k) [Ar]3 d 6 , (l) [Ar]3 d 10 4 s 2

(a) 1 s 2 2 s 2 2 p 6 3 s 2 3 p 1 ; Al 3+ : 1 s 2 2 s 2 2 p 6 ; (b) 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 10 4 s 2 4 p 5 ; 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 10 4 s 2 4 p 6 ; (c) 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 10 4 s 2 4 p 6 5 s 2 ; Sr 2+ : 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 10 4 s 2 4 p 6 ; (d) 1 s 2 2 s 1 ; Li + : 1 s 2 ; (e) 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 10 4 s 2 4 p 3 ; 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 10 4 s 2 4 p 6 ; (f) 1 s 2 2 s 2 2 p 6 3 s 2 3 p 4 ; 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6

NaCl consists of discrete ions arranged in a crystal lattice, not covalently bonded molecules.

ionic: (b), (d), (e), (g), and (i); covalent: (a), (c), (f), (h), (j), and (k)

(a) Cl; (b) O; (c) O; (d) S; (e) N; (f) P; (g) N

(a) H, C, N, O, F; (b) H, I, Br, Cl, F; (c) H, P, S, O, F; (d) Na, Al, H, P, O; (e) Ba, H, As, N, O

N, O, F, and Cl

(a) HF; (b) CO; (c) OH; (d) PCl; (e) NH; (f) PO; (g) CN

(a) eight electrons:

; (b) eight electrons:

; (c) no electrons Be 2+ ; (d) eight electrons:

; (e) no electrons Ga 3+ ; (f) no electrons Li + ; (g) eight electrons:

In this case, the Lewis structure is inadequate to depict the fact that experimental studies have shown two unpaired electrons in each oxygen molecule . (b)

(a) SeF 6 :

; (b) XeF 4 :

; (c) SeCl 3 + : SeCl 3 + :

; (d) Cl 2 BBCl 2 :

Two valence electrons per Pb atom are transferred to Cl atoms; the resulting Pb 2+ ion has a 6 s 2 valence shell configuration. Two of the valence electrons in the HCl molecule are shared, and the other six are located on the Cl atom as lone pairs of electrons.

Each bond includes a sharing of electrons between atoms. Two electrons are shared in a single bond; four electrons are shared in a double bond; and six electrons are shared in a triple bond.

CO has the strongest carbon-oxygen bond because there is a triple bond joining C and O. CO 2 has double bonds.

(a) H: 0, Cl: 0; (b) C: 0, F: 0; (c) P: 0, Cl 0; (d) P: 0, F: 0

Cl in Cl 2 : 0; Cl in BeCl 2 : 0; Cl in ClF 5 : 0

The structure that gives zero formal charges is consistent with the actual structure:

(a) −114 kJ; (b) 30 kJ; (c) −1055 kJ

The greater bond energy is in the figure on the left. It is the more stable form.

HCl ( g ) ⟶ 1 2 H 2 ( g ) + 1 2 Cl 2 ( g ) Δ H 1 ° = −Δ H f [ HCl ( g ) ] ° 1 2 H 2 ( g ) ⟶ H ( g ) Δ H 2 ° = Δ H f [ H ( g ) ] ° 1 2 Cl 2 ( g ) ⟶ Cl ( g ) Δ H 3 ° = Δ H f [ Cl ( g ) ] ° ¯ HCl ( g ) ⟶ H ( g ) + Cl ( g ) Δ H 298 ° = Δ H 1 ° + Δ H 2 ° + Δ H 3 ° HCl ( g ) ⟶ 1 2 H 2 ( g ) + 1 2 Cl 2 ( g ) Δ H 1 ° = −Δ H f [ HCl ( g ) ] ° 1 2 H 2 ( g ) ⟶ H ( g ) Δ H 2 ° = Δ H f [ H ( g ) ] ° 1 2 Cl 2 ( g ) ⟶ Cl ( g ) Δ H 3 ° = Δ H f [ Cl ( g ) ] ° ¯ HCl ( g ) ⟶ H ( g ) + Cl ( g ) Δ H 298 ° = Δ H 1 ° + Δ H 2 ° + Δ H 3 ° D HCl = Δ H 298 ° = Δ H f [ HCl ( g ) ] ° + Δ H f [ H ( g ) ] ° + Δ H f [ Cl ( g ) ] ° = − ( −92.307 kJ ) + 217.97 kJ + 121.3 kJ = 431.6 kJ D HCl = Δ H 298 ° = Δ H f [ HCl ( g ) ] ° + Δ H f [ H ( g ) ] ° + Δ H f [ Cl ( g ) ] ° = − ( −92.307 kJ ) + 217.97 kJ + 121.3 kJ = 431.6 kJ

The S–F bond in SF 4 is stronger.

