2. Mrs. Schultz Ch 8 Covalent bonding

3. In nature only the noble gas elements in Group 18 (8A) exist as _______ atoms. Considered to be MONATOMIC (single atoms)

4. Covalent Bonding – results from the sharing of electron pairs between two atoms (between nonmetals); forms molecules A molecule is a neutral group of atoms that are held together by covalent bonds A single molecule of a chemical compound is an individual unit capable of existing on its own.

5. A chemical compound whose simplest units are molecules is called a molecular compound. ex/ O 2 (oxygen gas); H 2 O (water) Molecular Formula Chemical formula for a compound that has covalent bonds.

6. Recall: What does the word part “di-” mean? Diatomic molecule – Two of the same atoms are bound covalently. (di = 2 ;Greek) Example: 7 diatomic gases *starting with element 7 (N) – “makes a seven on the Periodic Table” Br. HONClIF Br 2 H 2 O 2 N 2 Cl 2 I 2 F 2

7. octet rule: Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. atoms attain the electron configurations of noble gases. Ex/ Draw the electron dot structure: F 2 Ex/ Draw the e - dot structure, N 2

8. Recall: What is the trend for electronegativity? Draw the electron dot structures 1) NF 3 *Note: Least EN atom is central 2) SBr 2

9. Chapter 12 8 Electronegativity, Continued Electronegativity increases as you go left to right across a period. Electronegativity increases as you go from bottom to top in a family.

10. Polyatomic ions (many atoms w/charge + or -) Ex/ PO 4 3- phosphate ion and NH 4 + ammonium

11. Bond Dissociation Energies BDE: energy required to break the bond between two covalently bonded atoms. A large bond dissociation energy corresponds to a strong covalent bond. Look at page 236 table 8.3 Strong carbon-carbon bonds help explain the stability of carbon compounds. Write down the 3 strongest covalent bonds and the 3 weakest see pg. 236 chart

12. Bond strengths Strongest: C triple bond to O; C triple bond to C; C double bond to O Weakest: O single bond to O; N single bond to N, Cl single bond to Cl

13. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. When two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei, a sigma bond is formed.  Symbol – (Greek) sigma (σ). Molecular Orbitals – when atomic orbitals overlap to form orbitals that apply to the entire molecule Sigma Bonds s atomic orbital Bond axis Sigma-bonding molecular orbital  represents the nucleus

14. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. side-by-side overlap of p orbitals produces what are called pi molecular orbitals. Molecular Orbitals p atomic orbital Pi-bonding molecular orbital  represents the nucleus Pi Bonds

15. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. In a pi bond (Greek  ), the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms. Because orbitals in pi bonding overlap less than in sigma bonding, pi bonds tend to be weaker than sigma bonds. Molecular Orbitals Pi Bonds

16.  Polar covalent bond – bonded atoms have an unequal attraction for the shared electrons.  ex/ water H 2 O  Nonpolar covalent bond – bonding electrons are shared equally; balanced distribution of electrons  Ex/ hydrogen gas H 2  Note : Polar is a word that means uneven distribution of electrons.

17. Determination of Bond type? Bond type: difference in electronegativity? Ionic : greater than or equal to 2.0 Moderately Polar covalent :.4 – 1.0 Very polar covalent: 1.0 – 2.0 Nonpolar covalent: 0 -.4 p. 181 Table 6.2 What type of bonding would be expected between the following pairs of elements? a) Cl and K Ans. Ionic (EN diff = 3.0 -.8 = 2.2)

18. b) O and H Ans. 3.5 – 2.1 = 1.4 very polar covalent c) O and O Ans. 3.5 – 3.5 = 0 nonpolar covalent

19. CW: Fill in the following chart with a partner bonding between: EN diff.: Bond type: More EN atom: 1. Na and Br 2. Mg and O 3. N and N 4. C and O 5. Ba and O 6. P and O 7. Ca and Cl 8. Be and F 9. S and Cl 10. N and O bonding between: EN diff.: Bond type: More EN atom: 1. Na and Br 2. Mg and O 3. N and N 4. C and O 5. Ba and O 6. P and O 7. Ca and Cl 8. Be and F 9. S and Cl 10. N and O

20. Ch 8 VSEPR theory – stands for “valence- shell electron pair repulsion” all sets of valence electrons will repel each other and will be oriented as far apart as possible. This theory is used to determine shape of molecules (or MOLECULAR GEOMETRY)

21. Formula: e - dot model shape? polar? O 2 H 2 O CF 4 CH 2 Cl 2 NH 3 BBr 3

22. Chapter 12 21 Electronegativity, Continued Electronegativity increases as you go left to right across a period. Electronegativity increases as you go from bottom to top in a family.

23. Hybridization: Mixing of two or more atomic orbitals of similar energies on the same atom to produce new hybrid atom orbitals of equal energy. CH 4 1s 2 2s 2 2p 2 Draw orbital diagram: The paired 2s orbital and the 2p mix to form 4 sp 3 orbitals. (s,p,p,p) This is why carbon forms 4 bonds (hybridization)

24. Intermolecular forces Forces of attraction between molecules. Weaker than ionic or covalent bonds Hydrogen bonding – “Not a true bond” The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule. Ex/ water molecules:

25. van der Waals forces (2 types) Dipole interactions: Forces of attraction between polar molecules. Ex/ammonia Dispersion forces: forces of attraction between nonpolar atoms or molecules that are created by temporary dipoles. ex/ Noble gases

26. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. This table summarizes some of the characteristic differences between ionic and covalent (molecular) substances. Characteristics of Ionic and Molecular Compounds CharacteristicIonic CompoundMolecular Compound Representative unitFormula unitMolecule Bond formation Transfer of one or more electrons between atoms Sharing of electron parts between atoms Type of elementsMetallic and nonmetallicNonmetallic Physical stateSolidSolid, liquid, or gas Melting pointHigh (usually above 300°C)High (usually below 300°C) Solubility in waterUsually highHigh to low Electrical conductivity of aqueous solution Good conductorPoor to nonconducting Intermolecular Attractions and Molecular Properties