1. Molecules and Bonding Daniels Sims Fayola

2. How are molecules represented?  Chemical formula = symbols for the elements are used to indicate the types of atoms present and subscripts are used to indicate the relative numbers of atoms. EXAMPLE: CO2  Structural formula = individual bonds are shown indicated by lines. EXAMPLE: H-O-H

3. What are Chemical Bonds?  The forces that hold atoms together in compounds.  MUST follow the octet rule!!!  Types of Bonds…  Ionic  Metallic  Covalent  Polar Covalent  Non-polar Covalent (this one is like a covalent bond)

4. Ionic Bonds…  Closely packed, oppositely charged ions. (+) ion = cation and (-) ion = anion  Makes the compound electrically neutral and joined by electrostatic forces.  Formed when an atom that loses electrons easily (metals) reacts with an atom that has a whole lot of electrons (non-metals).  The most stable and strongest bond!  EXAMPLES… (NaCl, MgCl2, Na2O)

5. Properties of Ionic Compounds…  Most are 3-D crystalline solids  Each ion is strongly attracted to each other  Conduct electrical current in a molten state

6. Metallic Bonds…  Metals consist of closely packed cations (+ charges)  Cations are surrounded by mobile moving valence electrons  Consist of an attraction of free-floating valence electrons to the positively charges metal ions  Examples are Cu, Fe, Mg and precious metals

7. Lets Practice…  Write the symbols for each element  Draw an arrow to show the transfer of electrons from one atom to the other (lewis structures)  Determine the charge for each ion and write the chemical formula  1) Potassium + Fluorine  2) Magnesium + Iodide  3) Sodium + Oxygen  4) Sodium + Chlorine  5) Calcium + Chlorine  6) Aluminum + Chlorine

8. Covalent Bonds…  Electrons are shared by the 2 nuclei of the atoms.  Often non-metals form these bonds.  All organic molecules are covalent.  MOST COMMON BOND!!!  But there are 2 types of covalent bonds.

9. Polar Covalent Bonds…  Unequal sharing of electrons because 2 different atoms (different electronegativities)  Makes one atom look more (+) and the other atom look more (-) (dipole moments)  Any molecules that are different can show some polarity.  EXAMPLES… (H2O, ClF, CCl4, HF)

10. Non-Polar Covalent Bonds…  When atoms are shared equally because they are identical atoms.  Pretty much the same as regular covalent.  Doesn’t make one atom more + or more -  EXAMPLES…(H2, O2, N2, Cl2)

11. VSEPR Theory…  The Valence-Shell Electron-Pair Repulsion Theory  Because electron pairs repel the molecules adjust their shapes so that the valence-electron pairs are as far apart as possible. (why H2O is bent)  Molecular shapes: linear, trigonal planar, bent, pyramidal, tetrahedral, bypyrimidal

12. Covalent Bonding & Lewis Strucutes…  Single bonds: H2O, NH3, CH4, NF3  Double bonds: CO2, CH3COOH (acetic acid)  Triple bonds: N2  There can be some exceptions to the rules, but we won’t get into these

13. Let’s Practice…  Write the symbols for each element  Rearrange the electrons to pair up or “fit like a puzzle” (use different colors for each atom) AKA make the lewis structure  Draw circles to show the sharing of electrons  Write the chemical formula for each molecule  1) Hydrogen + Hydrogen  2) Hydrogen + Oxygen  3) Chlorine + Chlorine  4) Carbon + Oxygen  5) Carbon + Hydrogen