Unit 6 – Chemical Bonds

1. Unit 6 – Chemical Bonds

2. Ionic, Covalent, and Metallic Bonding

3. Valence Electrons – electrons in the outermost shell of the atom, these are electrons that can participate in forming chemical bonds

4. On the back of your handout!

5. Why do chemical bonds occur and what is the rule called? What group of elements will the bonded elements be like if they have a full outer shell? Chemical bonding occurs between atoms so they can have 8 valence electrons. This is called the octet rule. This allows atoms to have an electron configuration similar to the noble gases, which have 8 valence electrons.

6. Ionic Bonds – result from electrostatic attractions between cation and anion, electrons are transferred so that the octet is happy! • Remember, cations give electrons away making them positive and are usually formed by metals • Anions take electrons making them negative and are usually formed by non-metals • So, ionic bonds are generally metals and non-metals • a metal and a polyatomic ion (K+ and NO3-) • a polyatomic ion and a non-metal (NH4+ and F-) • a polyatomic ion and a polyatomic ion (NH4+ and SO42-) • a metalloid and a metal

7. Ionic bond between lithium and fluorine

9. Covalent Bonds • When atoms have covalent bonding, the atoms share the electrons in the outermost shells so both atoms have an octet. This type of bonding occurs when two atoms that are both nonmetals form a bond. This most often includes elements from Groups 14-18.

10. H2 F2 CO2 Most of the time, the center atom will be the one written first in the chemical formula CF4

11. •• •• • • •• F F •• •• •• •• •• •• •• •• F F F F •• •• Single Covalent Bonds • Two atoms share one pair of electrons. • 2 electrons. • One atom may have more than one single bond. •• • • • • H H O •• •• •• H H •• O •• Tro's Introductory Chemistry, Chapter 10

12. •• •• • • • • O O •• •• •• •• O •• •• •• •• O O O Double Covalent Bond • Two atoms sharing two pairs of electrons. • 4 electrons. • Shorter and stronger than single bond. Tro's Introductory Chemistry, Chapter 10

13. •• •• • • • • N N • • N N Triple Covalent Bond • Two atoms sharing 3 pairs of electrons. • 6 electrons. • Shorter and stronger than single or double bond. N N •• •• •• •• •• Tro's Introductory Chemistry, Chapter 10

14. Metallic Bonds- result when metal atoms donate their valence electrons to an “electron sea” that binds the atoms together

15. Metallic bonds The layers of atoms in metal are hard to pull apart because of the electrons holding them together, explaining why metals are tough and have high melting points. But individual atoms are not held to any other specific atom, so atoms slip easily past one another making metals ductile and malleable. Delocalized (free) electrons can move rapidly in response to electric fields, that is why metals are good conductors of electricity. Free electrons can transmit kinetic energy rapidly, that is why metals are good conductors of heat.

16. Lewis Dot Structures and VSEPR

17. Lewis Dot Structures for Ions in Compounds • The chemical symbol for the element is surrounded by the number of valence electrons present in the ion • The whole structure is then placed within square brackets, with a superscript to indicate the charge on the ion. • The overall charge on the compound must equal zero Example: Lithium, Fluoride

18. Example: Sodium and Chlorine Example: Magnesium and Oxygen Example: Magnesium and Chlorine

19. Practice! • CaCl2 KBr • Al2S3 Li2O • NaFMgI2

20. Steps to drawing Lewis Structures for Covalent Bonds Determine the type (covalent bonds are always non-metal) and number of atoms in the molecule. Write the Lewis structure for each type of atom in the molecule. Determine the total number of valence electrons in the atoms to be combined. Remember, each atom wants a full outer shell (2 e- in the first energy level or 8 e- in any other energy level) Arrange the atoms to form a skeleton structure for the molecule. If carbon is present, it is the central atom. Otherwise, it is usually the first atom written, except Hydrogen is never the center. Then connect the atoms by electron-pair bonds. 5. Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by eight electrons.

21. Example: Cl2 Example: CH4 Example: NH3 Example: H2S

22. Practice! • HI H2O • NCl3 O2 • CO2 N2

23. Molecular Shapes or Molecular Geometries • While Lewis dot structures can tell us how the atoms in molecules are bonded to each other, they don’t tell us the shape of the molecule. • The most important thing to remember when predicting the shape of a molecule is that the molecule will have the shape that most minimizes electron repulsion. Valence shell electron pair repulsion (VSEPR) theory is a model in chemistry used to predict the shape of individual molecules based upon the extent of electron-pair electrostatic repulsion.

24. Two types of electron sets must be considered: • 1. electrons can exist in bonding pairs, which are involved in creating a single or multiple covalent bond, or • 2. nonbonding pairs (lone pairs), which are pairs of electrons that are not involved in a bond, but are localized to a single atom. These nonbonding electrons do play an important role in determining the shape of a molecule. • AXE Method for determining structure • 1. Arepresents the central atom and usually there is only one of that atom. If there are three or more atoms, then the central atom is always C (carbon) or the least electronegative atom (but never hydrogen). • 2. X subscriptrepresents the number of atoms bonded to the central atom • 3. Esubscriptrepresents the number of lone pairs of electrons surrounding the central atom • **Look on the Molecular Geometry Table on following page to find the corresponding shape

26. Example: HF Example: CO2 Example: BF3 Example: CF4 Example: H2O

27. Practice! • O2 SiO2 • NH3 H2O • CCl4 CO2 • N2 BCl3