Bonding: General Concepts Dr. Walker DE Chemistry

Bonding Bond Energy – Energy required to break a chemical bond Atoms bond in a way that achieves the lowest possible energy for the system Bond length – distance between the nuclei of two bonded atoms Shorter bond length = stronger bond = higher bond energy!!

Ionic Bonding Electrons are transferred Metals react with nonmetals Ions paired together have lower energy (greater stability) than separated ions Ionic compounds conduct electricity in a molten state

Coulomb’s Law Shows energy of interaction between a pair of ions E = Energy, r = distance between ionic nuclei, Q = ionic charge Lower energy = stronger attraction, more favorable interaction

Covalent Bonding Electrons are shared by nuclei Pure covalent (non-polar covalent) Electrons are shared evenly Polar covalent bonds Electrons are shared unequally Atoms end up with fractional charges d+ or d- (partial charges)

Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. General Trend – Increases to the right, Increases going up on the Periodic Table

Electronegativity and Bonding The greater the electronegativity difference, the greater the chance of a bond being ionic This has been quantified, but (as usual), there are exceptions.

Bond Polarity and Dipole Moments Dipolar Molecules Molecules with a somewhat negative end and a somewhat positive end (a dipole moment) due to a difference in electronegativity Molecules with preferential orientation in an electric field (see next slide) All diatomic molecules with a polar covalent bond are dipolar

Physical Effect Of Dipole Moments

Polar Bonds, Nonpolar Molecules Covalent molecules can have multiple dipole moments, but can be nonpolar overall due to symmetry.

Ions: Electron Configurations and Sizes Ionic bonds Electrons are transferred until each species attains a noble gas electron configuration (ns2np6) This trend is primarily for main group elements and still has exceptions (Pb2+, Sn2+ among others) Formulas are predicted based on atoms required to form a neutral atom Covalent bonds Electrons are shared in order to complete the valence configurations of both atoms Exceptions can occur below period 2 Note: This SHOULD be review. I’m not sure why the book puts it here.

Ionic Size Anions are larger than the parent atom Cations are smaller than the parent atom Essentially, they “lose” a shell Ion size increases within a group (more energy levels) Isoelectronic ions Ions with the same number of electrons (for example, Cl-, S2-, and P3- all have 18 electrons) Size decreases as the nuclear charge (atomic number) increases

Ionic Size

Energy In Ionic Compounds Lattice Energy The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid Energy change is negative (-DH) as it is released to the surroundings

Lattice Energy Modified form of Coulomb’s Law, shows energy “loss” from lattice formation k = a proportionality constant dependent on the solid structure and the electron configuration Q1 and Q2 are charges on the ions r = shortest distance between centers of the cations and the anions Lattice energy increases as the ionic charge increases and the distance between anions and cations decreases More charge, closer ions = more energy “savings”

Example Of Lattice Energy The final step of lattice formation is what makes the process exothermic overall

Crystal Lattices Our previous visualization of ionic salts as “stand-alone” compounds is NOT correct. Ionic salts form lattice structures, shown on the left. This network formation ultimately makes the energy work out in favor of bond creation. Ionic formulas are actually empirical formulas, since we don’t count the atoms in an entire lattice, which is impossible

Bond Character Differences in electronegativity cause unequal sharing of electrons in covalent bonds (polarity) or complete transfer of electrons (ionic bonds) Some cases are ambiguous….it’s hard to predict what they are Bond character – most bonds are not 100% covalent or 100% ionic

Bond Character Bond Character measured by the above equation Ionic vs. Covalent Ionic compounds generally have greater than 50% ionic character Ionic compounds generally have electronegativity differences greater than 1.6 Percent ionic character is difficult to calculate for compounds containing polyatomic ions (each ion has covalent bonds)

Electronegativity Values

Bonding As A Model Models are a way to represent what we observe in nature, especially in what we can’t see with our eyes Models do not equal reality and are often wrong The way covalent bonds are represented are a model….. ….these have proven to be a good representation in most cases, but some molecules require us to think of them as a unit rather than a collection of “bonds” Keep this in mind as we progress through the chapter

Covalent Bonding – A Review Electron pairs are shared between two atoms in covalent bonding Single Bond – one pair of electrons being shared Double Bond – two pairs of electrons being shared Triple Bond – three pairs of electrons being shared More electrons being shared = higher bond energy = shorter bond

Bond Energies For Covalent Bonds These bond energies are endothermic – energy added to break the bonds Conversely, bond formation is exothermic (as previously mentioned)

Localized Electron Bonding Model Lone electron pairs Electrons localized on an atom (unshared) Bonding electron pairs Electrons found in the space between atoms (shared pairs) Localized Electron Model "A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms Lone Pair Bonding Pair

Localized Electron Bonding Model Derivations of the Localized Model Valence electron arrangement using Lewis structures Prediction of molecular geometry using VSEPR (valence shell electron pair repulsion) Description of the type of atomic orbitals used to share or hold lone pairs of electrons

Lewis Structures Lewis structures show how valence electrons are arranged among atoms in a molecule. Lewis structures reflect the central idea that stability of a compound relates to noble gas electron configuration (octet rule) Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - )

