1. Bonding Ionic and covalent

2. Key Terms 1  Chemical formula– the combination of chemical symbols and subscripts to indicate what the elements are in the compound and how many atoms of each element are in the compound  Example: H 2 O= two hydrogen atoms and 1 oxygen atom.

3. Key Terms 2  The octet rule– each atom wants to have 8 electron in its outer most energy level.  Atoms can share, take, or give away electrons to accomplish this.  Valence electrons– electrons in the outer most energy level that are responsible for the reactivity of that atom.

4. Key Terms 3  Lewis structure (electron-dot notation)– valence electrons are shown as dots around the element’s symbol. Label the # of Ve- on your periodic table. Only used for main block elements  Because each orbital can hold two electrons, electrons are grouped in pairs forming the shape of a box around the element’s symbol.  Paired electrons can also be represented by a dash instead of dots if they are being shared in a compound.

5. Key terms 4

6. Ions  Ion– any atom that has given up or taken electrons to create a positive or negative charge.  This is done to fill the highest energy level.  Cations (cat-ion) – any element that has given away its electrons to become a positively charged ion. Cations are metals.  Anions– any element that has taken electrons to become a negatively charged ion. Anions are nonmetals.

7. Atoms and charges  What happens when an atom gains or looses an electron?  The atom becomes charged!  Example: Copper has an atomic number of 29. This means copper has 29 electrons(-) and 29 protons(+). If copper were to loose two electrons, what would copper’s charge be? 29 Protons (+) 29 Protons (+) + 27 Electrons (-) 2 Protons left over, each proton has a positive charge so the charge of copper would be +2! 2 Protons left over, each proton has a positive charge so the charge of copper would be +2!

8. Practice Problem  Oxygen has an atomic number of 8 and an atomic mass of 16. How many neutrons does oxygen have?  Answer: 16= N+8, N=8  What would Oxygen’s charge be if it gained two electrons?  Answer: 8(+ protons) + 10 (- electrons)=- 2

9. Ionic Bonding

10.  Ionic bonding– any bond between metals and nonmetals (cations and anions)  Charges are based on how many Ve- are needed to fill the outer shell or drop to the previous full shell. Label this on your table.  The charges must cancel each other out.  Example: Na (+1) and Cl (-1)=NaCl (0)  Example: Ca (+2) and F (-1)=CaF 2 (0)  Ionic compounds are usually solids and in a crystal structure (crystal lattice).

11. Ionic Compounds  Both ions should have complete outer shells after bonding.  Both elements should have noble gas electron configurations  When naming, the first element always stays the same, but the last element should end with –ide  Ex. MgO= Magnesium Oxide instead of Magnesium Oxygen  Ex. CaCl 2 = Calcium Chloride vs Chlorine

12. Covalent Bond Key Terms 1  Molecule: a group of atoms held together by covalent bonds  Covalent bond: when atoms share electrons  Nonpolar covalent bond: electrons are equally shared by all atoms and the electrical charge is balanced  Polar covalent bond: electrons are not shared equally and there is an imbalance in the electrical charge surrounding the molecule.

13. Polar vs. nonpolar  Polarnonpolar

14. Key Terms 2  Polar bonds: when atoms in a molecule have an uneven electron distribution.  Bond length: average distance between two bonded atoms  Bond energy: energy required to break a chemical bond and form neutral isolated atoms.  When the bond length gets shorter, the bond energy gets higher

15. Key Terms 3  Lewis structures must be used to create covalent compounds (molecules). Sorry no short cuts this time.  Single bond: when only two electrons are being shared between two atoms  Double bond: when 4 electrons are being shared between two atoms  Triple bond: when 6 electrons are being shared between two atoms

16. Covalent Bonds

17. Covalent Bonding 2  There are no ions involved with covalent bonds which means no charges.  This is a bond between two nonmetals.  Electrons are shared.  The magic number is still 8.

18. Covalent properties  covalent bonds can produce solids, liquids, or gaseous molecules  they are poor conductors of electricity  They have low melting points and boiling points  They are usually very dull in appearance

19. Covalent Nomenclature  When naming covalent compounds, you MUST use prefixes for the first and second words  The only exception is if you only have one atom for the first element.  the less electronegative (furthest to the left on the p/t) element is given first and its full element name is written  -ide is still needed at the end of the second element as well as a prefix

20. Covalent molecule prefixes 1.Mono- 2.Di- 3.Tri- 4.Tetra- 5.Penta- 6.Hexa- 7.Hepta- 8.Octa- 9.Nona- 10.Deca-

21. Covalent nomenclature  Examples  Dihydrogen monoxide (H 2 O)  Sulfur trioxide (SO 3 )*there is only one sulfur so no prefix is needed.  Trisilicon tetranitride (Si 3 N 4 )  Dinitrogen Pentaoxide (N 2 O 5 )