GENERAL BONDING REVIEW

Transcription

1 GENERAL BONDING REVIEW Chapter 8 November 2, 2016

2 Questions to Consider 1. What is meant by the term chemical bond? 2. Why do atoms bond with each other to form compounds? 3. How do atoms bond with each other to form compounds?

3 A Chemical Bond No simple, and yet complete, way to define this. Forces that hold groups of atoms together and make them function as a unit. A bond will form if the energy of the aggregate is lower than that of the separated atoms.

4 Localized Electron Models A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Electron pairs are assumed to be localized on a particular atom or in the space between two atoms: Lone pairs pairs of electrons localized on an atom Bonding pairs pairs of electrons found in the space between the atoms

5 What Models Do We Use? 1. Lewis Dot Structures 2. VSEPR Structures & Geometry

6 Model 1: Lewis Structures Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration.

7 Review: Valence Electrons 1A Number of valence electrons is equal to the Group number. 8A 2A 3A 4A 5A 6A 7A

8 Electron Distribution Electron distribution is depicted with Lewis electron dot structures Electrons are distributed as: 1. Shared or BOND PAIRS 2. Unshared or LONE PAIRS.

9 Building a Lewis Structure - σbonds EXAMPLE: NH 3 Step 1: Determine the total number of valence electrons Step 2: Determine the central atom the least electronegative atom usually goes in the center Step 3: Draw bonds Step 4: Fill in the structure with the remaining valence electrons until everything has an octet

10 Building Lewis Structure Examples Methane Hydrosulfuric Acid Sulfite

11 Building a Lewis Structure - πbonds Sulfur Dioxide Nitrogen Gas NO - (NH 2 ) 2

12 Exceptions to the Octet Rules Odd number of e - ex. NO (more on this later) Boron (pairs each of its 3 à 6 valence) ex. BF 3 usually reacts with lone pairs Phosphorus and Sulfur (expands octet using empty valence d-orbitals) P pairs 5e - giving 10 valence e - S pairs 6e - giving 12 valence e -

13 Exception Rules C, N, O, and F should always be assumed to obey the octet rule. B and Be often have fewer than 8 electrons around them in their compounds. Second-row elements never exceed the octet rule. Third-row and heavier elements often satisfy the octet rule but can exceed the octet rule by using their empty valence d orbitals. When writing the Lewis structure for a molecule, satisfy the octet rule for the atoms first. If electrons remain after the octet rule has been satisfied, then place them on the elements having available d orbitals (elements in Period 3 or beyond).

14 Exceptions to the Rules BeCl 2 XeF 4 SF 4

15 Resonance Structures More than one valid Lewis structure can be written for a particular molecule. EXAMPLES: 1. NO 3-2. CO 3 2-

16 Formal Charges Used to evaluate nonequivalent Lewis structures. Atoms in molecules try to achieve formal charges as close to zero as possible. Any negative formal charges are expected to reside on the most electronegative atoms. EXAMPLE: SCN -

17 Review Problems: 11/08/ Draw the Lewis structure, determine the formal charge, and the VSEPR shape of NOCl. 2. Draw the Lewis structure, determine the formal charge, and the VSEPR shape for (HO)AsO 2.

18 COVALENT BONDING RECAP November 8, 2016

19 Types of Covalent Bonds The type of bond atoms form depends on electronegativity electronegativity increases going right and up the table Covalent bonds in which e - are not shared equally because of electronegativity differences are called polar covalent bonds Non polar covalent bonds are when the e - are shared equally.

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21 Categories of Bonds In general if the electronegativity difference between two bonded atoms is: 0, usually between identical nonmetal atoms, called nonpolar covalent <.4, mostly covalent.4 to 1.7, 2 different nonmetals called polar covalent > 1.7, usually nonmetals and reactive metals, is mostly ionic Note: there is no perfect ionic bond.

