Chapter 13: Phenomena

Chapter 13: Phenomena Phenomena: Scientists measured the bond angles of some common molecules. In the pictures below each line represents a bond that contains 2 electrons. If multiple lines are drawn together these are double or triple bonds and contain 4 and 6 electrons respectively. What patterns do you notice from the data? a) b) Cl c) H N H Cl C Cl H Cl Bond Angles: 107 d) O Bond Angles: 105 g) F B Bond Angles: 120 Bond Angles: 109.5 e) Bond Angles: 120 h) O C O Bond Angle: 117 H O N H N H C C H O + Bond Angles: 109.5 f) Bond Angles: 119 i) Bond Angles: 180 j) S H Bond Angles: 180

Chapter 13 Bonding: General Concepts Big Idea: Bonds are formed from the attraction between oppositely charged ions or by sharing electrons. Only the valence electrons participate in bonding. The shape of the molecules maximize the distance between areas of high electron density. o Types of Bonding o Electronegativity o Lewis Structures o Strength/Length of Covalent Bonds o Shapes of Molecules (VSEPR) o Polar Molecules 2

Types of Bonding Ionic Bonds: Formed when a lower energy can be achieved by the complete transfer of one or more electrons from the atoms of one element to those of another; the compound is then held together by electrostatic attraction between the ions. Covalent Bonds: Formed when the lowest energy structure can be achieved by sharing electrons. 3

Types of Bonding Ionic Bonds: Tend to be between a metal and a non metal. Metals: Usually lose their electrons. Nonmetals: Usually accept additional electrons. Note: In general, atoms gain or lose electrons until they have the same number of electrons as the nearest noble gas 4

Types of Bonding What ions do atoms form? Element S K I Electron Configuration (Atom) Gain/Lose Electrons Ion Formed Electron Configuration (Ion) 5

Types of Bonding Ionic Solids: Assembly of cations and anions stacked together in a regular array. Note: Ionic solids are example of crystalline arrays in which the overall charge on an ionic solid is neutral. Ionic Compounds are represented with formula units (lowest ratio of types of atoms in the compound). Ionic Compound + - CD - + 2CD - + 3CD Covalent Compound + - + - + - AB 2 2AB 2 3AB 2 6

Types of Bonding Steps to calculate the energy needed to form an ionic bond: Step 1: Standard states to gaseous single atom state Na(s) Na(g) ½F 2 (g) F(g) 97 kj mol 80. kj mol Step 2: Both atoms have to form ions Na(g) Na + (g) + e - (g) F(g) + e - (g) F - (g) 494 kj mol -323 kj mol (ionization energy) ((-)electron affinity) Step 3: The ions need to come together to form a crystal (Lattice Energy) Na + (g) + F - (g) NaF(s) -923 kj mol Total Reaction Na(s) + ½F 2 (g) NaF(s) 97 kj + 90. kj kj kj kj + 494 + 323 + 923 mol mol mol mol mol Note: When energy is released, the sign is negative because no work is needed to make the reaction happen. = 575 kj mol 7

Types of Bonding What is holding ionic solids together? Coulombic Potential Energy E P,12 = Z 1e Z2e 4πε r12 +Z Z e2 = = Z2 e 2 4πε d 4πε d Z 1 & Z 2 = charge of ions The total potential energy is the sum of all the potential energies E P = 1 4πε Z2 e 2 d +Z2 e 2 2d Z2 e 2 3d +Z2 e 2 4d = Z2 e 2 4πε d 1 1 2 +1 3 1 4 + Note: ln 2 = 1 1 2 + 1 3 1 4 + E P = ln 2 Z2 e 2 4πε d Need to multiply by 2 to account for the other half of the line. E P = 2ln 2 Z2 e 2 4πε d 8

Types of Bonding Once neighboring ions come into contact they start to repel each other. E P e d Τ d d=0 e dτd = 1 d>1 e dτd = 1 max repulsion repulsion decreases Note: d* is a constant that is commonly taken to be 34.5 pm The potential energy of an ionic solid is a combination of the favorable Coulombic interaction of the ions and the unfavorable exponential increase which results when the atoms touch. The ideal bond length occurs at the minimum potential energy. Energy Minimum Occurs: E p,min = N A Z 1Z2 4πε d 1 d d A 9

