Unit 3: Chemical Bonding and Molecular Structure Cartoon courtesy of NearingZero.net

Bonds Forces between nuclei and valence electrons of different atoms that binds the atoms together– bonds form because it LOWERS the potential energy Ionic bonds – transfer of electrons; smallest unit is called a “formula unit” Covalent bonds – sharing of electrons; smallest unit is a “molecule”

Properties of Ionic Compounds Structure: Crystalline solids Melting point: Generally high Boiling Point: Electrical Conductivity: Excellent conductors, molten and aqueous Solubility in water: Electrons: Elements: Generally soluble transferred Metal + nonmetal

Properties of Covalent Compounds Structure: Gases and some liquids Melting point: Generally low Boiling Point: Electrical Conductivity: Poor conductors, except aqueous acids Solubility in water: Elements: Two Types: Polar covalent are soluble Nonmetal + nonmetal Polar and nonpolar

Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling 1901 - 1994

Table of Electronegativities

Covalent Bonds Polar-Covalent bonds Nonpolar-Covalent bonds Electrons are unequally shared Electronegativity difference between .3 and 1.7 Nonpolar-Covalent bonds Electrons are equally shared Electronegativity difference of 0 to 0.3 IONIC Compounds have EN differences >1.7

Determining Bond Type: Bond EN Diff Type? More EN? S to H 2.5-2.1 Ca to Cl 1.0-3.0 Mg to O 1.2-3.5 F to F 4.0-4.0 Which of the above bonds has the LEAST ionic character? Which of the above bonds has the MOST ionic character?

The Octet Rule – Ionic Compounds Ionic compounds tend to form so that each atom, by gaining or losing electrons, has an octet of electrons in its highest occupied energy level.

Ionic Bonding: The Formation of Sodium Chloride Sodium has 1 valence electron Chlorine has 7 valence electrons An electron transferred gives each an octet Na 1s22s22p63s1 Cl 1s22s22p63s23p5

Ionic Bonding: The Formation of Sodium Chloride This transfer forms ions, each with an octet: Na+ 1s22s22p6 Cl- 1s22s22p63s23p6

Ionic Bonding: The Formation of Sodium Chloride The resulting ions come together due to electrostatic attraction (opposites attract): Na+ Cl- The net charge on the compound must equal zero

Examples of Ionic compounds All salts, which are composed of metals bonded to nonmetals, are ionic compounds and form ionic crystals. Examples: MgCl2 Na2O KI CaO BaS LiF

Monatomic Cations Name H+ Hydrogen Li+ Lithium Na+ Sodium K+ Potassium Mg2+ Magnesium Ca2+ Calcium Ba2+ Barium Al3+ Aluminum

Monatomic Anions Name F- Fluoride Cl- Chloride Br- Bromide I- Iodide O2- Oxide S2- Sulfide N3- Nitride P3- Phosphide

Sodium Chloride Crystal Lattice Ionic compounds form solids at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

Representation of Components in an Ionic Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

Ionic Bonds Electrons are transferred Electronegativity differences are generally greater than 1.7 The formation of ionic bonds is always exothermic!

Metallic Bonding The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons Electrons are “delocalized” Occurs between like metals in the D block Example: Cu to Cu

Properties of Metals Metals are good conductors of heat and electricity Metals are malleable Metals are ductile Metals have high tensile strength Metals have luster

Bond Energy Bonding Energy gives the strength of covalent compounds… Bonding Energy is the energy required to break a bond For ionic bonds, we measure Lattice Energy – the amount of energy it takes to form the bond

The Octet Rule – Covalent Compounds Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level. Diatomic Fluorine

Hydrogen Chloride by the Octet Rule

Electron Dot Notation

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.

Completing a Lewis Structure -CH3Cl Make carbon the central atom 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 Two pairs of shared electrons

Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons

Bond Length and Bond Energy Length (pm) Energy (kJ/mol) C - C 154 346 C=C 134 612 CC 120 835 C - N 147 305 C=N 132 615 CN 116 887 C - O 143 358 C=O 799 CO 113 1072 N - N 145 180 N=N 125 418 NN 110 942

Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures.

Resonance in Ozone Neither structure is correct.

VSEPR Model (Valence Shell Electron Pair Repulsion) The structure around a given atom is determined principally by minimizing electron pair repulsions.

Predicting a VSEPR Structure Draw Lewis structure. Put pairs as far apart as possible. Determine positions of atoms from the way electron pairs are shared. Determine the name of molecular structure from positions of the atoms.

Table – VSEPR Structures

Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

Intermolecular Forces Forces between molecules 3 Types: Hydrogen bonding Dipole-dipole forces London Disperson Forces

Hydrogen Bonding Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests.

Hydrogen Bonding in Water

Hydrogen Bonding between Ammonia and Water

Dipole-Dipole Attractions Attraction between oppositely charged regions of neighboring molecules.

The water dipole

London Dispersion Forces The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. London forces increase with the size of the molecules. Fritz London 1900-1954

Relative magnitudes of forces The types of bonding forces vary in their strength as measured by average bond energy. Strongest Weakest Covalent bonds (400 kcal) Hydrogen bonding (12-16 kcal ) Dipole-dipole interactions (2-0.5 kcal) London forces (less than 1 kcal)

The Blending of Orbitals Hybridization The Blending of Orbitals

Lets look at a molecule of methane, CH4. We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds. Lets look at a molecule of methane, CH4. Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it.

Carbon ground state configuration What is the expected orbital notation of carbon in its ground state? Can you see a problem with this? (Hint: How many unpaired electrons does this carbon atom have available for bonding?)

Carbon’s Bonding Problem You should conclude that carbon only has TWO electrons available for bonding. That is not not enough! How does carbon overcome this problem so that it may form four bonds?

Carbon’s Empty Orbital The first thought that chemists had was that carbon promotes one of its 2s electrons… …to the empty 2p orbital.

However, they quickly recognized a problem with such an arrangement… Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom. A Problem Arises

This would mean that three of the bonds in a methane molecule would be identical, because they would involve electron pairs of equal energy. But what about the fourth bond…? Unequal bond energy

The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron. Such a bond would have slightly less energy than the other bonds in a methane molecule. Unequal bond energy #2

This bond would be slightly different in character than the other three bonds in methane. This difference would be measurable to a chemist by determining the bond length and bond energy. But is this what they observe? Unequal bond energy #3

The simple answer is, “No”. Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane. Chemists have proposed an explanation – they call it Hybridization. Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy. Enter Hybridization

In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals. These new orbitals have slightly MORE energy than the 2s orbital… … and slightly LESS energy than the 2p orbitals. sp3 Hybrid Orbitals

sp3 Hybrid Orbitals Here is another way to look at the sp3 hybridization and energy profile…

sp Hybrid Orbitals While sp3 is the hybridization observed in methane, there are other types of hybridization that atoms undergo. These include sp hybridization, in which one s orbital combines with a single p orbital. Notice that this produces two hybrid orbitals, while leaving two normal p orbitals

sp2 Hybrid Orbitals Another hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel. Notice that one p orbital remains unchanged.