The C–C single bonds are longest.

(a) When two electrons are removed from the valence shell, the Ca radius loses the outermost energy level and reverts to the lower n = 3 level, which is much smaller in radius. (b) The +2 charge on calcium pulls the oxygen much closer compared with K, thereby increasing the lattice energy relative to a less charged ion. (c) Removal of the 4 s electron in Ca requires more energy than removal of the 4 s electron in K because of the stronger attraction of the nucleus and the extra energy required to break the pairing of the electrons. The second ionization energy for K requires that an electron be removed from a lower energy level, where the attraction is much stronger from the nucleus for the electron. In addition, energy is required to unpair two electrons in a full orbital. For Ca, the second ionization potential requires removing only a lone electron in the exposed outer energy level. (d) In Al, the removed electron is relatively unprotected and unpaired in a p orbital. The higher energy for Mg mainly reflects the unpairing of the 2 s electron.

4008 kJ/mol; both ions in MgO have twice the charge of the ions in LiF; the bond length is very similar and both have the same structure; a quadrupling of the energy is expected based on the equation for lattice energy

(a) Na 2 O; Na + has a smaller radius than K + ; (b) BaS; Ba has a larger charge than K; (c) BaS; Ba and S have larger charges; (d) BaS; S has a larger charge

The placement of the two sets of unpaired electrons in water forces the bonds to assume a tetrahedral arrangement, and the resulting HOH molecule is bent. The HBeH molecule (in which Be has only two electrons to bond with the two electrons from the hydrogens) must have the electron pairs as far from one another as possible and is therefore linear.

Space must be provided for each pair of electrons whether they are in a bond or are present as lone pairs. Electron-pair geometry considers the placement of all electrons. Molecular structure considers only the bonding-pair geometry.

As long as the polar bonds are compensated (for example. two identical atoms are found directly across the central atom from one another), the molecule can be nonpolar.

(a) Both the electron geometry and the molecular structure are octahedral. (b) Both the electron geometry and the molecular structure are trigonal bipyramid. (c) Both the electron geometry and the molecular structure are linear. (d) Both the electron geometry and the molecular structure are trigonal planar.

(a) electron-pair geometry: octahedral, molecular structure: square pyramidal; (b) electron-pair geometry: tetrahedral, molecular structure: bent; (c) electron-pair geometry: octahedral, molecular structure: square planar; (d) electron-pair geometry: tetrahedral, molecular structure: trigonal pyramidal; (e) electron-pair geometry: trigonal bypyramidal, molecular structure: seesaw; (f) electron-pair geometry: tetrahedral, molecular structure: bent (109°)

(a) electron-pair geometry: trigonal planar, molecular structure: bent (120°); (b) electron-pair geometry: linear, molecular structure: linear; (c) electron-pair geometry: trigonal planar, molecular structure: trigonal planar; (d) electron-pair geometry: tetrahedral, molecular structure: trigonal pyramidal; (e) electron-pair geometry: tetrahedral, molecular structure: tetrahedral; (f) electron-pair geometry: trigonal bipyramidal, molecular structure: seesaw; (g) electron-pair geometry: tetrahedral, molecular structure: trigonal pyramidal

All of these molecules and ions contain polar bonds. Only ClF 5 , ClO 2 − , ClO 2 − , PCl 3 , SeF 4 , and PH 2 − PH 2 − have dipole moments.

SeS 2 , CCl 2 F 2 , PCl 3 , and ClNO all have dipole moments.

(a) tetrahedral; (b) trigonal pyramidal; (c) bent (109°); (d) trigonal planar; (e) bent (109°); (f) bent (109°); (g) C H 3 CCH tetrahedral, CH 3 CC H linear; (h) tetrahedral; (i) H 2 C C CH 2 linear; H 2 C C C H 2 trigonal planar

; (d) CS 3 2− CS 3 2− includes three regions of electron density (all are bonds with no lone pairs); the shape is trigonal planar; CS 2 has only two regions of electron density (all bonds with no lone pairs); the shape is linear

The Lewis structure is made from three units, but the atoms must be rearranged:

The molecular dipole points away from the hydrogen atoms.

The structures are very similar. In the model mode, each electron group occupies the same amount of space, so the bond angle is shown as 109.5°. In the “real” mode, the lone pairs are larger, causing the hydrogens to be compressed. This leads to the smaller angle of 104.5°.

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• Authors: Paul Flowers, William R. Robinson, PhD, Richard Langley, Klaus Theopold
• Publisher/website: OpenStax
• Book title: Chemistry
• Publication date: Mar 11, 2015
• Location: Houston, Texas
• Book URL: https://openstax.org/books/chemistry/pages/1-introduction
• Section URL: https://openstax.org/books/chemistry/pages/chapter-7

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