Types of Bonding Covalent Bond: A pair of electrons shared between two atoms (occurs between two non metals) Note: In covalent bond formation, atoms go as far as possible toward completing their octets by sharing electron pairs. 10

Electronegativity Electronegativity (χ): The ability of an atom to attract electrons to itself when it is part of a compound Note: The atom with higher electronegativity has a stronger attractive power on electrons and pulls the electrons away from the atom with the lower electronegativity. 11

Electronegativity Difference in Electronegativity Type of Bond > 1.8 Mostly Ionic 0.4-1.8 Polar Covalent < 0.4 Mostly Covalent 0 Non-polar Covalent The dividing line between ionic and covalent bonds is hazy. 12 NaCl MgCl 2 AlCl 3 SiCl 4 PCl 3 S 2 Cl 2 Cl 2 ionic polar-covalent covalent

Electronegativity What makes covalent bonds partly ionic? Electric Dipole: A positive charge next to an equal but opposite negative charge. Electric Dipole Moment (μ): The magnitude of the electric dipole [units debye (D)]. Note: The dipole moment associated with H-Cl is about 1.1 D. Polar Covalent Bond: A covalent bond between atoms that have partial electric charges. 13

Electronegativity What makes ionic bonds partly covalent? Polarizability (α): The ease with which the electron cloud of a molecule can be distorted. As the cation s positive charge pulls on the anion s negative electrons, the spherical electron cloud of the anion becomes distorted in the direction of the cation. This causes the bond to have covalent bond properties. Note: The larger the anion the easier it is to distort the electron cloud. 14

Lewis Structures Lewis Symbols: The chemical symbol of an element, with a dot for each valence electron. Step 1: Determine the number of valence e - from electron configuration. Step 2: Place 1 dot around the element for each valence electron. 15

Lewis Structures Drawing Ionic Lewis Structures Step 1: Determine the electron configuration of the elements in the compound. Step 2: Determine the electron configuration of the ions that the elements form. Step 3: Draw the Lewis symbols for the ions. Do not forget to include charge. The cations should have no electrons around them and the anions should have 8 electrons around them. Step 4: Organize the Lewis symbols such that cations are next to anions. 16

Lewis Structures General Rules (Covalent Lewis Structures) All valence electrons of the atoms in the Lewis structures must be shown. Generally, electrons are paired. Except for odd electron molecules such as NO and NO 2. Generally, each atom has 8 electrons in its valence shell with the exception of H which only needs 2 valance electrons. Multiple bonds (double and triple bonds) can be formed. Show atoms by their chemical symbols (ex. H) Show covalent bonds by lines (ex. F F) Show lone pairs of electrons by pairs of dots (ex. :) 17

Lewis Structures 18 Drawing Covalent Lewis Structures (for structures that obey octet rule) Step 1: Count the number of valence electrons on each atom; for ions adjust the number of electrons to account for the charge. Step 2: Calculate the number of electrons that are needed to fill each atom s octet (or duplet, in the case of H). Step 3: Calculate the number of bonds: # Bonds = Wanted e Step 2 Valence e Step 1. 2 Step 4: Calculate the number of electrons left over: # e = Valence e 2 # Bonds. Step 5: Place bonds/electrons around elements so that octets/duplet are satisfied.

Lewis Structures Tips When Drawing Lewis Structures How to pick central atom: Choose the central atom to be the atom with the lowest ionization energy (atom closest to the lowest left hand corner of the periodic table) Arrange the atoms symmetrically around the central atom Examples: SO 2 would be arranged OSO not SOO Note: In simple formulas the central atom is often written first, followed by the atoms attached to it. Note: Acids are an exception to the rule because H is written first in acids. 19

Lewis Structures What is the Lewis Structure for SCN -? 1(S) 1(C) 1(N) 1(e - ) Total Valence e - 1(6) 1(4) 1(5) 1 16 Wanted e - 1(8) 1(8) 1(8) 24 # bonds = wanted e valence e 2 = 24 16 2 = 4 # e = valence e 2 # bonds = 16 2 4 = 8 Which structure is most likely? _ 20

