Covalent Bonding The Nature of Covalent Bonding Bonding Theories Polar Bonds and Molecules

Ch. 16.1 The Nature of Covalent Bonding Covalent bonds occur between atoms that are “sharing” electrons Form covalent compounds There is a kind of “tug of war” for the electrons There can be single, double and triple covalent bonds Single bond – a bond in which 2 atoms share a pair of electrons Double bond – bond that involves 2 shared pairs of e- Triple bond – bond that involves 3 shared pairs of e-

Ch. 16.1 The Nature of Covalent Bonding When writing formulas for covalent bonds: Pairs of electrons are represented by a dash Called structural formulas Chemical formulas that show the arrangement of atoms in molecules and polyatomic ions Each dash between atoms in structural formulas indicates a pair of electrons Dashes are never used to show ionic bonds Chemical formulas Chemical formulas of ionic cmpds describe formula units Chemical formulas of covalent cmpds describe molecules

Ch. 16.1 The Nature of Covalent Bonding Combinations of atoms of non-metallic atoms are likely to form covalent bonds Groups 4A, 5A, 6A, and 7A Summarized by G. Lewis in the octet rule sharing of e- occurs if atoms achieve noble gas config. H2 is an exception to this rule

Ch. 16.1 The Nature of Covalent Bonding Group trends Halogens form single covalent bonds in their diatomic molecules (ex: F – F ) Chalcogens form double covalent bonds in their diatomic molecules (ex: O = O) Phicogens form triple covalent bonds in their diatomic molecules ( N = N ) The Carbon group tends to form 4 bonds with other atoms

Ch. 16.1 The Nature of Covalent Bonding Covalent bonding can be explained using electron configurations and orbital boxes Show H2O, NH3 and CH4 on board Double and triple covalent bonds Oxygen forms a double bond in a diatomic molecule It is an exception to the octet rule, 2 unpaired e- (p. 442) Nitrogen forms a triple bond in a diatomic molecule Satisfies the octet rule, all e- are paired Multiple covalent bonds can form between unlike atoms (ex: CO2, CH3OH)

Ch. 16.1 The Nature of Covalent Bonding Coordinate covalent bonds Carbon monoxide (CO) is a covalent bond that cannot be explained by “normal” covalent bonding (p. 444) When C and O form a traditional double bond, the O atom satisfies the octet rule, but the C atom does not For both atoms to achieve a stable octet, oxygen donates one of its unshared pairs of e- for bonding Called a coordinate covalent bond The structural formula shows an arrow pointing from the atom donating the e- pair to the atom receiving the e- pair Once formed, it is like any other covalent bond

Ch. 16.1 The Nature of Covalent Bonding Coordinate covalent bonds Most polyatomic cations and anions contain covalent and coordinate covalent bonds The polyatomic ammonium ion (NH4+) has a coordinate covalent bond Polyatomic ions contain both covalent and ionic bonds (see pg. 446 H3O+)

Ch. 16.1 The Nature of Covalent Bonding Bond dissociation energies The release of heat in a bond suggests the product is more stable that the reactants The total energy required to break the bond between covalently bonded atoms is called the bond dissociation energy The higher the dissociation energy, the stronger the energy

Ch. 16.1 The Nature of Covalent Bonding Resonance Resonance structures are those that occur when it is possible to write 2 or more valid electron dot formulas that have the number of electron pairs for a molecule or ion Scientists once believed that the e- pairs flipped back and forth (resonated) between structures The actual bonding is a hybrid of the extremes represented by the resonance structures Double-headed arrows are used to connect resonance structures (see pg. 449 O3)

Ch. 16.1 The Nature of Covalent Bonding Exceptions to the octet rule For some molecules or ions, it is impossible to write structures that satisfy the octet rule This occurs when the total number of valence e- is an odd number NO2 is an example (see pg. 450) – an unpaired e- exists in each of the resonance structures

Ch. 16.1 The Nature of Covalent Bonding Magnetism and bonding Electrons are spinning charged particles with a magnetic field Paired e-, spinning in opposite directions, cancel out the magnetic field (diamagnetic substances have all electrons paired) These are weakly repelled by an external magnetic field Unpaired e- result in an external magnetic field (paramagnetic substances contain 1 or more unpaired e-) These are strongly attracted to an external magnetic filed

Ch. 16.1 The Nature of Covalent Bonding Ferromagnetism is the strong attraction of Fe, Co, and Ni to magnetic fields These ions all have unpaired electrons Electrons line up in magnetic domains, remain in that state permanently Covalent bond lengths Single bonds are longest, double bonds are intermediate, and triple bonds are the shortest Exceptions to the octet rule There are many molecules that do not follow the octet rule, yet exist as stable bonds A few atoms (esp. P and S) sometimes expand the octet to include 10 or 12 electrons (see pg. 451)

