Bonding

1. Bonding Chapter 6

2. Chemical Bonds • All types of chemical bonds form due to the mutual attraction between the nucleus of one atom and valence electrons of another atom.

3. Ionic Bonds • Ionic bonds form when electrons are transferred from one atom to another. • A metal will lose electrons and form a cation • A nonmetal will gain electrons and form an anion

4. Ion charges • When an ion forms it usually has an octet of electrons in its valence shell. • An octet (8 electrons) would result from a full s and p sublevel in the outermost energy level. What is the charge on the fluorine ion? What is the charge on the strontium ion? F- Sr2+

5. Ionic Compounds • Ionic compounds form when the cation and anion stick together due to the electrostatic attraction between the oppositely charged ions. • The overall compound will be neutral. NaBr AlCl3 Ca3N2

6. Ionic Crystals • Ionic solids contain cations and anions that are arranged in an organized three dimensional crystal lattice. • Within the lattice, each cation is surrounded by anions and vice-versa.

7. Ionic Bond Strength • Lattice energy is the energy released when one mole of an ionic compound is formed from gaseous ions. • The greater the lattice energy the stronger the ionic bond.

8. (Honors) • Lattice energy is affected by two variables; the ions’ charges and the ions’ radii. • The greater the charges, the greater the LE. LENaCl LEMgCl2 • The smaller the radii, the greater the LE. LENaCl LELiCl < <

9. Covalent Bonds • Covalent bonds result from the sharing of electrons between two atoms • Covalent bonding occurs when nonmetal atoms are bonded together • There are two types of covalent bonds; and nonpolar covalent bonds polar covalent bonds

10. Nonpolar Covalent Bonds • In a nonpolar covalent bond, electrons are shared equally between the atoms. • This results in a balanced electron distribution

11. Polar Covalent Bonds • In a polar covalent bond, electrons are shared unequally • This results in an uneven charge distribution (one end is positive and the other is negative)

12. Electronegativity • Def: the ability of an atom in a compound to attract electrons from another atom • Values given on page 161 in text

13. Electronegativity differences between atoms in a compound can help us predict the type of bond present Nonpolar ~ 0-0.3 Polar~ 0.3-1.7 ionic ~ 1.7 +

14. Lewis Dot Diagrams • Show how atoms are bonded together in a molecule. • A dot represents a valence electron • A pair of dots between two atom symbols represents a bond between the atoms

15. Drawing Lewis Diagrams • Step 1: determine the total number of valence electrons in the molecule • Step 2: Arrange element symbols to form a skeleton structure • If carbon is present, it should go in the center • Hydrogen atoms should always be on the edge of the molecule. Halogens usually are on an edge. • Otherwise, molecules tend to be somewhat symmetrical

16. Step 3: Put one pair of electrons between adjacent atoms • Step 4: Add unshared pairs (lone pairs) to the outside of each atom until each atom has a total of 8 electrons (an octet), except hydrogen will have 2 (a duet).

17. Step 5: Count all the electrons you have drawn • If that equals the amount that you had available from Step 1, you are done! • If you have drawn too many dots, you must make double and/or triple bonds until you have the proper number of electrons present. For every additional bond made, you will remove two lone pairs of electrons

18. HONORS – Exceptions to the Octet Rule • Some molecules will not obey the octet rule. (They often have a noble gas in them.) • If you complete Step 4 and in Step 5 you have fewer electrons shown in your diagram that you counted from Step 1, you will need to place all extra electrons (in pairs) around the central atom.

19. VSEPR • Valence Shell Electron Pair Repulsion Theory • States that regions of electrons around the central atom of a molecule will be arranged as far apart as possible • Determines the shapes of molecules

20. Hybridization • When a molecule is formed, the atomic orbitals (s, p, etc) from the individual atoms combine with each other to make new orbitals called hybrid orbitals • Hybridization must occur because the shapes of atomic orbitals does NOT explain the molecular shapes.

21. sp3 hybridization • If there are four regions of high electron density around the central atom, then 1 s-orbital has hybridized with 3 p-orbitals resulting in 4 sp3 orbitals arranged tetrahedrally

22. sp2 hybridization • If there are three regions of high electron density around the central atom, then 1 s-orbital has hybridized with 2 p-orbitals resulting in 3 sp2 orbitals arranged in a trigonal pyramid

23. sp hybridization • If there are two regions of high electron density around the central atom, then 1 s-orbital has hybridized with 1 p-orbital resulting in 2 sp orbitals arranged linearly

24. Sigma vs Pi Bonding • A sigma (s) bond occurs when electron density lies between the nuclei of the bonded atoms. • All single bonds are s-bonds.

25. Pi (p) bonding occurs when orbitals overlap above and below the line between the nuclei

26. All double bonds contain 1 s and 1 p bond • All triple bonds contain 1 s and 2 p bonds • The s-bond forms from the overlap of the hybrid orbitals while the p-bonds form from the overlap of the unhybridized p-orbitals

27. Molecular Polarity Cl H • Dipoles are polar molecules. • A dipole exists when there is an asymmetrical charge distribution in the molecule • For a molecule to be a dipole it must contain at least one polar bond and must have an asymmetrical shape O H H

28. Examples of Dipoles

29. Examples of Nonpolar Molecules

30. Intermolecular Forces • def: The forces of attraction that exist between molecules. • Intermolecular forces are weaker than a covalent or ionic bond • There are three types of intermolecular forces: • Dipole-Dipole Forces • Hydrogen Bonding • London Dispersion Forces

31. Dipole-Dipole Forces • Exists between two dipoles (polar molecules) • The positive end of one molecule is attracted to the negative end of another molecule • The greater the polarity of the molecule, the stronger the dipole-dipole force

32. Hydrogen Bonding • Def: a hydrogen atom bonded to a highly electronegative atom (F, O, or N) is attracted to an unshared pair of electrons on the F, O, or N of a nearby molecule. • H-bonding is a very strong type of dipole-dipole force.

33. London Dispersion Forces • A weak intermolecular force that results from a momentary (temporary) dipole. • All molecules both polar and nonpolar experience LDFs, but this is the ONLY type of intermolecular force that nonpolar molecules experience. • The larger the molecular mass, the stronger the LDFs. A temporary uneven electron distribution results in a momentary dipole one momentary dipole can cause another nearby molecule to also have a momentary dipole

34. Metallic Bonding • The bonding in metals consists of a crystal lattice of metal ions immersed in a sea of delocalized, mobile valence electrons • These mobile electrons are able to flow throughout the sample allowing metals to conduct electricity