22 Determining Bond Polarity PROBLEM: (a) Use a polar arrow to indicate the polarity of each bond: N-H, F-N, I-Cl. (b) Rank the following bonds in order of increasing polarity: H-N, H-O, H-C.

23 Covalent Bonds Bond length, bond energy, and bond order are closely related Higher bond order is shorter, and stronger for a given set of atoms With a constant bond order, longer bonds are usually weaker.

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25 Bond & Energy Trends Bond lengths should increase with increasing atomic radius Bond strength should decrease with increasing atomic radius Example: Using the periodic table, rank the bonds in each set in order of decreasing bond length and bond strength: (a) S - F, S - Br, S - Cl (b) C = O, C - O, C O

26 Review Problems

27 IONIC & METALLIC BONDING November 8, 2016

28 Ionic Bonding In ionic bonding, electrons are gained or lost, the resulting bonds are based on electrostatic attraction. How can we depict an ionic bond? Ex: Use partial orbital diagrams and Lewis symbols to depict the formation of Na + and O 2- ions from the atoms, and determine the formula of the compound.

29 Properties of Ionic Bonds Ionic compounds tend to be hard, rigid, and brittle, with high melting points. Ionic compounds do not conduct electricity in the solid state. In the solid state, the ions are fixed in place in the lattice and do not move. Ionic compounds conduct electricity when melted or dissolved. In the liquid state or in solution, the ions are free to move and carry a current.

30 Properties in Ionic Bonds Ionic compounds are hard, rigid, and brittle This is a result of ions being held in specific positions in a crystal. So a crystal retains it s shape until enough energy is applied to shift positions and crack the crystal.

31 Metallic Bonding Metallic bonding involves electron pooling and occurs when a metal bonds to another metal. Electron Sea Model All the metal atoms contribute their valence electrons to a delocalized pool of electrons. The metal cations are held together by attraction to the delocalized electrons.

32 Metallic Bond Properties Most are solid at RT, with moderate to high mp, and very high b.p. m.p. are not very high because the metallic bonds don t have to be broken to become liquid b.p. are very high because the cation and it s electrons must be separated from the others m.p. are higher for metals with more valence electrons cation charges are higher resulting in greater cation-electron sea attractions

33 Properties of Metals Metals are good conductors of electricity when solid, or liquid. The delocalized electrons are able to move under an electric field Metals are good conductors of heat. The delocalized electrons disperse heat more quickly Metals are malleable and ductile, not brittle The cations are able to slide past each other and still retain their attraction to the electron sea.

34 BONDING ENERGY

35 Bond Energies and ΔH º rxn The heat released or absorbed during a chemical change is due to differences between the bond energies of reactants and products. ΔH rxn = ΣΔH reactant bonds broken + ΣΔH product bonds formed

36 Using Bond Energies to Calculate ΔH º rxn Calculate ΔH rxn for the chlorination of methane to form chloroform. bonds broken ΣΔH positive bonds formed ΣΔH negative

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38 Energy in Ionic Bonding Lattice energy is the energy required to separate 1 mol of an ionic solid into gaseous ions. Lattice energy is a measure of the strength of the ionic bond.

39 Periodic Trends in Lattice Energy Lattice energy is affected by ionic size and ionic charge. As ionic size increases, lattice energy decreases. Lattice energy therefore decreases down a group on the periodic table. As ionic charge increases, lattice energy increases.

40 Calculating Lattice Energy Lattice energy cannot be directly measured, so it is found by using Hess Law. The enthalpy change for an overall reaction is the sum of the enthalpy changes of the reactions which make it up. Lattice energies are calculated by using a Born-Haber Cycle

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42 Big Idea 2 Check-In Ionic, metallic, and covalent Bonds Polarity and dipole moment à Intermolecular forces (IMF) Lewis dot structures Resonance forms Incomplete octets Expanded octets Formal charge Molecular geometry (VSEPR) 10. Bonding and phases Kinetic molecular theory The ideal gas equation Dalton s Law Deviations from the ideal Behavior Density Solutions: molarity and mole fraction