Covalent Bonding

1. Covalent Bonding Chapter 8

2. Covalent Bonding • Covalentbonds form when atoms share electrons • Atoms that lack the necessary electrons to form a stable octet are most likely to form covalent bonds. • Covalent bonds are formed between two nonmetals

3. Molecules and Molecular Compounds • A molecule is a neutral group of atoms joined together by covalent bonds. • Diatomic Elements—a molecule that contains two atoms of the same element. • Ex : Oxygen, Hydrogen, Nitrogen, Fluorine, Chlorine, Bromine, Iodine. • Molecules can also be made of atoms of different elements • A compound composed of molecules is called a molecular compound. • Water is an example of a molecular compound.

4. Molecular Formula • A molecular formula is the chemical formula of a molecular compound. It shows how many atoms of each element a substance contains.

5. Bonding and Interactions Molecular Compounds In molecular compounds, bonding occurs when atoms share electrons. Representative unit of a molecular compound is a molecule. Molecular compounds are composed of atoms of two or more nonmetals. Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds. Many molecular compounds are gases or liquids at room temperature. Ionic Compounds • In ionic compounds, bonding occurs when electrons are transferred between atoms. • The smallest representative unit , is a formula unit. • Ionic compounds are formed from a metal combined with a nonmetal. • Ionic Compounds tend to high melting points. • Ionic compounds are generally crystallinesolids at room temperature.

6. Comparing Molecular and Ionic Compounds Water, which is a molecular compound, and sodium chloride, which is an ionic compound, are compared here. Array of sodium ions and chloride ions Collection of water molecules Formula unit of sodium chloride Molecule of water Chemical formula H2O Chemical formula NaCl

7. Single Covalent Bonds • In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. • Single covalent bond :Two atoms held together by sharing one pair of electrons. • Structural formula represents the covalent bonds as dashes and shows the arrangement of covalently bonded atoms (H—H ). • A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair or a nonbondingpair. • In F2, each fluorine atom has three unshared pairs of electrons.

8. Single Covalent Bonds • Some Examples of Single Covalent Bonds

9. Double and Triple Covalent Bonds • Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two or three pairs of electrons.

10. Coordinate Covalent Bonds • Coordinate covalent bonds occur when one atom donates both of the electrons that are shared between two atoms • Coordinate covalent bonds are also called DativeBonds • Polyatomic ion is a tightly bound group of atoms that has a positive or negative charge and behaves as a unit ( H3O+, NH4+) .

11. Exceptions to the Octet Rule • The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. There are also molecules in which an atom has less, or more, than a complete octet of valence electrons. • Some molecules with an even number of valence electrons, such as some compounds of boron, also fail to follow the octet rule (BF3). • A few atoms, especially phosphorus and sulfur, expand the octet to ten or twelve electrons ( SF6, PCl5 ).

12. Bond Dissociation Energies • The total energy required to break the bond between 2 covalently bonded atoms • High dissociation energy usually means the chemical is relatively unreactive, because it takes a lot of energy to break it down • The units for this energy are often given in kJ/mol, which is the energy needed to break one mole of bonds.

13. Resonance • Resonance structures are structures that occur when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion. • Carbonate ion ( CO32-) • Benzene (C6H6) • Nitrate ion (NO3- )

14. Molecular Orbitals • The model for covalent bonding assumes the orbitals are those of the individual atoms = atomic orbital • Orbitals that apply to the overall molecule, due to atomic orbital overlap are the molecularorbitals • A bondingorbital is a molecular orbital that can be occupied by two electrons of a covalent bond • SigmaBond: When two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei. Its symbol is the Greek letter sigma (σ). • Pi Bond: When two atomic orbitals combine to form molecular orbital that is likely to be found above and below the bond axis (weaker than sigma). Its symbol is the Greek letter pi ().

15.  represents the nucleus p atomic orbital p atomic orbital Pi-bonding molecular orbital Molecular Orbitals

16. VESPR Theory • VSEPR theory states that the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible. • Valence Shell Electron Pair Repulsion • The name tells you the theory: • Valence shell = outside electrons. • Electron Pair repulsion = electron pairs try to get as far away as possible from each other. • Based on the number of pairs of valence electrons, both bonded and unbonded , VSEPRtheorypredicts the three dimensional shape of molecules. • Can determine the angles of bonds. • Please pick up the molecular geometry chart

18. Hybrid Orbitals • The VSEPR theory works well when accounting for molecular shapes, but it does not help much in describing the types of bonds formed. • Orbital hybridization provides information about both molecular bonding and molecular shape. • In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals • SP3 Hybridization : involves single bonds. Ex: Methane CH4 • SP2Hybridization : involves double bonds. Ex : Ethene • SPHybridization : involves triple bonds. Ex : Acetylene

19. Electronegativity Difference • If the difference in electronegativities is between: • 1.7 to 4.0: Ionic • 0.3 to 1.7: Polar Covalent • 0.0 to 0.3: Non-Polar Covalent • Example: NaCl • Na = 0.8, Cl = 3.0 • Difference is 2.2, so • this is an ionic bond!

20. H—Cl Bond Polarity • Consider HCl H = electronegativity of 2.1 ; Cl = electronegativity of 3.0 • the bond is polar • the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge • A molecule that has two poles is called a dipolar molecule, or dipole.

21. Molecular Polarity • Molecular Polarity depends on: • the relative electronegativities of the atoms in the molecule • The shape of the molecule • Molecules that have symmetrical charge distributions are usually non-polar • If there are equal polar bonds that balance each other around the central atom, then the overall molecule will be NONPOLAR with no dipole moment, even though the bonds within the molecule may be polar. - Polar bonds cancel - There is no dipole moment - Molecule is non-polar - Polar bonds do not cancel - There is a net dipole moment - The molecule is polar

22. Intermolecular Forces • Intermolecular forces are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature. • The weakest attractions is van der Waals forces and there are two types : 1. Dispersionforces : weakest of all, caused by motion of e- ; increases as # e- increases • halogens start as gases; bromine is liquid; iodine is solid – all in Group 7A 2. DipoleInteractions: Occurs when polar molecules are attracted to each other. • Slightly stronger than dispersion forces. • Opposites attract, but not completely hooked like in ionic solids.

23. Hydrogen bond Intermolecular Forces • HydrogenBonds : It is the attractive force caused by hydrogen bonded to N, O, F • N, O, and F are very electronegative, so this is a very strong dipole. • Hydrogen bonds are extremely important in determining the properties of water and biological molecules such as proteins. • This is the strongest of the intermolecular forces.