Lewis Structures Formal Charge: The electric charge of an atom in a molecule assigned on the assumption that the bonding is nonpolar covalent. Formal Charge = Valence e - e - Surrounding Atom Note: The formal charge on neutral molecules must add up to zero. Note: The formal charge on ions must add up to the charge on the ion. Generally, compounds with the lowest formal charges possible (charges closest to 0) are favored. 21

Lewis Structures Formal charge and oxidation numbers both give us information about the number of electrons around an atom in a compound. Formal Charge Exaggerates Covalent Character Assumes that the electrons are shared equally by all atoms. 0 0 0 22 Oxidation Number Exaggerates Ionic Character Assumes that octets are complete filled and all electrons must only belong to one atom. -2 4-2 2-4+ O C 2- O

Lewis Structures Resonance: A blend of Lewis structures into a single composite hybrid structure. Resonance Hybrid: The composite structure that results from a resonance. Delocalized Electrons: Electrons that are spread over several atoms in a molecule. 23

Lewis Structures What is the Lewis Structure for PO 4 3-? 1(P) 4(O) 1(e - ) Total Valence e - 1(5) 4(6) 1(3) 32 Wanted e - 1(8) 4(8) 40 # bonds = wanted e valence e 2 = 40 32 2 = 4 # e = valence e 2 # bonds = 40 2 4 = 32 3-24

Lewis Structures When the central atom in a molecule has empty d-orbitals, it may be able to accommodate 10, 12, or even more electrons, this is referred to as an expanded valence shell. Note: This only applies to nonmetal atoms in Period 3 and later Size also plays a role in how many atoms can fit around a given molecule. Examples: PCl 5 Known to exist NCl 5 Not known to exist (N is too small for the to fit 5 Cl atoms around it) Note: On homework problems only expand octets if there is no other way to accommodate electrons or if the problem tells you to minimize the formal charge. 25

Strength/Length of Covalent Bonds Enthalpy Change ( H): The amount of heat evolved or absorbed in a reaction carried out at constant pressure Dissociation Energy (D): The energy required to separate bonded atoms kj Avg. Bond Energies mol H-H 432 O-H 467 H-F 565 O-O 146 H-Cl 427 F-F 154 26 C-H 413 Cl-Cl 239 C-C 347 O=O 495 C-N 305 C=O* 745 C-O 358 N N 941 N-H 391 C C 839 N-N 160 C N 891 *C=O (CO 2 ) = 799

Strength/Length of Covalent Bonds Double bonds are not twice as strong as a single bonds. Triple bonds are not three times as strong as a single bond. 27

Strength/Length of Covalent Bonds Student Question Which of the following has the longest carbon-oxygen bond? Hint: You must draw the Lewis structures. a) CO b) CO 3 2- c) CO 2 d) CH 3 OH 28

Shapes of Molecules (VSEPR) VSEPR (Valence-Shell Electron-Pair Repulsion Model): Extends Lewis s theory of bonding to account for molecular shapes by adding rules that account for bond angle. Rule 1: Regions of high electron concentration (bonds and lone pairs on the central atom) repel one another and to minimize their repulsion, these regions move as far apart as possible while maintaining the same distance from the central atom. Rule 2: There is no distinction between single and multiple bonds: a multiple bond is treated as a single region of high electron concentration. 29

Shapes of Molecules (VSEPR) Rule 3: All regions of high electron density, lone pairs and bonds, are included in a description of the electronic arrangement, but only the positions of atoms are considered when reporting the shape of a molecule (molecular shape). Rule 4: The strengths of repulsion are in the order lone pair lone pair > lone pair atom > atom atom. 30

Shapes of Molecules (VSEPR) Assigning Shape and Bond Angles of Molecules Step 1: Draw the Lewis structure. Step 2: Assign the electronic arrangement around the central atom (linear, trigonal planer, tetrahedral) Note: Electronic arrangement includes all areas of electron density (lone pairs and bonds). Step 3: Identify the molecular shape (linear, bent, trigonal planer, trigonal pyramidal, tetrahedral) Note: Molecular shape includes only bonds. Step 4: Figure out the bond angle (allow for distortion) 31