HONC, HONC.. The HONC Rule Hydrogen (and Halogens) form one covalent bond Oxygen (and sulfur) form two covalent bonds One double bond, or two single bonds Nitrogen (and phosphorus) form three covalent bonds One triple bond, or three single bonds, or one double bond and a single bond Carbon (and silicon) form four covalent bonds. Two double bonds, or four single bonds, or a triple and a single, or a double and two singles

I Gotta Be Me…Exceptions To The Rules 2nd row elements C, N, O, F observe the octet rule (HONC rule as well). 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 3rd row (like S and P, especially) and heavier elements CAN exceed the octet rule using empty valence d orbitals. When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

Completing a Lewis Structure -CH3Cl Make the atom wanting the most bonds central Add up available valence electrons: C = 4, H = (3)(1), Cl = 7 Total = 14 Join peripheral atoms to the central atom with electron pairs. H .. .. Complete octets on atoms other than hydrogen with remaining electrons H .. C .. Cl .. .. .. H

Multiple Covalent Bonds: Double bonds Ethene Two pairs of shared electrons

Multiple Covalent Bonds: Triple bonds Ethyne Three pairs of shared electrons

Drawing Lewis Structures This differs from what you learned In previous class, as I GREATLY simplified it. Step 1: Count the number of valence electrons in the compound Step 2: Connect the atoms around the central atom by single bonds Step 3: Place remaining electrons on the outside atoms to fulfill octet rule Step 4: Place any remaining electrons on the central atom. Step 5: If the central atom does not have an octet, share pairs of electrons as bonds from outside atoms.

Practice

Answers

Practice - Exceptions

Answers - Exceptions

The Problem With Lewis The nitrate ion shows two types of bonds: 1 double bond and 2 single bonds Experiments show ONLY ONE BOND TYPE with a bond length BETWEEN that of a single and double bond This is not the only molecule exhibiting this behavior Presents a limitation of Lewis structures and the Localized Electron Model

Resonance Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule. The actual structure is an average of the resonance structures The bond lengths in the ring are identical and between those of single and double bonds Benzene, C6H6

More Resonance Ozone (O3) Bond lengths are identical, between that of a single and double bond Neither structure is Technically correct!

Resonance in Polyatomic Ions Resonance in a carbonate ion: Resonance in an acetate ion:

Odd-Electron Molecules Lewis structure models work best with molecules with even numbers of total electrons The localized electron model is based on pairs, so it doesn’t work so well with these Example NO, 11 total valence electrons

Formal Charge Method of handling ions/molecules with unusual arrangements to help determine proper Lewis structure

Assigning Formal Charge This pertains to a Lewis structure that you must draw first Number of valence electrons on the free atom minus the number of valence electrons assigned to the atom in the molecule Lone pair (unshared) electrons belong completely to the atom in question Shared electrons are divided equally between the sharing atoms The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species If the charge on an ion is -2, the sum of the formal charges must be -2

Assigning Formal Charge In a molecule with many possibilities, the representation with the lowest formal charges is most appropriate Bottom choices are best here with no charge on Xenon higher than +1

Formal Charge Practice

Formal Charge Examples A) P = 0, O = 0, Cl = -1 B) S = 0, 2 oxygens = 0, 2 oxygens = -1 C) Cl = 0, 3 oxygens = 0, 1 oxygen = -1 D) P = 0, 1 oxygen = 0, 3 oxygens = -1

VSEPR Valence Shell Electron Pair Repulsion The structure around a given atom is determined principally by minimizing electron-pair repulsions. Translation: all atoms and electrons want their own personal space and get as far away from everyone else as possible

Possible Arrangements X + E Overall Structure Forms 2 Linear AX2 3 Trigonal Planar AX3, AX2E 4 Tetrahedral AX4, AX3E, AX2E2 5 Trigonal bipyramidal AX5, AX4E, AX3E2, AX2E3 6 Octahedral AX6, AX5E, AX4E2 A = central atom X = atoms bonded to A E = nonbonding electron pairs on A

Linear Structure No unbonded electrons to deal with, chlorines as far apart as possible

Linear AX2 CO2

Trigonal Planar No unbonded electrons on central atom, molecule stays in a single plane

Trigonal Planar AX3 BF3 AX2E SnCl2

Tetrahedral When four pairs of electrons are present around an atom, they should always be arranged tetrahedrally. This can be in bonds OR in electron pairs

Tetrahedral AX4 CCl4 AX3E PCl3 AX2E2 Cl2O

Trigonal Bipyramidal Involves the arrangement of a compound with five electron pairs surrounding a central atom Notice the term “electron pairs” instead of “atoms”

Trigonal Bi-pyramidal AX5 PCl5 AX4E SF4 AX3E2 ClF3 AX2E3 I3-

Trigonal Bipyramidal Arrangements

Octahedral Six electron pairs arranged around a given atom with 90-degree angles

Octahedral AX6 SF6 AX5E BrF5 AX4E2 ICl4-

Electron Pairs Lone (unshared) electron pairs require more room than bonding pairs (they have greater repulsive forces) and tend to compress the angles between bonding pairs