Bonding Chapter 7

Unit Objectives 5.0 Define key terms and concepts 5.1 Draw Lewis Dot structures and line structures for simple chemical compounds. 5.2 Identify ionic and covalent compounds by their chemical formulas and their properties. 5.3 Determine the formal charge on an atom. 5.4 Explain the octet rule and exceptions to the rule. 5.5 Draw Resonance structures for chemical compounds. 5.6 Determine the lattice energy of a compound using the Born-Haber Cycle. 5.7 Calculate the bond enthalpy change in a reaction.

Stable Electron Configurations Valence electrons interact between atoms of elements to form bonds when compounds are formed This interaction involves either the sharing or transfer of electrons Octet Rule When electrons gain, lose, or share electrons to obtain 8 valence electrons (or a full outer shell) Basically every element is trying to mimic a Noble Gas with their electron configuration

Lewis Electron Dot Structures A way to represent the bonding and non-bonding electrons in a compound

Lewis Electron Dot Structures Bonding Electrons Involved in a bond Non-Bonding Electrons (Lone Pairs) Not involved in a bond Line Structure Replaces the bonding pair of electrons with a line

Lewis Electron Dot Structures He S Ca Na Al O C N B

Stable Electron Configuration Ionic Bonding Covalent Bonding

Ionic Bonding Bond formed when an electron is transferred from one atom to another resulting in the creation of two ions. Typically between a metal and a non-metal Ions Element with a charge Formed due to ionic bonding Ionic Charge An electrical charge on an element due to the lose/gain of an electron This charge difference creates an electrostatic force, holding the two ions together. The overall compound has a neutral charge.

Ions Cations Anions Ions with a positive charge Typically metals Na+1, Al+3, C+4 Anions Ions with a negative charge Typically nonmetals O-2, Cl-1, N-3

Lewis Dot Structures – Ionic Compounds

Draw Lewis Dot and line structures of the following compounds. Be2O HBr MgF2 K2S

Covalent Bonds Formed when electrons are shared between two elements. Created when two nonmetals bond together Polar Covalent Electrons are held slightly closer to one atom than another, creating an imbalance of electrons, resulting in a dipole moment Dipole Separation of positive and negative charges in a bond Non-Polar Covalent Equal sharing of electrons Does not have a dipole moment

Covalent Bonds

Covalent Bonds

Covalent Bonds Diatomic elements Formed when an 2 of the same elements bond covalently Are gaseous Non-polar

Covalent Bonds

Multiple Bonds Single Bond Double Bonds Triple Bonds 1 pair of electrons is involved in bonding Double Bonds Sharing of two electron pairs Triple Bonds Sharing of three electron pairs Takes more energy to hold a triple bond that a single bond.

Lewis Dot Structures

Lewis Dot Structures

Draw Lewis Dot and line structures of the following compounds Draw Lewis Dot and line structures of the following compounds. H2 SH2 CBr4 PCl3

Covalent Bonds Bond Length The distance between the nuclei of two covalently bonded atoms Single bonds have the longest bond length while triple bonds have the shortest Bond length increases with atomic radius Bond Type Bond Length (pm) C-C 154 C=C 133 C≡C 120

Electronegativity The tendency of an atom to attract electrons Increases to the right and up on the periodic table The greater the difference in charges between to atoms, the more the electrons are attracted to one atom over another

Electronegativity Nonpolar Covalent Bond Polar Covalent Bond EN difference less than 0.5 Polar Covalent Bond EN difference from 0.5 to 1.7 Ionic Bond EN difference greater than 1.7

What Are Your Questions?

Formal Charge of an atom The electrical charge difference between the valence electrons of an atom and the number of electrons assigned to that same atom in the Lewis structure. To determine the formal charge of an atom Assign all of the non-bonding electrons to the atom they are attached to Assign half of the bonding electrons to each of the two atoms involved in a bond Writing formal charges The sum of the charges must equal zero for molecules The sum of the charges must be a positive charge for cations The sum of the charges must be a negative number for anions

Formal Charge of an atom Determining the most plausible Lewis structure for a molecule Having no charge is preferable to having a charge Structures with small formal charges are preferable to those with large formal charges Charges are usually limited to -1, 0, or +1 unless a metal is present It is more plausible for the negative charges to be placed on the more electronegative atom, etc.

Exceptions to the Octet rule An Incomplete Octet Applies to hydrogen to carbon Aluminum Some atoms with a formal charge

Exceptions to the Octet rule Odd number of electrons Radical Has an unpaired electron Expanded Octet More than 8 valence electrons Can occur in elements in the third period and beyond

Draw the Lewis dot structure and determine the formal charge of the elements in each of the following compounds: H3O+ NO3-

Draw the Lewis dot structure and determine the formal charge of the elements in each of the following compounds: CN- H2SO4

Resonance Structures When more than one Lewis structure can be drawn for a given compound. Typically happens when a molecule contains 1 or more multiple bonds. Common on ions

Draw the resonance structures for the following compounds. NO3- PO43- HSO4-

What Are Your Questions?

∆H° = ΣBE(reactants) – ΣBE(products) Bond Enthalpy The energy required to break a bond in 1 mole of gaseous molecules. The bond enthalpies of liquids and solids are effected by the surrounding molecules. Where BE stands for average bond energy This applies only to the gas phase Turn to Table 9.4, pg. 401. ∆H° = ΣBE(reactants) – ΣBE(products)

Estimate the enthalpy change for the following reaction: N2(g) + H2(g)  NH3(g)

Estimate the enthalpy change for the following reaction: CH4(g) + O2(g)  H2O(g) + CO2(g)

Estimate the enthalpy change for the following reaction: C2H4(g) + H2(g)  C2H6(g)

Ionic Compounds Lattice Energy Coulomb’s Law Born-Haber Cycle The energy required to completely separate one mole of a solid ionic compound into gaseous ions. Coulomb’s Law The potential energy (E) between two ions is directly proportional to the product of their charges and inversely proportional to the distance between the two ions. Born-Haber Cycle Relates lattice energies of ionic compounds to other atomic and molecular properties, such as ionization energy and electron affinity. Based on Hess’s Law

The Born-Haber Cycle

Calculate the enthalpy change for following reaction using the Born-Haber Cycle. Mg(s) + ½O2(g)  MgO(s) ∆Hatomization O = 249kJ/mole ∆Hatomization Mg = 148kJ/mole ∆H1st ionization energy Mg = 738kJ/mole ∆H2nd ionization energy Mg = 1451kJ/mole ∆H1st electron affinity O = -141kJ/mole ∆H2nd electron affinity O = 798kJ/mole ∆Hlattice energy MgO = -3791kJ/mole

Calculate the enthalpy change for following reaction using the Born-Haber Cycle. 2Li(s) + ½O2(g)  Li2O(s) ∆Hatomization O = 249kJ/mole ∆Hatomization Li = 162kJ/mole ∆H1st ionization energy Li = 540kJ/mole ∆H1st electron affinity O = -141kJ/mole ∆H2nd electron affinity O = 798kJ/mole ∆Hlattice energy Li2O = -2800kJ/mole

Calculate the enthalpy change for following reaction using the Born-Haber Cycle. Mg(s) + Cl2(g)  MgCl2(s) ∆Hatomization Cl = 122kJ/mole ∆Hatomization Mg = 148kJ/mole ∆H1st ionization energy Mg = 738kJ/mole ∆H2nd ionization energy Mg = 1451kJ/mole ∆Hformation MgCl2 = -641kJ/mole ∆Hlattice energy MgCl2 = -25261kJ/mole

What Are Your Questions?