Ch a p t e r s 8 a n d 9

Size: px
Start display at page:

Download "Ch a p t e r s 8 a n d 9"

Transcription

1 Ch a p t e r s 8 a n d 9 Covalent Bonding and Molecular Structures Objectives You will be able to: 1. Write a description of the formation of the covalent bond between two hydrogen atoms to form a hydrogen molecule. Your description should include mention of overlapping atomic orbitals to form molecular orbitals. It should also include rough pictures of the atomic orbitals and the molecular orbitals. 2. Write a description of the assumptions made in the Linear Combination of Atomic Orbitals (LCAO) approximation for the molecular orbital approach to covalent bonding. 3. Write an explanation for why bonding molecular orbitals are more stable and why antibonding molecular orbitals are less stable than the separate atomic orbitals that form them. 4. Write a description of how the σ, σ*, σ 2s, σ* 2s, σ, σ*, π*, and π* are formed from the, 2s and atomic orbitals. 5. Draw sketches of the σ, σ*, σ, σ*, π, and π* molecular orbitals. 6. Write an explanation for why the σ molecular orbital of O 2 is lower energy than the π molecular orbitals and why the σ* is higher energy than the π*. 7. Draw molecular orbital diagrams for 2, O 2, 2, and Ne rom a molecular orbital diagram of a molecule, predict its bond order. 9. rom a molecular orbital diagram for a molecule, predict whether it is paramagnetic or diamagnetic. 10. Draw the molecular orbital diagrams for CO and NO and use them to predict the bond order of each and to explain why each is relatively stable despite their strange Lewis Structures. 11. Describe the advantages and disadvantages of scientific models. 12. Write a list of the basic assumptions of the valence bond model of covalent bonding. 13. Write or identify the number of valence electrons for each of the Representative elements. 14. With reference to the valence bond model, describe how the covalent bond forms between fluorine atoms in Write or identify a description of the information given in a Lewis structure. 16. With reference to the valence bond model, describe how the covalent bond forms between hydrogen atoms in Explain how the carbon atom in C 4 is able to form four equivalent covalent bonds to hydrogen atoms with bond angles of Describe (1) how sp 3 hybrid orbitals form, (2) their shapes, and (3) how they are arranged in space. 19. Write or identify a description of the tetrahedral electron group geometry. 157

2 158 Chapter 8 & 9 Covalent Bonding and Molecular Structures 20. or ethene (ethylene) molecules, C 2 4 : (1) explain how the carbon atoms are able to form four equivalent covalent bonds to hydrogen atoms, (2) explain why all the bond angles are about 120, (3) with reference to the valence bond model, explain how the double bond between carbon atoms forms, and (4) explain why one bond in the double bond is weaker than the other. 21. Describe (1) how sp 2 hybrid orbitals form, (2) their shapes, and (3) how they are arranged in space. 22. Write or identify a description of the trigonal planar electron group geometry. 23. or ethyne (acetylene) molecules, C 2 2 : (1) explain how the carbon atoms are able to form two equivalent covalent bonds to hydrogen atoms, (2) explain why the bond angles are 180, (3) with reference to the valence bond model, explain how the triple bond between carbon atoms forms, and (4) explain why two bonds in the triple bond are weaker than the other one. 24. Describe (1) how sp hybrid orbitals form, (2) their shapes, and (3) how they are arranged in space. 25. Write or identify a description of the linear electron group geometry. 26. Explain how the nitrogen atom in N 3 is able to form three equivalent covalent bonds to hydrogen atoms with bond angles of Write or identify a description of the trigonal pyramid molecular geometry. 28. Explain how the nitrogen atom in N 4 + is able to form four equivalent covalent bonds to hydrogen atoms with bond angles of Explain how the oxygen atom in 2 O is able to form two equivalent covalent bonds to hydrogen atoms with bond angles of Write or identify a description of the bent molecular geometry. 31. With reference to the valence bond model, explain how the triple bond forms in the carbon monoxide, CO, molecule. 32. Explain how the boron atom in B 3 is able to form three equivalent covalent bonds to fluorine atoms with bond angles of Explain how the phosphorus atom in P 5 is able to form five equivalent covalent bonds to fluorine atoms with bond angles of, 120, and Write or identify a description of the trigonal bipyramid electron group geometry. 35. Explain how the sulfur atom in S 4 is able to form four covalent bonds to fluorine atoms with bond angles of, 120, and Write or identify a description of the see-saw molecular geometry. 37. Explain how the iodine atom in I 3 is able to form three covalent bonds to fluorine atoms with bond angles of and Write or identify a description of the T-shaped molecular geometry. 39. Explain how the xenon atom in Xe 2 is able to form two covalent bonds to fluorine atoms with a bond angle of Write or identify a description of the linear molecular geometry. 41. Explain how the sulfur atom in S 6 is able to form six equivalent covalent bonds to fluorine atoms with bond angles of and Write or identify a description of the octahedral electron geometry. 43. Explain how the iodine atom in I 5 is able to form five equivalent covalent bonds to fluorine atoms with bond angles of and Write or identify a description of the square pyramid molecular group geometry.

3 45. Explain how the xenon atom in Xe 4 is able to form four covalent bonds to fluorine atoms with bond angles of and Write or identify a description of the square planar molecular geometry. 47. Write an explanation for why the elements listed on the Table 8,9.3 can have each of the bonding patterns listed. 48. Given a Lewis structure or enough information to get one, do the following: (1) list the hybridization for each atom in the structure, (2) write a description of the formation of each bond in terms of the overlap of atomic orbitals, (3) describe the electron group geometry around each atom that has two or more atoms attached to it, (4) draw the geometric sketch of the molecule, including bond angles, and (5) describe the molecular geometry around each atom that has two or more atoms attached to it. 49. Given a Lewis structure or enough information to get one, determine the formal charge for each atom in the structure. 50. Given a formula for a molecule or polyatomic ion, draw a reasonable Lewis structure that corresponds to the formula. 51. Given a Lewis structure or enough information to get one, do the following: (1) identify whether it is best described as having resonance or not, (2) if can be seen as having resonance, draw the Lewis Structures for all of the reasonable resonance forms, and (3) draw the resonance hybrid for the formula. 52. Draw or identify cis and trans isomers for molecules with carbon-carbon double bonds. 53. Given a structure for a triglyceride, identify it as saturated or unsaturated, and identify whether it is more likely to be solid or liquid at room temperature. 54. Given a formula for a molecular compound, predict whether or not it represents a polar molecule. 55. Write a description of how London forces come to be. 56. Write an explanation for why larger molecules have stronger London forces than smaller molecules. 57. Write a description of how London forces play a part in the attractions between polar molecules. 58. Given a formula for a pure substance, identify which of the following types of substances it represents: ionic compound, metallic substance, polar molecular compound with hydrogen bonds, polar molecular substance without hydrogen bonds, or nonpolar molecular substance. 59. Identify carbon in the diamond form, C(dia), and silicon dioxide, SiO 2 as examples of network crystals. 60. Given a formula for a pure substance, identify which of the following attractions is the predominant force holding its particles in the liquid and solid state: ionic bonds, metallic bonds, covalent bonds, dipole-dipole attractions, hydrogen bonds, or London forces. 61. Given chemical formulas for two substances that are about the same size but have different types of attractions between the particles, predict which substance would have the strongest attractions between the particles in the liquid and solid state. 159

4 160 Chapter 8 & 9 Covalent Bonding and Molecular Structures 62. Given chemical formulas for two substances with the same type of attraction between the particles but different sizes, predict which substance would have the stronger attractions between the particles in the liquid and solid state. 63. Given descriptions of two compounds that differ in the percentages of their structures that are polar and nonpolar (hydrophilic and hydrophobic), predict their relative solubilities in water and in hexane. 64. Explain why amphetamine can pass through the blood brain barrier more easily than epinephrine. 65. Explain why methamphetamine is often converted to methamphetamine hydrochloride. 66. Convert between the definition and the term for the following words or phrases. Chapters 8 and 9 Glossary Molecular orbital A volume that contains a high percentage of the electron charge for an electron in a molecule or a volume within which an electron in a molecule has a high probability of being found. Bonding molecular orbital ormed from in-phase interaction of two atomic orbitals. This leads to an increase in negative charge between two nuclei where the atomic orbitals overlap and leads to more +/- attraction between the negative charge generated by the electrons and the nuclei. Antibonding molecular orbital ormed from out-of-phase interaction of two atomic orbitals. This leads to a decrease in negative charge between two nuclei where the atomic orbitals overlap and leads to less +/- attraction between the negative charge generated by the electrons and the nuclei. Valence bond model The model for covalent bonding that includes the assumptions that (1) only the highest-energy (valence) electrons participate in bonding, (2) covalent bonds arise due to the overlap of atomic orbitals on adjacent atoms, forming molecular orbitals, and (3) covalent bonds often form to pair unpaired electrons. Valence electrons The highest energy s and p electrons for an atom. The electrons that are most important in the formation of chemical bonds. The number of valence electrons for the atoms of an element is equal to the element s A-group number on the periodic table. Electron-dot symbol A representation of an atom that consists of its elemental symbol surrounded by dots representing its valence electrons. Lone pair Two electrons that are not involved in the covalent bonds between atoms but are important for explaining the arrangement of atoms in molecules. They are represented by pairs of dots in Lewis structures. Lewis structure A representation of a molecule that consists of the elemental symbol for each atom in the molecule, lines to show covalent bonds, and pairs of dots to indicate lone pairs. ybrid orbital An atomic orbital formed from a blend of atomic orbitals, e.g. four sp 3 hybrid orbitals are formed from the blend of one s orbital and three p orbitals. Bond angle The angle formed by straight lines (representing bonds) connecting the nuclei of three adjacent atoms.

