Chemical Bonding

1. Ch. 6 – Part 1: Covalent Bonding Chemical Bonding

2. What is a chemical bond? • mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms together • Why don’t noble gases do this? • Already have filled s and p orbitals • stable octet: 8 valence e- (or 2 if, you’re helium) Atoms that don’t have a stable octet are more reactive

3. Key Point #1: By forming bonds with each other, most atoms reducetheir potential energy, becoming more stable. • Because forming bonds involves the rearrangement of atoms to recreate new substances, this is considered a chemical change. All chemical changes involve energy changes.

4. Types of Bonds • What type of bonds can be formed? • Ionic bond • Covalent bond • Nonpolar covalent • Polar covalent • Ionic bonding: bonds that result from electrical attractions between cations and anions • 1 atom losses electrons • 1 atom gains electrons • Covalent bonding: sharing of electrons between 2 or more atoms

5. Key Point 2: Rarely is bonding between atoms purely ionicor purely covalent. • Instead, it usually falls somewhere between the two extremes. Why? • Key Point 3: The extent of ionic or covalent bonding between two atoms can be estimated by calculating the differencein each elements’ electronegativity.

6. Covalent Bonding • Large difference in E.N.: bond has more ionic character • Small difference in E.N: bond has more covalent character Types of Covalent Bonds • Non-polar covalent bonding: both electrons equally shared between atoms • Polar covalent bonding: unequal attraction for the shared electrons

7. 6.1 Practice Worksheet Part 1 The property of electronegativity, which is the measure of an atom’s ability to attract electrons, can be used to predict the degree to which the bonding between atoms of two elements is ionic or covalent. The greaterthe electronegativity difference, the more ionicthe bonding is.

8. Increasing difference in electronegativity Nonpolar Covalent Polar Covalent share e- partial transfer of e- • If the calculated electronegative difference is… • > 1.7 : ionic bond is formed • > 0.3 , < 1.7 : polar-covalent bond • 0 – 0.3 : non-polar covalent bond Ionic transfer e-

9. 3.0 1.8 Ionic 1.2 3.5 Polar covalent 1.4 2.1 3.0 .5 2.5 Polar covalent 3.0 2.1 .9 Polar covalent 2.5 2.5 Non-polar covalent 0 .8 4.0 3.2 Ionic .9 3.0 Ionic 2.1 2.1 2.1 0 Non-polar covalent

10. Polyatomic Ions • It is also possible if a compound contains polyatomic ions, for both types of bonding to be present. • Monatomic Ions: Fe2+ , Na+, Cl- • Polyatomic Ions: PO43-, NH4+ , NO-1 • Groups of atoms are bonded covalent together, but because of few or more than expected valence electrons they have an overall charge (so they can also bond ionically with other ions) • Ex: Ca2+ and SO42-  CaSO4 (metal & diff. nonmetals)

11. Classify the following as ionic, covalent, or both Both Ionic 1. CaCl2 = __________ (metal & nonmetal) 2. CO2 = __________ (nonmetal & nonmetal) 3. MgO = __________ (metal & nonmetal) 4. HCl = ___________ (nonmetal & nonmetal) 5. BaSO4 = ___________ (metal & diff. nonmetals) 6. H2O = ____________ (nonmetal & nonmetal) 7. SO3 = ___________ (nonmetal & nonmetal) 8. AlPO4 = ___________ (metal & diff. nonmetals) Covalent Covalent Ionic Covalent Covalent Both

12. Section 6.2 Covalent Bonding

13. What is a molecule? • Neutral group of atoms that are held together by covalent bonds. • Chemical formula: indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.

14. Formation of Covalent Bonds • The electrons of one atom and protons of the other atom attracteach another. • The two nuclei and two electrons repel each other. • These two forces cancel out to form a covalent bond at a length where the potential energy is at a minimum.

16. Bond Length vs. Bond Energy • Bond length (pm):distance between two bonded atoms at their minimum potential energy • Bond energy (KJ/mol): energy required to break a chemical bond and form neutral isolated atoms. • Breaking bonds: absorbs (requires) energy • Forming bonds: releases energy • A few general principles – • Bond length increases with atomic size • Bond length and bond energy are inverses • As you increase the number of bonds (between 2 atoms), bond energy increases while bond length decreases

17. Bond Energies & Bond Lengths

18. Octet Rule • Octet Rule: chemical compounds tend to form so that each atom has an octet of e-’s in its highest occupied energy level • Exceptions to the octet rule: • Atoms that cannot fit eight electrons • Atoms that can fit more than eight electrons • Hydrogen forms bonds surrounded by only two electrons. • Boron has just three valence electrons, so it tends to form bonds in which it is surrounded by six electrons. • Phosphorus, Sulfur, & Xenon can form bonds with expanded valence, involving more than eight electrons.

19. Writing Lewis Structures • Lewis Structures: formulas in which atomic symbols represent nuclei and inner shells, dot-pairs/dashes represent shared electron pairs, and dots adjacent to only one atomic symbol represent unshared pairs of electrons. • Structural formulas: represent kind, number, arrangement, and bonds in a molecule but not does not include unshared pairs of electrons.