1. Chemical Bonds and Molecular Geometry

2. Electrons  Valence electrons  Those electrons that are important in chemical bonding. For main-group elements, the valence electrons are those in the outermost principal energy level  Electron orbitals  A probability distribution map, based on the quantum mechanical model of the atom, used to describe the likely position of an electron in an atom; also an allowed energy state for an electron  Energy levels  The possible locations around an atom where electrons with particular energy values may be found

3. Valence Electrons

4. Electron Configuration  Electron configuration  A notation that shows the particular orbitals that are occupied by electrons in an atom  Each element has its own unique electron configuration  Electron configuration will change for ions of an element

5. Orbitals and Energy Levels  Electrons are found in shells around nucleus, numbered 1, 2, 3 etc  Each shell has sub-shells, labeled s, p, d, or f  Number of sub-shells are equal to the number of the shell

6. Sub-Shells  Each sub-shell contains a particular number of orbitals, where the electrons are located  Number of orbitals depends on the type of sub-shell  Each orbital can hold a maximum of 2 electrons Type of sub-shell (sub-level) Number of orbitalsMaximum number of electrons in the sub-shell (sub-level) s12 p36 d510 f714

7. Electron Configuration and the Periodic Table

8. Filling Electron Orbitals  When determining the electron configuration for an element, the lower energy levels will be filled first (aufbau principle)

9. Atomic Orbitals and Electrons in Principle Energy Levels Principle energy level Type of sublevel Number of orbitals in sublevel Maximum number of electrons 1s12 2s, p1+3=48 3s, p, d1+3+5=918 4s, p, d, f1+3+5+7=1632

10. Electron Configuration  Determine the electron configuration for the following elements 1. Li 2. Be 3. C 4. N 5. Na 6. Mg 7. Al 8. Si 9. P 10. B

11. History of Chemistry  Wolfgang Pauli, an Austrian theoretical physicist, proposed in 1924 the Pauli exclusion principle  No two electrons could exist in the same quantum state of an atom

12. Orbital diagrams  A diagram which gives information similar to an electron configuration, but symbolizes an electron as an arrow in a box representing an orbital, with an arrow’s direction denoting the electron’s spin

13. Filling Orbitals  Hund’s Rule

14. History of Chemistry  Friedrich Hund was a German physicist who developed Hund’s Rule for electron configuration  Hund’s Rule is that orbitals of sub-shells will be filled singly until all sub-shells have a single electron before filling doubly

15. Oxidation number  A positive or negative whole number that represents the “charge” an atom in the compound would have if all shared electrons were assigned to the atom with a greater attraction for those electrons  Label your periodic table with appropriate oxidation numbers

16. Quick Review  Ion  An atom or molecule with a net charge caused by the loss or gain of electrons  Electronegativity  The ability of an atom to attract electrons to itself  Ionization energy  The energy required to remove an electron from an atom or ion in its gaseous state  Electron affinity  The energy change associated with the gaining of an electron by an atom in its gaseous state

17. Lewis dot structures  A drawing that represents chemical bonds between atoms as shared to transferred electrons; the valence electrons of atoms are represented as dots  Ex.

18. Lewis Structures  Only valence electrons are represented  Families have the same number of valence electrons  The first four electrons are arranged one per side of the atomic symbol  Each additional dot is paired with a previous dot

19. Lewis Structure Assignment  Write a complete Lewis diagram for each:  Alkali metal  Alkaline earth metal  Families 13-18  Write them in order

20. Covalent Bonds  A chemical bond in which two atoms share electrons that interact with the nuclei of both atoms, lowering the potential energy of each through electrostatic interactions ; between two nonmetals

21. Ionic bond  A chemical bond formed between two oppositely charged ions, generally a metallic cation and a nonmetallic anion, that are attracted to one another by electrostatic forces

22. Formula Unit  Lowest whole number ratio of elements in a compound  Ex.  NaCl  CaCl 2  P 2 O 5  Must know the oxidation number before the formula unit can be determined

23. Metallic bond  The type of bonding that occurs in metal crystals, in which metal atoms donate their electrons to an electron sea, delocalized over the entire crystal lattice

24. Metallic Bond Theory and Metallic Properties  Most metals good conductors of thermal energy  cations held in crystal lattice form makes them move rapidly distributing the energy across the solid  Good conductors of electrical energy  Because electrons flow freely and move in response to electric potential energy  Metals are malleable and ductile  the cations move through the sea of electrons to reshape without breaking their bonds  High boiling points  much energy is required to break the tight metallic bonds

25. Alloys  Mixtures composed of two or more elements, at least one is a metal  Their properties are superior to the elements composing them  Common alloys   Brass (copper and zinc)  Bronze (copper and tin)  Stainless steel (iron and chromium)  Pewter (Tin and copper)  Sterling silver (silver and copper or zinc or other metal)

26. Molecular Geometry  The geometrical arrangement of atoms in a molecule  VSEPR theory (Valence shell electron pair repulsion)  A theory that allows prediction of the shapes of molecules based on the idea that electrons either as lone pairs or bonding pairs repel one another

27. Polarity  Polar  A compound with a partially positive pole and a partially negative pole; resulting from unequal electron distribution (Ex. Water)  Non-polar  A compound that does not dissociate into ions due to uniform electron distribution throughout, causing a stable compound