5 161 Electron group geometry A description of the arrangement of all the electron groups around a central atom in a molecule or polyatomic ion, including the lone pairs. Molecular geometry The description of the arrangement of all the atoms around a central atom in a molecule or polyatomic ion. This description does not consider lone pairs. Isomers Compounds that have the same molecular formula but different molecular structures. ormal charge A tool for evaluating Lewis structures, calculated from the formula: A-group number minus number of bonds minus number of electrons in lone pairs. Resonance structures Two or more Lewis structures for a single molecule or polyatomic ion that differ in the positions of lone pairs and multiple bonds but not in the positions of the atoms in the structure. It is as if the molecule or ion were able to shift from one of these structures to another by shifting pairs of electrons from one position to another. Resonance The hypothetical switching from one resonance structure to another. Resonance hybrid A structure that represents the average of the resonance structures for a molecule or polyatomic ion. Delocalized electrons Electrons that are shared among three or more atoms. Delocalized pi system A system of overlapping p orbitals used to describe the bonding in resonance hybrids. Cis isomer A structure that has like groups on different carbons (which are linked by a double bond) and on the same side of the double bond. Trans isomer A structure that has like groups on different carbons (which are linked by a double bond) and on different sides of the double bond. Triglyceride A compound with three hydrocarbon groups attached to a three carbon backbone by ester functional groups. Saturated fat A triglyceride with single bonds between all of the carbon atoms. Unsaturated fat A triglyceride that has one or more carbon carbon double bonds. ydrogenation A process by which hydrogen is added to an unsaturated triglyceride to convert double bonds to single bonds. This can be done by combining the unsaturated triglyceride with hydrogen gas and a platinum catalyst. Noncovalent interaction All forces of attraction between particles other than covalent, ionic, or metallic bonds. Intermolecular forces Attractions between molecules. Dipole-dipole attraction The intermolecular attraction between the partial negative end of one polar molecule and the partial positive end of another polar molecule. ydrogen bond The intermolecular attraction between a nitrogen, oxygen, or fluorine atom of one molecule and a hydrogen atom bonded to a nitrogen, oxygen, or fluorine atom in another molecule. London dispersion forces, London forces, or dispersion forces The attractions produced between molecules by instantaneous and induced dipoles. ydrophilic ( water loving ): A polar molecule or ion (or a portion of a molecule or polyatomic ion) that is attracted to water. ydrophobic ( water fearing ): A nonpolar molecule (or a portion of a molecule or polyatomic ion) that is not expected to mix with water. Lipid bilayer Cell membrane.

6 162 Chapter 8 & 9 Covalent Bonding and Molecular Structures Molecular Orbital Theory The goal of molecular orbital theory is to describe molecules in a similar way to how we describe atoms, that is, in terms of orbitals, orbital diagrams, and electron configurations. or example, to give you a glimpse at where we are headed, the following are orbital diagrams for O 2 and O. π π σ π π 2s σ 2s σ 2s O 2 O Each line in the molecular orbital diagram represents a molecular orbital, which is the volume within which a high percentage of the negative charge generated by the electron is found. The molecular orbital volume encompasses the whole molecule. We assume that the electrons would fill the molecular orbitals of molecules like electrons fill atomic orbitals in atoms. The molecular orbitals are filled in a way that yields the lowest potential energy for the molecule. The maximum number of electrons in each molecular orbital is two. (We follow the Pauli exclusion principle.) Orbitals of equal energy are half filled with parallel spin before they begin to pair up. (We follow und s Rule.) Before we continue with a description of a model used to generate molecular orbital diagrams, let s get a review of light and electron waves and how two waves can interact. The wave description of light describes the effect that the light has on the space around it. This effect is to generate an oscillating electric and magnetic fields. These fields can vary in intensity, which is reflected in varying brightness of light.

7 163 Source Wavelength, λ, the distance between two peaks Electric field, perpendicular to magnetic field igure 8,9.1 Electromagnetic Waves Magnetic field, perpendicular to electric field Radiant energy The wave description of the electron describes the variation in the intensity of negative charge generated by the electron. Nucleus, about the diameter of the atom The negative charge is most intense at the nucleus and decreases in intensity with distance outward. igure 8,9.2 Orbital Light waves can interact in-phase, which leads to an increase in the intensity of the light (brighter) and out-of-phase, which leads to a decrease in the intensity of the light (less bright). igure 8,9.3 Wave Interference In-phase Brighter Dark Out-of-phase Electron waves can also interact in-phase and out-of-phase. In-phase interaction leads to an increase in the intensity of the negative charge. Out-of-phase interaction leads to a decrease in the intensity of the negative charge. One common approximation that allows us to generate molecular orbital diagrams for some small diatomic molecules is called the Linear Combination of Atomic Orbitals (LCAO) approach. The following assumptions lie at the core of this model. Molecular orbitals are formed from the overlap of atomic orbitals. Only atomic orbitals of about the same energy interact to a significant degree. When two atomic orbitals overlap, they interact in two extreme ways to form two molecular orbitals, a bonding molecular orbital and an antibonding molecular orbital.

8 164 Chapter 8 & 9 Covalent Bonding and Molecular Structures or example, our model assumes that two atomic orbitals can overlap in two extreme ways to form two molecular orbitals. One of the ways the atomic orbitals interact is in-phase, which leads to wave enhancement similar to the enhancement of two in-phase light waves. Where the atomic orbitals overlap, the in phase interaction leads to an increase in the intensity of the negative charge in the region where they overlap. This creates an increase in negative charge between the nuclei and an increase in the plus minus attraction between the electron charge and the nuclei for the atoms in the bond. The greater attraction leads to greater stability and lower potential energy. Because electrons in the molecular orbital are lower potential energy than in separate atomic orbitals, energy would be required to shift the electrons back into the orbitals of separate atoms. This keeps the atoms together in the molecule, so we call this orbital a bonding molecular orbital. The molecular orbital formed is symmetrical about the axis of the bond. Symmetrical molecular orbitals are called sigma, σ, molecular orbitals. The symbol σ is used to describe the bonding molecular orbital formed from two atomic orbitals. The second way that two atomic orbitals interact is out of phase. Where the atomic orbitals overlap, the out-of phase interaction leads to a decrease in the intensity of the negative charge. This creates a decrease in negative charge between the nuclei and a decrease in the plus-minus attraction between the electron charge and the nuclei for the atoms in the bond. The lesser attraction leads to lower stability and higher potential energy. The electrons are more stable in the atomic orbitals of separate atoms, so electrons in this type of molecular orbital destabilize the bond between atoms. We call this type of molecular orbital an antibonding molecular orbital. The molecular orbital formed is symmetrical about the axis of the bond, so it is a sigma molecular orbital with a symbol of σ*. The asterisk indicates an antibonding molecular orbital. The following diagram shows the bonding and antibonding molecular orbitals formed from the interaction of two atomic orbitals.

9 165 Decreased negative charge between the nuclei leads to decreased attractions between the negative charge from the electrons and the positively charged nuclei. This makes the sigma anitbonding molecular orbital higher potential energy than the separate atomic orbitals. σ out-ofphase in-phase nuclei σ Increased negative charge between the nuclei leads to increased attractions between the negative charge from the electrons and the positively charged nuclei. This makes the sigma bonding molecular orbital lower potential energy than the separate atomic orbitals. σ igure 8,9.4 ormation of Sigma Molecular Orbitals When two larger atoms combine to form a diatomic molecules (such as O 2, 2, or Ne 2 ), more atomic orbitals interact. The LCAO approximation assumes that only the atomic orbitals of about the same energy interact. or O 2, 2, and Ne 2, the orbital energies are different enough in energy so only orbitals of the same energy interact to a significant degree. Like for hydrogen, the from one atom overlaps the from the other atom to form a σ bonding molecular orbital and a σ* antibonding molecular orbital. The shapes would be similar to those formed from the orbitals for hydrogen. The 2s atomic orbital from one atom overlaps the 2s atomic orbital from the other atom to form a σ 2s bonding molecular orbital and a σ* 2s antibonding molecular orbital. The shapes of these molecular orbitals would be similar to those for the σ and σ* molecular orbitals. Both σ 2s and σ* 2s molecular orbitals are higher energy and larger than the σ and σ* molecular orbitals. The p atomic orbitals of the two atoms can interact in two different ways, parallel or end-on. The molecular orbitals are different for each type of interaction. The end-on interaction between two x atomic orbitals yields sigma molecular orbitals, which are symmetrical about the axis of the bond.

10 166 Chapter 8 & 9 Covalent Bonding and Molecular Structures igure 8,9.5 ormation of Sigma Molecular Orbitals out-ofphase nuclei σ* in-phase x x σ The two y atomic orbitals overlap in parallel and form two pi molecular orbitals. Pi molecular orbitals are asymmetrical about the axis of the bond. igure 8,9.6 ormation of Pi Molecular Orbitals out-ofphase in-phase nuclei π * y y π There is less overlap between parallel p orbitals than between two p orbitals overlapping end-on. The z - z overlap generates another pair of π and π* molecular orbitals. The z - z overlap is similar to the y - y overlap. To visualize this overlap, picture all of the orbitals in the image above rotated, so that the axes that run through each atomic and molecular orbitals are perpendicular to the paper. The molecular orbitals formed have the same potential energy as the molecular orbitals formed from the y - y overlap. There is less overlap for the parallel atomic orbitals. When the interaction is in-phase, less overlap leads to less electron charge enhancement between the nuclei. This leads to less electron charge between the nuclei for the pi bonding molecular orbital than for the sigma bonding molecular orbital. Less electron character between the nuclei means less plus-minus attraction, less stabilization, and higher potential energy for the pi bonding molecular orbital compared to the sigma bonding molecular orbital. When the interaction is out-of-phase, less overlap leads to less shift of electron charge from between the nuclei. This leads to more electron charge between the nuclei for the pi antibonding molecular orbital than for the sigma antibonding molecular orbital. More electron charge between the nuclei means more plus-minus attraction and lower potential energy for the pi antibonding molecular orbital compared to the sigma antibonding molecular orbital.

11 167 The expected molecular orbital diagram from the overlap of, 2s, and atomic orbitals is as follows. We will use this diagram to describe O 2, 2, Ne 2, CO, and NO. π π π π σ 2s σ 2s σ We assume that the electrons would fill the molecular orbitals of molecules like electrons fill atomic orbitals in atoms. We use the following procedure when drawing molecular orbital diagrams. Determine the number of electrons in the molecule. We get the number of electrons per atom from their atomic number on the periodic table. (Remember to determine the total number of electrons, not just the valence electrons.) ill the molecular orbitals from bottom to top until all the electrons are added. Describe the electrons with arrows. Put two arrows in each molecular orbital, with the first arrow pointing up and the second pointing down. Orbitals of equal energy are half filled with parallel spin before they begin to pair up. We describe the stability of the molecule with bond order. bond order = 1/2 (#e in bonding MO s #e in antibonding MO s) We use bond orders to predict the stability of molecules. If the bond order for a molecule is equal to zero, the molecule is unstable. A bond order of greater than zero suggests a stable molecule. The higher the bond order is, the more stable the bond. We can use the molecular orbital diagram to predict whether the molecule is paramagnetic or diamagnetic. If all the electrons are paired, the molecule is diamagnetic. If one or more electrons are unpaired, the molecule is paramagnetic.

