CH 8: Bonding General Concepts

Chapter Outline – Part I Types of chemical bonds (8.1) Electronegativity and bond polarity (8.2/3 Ions (8.4) Energy changes when a binary ionic compound forms (8.5) Ionic character of covalent bonds (8.6)

Introduction to Bonding Chemical bond – force that holds atoms together so that they function as a unit. Consider 2 classes of bonds: Ionic bonding Covalent bonding

Bond Types Ionic bonds – attractive forces among oppositely charged ions Forms when a metal loses electron(s) to a nonmetal. Bond strength can be calculated using Coulomb’s law

Ionic Bonds Strength of the attraction between the ions can be calculated using Coulomb’s law. E = (2.31 x 10-19 J nm) (Q1Q2/r) Q1 and Q2 are the charges on the ions. r = distance between ion centers in nm

Using Coulomb’s Law E = (2.31 x 10-19 J nm) (Q1Q2/r) Sign on E??? The more negative E, the stronger the attractive force between the ions.

Using Coulomb’s Law E = (2.31 x 10-19 J nm) (Q1Q2/r) Magnitude of E. E is more negative when:

Covalent Bonds Covalent bond – bonded atoms share pairs of valence electrons Covalent bonding results in formation of a molecule. Covalent bonding occurs between nonmetals.

Types of Covalent Bonds “Pure” covalent bond – electrons are shared by like nonmetals E.g. diatomic molecules Results in equal sharing of the electrons Aka – nonpolar covalent bond

Types of Covalent Bonds Polar covalent bond – unequal sharing of electrons by the bonded atoms bond between different nonmetals each with its own ability to attract the shared electrons

Polar Covalent Bonds Showing bond polarity: Consider the HF molecule. See board and/or page 290. Experimental determination of bond polarity, page 289

Bond Polarity To predict bond polarity…consider the electronegativity (EN) of the bonded atoms. EN – the ability of an atom in a molecule to attract shared electrons.

EN Values The higher the EN the greater the atom’s ability to attract shared electrons. EN values and the periodic table EN ________ down a group. EN ________ across a period. See “back” of the periodic table.

EN and Bond Polarity As the difference in EN between bonded atoms increases so does the polarity of the bond. Can also say that the ionic character of the bond is increasing. See table 8.1 on page 289.

Bond Polarity and Dipoles Polar molecules have a preferred orientations when placed in an electric field. Said to have a dipole moment. Dipole moment – molecule has a center of positive charge and a center of negative charge

Bond Polarity and Dipoles Not all molecule with polar bonds have dipole moments! Bond polarities cancel each other in molecules with symmetrical dipoles. Molecules with equal, opposing dipoles. See page 291and 8.2 on page 292 Dog walking example!

Compound Formation Atoms gain, lose, or share enough electrons to achieve the same stable electron configuration as a noble gas Nonmetals share electrons Form molecules with covalent bonds Representative metals lose electrons to nonmetals in ionic compounds Ions are isoelectronic to noble gases

8.4 Ion Formation Binary ionic compounds The metal loses electron(s) to a nonmetal Focus on represenative metals The atoms lose/gain enough electrons to obtain a noble gas electron configuration.

Cations Group IA metals form ions with a _____ charge. Na atom Na ion Isoelectronic to: ____________________

Anions Group VIA elements form ions with a ______ charge. Sulfur atom Sulfur ion (called _________________) Isoelectronic to: ____________________

Ionic Compound Consider the compound formed between sodium and sulfur. Each sodium atom loses 1 electron. Each sulfur atom needs 2 electrons. Formula for compound:

Ion Size Cations are smaller than their parent atom. Atoms lose their valence shell when the ion forms. Na 1s22s22p63s1 ____ protons Na+ 1s22s22p6 ____ protons

Ion Size Anions are larger than their parent atom. Atoms add electrons to their valence shell when the ion forms – proton # remains the same. F 1s22s22p5 ____ protons F1- 1s22s22p6 ____ protons

Ion Size Isoelectronic ions decrease in size as the number of protons increases. Example: ions with 10 electrons 10 e O2- F1- Na1+ Mg2+ Al3+ # p 8 9 11 12 13

Isoelectronic Ions 10 e O2- F1- Na1+ Mg2+ Al3+ # p 8 9 11 12 13 Radius* 140 136 95 65 50 * picometers The diagram on page 296 should make sense.

8.5 Energy in Binary Ionic Compounds Lattice energy – change in energy when separated gaseous ions form an ionic solid. M+(g) + X-(g)  MX(s) LE < 0

Lattice Energy LE = k (Q1Q2)/r K is the proportionality constant Q1 and Q2 are the charges on the ions r is the ionic radius

Lattice Energy LE becomes more exothermic as the ion charges increase and the ion radius decreases. Small highly charged ions have more exothermic LE See board for examples.

Formation of ionic compounds. Consider energy changes associated with formation of a binary ionic compound. 5 step process, page 297/298 Most common series of steps is shown on the next slide.

