General Bonding Concepts Chapter 8 Sections 1-12 only

Objectives Define 3 different types of bonds based on electron distribution Use Coulomb's law equation to determine attraction/repulsion

Chemical Bonds What is a bond? Why do bonds form? What is bond energy? The forces that cause atoms to act in unison Why do bonds form? to use or achieve the lowest possible energy What is bond energy? the amount of energy required to BREAK the bond

Types of Bonds Ionic: Electrons are transferred Polar Covalent: Electrons are shared unequally Covalent: Electrons are shared equally

Ionic Bonding What types of substances react to form ionic bonds? Answer: Metals and non metals Energy of interaction between a pair of ions can be calculated using COULOMB'S LAW. The formula:

Coulomb’s Law Equation Constant: Joules= energy unit, nm = distance unit Charge on Ions Distance between atoms (in nm) In Joules

Using Coulomb's Law In NaCl the distance between ions is 0.276 nm, What is the energy of this bond interaction? What equation? Set it up: Answer: E = -8.37x10-19 J

Coulomb's Law A Negative Coulomb's Law answer indicates an attractive force A Positive Coulomb's Law answer indicates a repulsive force

Coulomb’s Law Practice

Objectives Discuss relationship between bond energy and bond length Discuss how electronegativity impacts bonding Determine if a bond will be polar, if so show dipole moment Define bond, bond energy, bond length, dipole moment Explain the three forces at work when atoms bond

Relate Bond Energy and Bond Length Finish this statement: as bond length decreases bond energy… INCREASES Defend using Coulomb's Law:

What about two same type atoms? Example H:H A bond will form if the energy of the aggregate is lower than that of the separated ions or atoms (p. 331) Picture from: chemkb.cs.iupui.edu

Bonding between same type ions These are covalent bonds. Bond Length is defined as the distance where the lowest energy is achieved. Several forces are involved to establish bond length

The Forces at Work When same type atoms interact you get a covalent bond. The following forces must be balanced for a covalent bond to occur: What is this trying to do? Type of interaction: Proton vs Proton Proton vs Electron Electron vs Electron Repel Attract

Key points of Covalent Bonding Shared Electrons Electrons shared in area between nuclei Bonds created to increase stability

Key points of Covalent Bonding Unequal sharing of electrons Based on electronegativity difference Electronegativity Differences and bond type 0-1 "covalent" 1-2 "polar covalent" >2 "ionic“ Actually more like a sliding scale Most Covalent Most Ionic 0 4

Consequences of Unequal Sharing Dipole Moment: partial charges caused by the unequal sharing of electrons Polar bond does not always equal polar molecule Dipoles can cancel due to molecular structure Two ways to show dipoles: H F H F d+ d-

Assignment Page 383 # 24, 26, 30, 32

Show the bond polarity for the following bonded atoms: C-O Se-S P-H Cl-I H-Cl Br-Te Br-Br O-P Si-S

Objectives: Describe what it means for an atom/ion to be stable Use electron configurations to predict formulas of ionic compounds Describe the relationship between parent atom size and ion size for cations and anions Use a periodic trend to determine relative ion size in a group Define an isoelectronic ion group Discuss the trend for ions size in an isoelectronic group

Stable Atoms Quantum mechanical model has helped to show that atoms in a stable compound have achieved noble gas configuration (full energy level) by either sharing electrons or by forming ions by either the loss or gain of electrons.

Stable covalent compounds are achieved by atoms sharing electrons so that both complete their valence shells to attain noble gas configuration Stable ionic compounds are formed when a non-metal removes the valence electrons from the metal so that it fills its valence shell to achieve noble gas configuration, the metal reverts back to the last noble gas configuration.