12 168 Chapter 8 & 9 Covalent Bonding and Molecular Structures EXAMPLES: 1. The molecular orbital diagram for a diatomic hydrogen molecule, 2, is σ The bond order is 1. Bond Order = 1 2 (2-0) = 1 The bond order above zero suggests that 2 is stable. Because there are no unpaired electrons, 2 is diamagnetic. 2. The molecular orbital diagram for a diatomic helium molecule, e 2, shows the following. σ The bond order is 0 for e 2. Bond Order = 1 2 (2-2) = 0 The zero bond order for e 2 suggests that e 2 is unstable. If e 2 did form, it would be diamagnetic. 3. The molecular orbital diagram for a diatomic oxygen molecule, O 2, is π π π π σ 2s σ 2s σ O 2 has a bond order of 2. Bond Order = 1 2 (10-6) = 2 The bond order of two suggests that the oxygen molecule is stable. The two unpaired electrons show that O 2 is paramagnetic.

13 The molecular orbital diagram for a diatomic fluorine molecule, 2, is π π σ π π 2s σ 2s σ 2 has a bond order of 1. Bond Order = 1 2 (10-8) = 1 The bond order of one suggests that the fluorine molecule is stable. Because all of the electrons are paired 2 is diamagnetic. 5. The molecular orbital diagram for a diatomic neon molecule, Ne 2, is π π σ π π 2s σ 2s σ Ne 2 has a bond order of 0. Bond Order = 1 2 (10-10) = 0 The bond order zero for Ne 2 suggests that Ne 2 is unstable. If Ne 2 did form, it would be diamagnetic.

14 170 Chapter 8 & 9 Covalent Bonding and Molecular Structures We can describe diatomic molecules composed of atoms of different elements in a similar way. The bond between the carbon and oxygen in carbon monoxide is very strong despite what looks like a strange and perhaps unstable Lewis Structure. C O The plus formal charge on the more electronegative oxygen and the minus formal charge on the less electronegative carbon would suggest instability. The molecular orbital diagram predicts CO to be very stable with a bond order of three. π π π π σ 2s σ 2s σ We predict the nitrogen monoxide molecule to be unstable according to the Lewis approach to bonding. N O The unpaired electron and the lack of an octet of electrons around nitrogen would suggest an unstable molecule. NO is actually quite stable. The molecular orbital diagram predicts this by showing the molecule to have a bond order of 2.5. π π π π σ 2s σ 2s σ

15 171 Electrons in a sigma molecular orbital lead to a sigma bond. igure 8,9.7 Bonding in 2 nuclei x x σ Electrons in a sigma molecular orbital lead to a sigma bond. igure 8,9.8 Bonding in 2 nuclei σ ormation of sp 3 hybrid orbitals One 2s and three atomic orbitals blend 2s x y z to form four equivalent sp 3 hybrid orbitals, which are arranged in a tetrahedral geometry with angles of sp sp sp 3 sp 3 sp 3 sp 3 sp sp 3 igure 8,9.9 ormation of sp 3 ybrid Atomic Orbitals

16 172 Chapter 8 & 9 Covalent Bonding and Molecular Structures Overlap of atomic orbitals in methane, C C igure 8,9.10 Bonding in C 4 C sp C sp 3 C sp our C- sigma bonds due to sp 3 - overlap sp 3 C ormation of sp 2 hybrid orbitals One 2s and two atomic orbitals blend 2s x y z to form three equivalent sp 2 hybrid orbitals, which are arranged in a trigonal planar geometry with angles of 120. This leaves one atomic orbital unhybridized. The axis through the orbital is from the axes through the sp 2 hybrid orbitals. 120 sp 2 sp sp 2 2 sp 2 sp 2 igure 8,9.11 ormation of sp 2 ybrid Atomic Orbitals sp 2

17 173 igure 8,9.12 Bonding in C 2 4 Overlap of atomic orbitals in ethene (ethylene), C 2 4 sp 2 sp 2 sp 2 sp 2 sp 2 sp 2 C (with stylized orbitals) C (with stylized orbitals) sp 2 sp 2 sp 2 sp2 sp 2 sp 2 One C-C sigma bond due to sp 2 -sp 2 overlap and one C-C pi bond due to - overlap 4 C- sigma bonds due to sp 2 - overlap

18 174 Chapter 8 & 9 Covalent Bonding and Molecular Structures igure 8,9.13 Bonding in C 2 2 Overlap of atomic orbitals in ethyne (acetylene), C 2 2 x y sp z y sp sp sp C (with stylized orbitals) C (with stylized orbitals) C sp sp C z C z C y y C C One C-C sigma bond due to sp-sp overlap One C-C pi bond due to z - z overlap One C-C pi bond due to y - y overlap sp sp C C Two C- sigma bonds due to sp- overlap igure 8,9.14 Bonding in N 3 Overlap of atomic orbitals in ammonia, N 3 lone pair N 107 The lone pair is more repulsive than the bond groups, so the hydrogen atoms are pushed closer together than we would predict from sp 3 hybridization. The angle is about 107 instead of N sp 3 sp 3 N sp 3 Three N- sigma bonds due to sp 3 - overlap N 107

19 175 igure 8,9.15 Bonding in N 4 + Overlap of atomic orbitals in ammonium, N N N sp N sp 3 N sp sp 3 N our N- sigma bonds due to sp 3 - overlap igure 8,9.16 Bonding in 2 O Overlap of atomic orbitals in water, 2 O lone pairs O 105 The lone pairs are more repulsive than the bond groups, so the hydrogen atoms are pushed closer together than we would predict from sp 3 hybridization. The angle is about 105 instead of Two O- sigma bonds due to sp 3 - overlap sp 3 O 105 O sp 3 105

20 176 Chapter 8 & 9 Covalent Bonding and Molecular Structures igure 8,9.17 Bonding in B 3 Overlap of atomic orbitals in boron trifluoride, B 3 sp 3 sp 2 sp 3 sp 2 B sp sp 3 Three B- sigma bonds due to sp 2 -sp 3 overlap ormation of sp 3 d hybrid orbitals One s, three p, and one d atomic orbitals blend 3s 3p 3p 3p 3d z 2 3d xz 3d yz 3d xy 3d x 2- y2 igure 8,9.18 ormation of sp 3 d ybrid Atomic Orbitals to form five equivalent sp 3 d hybrid orbitals, which are arranged in a trigonal bipyramid geometry with angles of, 120, and 180. This leaves four d atomic orbitals unhybridized. axial 180 sp 3 d 120 sp 3 d 180 axial 120 equatorial sp 3 d 120 sp 3 d 120 sp 3 d 120 equatorial equatorial

21 177 igure 8,9.19 Bonding in P 5 Overlap of atomic orbitals in phosphorus pentafluoride, P 5 sp 3 d-sp 3 overlap 120 P 180 sp 3 d-sp 3 overlap sp 3 d-sp 3 overlap 120 sp 3 d-sp 3 overlap sp 3 d-sp 3 overlap ive P- sigma bonds due to sp 3 d-sp 3 overlap igure 8,9.20 Bonding in S 4 Overlap of atomic orbitals in sulfur tetrafluoride, S 4 sp 3 d-sp 3 overlap sp 3 d-sp 3 overlap 120 S 180 lone pair sp 3 d-sp 3 overlap sp 3 d-sp 3 overlap our S- sigma bonds due to sp 3 d-sp 3 overlap

22 178 Chapter 8 & 9 Covalent Bonding and Molecular Structures igure 8,9.21 Bonding in I 3 Overlap of atomic orbitals in iodine trifluoride, I 3 sp 3 d-sp 3 overlap sp 3 d-sp 3 overlap lone pairs I 180 sp 3 d-sp 3 overlap Three I- sigma bonds due to sp 3 d-sp 3 overlap 3 igure 8,9.22 Bonding in Xe 2 Overlap of atomic orbitals in xenon difluoride, Xe 2 sp 3 d-sp 3 overlap 180 lone pairs Xe lone pair sp 3 d-sp 3 overlap Two Xe- sigma bonds due to sp 3 d-sp 3 overlap

23 179 ormation of sp 3 d 2 hybrid orbitals One s, three p, and two d atomic orbitals blend 3s 3p x 3p y 3p z igure 8,9.23 ormation of sp 3 d 2 ybrid Atomic Orbitals 3d z 2 3d xz 3d yz 3d xy 3d x 2- y2 to form six equivalent sp 3 d 2 hybrid orbitals, which are arranged in an octahedral geometry with angles of and 180. This leaves three d atomic orbitals unhybridized. sp 3 d sp 3 d sp 3 d 2 sp 3 d sp 3 d 2 sp 3 d 2 sp 3 d sp 3 d 2 sp 3 d 2 sp 3 d 2 sp 3 d sp 3 d Overlap of atomic orbitals in sulfur hexafluoride, S 6 sp 3 d 2 -sp 3 overlap sp 3 d 2 -sp 3 overlap sp 3 d 2 -sp 3 overlap igure 8,9.24 Bonding in S 6 S sp 3 d 2 -sp 3 overlap sp 3 d 2 -sp 3 overlap sp 3 d 2 -sp 3 overlap Six S- sigma bonds due to sp 3 d 2 -sp 3 overlap

24 180 Chapter 8 & 9 Covalent Bonding and Molecular Structures igure 8,9.25 Bonding in I 5 Overlap of atomic orbitals in iodine pentafluoride, I 5 sp 3 d 2 -sp 3 overlap sp 3 d 2 -sp overlap sp 3 d 2 -sp 3 overlap I lone pair sp 3 d 2 -sp 3 overlap sp 3 d 2 -sp 3 overlap ive I- sigma bonds due to sp 3 d 2 -sp 3 overlap igure 8,9.26 Bonding in Xe 4 Overlap of atomic orbitals in xenon tetrafluoride, Xe 4 sp 3 d 2 -sp 3 overlap lone pair Xe lone pair sp 3 d 2 -sp 3 overlap sp 3 d 2 -sp 3 overlap sp 3 d 2 -sp 3 overlap our Xe- sigma bonds due to sp 3 d 2 -sp 3 overlap