Ch. 16.2 Bonding Theories Molecular orbitals When 2 atoms combine, their atomic orbitals overlap to produce molecular orbitals (orbitals that apply to the entire molecule) The molecular orbital model of bonding requires that the number of molecular orbitals equal the number of overlapping atomic orbitals When 2 atomic orbitals overlap, 2 molecular orbitals are created One is called a bonding orbital, the other is called an anti- bonding orbital The anti-bonding orbital has a higher energy that the atomic orbitals from which it formed

Ch. 16.2 Bonding Theories Molecular orbitals When H2 forms, the 1s atomic orbitals overlap 2 electrons are available for bonding (see pg. 452) The energy of the e- in the bonding molecular orbital is lower than the e- in the atomic orbitals of the separate H atoms Electrons seek the lowest energy level, so they fill the bonding molecular orbital This makes a stable covalent bond between the H atoms The anti-bonding orbital is empty Sigma and pi bonds are caused by the overlapping of “s” and “p” orbitals

Ch. 16.2 Bonding Theories VSEPR Theory The valence-shell electron-pair repulsion theory States that because e- pairs repel, molecular shape adjusts so that the valence-electron pairs are as far apart as possible Unshared pairs are important when you are trying to predict the shapes of molecules Unshared pairs are held more closely to the atom Unshared pairs strongly repel bonding pairs of e-, pushing them closer together than might be expected

Ch. 16.2 Bonding Theories Molecular geometries Tetrahedral Pyramidal all 4 bond angles are 109.5o CH4 is an example Pyramidal bond angles between shared pairs is 107o NH3 is an example Bent bond angles are 105o H2O is an example

Ch. 16.2 Bonding Theories Molecular geometries Trigonal planar Linear Bond angles are 120o An example is BF3 Linear Bond angles are 180o CO2 is an example Trigonal bipyramidal Bond angles of 120o and 90o An example is PF5

Ch. 16.2 Bonding Theories Hybrid orbitals In orbital hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals One 2s and three 2p orbitals of a carbon atom overlap to form an sp3 hybrid orbital These are at the tetrahedral angle of 109.5o Four sp3 orbitals of carbon overlap with the 1s orbitals of the four hydrogen atoms This allows for a great deal of overlap, which results in the formation of 4 C-H sigma bonds These are unusually strong covalent bonds

Ch. 16.3 Polar Bonds and Molecules Bond polarity Covalent bonds involve the sharing of electrons However, they can differ in how the bonds are shared Depends on the kind and number of atoms joined together When electrons are shared equally, a nonpolar covalent bond is formed When the atoms share the electron unequally, a polar covalent bond is formed The more electronegative element will have the stronger electron attraction and will acquire a slightly negative charge The less electronegative element will acquire a slightly positive charge

Ch. 16.3 Polar Bonds and Molecules Bond polarity The lower case Greek letter delta is used to show partial positive and negative charges (p.461) Table 16.4 (p. 462) shows how the size of the electronegativity difference between bonded atoms can predict the type of bond that will form These are the most probable bonds that will form The boundary between covalent and ionic is not always distinct Polar molecules have one end that is slightly positive and one end that is slightly negative A molecule with 2 different poles is called a dipole

Ch. 16.3 Polar Bonds and Molecules Attractions between molecules Molecules are attracted to each other by a variety of forces These are intermolecular forces They are weaker than either ionic or covalent bonds They are largely responsible for determining whether a substance is a solid, liquid or gas There are several kinds of intermolecular forces Van der Waals forces Dispersion forces Dipole interactions Hydrogen bonds

Ch. 16.3 Polar Bonds and Molecules Intermolecular forces Van der Waals forces The weakest attractions are known collectively as van der Waals forced Named after Johannes van der Waals Consist of two possible types Dispersion forces The weakest of all intermolecular forces Caused by the motion of electrons Strength increases as the number of electrons increases Dipole interactions Occurs when polar molecules are attracted to each other Slightly negative regions are attracted to slightly positive ones

Ch. 16.3 Polar Bonds and Molecules Intermolecular forces Hydrogen bonds Attractive forces in which a hydrogen covalently bonded to a to a very electronegative atom is also weakly bonded to to an unshared pair of another electronegative atom Always involves hydrogen The strongest of all the intermolecular forces Extremely important in determining the properties of water and biological molecules such as proteins

Ch. 16.3 Polar Bonds and Molecules Intermolecular attractions and molecular properties The physical properties of a compound depend on the type of bonding it has whether the bonding is ionic or covalent Table 16.5 (p. 465) Lists the characteristics of ionic and covalent compounds Melting and boiling points of most covalent compounds are low compared to ionic compounds A few molecular solids are very stable substances - network solids in which all atoms are covalently bonded to one another An example is diamond These compounds have extremely high melting and boiling points, or may simply vaporize above a certain temperature