Shapes of Molecules (VSEPR) Possible Electronic Arrangements 32

Shapes of Molecules (VSEPR) Possible Molecular Arrangements The names of the shapes of simple molecules and their bond angles. Lone pairs of electrons are not shown. 33

Shapes of Molecules (VSEPR) 2 Areas of Electron Density 3 Areas of Electron Density No Lone Pairs Linear No Lone Pairs Trigonal Planar 1 Lone Pair Bent 4 Areas of Electron Density No Lone Pairs Tetrahedral 1 Lone Pair Trigonal Pyramidal 2 Lone Pairs Bent 34

Shapes of Molecules (VSEPR) 5 Areas of Electron Density No Lone Pairs Trigonal Bipyramidal 1 Lone Pair Seesaw 2 Lone Pairs T-Shaped 3 Lone Pair Linear 35

Shapes of Molecules (VSEPR) 6 Areas of Electron Density No Lone Pairs Octahedral 1 Lone Pair Square Pyramidal 2 Lone Pairs Square Planer 3 Lone Pairs T-Shaped 4 Lone Pair Linear 36

Shapes of Molecules (VSEPR) Step 1: Draw Lewis Structure CH 4 1(C) 4(H) Total Valence e - Wanted e - Determine number of bonds # bonds = wanted valence 2 Determine number of electrons # e = valence 2 bonds Step 2: Determine Electronic Shape Step 3: Determine Molecular Shape Step 4: Determine Angle 37

Shapes of Molecules (VSEPR) Step 1: Draw Lewis Structure (Obeys Octet Rule) SO 3 2-1(S) 3(O) 2(e - ) Total Valence e - Wanted e - Determine number of bonds # bonds = wanted valence 2 Determine number of electrons # e = valence 2 bonds Step 2: Determine Electronic Shape Step 3: Determine Molecular Shape Step 4: Determine Angle 38

Shapes of Molecules (VSEPR) Student Question What is the most likely shape of ICl 4-? Helpful Hint: Make sure that your formal charges are minimized. a) Octahedral b) Trigonal Planar c) Seesaw d) Tetrahedral e) None of the Above 39

Polar Molecules Student Question Is PCl 4- polar or nonpolar? Helpful Hint: Make sure that your formal charges are minimized. a) Polar b) Nonpolar 40

Take Away From Chapter 13 Big Idea: Bonds are formed from the attraction between oppositely charged ions or by sharing electrons. Only the valence electrons participate in bonding. The shape of the molecules maximize the distance between areas of high electron density. Types of Bonding Ionic Bonds (metal/non metal) Be able to write electron configuration of ions. (30,31,32, 33,34) Be able to predict size of ions.(27, 28, 29) Be able to predict formula unit ionic compound.(37) Covalent Bonds (non metal/non metal) 41 Numbers correspond to end of chapter questions.

42 Take Away From Chapter 13 Electronegativity Know the general electronegativity trend. (17) Know that covalent bonds can have ionic because of dipole moments Be able to identify the most polar bond. (18,22,23) Know that ionic bonds can have covalent character because of polarizablity Lewis Structures Be able to draw Lewis symbols (atoms). Be able to draw Lewis structures of ionic compounds. Be able to draw Lewis structures of covalent compounds. (63,64) Know how to calculate formal charges.(85) Identification of most likely Lewis structure. Know when multiple resonance structures are possible for a compound. (66,67,71) Know when atoms can expand their octets (group 3 and greater). (86) Numbers correspond to end of chapter questions.

Take Away From Chapter 13 Strength/Length of Covalent Bonds Know how to calculate ΔH from bond dissociation energies. H = σ D broken σ D formed (47,49,52) Know how to estimate the length of bonds. Triple < Double < Single (79,80,84) Shape of Molecules (VSEPR) (96,99,100) Know how to determine electronic shape. Linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral. Know how to determine molecular shape. Linear, angular, trigonal planar, trigonal pyramidal, T- shaped, tetrahedral, seesaw, square planar, trigonal bipyramidal, square pyramidal, or octahedral. Know how to determine bond angles. 43 Numbers correspond to end of chapter questions.

Take Away From Chapter 13 Polar Molecules Be able to determine if a molecule is polar or non polar. (101,102) 44 Numbers correspond to end of chapter questions.