25 181 The following assumptions lie at the core of the valence bond model. Only the highest energy electrons participate in bonding. Covalent bonds usually form to pair unpaired electrons. Covalent bonds arise due to the overlap of atomic orbitals on adjacent atoms. When the columns on the periodic table are numbered by the A group convention, the number of valence electrons in each atom of a representative element is equal to the element s group number on the periodic table. One valence electron 1 3A Number of valence electrons equals the A-group number 4A 5A 6A 7A B C N O P S Cl 8A 2 e 10 Ne 18 Ar As Se Br Te I 36 Kr 54 Xe igure 8,9.27 Valence Electrons Table 8,9.1 Most Common Covalent Bonding Patterns Element requency of pattern Number of bonds Number of lone pairs Example always 1 0 B usually 3 0 B C usually 4 0 C or C or C rarely 3 1 C N, P, & As usually 3 1 N commonly 4 0 N O, S, & Se usually 2 2 O or O commonly 1 3 O rarely 3 1 O, Cl, Br, & I usually 1 3 X

26 182 Chapter 8 & 9 Covalent Bonding and Molecular Structures ormal Charge = Group Number Number of lines Number of dots ormal Charge = (The number of valence electrons necessary to be uncharged) (the number of electrons from bonds assuming equal sharing) (the number of electrons from lone pairs) Sample Study Sheet 8,9.1: Explanation of Bonding Patterns - Valence Bond Model Tip-off: You will be given a bonding pattern for a nonmetal atoms and asked to explain how it is possible in terms of the valence bond model. General Procedure Step 1 Write, Only the highest energy electrons participate in bonding. Step 2 Draw the orbital diagram for the valence electrons. Include the empty d orbitals for the third period nonmetals and below. Step 3 If there is a formal charge, add or subtract electrons. a. Add one electron for -1. b. Subtract one electron for +1. Step 4 If necessary, promote one or more electrons from a pair to an empty orbital to get the number of unpaired electrons equal to the number of bonds to be explained. Step 5 Rewrite the orbital diagram, showing the predicted hybrid orbitals. Step 6 Write, Covalent bonds form in order to pair unpaired electrons. Step 7 Indicate that the unpaired electrons form the bonds. Step 8 Indicate that the paired electrons are lone pairs. Exercise 8,9.1 - Explanation of Bonding Patterns Explain the following bonding patterns a. C 4 bonds, no lone pairs, and no formal charge b. N 4 bonds, no lone pairs, and a +1 formal charge c. O 1 bond, 3 lone pairs, and a 1 formal charge d. S 4 bonds, 1 lone pair, and no formal charge e. I 5 bonds, 1 lone pair, and no formal charge

27 183 Tip-off: You will be given a Lewis structure, and you will be asked the following. What is the hybridization for each atom in the structure? Write a description of the formation of each bond in terms of the overlap of atomic orbitals. Describe the electron group geometry around each atom that has two or more atoms attached to it. Draw the geometric sketch of the molecule, including bond angles. Describe the molecular geometry around each atom that has two or more atoms attached to it. General Steps for predicting hybridization for each atom in a structure. Step 1 Count the number of electron groups around each atom. An electron group is any one of the following. single bond double or triple bond (counts as one group) lone pair Step 2 Apply the following guidelines. See Table 8, group - No hybridization (ydrogen atoms are the only atoms that have only one group around them in a Lewis structure. The orbital is the only occupied atomic orbital for hydrogen atoms. Because there is no other orbital with which to blend, there is no hybridization. ydrogen atoms use the atomic orbital to form bonds.) 2 groups - sp hybridization 3 groups - sp 2 hybridization 4 groups - sp 3 hybridization 5 groups - sp 3 d hybridization 6 groups - sp 3 d 2 hybridization General Steps for describing the bonds in terms of the overlap of atomic orbitals. Step 1 Describe single bonds that do not involve hydrogen as due to the overlapping of two hybrid orbitals, one from each atom. Step 2 Describe single bonds that involve hydrogen as due to the overlapping a hybrid orbital with a orbital for the hydrogen atom. Step 3 Describe double bonds as one bond forming from the overlap of two hybrid orbitals, one from each atom. Describe the second bond as due to the overlap of two p orbitals, one from each atom. Step 4 Describe triple bonds as one bond forming from the overlap of two hybrid orbitals, one from each atom. Describe each of the other two bonds as due to the overlap of two p orbitals, one from each atom. Sample Study Sheet 8,9.2: rom Lewis Structures

28 184 Chapter 8 & 9 Covalent Bonding and Molecular Structures To determine the electron group geometry around each atom that is attached to two or more atoms, use the number of electron groups around each atom and the guidelines found on Table 8,9.2. Use one or more of the following geometric sketches shown on Table 8,9.2 for the geometric sketch of you molecule. To describe the molecular geometry around each atom that has two or more atoms attached to it, count the number of bond groups and apply the guidelines found on Table 8,9.2. (Single, double, and triple bonds all count as one bond group.) Note that if all of the electron groups attached to the atom are bond groups (no lone pairs), the molecular geometry is the same as the electron group geometry.

29 185 Table 8,9.2 ybridization and Geometry e ybrid General Geometric Sketch e group bond bond lone molecular groups Orbitals geometry angles groups pairs geometry 2 sp 180 linear linear 3 sp trigonal planar trigonal planar 4 sp tetrahedral bent 5 sp 3 d trigonal bipyramid, 120, tetrahedral 3 1 trigonal pyramid 2 2 bent sp 3 d octahedral, trigonal bipyramid 4 1 see-saw 3 2 T-shaped 2 3 linear octahedral 5 1 square pyramid 4 2 square planar

30 186 Chapter 8 & 9 Covalent Bonding and Molecular Structures Table 8,9.3 Common Covalent Bonding Patterns for the Nonmetal Elements Element Number of bonds Number of lone pairs ormal charge always C usually rarely N usually commonly O usually commonly rarely always P and As usually commonly commonly S and Se usually commonly commonly commonly possibly halogens except usually Xe possibly possibly possibly usually none rarely rarely Table 8,9.4 The Most Common Bonding Pattern for Each Nonmetal Element Group Number of Covalent Bonds Number of Lone Pairs 4A (C) 4 0 5A (N, P, As) 3 1 6A (O, S, Se) 2 2 7A (halogens) 1 3 8A (noble gases) none 4

31 187 O O O O N O O N O O N O Tip-off: You are asked to draw a resonance hybrid. General Steps Draw the skeletal structure with solid lines for the bonds that are found in all of the resonance forms. Where there is sometimes a multiple bond and sometimes not, draw a dotted line. Draw in all of the lone pairs that are found on every one of the resonance forms. (Leave off the lone pairs that are on one or more resonance form but not on all of them.) Put on full formal charges on those atoms that have the formal charge in all of the resonance forms. Put on partial formal charges on the atoms that have formal charges in one or more of the resonance forms but not in all of them. The resonance hybrid for the nitrate polyatomic ion is below. Sample Study Sheet 8,9.3: Resonance ybrids A bond found in at least one but not all the resonance structures O O N O A bond found in all the resonance structures A lone pair found in all the resonance structures igure 8,9.28 Resonance ybrid

32 188 Chapter 8 & 9 Covalent Bonding and Molecular Structures Sample Study Sheet 8,9.4: Resonance and Resonance forms Tip-off - You are asked to draw a Lewis structure for a molecule or a polyatomic ion, and the structure drawn has the following form. X Y Z Z can have more than one lone pair. X and Y can have lone pairs. The X Y bond can be a triple bond. The Y Z bond can be a double bond. Z cannot be with one bond and three lone pairs or O with two bonds and two lone pairs. General Steps: ollow these steps when writing the resonance forms. Shift one of the lone pairs on one adjacent atom down to form a multiple bond. Shift one of the multiple bonds up to form a lone pair. You might find it useful to draw the arrows to indicate the hypothetical shift of electrons. Repeat this process for each adjacent atom with a lone pair. X Y Z X Y Z Sample Study Sheet 8,9.5: Drawing Lewis Structures Tip-off: You are asked to draw a Lewis structure from a chemical formula for a molecule or a polyatomic ion. General Steps (As you become more experienced with the process of drawing Lewis structures, you will see shortcuts that will allow you to bypass some or all of the following steps.) Step 1 Determine the number of valence electrons for the molecule or polyatomic ion. or neutral molecules, the total number of valence electrons is the sum of the valence electrons of each atom. Remember that the number of valence electrons for any one of the representative elements is equal to its group number, using the A group convention for numbering the groups. or example, chlorine, Cl, is in group 7A, so it has seven valence electrons. ydrogen has one valence electron. or polyatomic cations, the total number of valence electrons is the sum of the valence electrons for each atom minus the charge. or polyatomic anions, the total number of valence electrons is the sum of the valence electrons for each atom plus the charge.

33 189 Step 2 Draw a reasonable skeletal structure joining the atoms with single bonds. One or more of the following guidelines might help you with this step. Try to set the atoms in space to gain the most common number of bonds for each atom. Table 8,9.3 lists the most common bonding patterns for the nonmetal elements. Apply the following guidelines in deciding what goes in the center of your structure. ydrogen and fluorine atoms are never in the center. Oxygen atoms are rarely in the center. The element with the fewest atoms in the formula is usually in the center. The atom that is capable of making the most bonds is usually in the center. See Table 8,9.3. The atom with the lowest electronegativity is often in the center. Oxygen atoms rarely bond to other oxygen atoms. The molecular formula often reflects the structural formula. Carbon atoms commonly bond to other carbon atoms. Step 3 Subtract 2 electrons from the total for each of the single bonds (lines) described in Step 2 above. Step 4 Try to distribute the remaining electrons as lone pairs to get eight total electrons around each atom except hydrogen, beryllium, and boron. When an atom has a total of eight electrons around it, we say it has an octet of electrons. Atoms of the noble gases other than helium have an octet of electrons around them. The orbital diagram and electron dot structure for the valence electrons of neon show this. 2s Ne The stability of the octet is reflected in the fact that atoms in reasonable Lewis structures often have this octet. Carbon, nitrogen, oxygen and fluorine always have 8 electrons around them in a reasonable Lewis structure. Beryllium and boron can have less than eight electrons but never more than eight. The nonmetal elements below the second period (P, S, Cl, Se, Br, and I) usually have eight electrons around them but often have more. The metalloids arsenic, As, and tellurium, Te, form covalent bonds in similar ways to phosphorus, P, and sulfur, S, so they usually have eight electrons around them but often have more. ydrogen will always have two electrons total from its one bond.