Predicting Ion Formation You can use electron configurations to predict ion formation Ca: [Ar]4s2 Ca wants to lose 2 electrons and achieve the configuration of Argon O: [He] 2s22p4 O wants to gain 2 electrons and achieve the configuration of Neon From this we predict: The calcium ion The oxide ion The compound calcium oxide Ca2+ O2- CaO

Ion Prediction and Compound Formation What happens when aluminum and oxygen react? Look at electron configurations: Al = [Ne] 3s23p1 O = [He] 2s22p4 What does each atom "want"? Al "wants" to lose 3 electrons, O "wants" to gain 2 electrons What formula would you predict? Why? Al2O3 all atoms are "satisfied"

Ion Size Important when determining stability, structure, and properties of ionic solids and aqueous ions. Impossible to precisely define ion size There are several ways to look at trends for ion size

Ion Size and The Parent Atom Cations are smaller than their parent atom. Why? loss of electrons results in the loss of an entire energy level which decreases the size of the entity Anions are larger than their parent atom. Why? gain of electrons results in a fuller outer energy level which increases the size of the entity

Ion Size and Periodic Groups Ion size increases down a group. Why? as you move down a group, there are more filled energy levels= larger ion There is not a trend for ion size across a period like there is for other types of periodic trends. as you move across a period, there is a change over from metals (cations) to non metals (anions) so in the middle of a period, it will change from smaller ions to large ones

Ion Size and Isoelectronic Ions Define Isoelectronic: ions containing the same number of electrons Example: O-2, F-1, Na+1, Mg+2, Al+3 What is the electron configuration of each ? They all have the configuration of NEON! 10 total electrons are in each ion! This is an isoelectronic set of ions

Ion Size and Isoelectronic Ions ... continued The trend for isoelectronic ions is that ion size decreases as atomic number (nuclear charge, number of protons, z) increases. Why?? Isoelectronic ions by definition have the same number of electrons, therefore, the greater the number of protons (atomic number) the greater the positive force drawing those electrons towards the nucleus (making the ion smaller by drawing the electrons closer)

Ion Size and Isoelectronic Ions... Practice Arrange the following in order of decreasing ion size Br-1, Rb+1, Se-2, Sr+2 Se-2, Br-1, Rb+1, Sr+2 Choose the smallest ion from the following sets Li+1, Na+1, K+1, Rb+1 Li+1 Ba+2, Cs+1, I-1, Te-2 Ba+2

Assignment: Page 383 # 35, 37, 39, 41

Answers: #35 a. Sc+3 , b. Te-2 , c. Ce+4 and Ti+4, d. Ba+2 #37. La+3 , Ba+2 , Cs+1 , I-1 , Te-2 #39 a. Cu > Cu+1 > Cu+2 b. Pt+2 >Pd+2 >Ni+2 c. O-2 , O-1 , O d. La+3 , Eu+3 , Gd+3 , Yb+3 e . Te-2 , I -1 , Cs+1 , Ba+2 , La+3 #41 a. Al2S3 – aluminum sulfide b. K3N – potassium nitride c. MgCl2 – magnesium chloride d. CsBr – cesium bromide

Objectives: Define lattice energy Determine relative lattice energy among a group of compounds Use lattice energy to calculate reaction energy (5 step)

Na+1 (g) + Cl-1 (g) --> NaCl (s) Lattice Energy Define Lattice Energy: the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid Na+1 (g) + Cl-1 (g) --> NaCl (s) change in energy from here to here is lattice energy !

The Textbooks's Perspective on Energy Viewed from the system's point of view Negative energy if the process is exothermic Positive energy if the process is endothermic This perspective gives lattice energy a negative value!

Lattice Energy Calculation Lattice Energy is Calculated by a modified form of Coulomb's Law. Where k is a constant that depends on the structure of the solid and the electron configurations of the ions Note: you will NOT be expected to calculate lattice energy, just to know how it is calculated and to use lattice energy to calculate total reaction energy

Estimating Relative Lattice Energy Lattice energy increases as ion charge (Q1 and Q2) increases Lattice energy decreases as the distance between ions (r) increases

Practice: Which compound in the following pairs has the most exothermic lattice energy? Why? NaCl or KCl Mg(OH)2 or MgO LiF or LiCl Fe(OH)2 or Fe(OH)3 NaCl or Na2O MgO or BaS NaCl : smaller r MgO : larger Q LiF : smaller r Fe(OH)3 : larger Q Na2O : Larger Q MgO : smaller r

Ionic Solid Formation Energy Calculation Li (s) + 1/2 F2 (g) --> LiF (s) Note: must have balanced equation Use lattice energy to determine the energy change experienced in this reaction. The calculation is broken into steps based on states of matter Sublimation Ionization Dissociation Ion formation Solid formation (lattice energy**) ** remember to use lattice energy you need separate, gaseous ions. The other 4 steps just get you to this point.