Formation of ionic compounds. Sublime the metal. Ionize the gaseous metal atoms. ionization energy(ies) Dissociate the nonmetal (if diatomic). Bond energy Ionize the gaseous nonmetal atoms. Electron affinity Form the solid from the gaseous ions LE

Born Haber Practice… NaF MgF2 IE – page 272; EA – page 275 Bond energies – page 306 Sublimation: Na 109 kJ; Mg 147 kJ Lattice energy: NaF -923kJ; MgF2 2913

8.6 Partial Ionic Character When atoms with different EN bond the result is either a polar covalent or an ionic bond. There’s evidence that some level of electron sharing occurs in all bonds. Even in what we consider as ionic bonds.

8.6 Partial Ionic Character Classify a bond as ionic if it conducts electricity when melted. Essentially all compounds with metals meet this criteria. These compounds generally have more than 50% ionic character.

8.8 Covalent Bond Energies Strength of a given bond depends upon the compound. Not all C-H bonds are of the same energy! See page 305. Bond energies given in tables are averages based on experimental data.

Bond Energies Consider the bond energies on page 306. Compare the bond energies and bond length associated with single, double, and triple bonds between a given pair of atoms.

8.8 Using Bond Energies The DH for a reaction can be estimated from bond energies. DH = energy needed to break bonds of reactants – energy released when product bonds form

8.8 Using Bond Energies Estimate the DH for.

CH 8 Part II: Bonding Models Introduction to models (8.7) Localized Electron (LE) Bonding Model (8.9) Lewis Structure (8.10) Resonance (8.12) Exceptions to the Octet Rule(8.11) VSEPR Theory (8.13) Key pages: 326/27

8.7 Models Read “Fundamental Properties of Models” on page 350.

8.9 LE Bonding Model Localized electron bonding model Assumes a molecule is made of atoms bound together by sharing pairs of electrons using the orbitals of the bonding atoms.

8.9 LE Bonding Model Localized electron bonding model Shared electrons are pictures to be localized in the space between the atoms Called bonding pairs Non-bonding valence electrons are pictured to be localized on the parent atom. Called lone pairs Consider HCl

8.10 Lewis Structures Lewis structures show the arrangement of the valence electrons in molecules (and ions). Representative atoms will have the same number of valence electrons as one of the noble gases 2 electrons to be like H 8 electrons to be like all other noble gases

Lewis Structures Lewis structures illustrate LE bonding model. Show the bonding electrons and the lone pairs. Lewis structures can be used to predict the 3D geometry of a molecule. Requires application of VSEPR Theory More to come on this…..

1st Goal: To Write Lewis Structures Sum the valence electrons. Use a pair of electrons to form a bond between each of the bonded atoms. Put the atom that needs the most electrons in the center when the molecule contains more than 2 atoms. Arrange the remaining electrons to satisfy the duet rule for H and the octet rule for elements in the 2nd row of elements.

Writing Lewis Structures Practice! H2O O2 HCN NO31- PH3

8.12 Resonance More than one valid Lewis structure can often be drawn for molecules with multiple bonds (double, triple..) Consider NO21- 2 valid Lewis structures can be drawn.

Resonance Structures Lewis structure just drawn indicate 2 types of bonds in NO21- -- single bond and a double bond However….the data shows that both bonds in NO21- are of the same energy and bond length Both bonds are stronger and shorter than a single bond, but not as strong or short as a double bond!

Exceptions to the Octet Rule Less than an octet. Be and B More than an octet 3rd period elements and “up” Odd number of electrons

Exceptions to the Octet Rule Less than an octet. Be - satisfied/stable with 4 electrons B - satisfied/stable with 6 electrons

Exceptions to the Octet Rule 2. More than an octet Atoms in the 3rd period and “up*” can use their unfilled d orbitals to accommodate more than 8 electrons Commonly see 10 electrons and 12 electrons around the central atom. “Up” refers to periods 4, 5,6,…

More than an octet ICl3 PF5

Exceptions to the Octet Rule Odd number of electrons A small number of molecules have an odd number of electrons Called “free radicals” Molecules are highly reactive/unstable “steal” an electron from other molecules Example: NO

VSEPR Theory Valence Shell Electron Pair Repulsion Theory Structure around a given atom is determined by minimizing e-pair repulsions Atoms arrange themselves in 3D space in a manner that minimizes electron pair repulsive forces

VSEPR Theory Predictions based on VSEPR theory agree closely with experimental data. CO2 BF3 SO2

Goals – much practice in lab Applying LE and VSEPR Theory From Lewis structure to: Electron pair geometry (electronic geometry, EG) molecular geometry (MG) bond angles Molecular polarity Hybridization (CH 9) Bond type – sigma, pi

Sigma and Pi Bonds Sigma bonds – bonding electrons overlap (bond) on the internuclear axis Pi bonds – unhybridized p orbitals overlap (bond) above and below the inter nuclear axis

Single bond = 1 sigma bond Double bond = 1 sigma bond, 1 pi bond Triple bond = 1 sigma bond, 2 pi bonds

O=C=O

Count # electron densities around the central atom Write Lewis structure. Count # electron densities around the central atom Determines electronic geometry (EG) Also determines hybridization of central atom Count # lone pairs around central atom. This combined with #2 determines molecular geometry (MG) Determine polarity of each bond. From this determine polarity of molecule. Other….Label sigma and pi bonds, identify orbitals involved.

Lab – when polarities differ