34 190 Chapter 8 & 9 Covalent Bonding and Molecular Structures Step 5 Do one of the following. If in Step 4 you were able to get eight electrons around each atom other than those of hydrogen, beryllium, and boron and if you used all of the remaining valence electrons, go to Step 6. If you have electrons remaining after each of the atoms other than those of hydrogen, beryllium, and boron have their octet, you can put more than eight electrons around the elements in the third period and below. If you do not have enough electrons to get octets of electrons around each atom (other than hydrogen, beryllium, and boron), convert one lone pair into a multiple bond for each 2 electrons that you are short. If you would need two more electrons to get octets, convert one lone pair in your structure to a second bond between two atoms. If you would need four more electrons to get octets, convert two lone pairs into bonds. This could mean creating a second bond in two different places or creating a triple bond in one place. If you would need six more electrons to get octets, convert three lone pairs into bonds. Etc. Step 6 Check your structure. irst check to see if all of the atoms have their most common bonding pattern. If each atom has its most common bonding pattern, your structure is a reasonable structure. Skip Step 7, and proceed to Step 8. Determine the formal charge for each atom that has other than its most common bonding pattern. (Atoms with their most common bonding pattern always have a zero formal charge.) ormal Charge = Group Number Number of lines Number of dots If all of the formal charges are zero, skip Step 7, and go to Step 8. If you have formal charges, continue with Step 7. Step 7 Try to rearrange your structure to eliminate or at least reduce any formal charges. (This step is unnecessary if there are no formal charges.) One way to eliminate formal charges is to return to Step 2 and try another skeleton. ormal charges can also be eliminated by converting lone pairs into bonds. This process leads to placing more than eight electrons around some atoms. ere is where we hit controversy. Chemists are not in agreement about whether it is better to emphasize the importance of octets in reasonable Lewis structures or better to minimize formal charges.

35 191 Step 8: Once you have a reasonable Lewis structure, consider the possibility of resonance. If resonance is possible, write all of the reasonable resonance structures for the original Lewis structure. Draw the resonance hybrid from the resonance forms. It is possible to generate more than one reasonable Lewis structure for a given formula. When this happens, the following rules will help you decide which is most likely. The more usual or common the bond is, the more stable the structure. See Table 8,9.3. A structure with formal charges is less stable than one without formal charges. + formal charges are more likely on the less electronegative elements and - formal charges are more likely on the more electronegative elements. It is possible to have two Lewis structures for the same molecular formula, which are considered equally reasonable according to the above criteria. The substances represented by the Lewis structures are isomers, molecules that have the same molecular formula but different structural formulas. The dimethyl ether and ethanol are isomers of C 2 6 O. C O C C C O dimethyl ether ethanol

36 192 Chapter 8 & 9 Covalent Bonding and Molecular Structures TE OLLOWING IS A SUMMARY O TE STEPS OR DRAWING LEWIS STRUCTURES ROM CEMICAL ORMULAS. Step 1 Determine the number of valence electrons for the molecule or polyatomic ion. Step 2 Draw a reasonable skeletal structure joining the atoms with single bonds. Step 3 Subtract 2 electrons from the total for each of the single bonds described in Step 2 above. Step 4 Try to distribute the remaining electrons as lone pairs to get eight total electrons around each atom except hydrogen, beryllium, and boron. Step 5 Do one of the following. If in Step 4 you were able to get eight electrons around each atom other than those of hydrogen, beryllium, and boron and if you used all of the remaining valence electrons, go to Step 6. If you have electrons remaining after the atoms other than those of hydrogen, beryllium, and boron have their octet, you can put more than eight electrons around the elements in the third period and below. If you do not have enough electrons to get octets of electrons around each atom (other than those of hydrogen, beryllium, and boron), convert one lone pair into a multiple bond for each 2 electrons that you are short. Step 6 Check your structure. Step 7 Try to rearrange your structure to eliminate or at least reduce any formal charges. Step 8 Once you have a reasonable Lewis structure, consider the possibility of resonance. If resonance is possible, draw the reasonable resonance structures and the resonance hybrid for the structure. Exercise 8,9.2 - Drawing Lewis Structures Draw a reasonable Lewis structure for each of the following. If the structure has resonance, draw all of the reasonable resonance structures and the resonance hybrid. a. C 3 Br h. C 3 7 O b. Cl 3 i. CO 2 c. C 2 O j. C 2 C d. CN k. CO 2 e. C 3 CCl 2 l. N 2 COC 3 f. C 2 6 O m. SO 2 4 g. C n. PO 4 3

Chapter 5 Chemical Compounds. An Introduction to Chemistry by Mark Bishop

Chapter 5 Chemical Compounds. An Introduction to Chemistry by Mark Bishop Chapter 5 Chemical Compounds An Introduction to Chemistry by Mark Bishop Chapter Map Elements, Compounds, and Mixtures Element: A substance that cannot be chemically converted into simpler substances;

More information

CHAPTER NOTES CHAPTER 16. Covalent Bonding

CHAPTER NOTES CHAPTER 16. Covalent Bonding CHAPTER NOTES CHAPTER 16 Covalent Bonding Goals : To gain an understanding of : NOTES: 1. Valence electron and electron dot notation. 2. Stable electron configurations. 3. Covalent bonding. 4. Polarity

More information

Chapter 10: Chemical Bonding II: Molecular Shapes; VSEPR, Valence Bond and Molecular Orbital Theories

Chapter 10: Chemical Bonding II: Molecular Shapes; VSEPR, Valence Bond and Molecular Orbital Theories C h e m i s t r y 1 A : C h a p t e r 1 0 P a g e 1 Chapter 10: Chemical Bonding II: Molecular Shapes; VSEPR, Valence Bond and Molecular Orbital Theories Homework: Read Chapter 10: Work out sample/practice

More information

Effect of unshared pairs on molecular geometry

Effect of unshared pairs on molecular geometry Chapter 7 covalent bonding Introduction Lewis dot structures are misleading, for example, could easily represent that the electrons are in a fixed position between the 2 nuclei. The more correct designation

More information

SHAPES OF MOLECULES (VSEPR MODEL)

SHAPES OF MOLECULES (VSEPR MODEL) 1 SAPES MLEULES (VSEPR MDEL) Valence Shell Electron-Pair Repulsion model - Electron pairs surrounding atom spread out as to minimize repulsion. - Electron pairs can be bonding pairs (including multiple

More information

Valence shell electrons repel each other Valence shell electrons are arranged geometrically around the central atom to

Valence shell electrons repel each other Valence shell electrons are arranged geometrically around the central atom to Molecular Geometry (Valence Shell Electron Pair Repulsion -VSEPR) & Hybridization of Atomic Orbitals (Valance Bond Theory) Chapter 10 Valence Shell Electron Pair Repulsion (VSEPR) Valence shell electrons

More information

The Lewis electron dot structures below indicate the valence electrons for elements in Groups 1-2 and Groups 13-18

The Lewis electron dot structures below indicate the valence electrons for elements in Groups 1-2 and Groups 13-18 AP EMISTRY APTER REVIEW APTER 7: VALENT BNDING You should understand the nature of the covalent bond. You should be able to draw the Lewis electron-dot structure for any atom, molecule, or polyatomic ion.

More information

Chemical Bonding and Molecular Structure (Chapter 10)

Chemical Bonding and Molecular Structure (Chapter 10) Chemical Bonding and Molecular Structure (Chapter 10) Molecular Structure 1. General Summary -- Structure and Bonding Concepts Electronic Configuration of Atoms Octet Rule Lewis Electron Dot ormula of

More information

A pure covalent bond is an equal sharing of shared electron pair(s) in a bond. A polar covalent bond is an unequal sharing.

A pure covalent bond is an equal sharing of shared electron pair(s) in a bond. A polar covalent bond is an unequal sharing. CHAPTER EIGHT BNDING: GENERAL CNCEPT or Review 1. Electronegativity is the ability of an atom in a molecule to attract electrons to itself. Electronegativity is a bonding term. Electron affinity is the

More information

Chemical Bonds. a. Duet Rule: 2 electrons needed to satisfy valence shell. i. What follows this rule? Hydrogen and Helium

Chemical Bonds. a. Duet Rule: 2 electrons needed to satisfy valence shell. i. What follows this rule? Hydrogen and Helium Chemical Bonds 1. Important points about Lewis Dot: a. Duet Rule: 2 electrons needed to satisfy valence shell. i. What follows this rule? Hydrogen and Helium b. Octet Rule: 8 electrons needed to satisfy

More information

Lewis Structures. Molecular Shape. VSEPR Model (Valence Shell Electron Pair Repulsion Theory)

Lewis Structures. Molecular Shape. VSEPR Model (Valence Shell Electron Pair Repulsion Theory) Lewis Structures Molecular Shape VSEPR Model (Valence Shell Electron Pair Repulsion Theory) PART 1: Ionic Compounds Complete the table of Part 1 by writing: The Lewis dot structures for each metallic and

More information

11 Chemical Bonds: The Formation of Compounds from Atoms. Chapter Outline. Periodic Trends in Atomic Properties. Periodic Trends in Atomic Properties

11 Chemical Bonds: The Formation of Compounds from Atoms. Chapter Outline. Periodic Trends in Atomic Properties. Periodic Trends in Atomic Properties 11 Chemical Bonds The Formation of Compounds from Atoms Chapter Outline 11.1 11.2 Lewis Structures of Atoms 11.3 The Ionic Bond Transfer of Electrons from One Atom to Another 11.4 Predicting Formulas of

More information

CHAPTER 9 COVALENT BONDING: ORBITALS. Questions

CHAPTER 9 COVALENT BONDING: ORBITALS. Questions APTER 9 VALET BDIG: RBITALS Questions 9. In hybrid orbital theory, some or all of the valence atomic orbitals of the central atom in a molecule are mixed together to form hybrid orbitals; these hybrid

More information

Chapter 11. Chemical Bonds: The Formation of Compounds from Atoms

Chapter 11. Chemical Bonds: The Formation of Compounds from Atoms Chapter 11 Chemical Bonds: The Formation of Compounds from Atoms 1 11.1 Periodic Trends in atomic properties 11.1 Periodic Trends in atomic properties design of periodic table is based on observing properties

More information

MULTIPLE CHOICE. Choose the one alternative that best completes the statement or answers the question.