An Example Step Process Energy change Sublimation Li (s) --> Li (g) 161 kJ/mol Ionization Li (g) --> Li+1 + 1 e- 520 kJ/mol Dissociation 1/2 F2 (g) --> F (g) 154 kJ/mol Ion Formation F (g) + e- --> F-1 (g) -328 kJ/mol Solid Formation Li+1(g) + F-1 (g) --> LiF (s) -1047 kJ/mol Total Energy Li (s) + 1/2 F2 (g) --> LiF (s) -617 kJ the negative sign tells that energy is being released (exothermic)

Assignment: Page 383-4 # 44, 45, 46, 49

Answers:

A better definition of ionic compounds: Any compound that conducts electricity when melted is ionic Avoids confusion when polyatomic ions, which are often held together covalently are part of the compound

Use of Models Models attempt to explain how nature operates Models are human inventions: Models do not equal reality Models can be wrong. They are oversimplifications. Exceptions deal with items that do not meet standards due to oversimplifications. Need to understand strengths and weaknesses When models are wrong, we can learn

Using Bond Energy to Calculate Reaction Energy Bond Energy depends on the environment Example: C-H bond in HCCl3 = 380 kJ/mol C-H bond in C2H6 = 410 kJ/mol Average Bond Energy is still useful

Objectives: Discuss a "bond" as a model Explain why models are used in science Discuss 5 properties of models Define single, double and triple bonds- # of e- shared, bond energy & length Use bond energies to calculate ΔH Explain the relationship between number of shared electrons and bond strength and/or bond length Write the expression for ΔH when using bond energy to calculate the enthalpy of the reaction

Ways to share electrons Single Bonds - single pair of electrons shared C-C - two total shared electrons Double Bonds - two pairs of electrons shared C=C - four total shared electrons Triple Bonds - three pairs of electrons shared C=C - six total shared electrons

Using bond energy to calculate Enthalpy H2 (g) + F2 (g) --> 2 HF (g) This reaction requires us to break a H-H bond and a F-F bond on the reactant side of the reaction It also requires that we make two H-F bonds on the product side of the reaction. Breaking bonds is an endothermic process- positive numbers Making bonds is an exothermic process- negative numbers Average bond energy Table 8.4 is on page 351.

How it is done . . . Enthalpy change for a reaction is calculating by the addition of the energy required to break old bonds and the energy released by the formation of new bonds. ΔH = ΣD (bonds broken) - ΣD (bonds formed) Note: you are actually adding the energy from "bonds formed", however, since these are releasing energy the sign on the numbers is negative leading to the negative sign in the equation.

Back to the problem . . . H2 (g) + F2 (g) --> 2 HF (g) Breaking one mol H-H @ 432 kJ/mol = 432 kJ Breaking one mol F-F @ 154 kJ/mol = 154 kJ Making two moles H-F @ 565 kJ/mol = 1130 kJ Bonds broken total= 586 kJ Bonds formed total= 1130 kJ 586 - 1130 = -544 kJ Average bond energy Table 8.4 is on page 351.