MULTIPLE CHOICE. Choose the one alternative that best completes the statement or answers the question. MULTIPLE CHOICE. Choose the one alternative that best completes the statement or answers the question. 1) A chemical bond formed between two identical atoms is a(an) bond. A) covalent B) ionic C) molecular

More information

Chapter 12 Molecular Structure

Chapter 12 Molecular Structure hapter 12 Molecular Structure t s Monday morning, and you d like a cup of coffee, but when you try cranking up the stove to reheat yesterday s brew, nothing happens. Apparently, the city gas line has sprung

More information

Lab Manual Supplement

Lab Manual Supplement Objectives 1. Learn about the structures of covalent compounds and polyatomic ions. 2. Draw Lewis structures based on valence electrons and the octet rule. 3. Construct 3-dimensional models of molecules

More information

CHEM 2323 Unit 1 General Chemistry Review

CHEM 2323 Unit 1 General Chemistry Review EM 2323 Unit 1 General hemistry Review I. Atoms A. The Structure of the Atom B. Electron onfigurations. Lewis Dot Structures II. Bonding A. Electronegativity B. Ionic Bonds. ovalent Bonds D. Bond Polarity

More information

Covalent Bonding And Molecular Geometry

Covalent Bonding And Molecular Geometry ovalent Bonding And Molecular Geometry Questions: 1.ow can the valence electrons of an atom be represented? 2.ow do atoms achieve an octet? 3.ow are electrons shared in a molecule? 4.ow can the geometries

More information

CHEM 110 Exam 2 - Practice Test 1 - Solutions

CHEM 110 Exam 2 - Practice Test 1 - Solutions CHEM 110 Exam 2 - Practice Test 1 - Solutions 1D 1 has a triple bond. 2 has a double bond. 3 and 4 have single bonds. The stronger the bond, the shorter the length. 2A A 1:1 ratio means there must be the

More information

5. Structure, Geometry, and Polarity of Molecules

5. Structure, Geometry, and Polarity of Molecules 5. Structure, Geometry, and Polarity of Molecules What you will accomplish in this experiment This experiment will give you an opportunity to draw Lewis structures of covalent compounds, then use those

More information

Chemistry 4th Edition McMurry/Fay

Chemistry 4th Edition McMurry/Fay 7 Chapter Covalent Bonding Chemistry 4th Edition McMurry/Fay Dr. Paul Charlesworth Michigan Technological University The Covalent Bond 01 Covalent bonds are formed by sharing at least one pair of electrons.

More information

Section 1: Organic Structure and Bonding

Section 1: Organic Structure and Bonding Section 1: Organic Structure and Bonding What is Organic Chemistry? Compounds containing only carbon and hydrogen, also known as, are the simplest form of organic compounds. Examples: C C C C C C Atoms

More information

Chapter 9 Molecular Geometry and Bonding Theories

Chapter 9 Molecular Geometry and Bonding Theories Chapter 9 Molecular Geometry and Bonding Theories 1. or a molecule with the formula AB 2 the molecular shape is. (a). linear or trigonal planar (b). linear or bent (c). linear or T-shaped (d). T-shaped

More information

Chapter 9 - Covalent Bonding: Orbitals

Chapter 9 - Covalent Bonding: Orbitals Chapter 9 - Covalent Bonding: Orbitals 9.1 Hybridization and the Localized Electron Model A. Hybridization 1. The mixing of two or more atomic orbitals of similar energies on the same atom to produce new

More information

Chapter 10 Molecular Geometry and Chemical Bonding Theory

Chapter 10 Molecular Geometry and Chemical Bonding Theory Chem 1: Chapter 10 Page 1 Chapter 10 Molecular Geometry and Chemical Bonding Theory I) VSEPR Model Valence-Shell Electron-Pair Repulsion Model A) Model predicts Predicts electron arrangement and molecular

More information

Molecular Geometry and VSEPR We gratefully acknowledge Portland Community College for the use of this experiment.

Molecular Geometry and VSEPR We gratefully acknowledge Portland Community College for the use of this experiment. Molecular and VSEPR We gratefully acknowledge Portland ommunity ollege for the use of this experiment. Objectives To construct molecular models for covalently bonded atoms in molecules and polyatomic ions

More information

Chapter 8: Bonding General Concepts. Valence Electrons. Representative Elements & Lewis Dot Structures

Chapter 8: Bonding General Concepts. Valence Electrons. Representative Elements & Lewis Dot Structures Chapter 8: Bonding General Concepts Valence Electrons 8.1 Chemical Bond Formation 8.2 Covalent Bonding (Lewis Dot Structures) 8.3 Charge Distribution in Covalent Compounds 8.4 Resonance 8.5 Molecular Shapes

More information

A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together is called a(n)

A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together is called a(n) Chemistry I ATOMIC BONDING PRACTICE QUIZ Mr. Scott Select the best answer. 1) A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together is

More information

8/19/2011. Periodic Trends and Lewis Dot Structures. Review PERIODIC Table

8/19/2011. Periodic Trends and Lewis Dot Structures. Review PERIODIC Table Periodic Trends and Lewis Dot Structures Chapter 11 Review PERIODIC Table Recall, Mendeleev and Meyer organized the ordering the periodic table based on a combination of three components: 1. Atomic Number

More information

Bonding Models. Bonding Models (Lewis) Bonding Models (Lewis) Resonance Structures. Section 2 (Chapter 3, M&T) Chemical Bonding

Bonding Models. Bonding Models (Lewis) Bonding Models (Lewis) Resonance Structures. Section 2 (Chapter 3, M&T) Chemical Bonding Bonding Models Section (Chapter, M&T) Chemical Bonding We will look at three models of bonding: Lewis model Valence Bond model M theory Bonding Models (Lewis) Bonding Models (Lewis) Lewis model of bonding

More information

Chapter 7. Comparing Ionic and Covalent Bonds. Ionic Bonds. Types of Bonds. Quick Review of Bond Types. Covalent Bonds

Chapter 7. Comparing Ionic and Covalent Bonds. Ionic Bonds. Types of Bonds. Quick Review of Bond Types. Covalent Bonds Comparing Ionic and Covalent Bonds Chapter 7 Covalent Bonds and Molecular Structure Intermolecular forces (much weaker than bonds) must be broken Ionic bonds must be broken 1 Ionic Bonds Covalent Bonds

More information

Valence Bond Theory - Description

Valence Bond Theory - Description Bonding and Molecular Structure - PART 2 - Valence Bond Theory and Hybridization 1. Understand and be able to describe the Valence Bond Theory description of covalent bond formation. 2. Understand and

More information

Chapter 9. Chemical reactivity of molecules depends on the nature of the bonds between the atoms as well on its 3D structure

Chapter 9. Chemical reactivity of molecules depends on the nature of the bonds between the atoms as well on its 3D structure Chapter 9 Molecular Geometry & Bonding Theories I) Molecular Geometry (Shapes) Chemical reactivity of molecules depends on the nature of the bonds between the atoms as well on its 3D structure Molecular

More information

AP Chemistry A. Allan Chapter 8 Notes - Bonding: General Concepts

AP Chemistry A. Allan Chapter 8 Notes - Bonding: General Concepts AP Chemistry A. Allan Chapter 8 Notes - Bonding: General Concepts 8.1 Types of Chemical Bonds A. Ionic Bonding 1. Electrons are transferred 2. Metals react with nonmetals 3. Ions paired have lower energy

More information

Lewis Structure Exercise

Lewis Structure Exercise Lewis Structure Exercise A Lewis structure shows how the valence electrons are arranged and indicates the bonding between atoms in a molecule. We represent the elements by their symbols. The shared electron

More information

Chemistry Workbook 2: Problems For Exam 2

Chemistry Workbook 2: Problems For Exam 2 Chem 1A Dr. White Updated /5/1 1 Chemistry Workbook 2: Problems For Exam 2 Section 2-1: Covalent Bonding 1. On a potential energy diagram, the most stable state has the highest/lowest potential energy.

More information

Covalent Bonds. A group of atoms held together by covalent bonds is called a molecule.

Covalent Bonds. A group of atoms held together by covalent bonds is called a molecule. Covalent Bonds The bond formed when atoms share electrons is called a covalent bond. (Unlike ionic bonds, which involve the complete transfer of electrons). A group of atoms held together by covalent bonds

More information

Laboratory 11: Molecular Compounds and Lewis Structures

Laboratory 11: Molecular Compounds and Lewis Structures Introduction Laboratory 11: Molecular Compounds and Lewis Structures Molecular compounds are formed by sharing electrons between non-metal atoms. A useful theory for understanding the formation of molecular

More information

3.4 Covalent Bonds and Lewis Structures

3.4 Covalent Bonds and Lewis Structures 3.4 Covalent Bonds and Lewis Structures The Lewis Model of Chemical Bonding In 1916 G. N. Lewis proposed that atoms combine in order to achieve a more stable electron configuration. Maximum stability results

More information

10. Geometry Hybridization Unhybridized p atomic orbitals

10. Geometry Hybridization Unhybridized p atomic orbitals APTER 14 VALET BDIG: RBITALS The Localized Electron Model and ybrid rbitals 9. The valence orbitals of the nonmetals are the s and p orbitals. The lobes of the p orbitals are 90E and 180E apart from each

More information

EXPERIMENT 14: COMPARISONS OF THE SHAPES OF MOLECULES AND IONS USING MODELS

EXPERIMENT 14: COMPARISONS OF THE SHAPES OF MOLECULES AND IONS USING MODELS EXPERIMENT 14: CMPARISNS F TE SAPES F MLECULES AND INS USING MDELS PURPSE Models of various molecules and ions will be constructed and their shapes and geometries will be compared. BACKGRUND LEWIS STRUCTURES

More information

CHAPTER 6 REVIEW. Chemical Bonding. Answer the following questions in the space provided.

CHAPTER 6 REVIEW. Chemical Bonding. Answer the following questions in the space provided. Name Date lass APTER 6 REVIEW hemical Bonding SETIN 1 SRT ANSWER Answer the following questions in the space provided. 1. a A chemical bond between atoms results from the attraction between the valence

More information

George Mason University General Chemistry 211 Chapter 10 The Shapes (Geometry) of Molecules

George Mason University General Chemistry 211 Chapter 10 The Shapes (Geometry) of Molecules Acknowledgements George Mason University General Chemistry 211 Chapter 10 The Shapes (Geometry) of Molecules Course Text Chemistry the Molecular Nature of Matter and Change, 7 th edition, 2011, McGraw-Hill

More information

CHEMISTRY NOTES: Structures, Shapes, Polarity and IMF s

CHEMISTRY NOTES: Structures, Shapes, Polarity and IMF s CHEMISTRY NOTES: Structures, Shapes, Polarity and IMF s DRAWING LEWIS STRUCTURES: RULES 1) Draw the skeleton structure for the molecule. The central atom will generally be the least electronegative element

More information

Illustrating Bonds - Lewis Dot Structures

Illustrating Bonds - Lewis Dot Structures Illustrating Bonds - Lewis Dot Structures Lewis Dot structures are also known as electron dot diagrams These diagrams illustrate valence electrons and subsequent bonding A line shows each shared electron

More information

Topic 4. Chemical bonding and structure

Topic 4. Chemical bonding and structure Topic 4. Chemical bonding and structure There are three types of strong bonds: Ionic Covalent Metallic Some substances contain both covalent and ionic bonding or an intermediate. 4.1 Ionic bonding Ionic

More information

Laboratory 20: Review of Lewis Dot Structures

Laboratory 20: Review of Lewis Dot Structures Introduction The purpose of the laboratory exercise is to review Lewis dot structures and expand on topics discussed in class. Additional topics covered are the general shapes and bond angles of different

More information

Theme 3: Bonding and Molecular Structure. (Chapter 8)

Theme 3: Bonding and Molecular Structure. (Chapter 8) Theme 3: Bonding and Molecular Structure. (Chapter 8) End of Chapter questions: 5, 7, 9, 12, 15, 18, 23, 27, 28, 32, 33, 39, 43, 46, 67, 77 Chemical reaction valence electrons of atoms rearranged (lost,

More information

Lewis Dot Structure Answer Key

Lewis Dot Structure Answer Key Lewis Dot Structure Answer Key 1) Nitrogen is the central atom in each of the following species: N2 N2 - N2 + Nitrogen can also form electron deficient compounds with a single unpaired electron on the

More information

Molecular Structures. Chapter 9 Molecular Structures. Using Molecular Models. Using Molecular Models. C 2 H 6 O structural isomers: .. H C C O..

Molecular Structures. Chapter 9 Molecular Structures. Using Molecular Models. Using Molecular Models. C 2 H 6 O structural isomers: .. H C C O.. John W. Moore onrad L. Stanitski Peter. Jurs http://academic.cengage.com/chemistry/moore hapter 9 Molecular Structures Stephen. oster Mississippi State University Molecular Structures 2 6 structural isomers:

More information

A REVIEW OF GENERAL CHEMISTRY: ELECTRONS, BONDS AND MOLECULAR PROPERTIES

A REVIEW OF GENERAL CHEMISTRY: ELECTRONS, BONDS AND MOLECULAR PROPERTIES A REVIEW OF GENERAL CEMISTRY: ELECTRONS, BONDS AND MOLECULAR PROPERTIES A STUDENT SOULD BE ABLE TO: 1. Draw Lewis (electron dot and line) structural formulas for simple compounds and ions from molecular

More information

Unit 8: Drawing Molecules

Unit 8: Drawing Molecules Unit 8: Drawing Molecules bjectives Topic 1: Lewis Dot Diagrams & Ionic Bonding 1. Draw a Lewis dot diagram of any representative element. 2. Draw a Lewis dot diagram of any ionic compound. A Lewis structure

More information

Lewis Structures & the VSEPR Model

Lewis Structures & the VSEPR Model Lewis Structures & the VSEPR Model A Directed Learning Activity for Hartnell College Chemistry 1 Funded by the Title V STEM Grant #P031S090007 through Hartnell College For information contact lyee@hartnell.edu

More information

EXPERIMENT 17 : Lewis Dot Structure / VSEPR Theory

EXPERIMENT 17 : Lewis Dot Structure / VSEPR Theory EXPERIMENT 17 : Lewis Dot Structure / VSEPR Theory Materials: Molecular Model Kit INTRODUCTION Although it has recently become possible to image molecules and even atoms using a high-resolution microscope,

More information

Valence Electrons. core and CHAPTER 9. Introduction. Bonds - Attractive forces that hold atoms together in compounds

Valence Electrons. core and CHAPTER 9. Introduction. Bonds - Attractive forces that hold atoms together in compounds Structure and Molecular Bonding CAPTER 9 1 Introduction Bonds - Attractive forces that hold atoms together in compounds Valence Electrons - The electrons involved in bonding are in the outermost (valence)

More information

Drawing Lewis Structures

Drawing Lewis Structures Drawing Lewis Structures 1. Add up all of the valence electrons for the atoms involved in bonding 2. Write the symbols for the elements and show connectivity with single bonds (2 electrons shared). a.

More information

GRADE 11 PHYSICAL SCIENCES SESSION 3: CHEMICAL BONDING. Key Concepts. X-planation

GRADE 11 PHYSICAL SCIENCES SESSION 3: CHEMICAL BONDING. Key Concepts. X-planation GRADE 11 PHYSICAL SCIENCES SESSION 3: CHEMICAL BONDING Key Concepts In this session we will focus on summarising what you need to know about: Bonding Covalent bonding Electronegativity in covalent bonding

More information

Bonding Practice Problems

Bonding Practice Problems NAME 1. When compared to H 2 S, H 2 O has a higher 8. Given the Lewis electron-dot diagram: boiling point because H 2 O contains stronger metallic bonds covalent bonds ionic bonds hydrogen bonds 2. Which

More information

Bonding & Molecular Shape Ron Robertson

Bonding & Molecular Shape Ron Robertson Bonding & Molecular Shape Ron Robertson r2 n:\files\courses\1110-20\2010 possible slides for web\00bondingtrans.doc The Nature of Bonding Types 1. Ionic 2. Covalent 3. Metallic 4. Coordinate covalent Driving

More information

CHAPTER 6 Chemical Bonding

CHAPTER 6 Chemical Bonding CHAPTER 6 Chemical Bonding SECTION 1 Introduction to Chemical Bonding OBJECTIVES 1. Define Chemical bond. 2. Explain why most atoms form chemical bonds. 3. Describe ionic and covalent bonding.. 4. Explain

More information

Models. Localized Electron Model. Localized Electron Model. Fundamental Properties of Models

Models. Localized Electron Model. Localized Electron Model. Fundamental Properties of Models Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world. Fundamental Properties of Models A model does not equal reality. Models

More information

Lewis Dot Symbols for Representative Elements

Lewis Dot Symbols for Representative Elements CHEM 110 - Section 4 Guest Instructor: Prof. Elizabeth Gaillard Fall 2011 Lewis Dot Symbols for Representative Elements Principal Types of Chemical Bonds: Ionic and Covalent Ionic bond - a transfer of

More information

Chapter 8. Chemical Bonding. Introduction. Molecular and Ionic Compounds. Chapter 8 Topics. Ionic and Covalent. Ionic and Covalent

Chapter 8. Chemical Bonding. Introduction. Molecular and Ionic Compounds. Chapter 8 Topics. Ionic and Covalent. Ionic and Covalent Introduction Chapter 8 Chemical Bonding How and why to atoms come together (bond) to form compounds? Why do different compounds have such different properties? What do molecules look like in 3 dimensions?

More information

UNIT TEST Atomic & Molecular Structure. Name: Date:

UNIT TEST Atomic & Molecular Structure. Name: Date: SCH4U UNIT TEST Atomic & Molecular Structure Name: _ Date: Part A - Multiple Choice Identify the letter of the choice that best completes the statement or answers the question. 1. Who postulated that electrons

More information

Assignment 9 Solutions. Chapter 8, #8.32, 36, 38, 42, 54, 56, 72, 100, 102, Chapter 10, #10.24, 40, 55, 63. Number of e in Valence Shell

Assignment 9 Solutions. Chapter 8, #8.32, 36, 38, 42, 54, 56, 72, 100, 102, Chapter 10, #10.24, 40, 55, 63. Number of e in Valence Shell Assignment 9 Solutions Chapter 8, #8.32, 36, 38, 42, 54, 56, 72, 100, 102, Chapter 10, #10.24, 40, 55, 63. 8.32. Collect and Organize Of B 3+, I, Ca 2+, and Pb 2+ we are to identify which have a complete

More information

EXPERIMENT 9 Dot Structures and Geometries of Molecules

EXPERIMENT 9 Dot Structures and Geometries of Molecules EXPERIMENT 9 Dot Structures and Geometries of Molecules INTRODUCTION Lewis dot structures are our first tier in drawing molecules and representing bonds between the atoms. The method was first published

More information

Prepare well for this topic!!