Try this one on your own . . . carbon is central atom CH4 + 2 Cl2 + 2 F2 --> CF2Cl2 + 2HF + 2HCl Answer: -1194 kJ

Assignment: Page 384 # 53,55,57,59,61 Answers

Answers Page 384 # 53,55,57,59,61

Objectives Explain the Localized Energy Model Define lone pair Define bonding pair Define Lewis structure Define duet rule List elements that obey the duet rule Define Octet rule List elements that obey the octet rule

Objectives Draw Lewis structures for atoms Draw Lewis structures for molecules Explain how there are exceptions to the octet rule List some elements that are exceptions to the octet rule Draw Lewis structures for compounds where elements exceed the octet rule

The Localized Electron Model **Assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms

The LE Model Continued Electrons are assumed to be localized on an atom or in the space between two atoms. Lone Pair : Localized on one atom Bonding Pair: Found in space between two atoms

Three Parts of LE Model Description of valence electron arrangement with Lewis Structures Prediction of Molecular Geometry using VSEPR Model Description of the type of atomic orbitals used by atoms to share electrons. For this unit we are only going to deal with the first part, the second and third will be worked with next unit.

Lewis Structures Shows how valence electrons are arranged among atoms in a molecule. Most important requirement for a stable compound is the achievement of noble gas configuration.

Two Basic Rules Duet Rule: a stable molecule is formed when two electrons are shared Octet Rule: a stable molecule is formed when electrons are shared so that each atom is surrounded by eight electrons

Steps for Writing Lewis Structures Sum the VALENCE electrons for all atoms in the molecule, it does not matter how many come from each, just the total Use a pair of electrons to form a bond between each pair of bond atoms Arrange remaining atoms to satisfy the duet rule for hydrogen and the octet rule for second row elements.

An Example: Water: H2O 1. Sum valence electrons: Total= 8 valence electrons H= 1 valence electron O= 6 valence electrons 2. Use a pair to form bond between atoms 3. Arrange remaining to satisfy duet/octet rule

Another Example: Carbon Dioxide: CO2 Sum valence electrons Total: 16 valence electrons C= 4 valence electrons O= 6 valence electrons 2. Use pairs to bond atoms 3. Arrange remaining to meet duet/octet

Assignment Page 385 # 67 and 68 Answers

Answers Page 385 # 67 and 68

Violations of the Octet Rule Be, B are often electron deficient Third row and beyond elements are able to exceed the octet rule by placing extra electrons in their "d" orbitals

Examples: BF3 SF6 PCl5

Assignment Page 385 69, 71,72 Answers

Answers Page 385 69, 71,72

Objectives Define Resonance Define Resonance structure Explain why odd electron molecules cannot be shown using the Localized Energy Model Define formal charge Explain the “problem” with oxidation state assignments Explain how formal charge is different from oxidation state Use formal charge to determine the most probable Lewis structure for a molecule

Resonance Occurs when there is more than one equivalent and valid Lewis structure for a molecule. Is necessary to compensate for the incorrect assumption that electrons are localized In reality, electrons are delocalized- meaning that they move around the entire molecule LE model is still useful, so the exception of resonance is added to accommodate certain molecules rather than tossing the whole model

An example: NO3-1

Another Example NO2-1

Assignment Page 385 # 73,74,75 Answers

Answers Page 385 # 73,74,75

Odd - Electron Molecules Only a few molecules form containing odd numbers of electrons Examples: NO and NO2 The Localized Electron Model only works with PAIRS of electrons so this model is not able to accomodate odd electron molecules

Formal Charge Formal charge works with nonequivalent Lewis structures often found in molecules and polyatomic ions that have atoms that can exceed the octet rule. Charges are utilized to determine the most appropriate structure(s)

Charge Determination Oxidation states: Both shared electrons count for the more electronegative atom Results in exaggerated charge estimates Not good for determining proper Lewis structures

Charge Determination Formal Charge: The difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule When calculating formal charge: Lone pair electrons belong to the atom on which they are located. Shared pair electrons are split equally among sharing atoms.

Rules for working with Formal Charge: Atoms in molecules want to have formal charges as close to zero as possible. Negative formal charges should occur on the more electronegative atom.

An Example: SO4-2

Another Example XeO3 Hint: there are 8 possible Lewis structures

Assignment Page 386 # 81 and 82 Answers

Answers Page 386 # 81 and 82

End of Unit 1 Chapter 8 Sections 1-12.