Prepare well for this topic!! 1 Bonding There is almost always a chemical bonding question. Since this is such an important topic, you should be prepared for any and all bonding questions. Prepare well for this topic!! Some suggestions

More information

Bonding Web Practice. Trupia

Bonding Web Practice. Trupia 1. If the electronegativity difference between the elements in compound NaX is 2.1, what is element X? bromine fluorine chlorine oxygen 2. Which bond has the greatest degree of ionic character? H Cl Cl

More information

Copyright Page 1

Copyright Page 1 1 Introduction to Bonding: Two-Dimensional Lewis tructures Key Terms: abbreviated electron configuration combines the inert, noble core electrons (abbreviated via a noble element) with the outermost or

More information

Periodic Table Trends

Periodic Table Trends Name Date Period Periodic Table Trends (Ionization Energy and Electronegativity) Ionization Energy The required to an electron from a gaseous atom or ion. Period Trend: As the atomic number increases,

More information

Unit 3: Quantum Theory, Periodicity and Chemical Bonding

Unit 3: Quantum Theory, Periodicity and Chemical Bonding Selected Honour Chemistry Assignment Answers pg. 9 Unit 3: Quantum Theory, Periodicity and Chemical Bonding Chapter 7: The Electronic Structure of Atoms (pg. 240 to 241) 48. The shape of an s-orbital is

More information

Chemistry 105, Chapter 7 Exercises

Chemistry 105, Chapter 7 Exercises hemistry 15, hapter 7 Exercises Types of Bonds 1. Using the periodic table classify the bonds in the following compounds as ionic or covalent. If covalent, classify the bond as polar or not. Mg2 4 i2 a(3)2

More information

Molecular Compounds. Chapter 5. Covalent (Molecular) Compounds

Molecular Compounds. Chapter 5. Covalent (Molecular) Compounds Molecular Compounds Chapter 5 Covalent (Molecular) Compounds Covalent Compound- a compound that contains atoms that are held together by covalent bonds Covalent Bond- the force of attraction between atoms

More information

Homework. Chapter 9. Chapter 9. Sigma Bond Formation by Orbital Overlap. Valence Bond Theory VALENCE BOND THEORY

Homework. Chapter 9. Chapter 9. Sigma Bond Formation by Orbital Overlap. Valence Bond Theory VALENCE BOND THEORY Homework Chapter 9 Chapter 9 11, 21, 25, 27, 29, 31, 35, 39, 45, 51, 65 Bonding and Molecular Structure: Orbital Hybridization and Molecular Chapter 9 Broken into two different sections discussing two

More information

2. Atoms with very similar electronegativity values are expected to form

2. Atoms with very similar electronegativity values are expected to form AP hemistry Practice Test #6 hapter 8 and 9 1. Which of the following statements is incorrect? a. Ionic bonding results from the transfer of electrons from one atom to another. b. Dipole moments result

More information

CHAPTER 10 THE SHAPES OF MOLECULES

CHAPTER 10 THE SHAPES OF MOLECULES ATER 10 TE AE MLEULE 10.1 To be the central atom in a compound, the atom must be able to simultaneously bond to at least two other atoms. e,, and cannot serve as central atoms in a Lewis structure. elium

More information

Question 4.2: Write Lewis dot symbols for atoms of the following elements: Mg, Na, B, O, N, Br.

Question 4.2: Write Lewis dot symbols for atoms of the following elements: Mg, Na, B, O, N, Br. Question 4.1: Explain the formation of a chemical bond. A chemical bond is defined as an attractive force that holds the constituents (atoms, ions etc.) together in a chemical species. Various theories

More information

CHEM 200/202 Exam 3 November 15, 2014

CHEM 200/202 Exam 3 November 15, 2014 HM 200/202 xam 3 ovember 15, 2014 ame: Lab Section #: Please mark your answers on the scantron sheet using a #2 pencil and also mark your answers on the exam itself. Mark Test rom A on your scantron. 1.

More information

Name: Class: Date: 3) The bond angles marked a, b, and c in the molecule below are about,, and, respectively.

Name: Class: Date: 3) The bond angles marked a, b, and c in the molecule below are about,, and, respectively. Name: Class: Date: Unit 9 Practice Multiple Choice Identify the choice that best completes the statement or answers the question. 1) The basis of the VSEPR model of molecular bonding is. A) regions of

More information

EXPERIMENT - 1. Molecular Geometry- Lewis Dot structures

EXPERIMENT - 1. Molecular Geometry- Lewis Dot structures EXPERIMENT - 1 Molecular Geometry- Lewis Dot structures INTRODUCTION Although it has recently become possible to image molecules and even atoms using a high-resolution microscope, most of our information

More information

Principal energy levels are divided into sublevels following a distinctive pattern, shown in Table 5.1 below.

Principal energy levels are divided into sublevels following a distinctive pattern, shown in Table 5.1 below. 56 Chapter 5: Electron Configuration, Lewis Dot Structure, and Molecular Shape Electron configuration. The outermost electrons surrounding an atom (the valence electrons) are responsible for the number

More information

Worked solutions to student book questions Chapter 7 Covalent molecules, networks and layers

Worked solutions to student book questions Chapter 7 Covalent molecules, networks and layers E1. a Give the electronic configuration for an atom of beryllium. b How many electrons are in the outer shell of an atom of beryllium in the molecule BeH 2? AE1. a 1s 2 2s 2 b 4 E2. The noble gases helium

More information

Lewis Structures. X } Lone Pair (unshared pair) } Localized Electron Model. Valence Bond Theory. Bonding electron (unpaired electron)

Lewis Structures. X } Lone Pair (unshared pair) } Localized Electron Model. Valence Bond Theory. Bonding electron (unpaired electron) G. N. Lewis 1875-1946 Lewis Structures (The Localized Electron Model) Localized Electron Model Using electron-dot symbols, G. N. Lewis developed the Localized Electron Model of chemical bonding (1916)

More information

VSEPR Model. The Valence-Shell Electron Pair Repulsion Model. Predicting Molecular Geometry

VSEPR Model. The Valence-Shell Electron Pair Repulsion Model. Predicting Molecular Geometry VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions. The Valence-Shell Electron Pair Repulsion Model The valence-shell electron pair repulsion

More information

Sample Exercise 8.1 Magnitudes of Lattice Energies

Sample Exercise 8.1 Magnitudes of Lattice Energies Sample Exercise 8.1 Magnitudes of Lattice Energies Without consulting Table 8.2, arrange the ionic compounds NaF, CsI, and CaO in order of increasing lattice energy. Analyze From the formulas for three

More information

Unit 28 Molecular Geometry

Unit 28 Molecular Geometry Unit 28 Molecular Geometry There are two concepts in the study of molecular geometry. One is called the Valence Shell Electron Pair Repulsion (VSEPR) model. The other is electron orbital hybridization.

More information

Which substance contains positive ions immersed in a sea of mobile electrons? A) O2(s) B) Cu(s) C) CuO(s) D) SiO2(s)

Which substance contains positive ions immersed in a sea of mobile electrons? A) O2(s) B) Cu(s) C) CuO(s) D) SiO2(s) BONDING MIDTERM REVIEW 7546-1 - Page 1 1) Which substance contains positive ions immersed in a sea of mobile electrons? A) O2(s) B) Cu(s) C) CuO(s) D) SiO2(s) 2) The bond between hydrogen and oxygen in

More information

Copyright 2014 Edmentum - All rights reserved. Chemistry Chemical bonding, molecular structure and Gases Blizzard Bag 2014-2015

Copyright 2014 Edmentum - All rights reserved. Chemistry Chemical bonding, molecular structure and Gases Blizzard Bag 2014-2015 Copyright 2014 Edmentum - All rights reserved. Chemistry Chemical bonding, molecular structure and Gases Blizzard Bag 2014-2015 1. Which of the following is a unit of pressure? A. newton-meters per second

More information

CHAPTER 10 THE SHAPES OF MOLECULES

CHAPTER 10 THE SHAPES OF MOLECULES ATER 10 TE AE MLEULE EMIAL ETI BED READIG RBLEM B10.1 lan: Examine the Lewis structure, noting the number of regions of electron density around the carbon and nitrogen atoms in the two resonance structures.

More information

CHAPTER 10 THE SHAPES OF MOLECULES

CHAPTER 10 THE SHAPES OF MOLECULES ATER 10 TE AE MLEULE 10.1 To be the central atom in a compound, the atom must be able to simultaneously bond to at least two other atoms. e,, and cannot serve as central atoms in a Lewis structure. elium

More information

Ionic and Covalent Bonds

Ionic and Covalent Bonds Ionic and Covalent Bonds Ionic Bonds Transfer of Electrons When metals bond with nonmetals, electrons are from the metal to the nonmetal The becomes a cation and the becomes an anion. The between the cation

More information

CHEM 101 Exam 4. Page 1

CHEM 101 Exam 4. Page 1 CEM 101 Exam 4 Form 1 (White) November 30, 2001 Page 1 Section This exam consists of 8 pages. When the exam begins make sure you have one of each. Print your name at the top of each page now. Show your

More information

ch9 and 10 practice test

ch9 and 10 practice test 1. Which of the following covalent bonds is the most polar (highest percent ionic character)? A. Al I B. Si I C. Al Cl D. Si Cl E. Si P 2. What is the hybridization of the central atom in ClO 3? A. sp

More information

3) Of the following, radiation has the shortest wavelength. A) X-ray B) radio C) microwave D) ultraviolet E) infrared Answer: A

3) Of the following, radiation has the shortest wavelength. A) X-ray B) radio C) microwave D) ultraviolet E) infrared Answer: A 1) Which one of the following is correct? A) ν + λ = c B) ν λ = c C) ν = cλ D) λ = c ν E) νλ = c Answer: E 2) The wavelength of light emitted from a traffic light having a frequency of 5.75 1014 Hz is.

More information

COVALENT BONDING. [MH5; Chapter 7]

COVALENT BONDING. [MH5; Chapter 7] COVALENT BONDING [MH5; Chapter 7] Covalent bonds occur when electrons are equally shared between two atoms. The electrons are not always equally shared by both atoms; these bonds are said to be polar covalent.

More information

Packet 4: Bonding. Play song: (One of Mrs. Stampfel s favorite songs)

Packet 4: Bonding. Play song:  (One of Mrs. Stampfel s favorite songs) Most atoms are not Packet 4: Bonding Atoms will, or share electrons in order to achieve a stable. Octet means that the atom has in its level. If an atom achieves a stable octet it will have the same electron

More information

Health Science Chemistry I CHEM-1180 Experiment No. 15 Molecular Models (Revised 05/22/2015)

Health Science Chemistry I CHEM-1180 Experiment No. 15 Molecular Models (Revised 05/22/2015) (Revised 05/22/2015) Introduction In the early 1900s, the chemist G. N. Lewis proposed that bonds between atoms consist of two electrons apiece and that most atoms are able to accommodate eight electrons

More information

Chemical Bonds: A Preview Chapter 9 Section 1.1 Forces called chemical bonds hold atoms together in molecules and keep ions in place in solid ionic

Chemical Bonds: A Preview Chapter 9 Section 1.1 Forces called chemical bonds hold atoms together in molecules and keep ions in place in solid ionic Chemical Bonds: A Preview Chapter 9 Section 1.1 Forces called chemical bonds hold atoms together in molecules and keep ions in place in solid ionic compounds. Chemical bonds are electrostatic forces; they

More information