1. Structure of Matter Chapter 6 Section 1 Compounds and Molecules

2. Let’s Review! An element is made up of entirely the same atom  It is something that can not be broken down any further. Compounds are pure substances composed of the same ions and/or molecules.  Compounds can be broken down into other substances.

3. Chemical Bonding Joining atoms to form new substances (molecules) This is a chemical change, the new substances will have different properties than the original elements. Chemical bonding involves valence electrons

4. Why do Atoms Bond? Atoms are most stable (happy) with a full valence shell. A full valence is normally 8 e - (for H, He, Li and Be it is 2) Bonding is when atoms share or transfer electrons in order to get to a full valence shell.

5. Chemical Structure Chemical Structure: how the atoms are bonded together to make the compound Water’s chemical formula tells us what atoms make up water, but it doesn’t tell us anything about how it is put together or how it acts.

6. Bond Length and Angle Bond Length: the distance between nuclei of two bonded atoms Bond Angle: the angle formed by two bonds to the same atom, tells which way these atoms point.

7. Chemical Compounds Section 2

8. Let’s Recall… Why do atoms bond?  To have a stable electron configuration

9. Types of Bonds Two kinds of chemical bonds  Ionic bonding  Covalent bonding The way a compound bonds determines many of its properties

10. Ionic Bonds Form from the attraction of oppositely charged ions Ionic Bonding: a transfer of electrons (something gains, something loses) Ionic bonds occur between metals (cations) and nonmetals (anions)

11. Ionic Compounds Ionic bonds are held together by electromagnetic force (opposites attract) Compounds with ionic bonds are ionic compounds Form networks, not molecules  For example  Na + is attracted to Cl -  when large amounts get together, they stack in a crystal arrangement

13. Properties of Ionic Compounds Examples of ionic compounds  Salts, baking soda, and rust Brittle – tend to shatter when hit High Melting points – most don’t melt until they are very hot Solubility – many tend to dissolve in water.

14. Dissolving Ions If you dissolve ions in water, they will break into parts, some will have a positive charge, some will have a negative charge Because of this difference in charge, IONIC COMPOUNDS conduct electricity

15. Covalent Bonds Covalent Bonding: a sharing of electrons Usually between nonmetals

16. Covalent Bonds Structural formula for a covalent bond is shown as a single line More than one pair of electrons can be shared, and there will be an additional line for each pair shared

17. Covalent Bonds Covalent bonds can also be shown with a Lewis Dot diagram The valence electrons of the atom are drawn around the atom symbol Bonded molecules are drawn with the dots being shared

18. Covalent Compound Covalent Compounds are made of molecules.  Molecules are atoms that share electrons in a bond. Covalent compounds can be solids, liquids or gases

19. Properties of Covalent Compounds Examples  Glass, rubbing alcohol, nitroglycerin, and natural gas Low melting points – several exist as gases or liquids at room temperature. Most covalent compounds do not dissolve in water or produce positive and negative particles if they dissolve.  Therefore most covalent compounds do not conduct electricity by themselves.

20. Polarity Some atoms don’t share electrons equally Sometimes the electron is more attracted to one side of the molecule than the other causing the compound to be POLAR If the electrons are shared equally, they are said to be NONPOLAR COVALENT BONDS

21. Water and Conductivity Pure water (highly polar) does not conduct electricity. In order to conduct electricity, water must have positive and negative particles in it. Dissolving ionic compounds puts positive and negative particles in the water, increasing the conductivity of water. Electrolyte: anything dissolved in water that increases conductivity

22. Metallic Bonding Bonds between metals Sharing and transfer of electrons Metallic bonds only occur with the same metal, not with others  Ca can bond with other Ca atoms, but not Ba

23. Metallic Bond In metallic bonds, the valence electrons become community property, traveling anywhere they want to throughout the metal’s packed structure.  The outermost energy levels overlap, and the electrons are free to move from atom to atom This “ Sea of Electrons ” is why metals are such good conductors of electricity and heat. Sea of Electrons

24. Polyatomic Ions Some compounds have both covalent and ionic bonds Polyatomic Ion: groups of covalently bonded atoms that have an overall charge as a group

25. Compound Names and Formulas Section 3

26. Remember: Valence electrons are the ones that want to react Metals have fewer valence electrons, so they will give up 1, 2, or 3 electrons (cations) Nonmetals have more valence electrons, they will gain 1, 2, or 3 valence electrons (anions)

28. Naming Ionic Compounds Ionic bonds are between metals and nonmetals The name of an ionic compound consists of the names of the ions in the compound The metal is always named FIRST, and the nonmetal LAST

29. Naming Ionic Compounds The cation (metal) name will remain the same The anion (nonmetal) name will change  drop the ending and add “-ide” For example – F-F-  fluoride  Cl -, O 2-, C 4-  chloride, oxide and carbide Ions of chlorine and sodium give you  sodium chloride (metal) (nonmetal)

30. Determining Formula of Ions Ions have different charges Ionic compounds want to have an overall charge of 0 (this makes them neutral and stable)  Total positive charge = total negative charge For example:  Na + and O 2-  2 sodium for every one oxygen  Na 2 O

31. Determining Formula of Ions Example: BORON OXIDE 1. Write the symbols of both elements (cation 1 st, anion 2 nd ) 2. Write the valence of each as a superscript 3. Drop the positive and negative signs 4. Crisscross the superscripts so they become subscripts 5. Reduce when possible (not possible here) B O B 3+ O 2- B 3 O 2 B2O3B2O3

32. Let’s Practice! Al 3+ and O 2-  Al 2 O 3  Aluminum oxide K + and Cl -  KCl  Potassium chloride Sr 2+ O 2-  Sr 2 O 2  SrO  Strontium Oxide The subscripts don’t effect the name if there is only one possibility

33. Determining Formula of Ions Some metals form different cations with different charges These are all metals that aren’t in group 1, 2 or aluminum – the transition metals! The cation charge follows the symbol as a roman numeral For example  iron can form Fe 2+ or Fe 3+  These are said as iron (II) and iron (III) Cu + and Cu 2+  Copper (I) and Copper (II)

34. Covalent Naming Covalent bonds involve shared electrons, so there are no charges You still drop the ending of the second atom and replace it with the suffix “-ide”. Ionic names ignored the subscript numbers. Covalent does not  Prefixes are used in the name, they tell you how many atoms of the element are in the compound You cannot reduce these formulas!

35. Prefixes prefixmeaningprefixmeaning *mono-1hexa-6 di-2hepta-7 tri-3octa-8 tetra-4nona-9 penta-5deca-10 *“mono-”, it just keeps its original name!

36. Examples CO  carbon monoxide CO 2  carbon dioxide NI 3  nitrogen triiodide P 4 O 6  tetraphosphorus hexoxide

37. Continuing I 4 O 9  tetriodine nonoxide S 2 F 10  disulfur decafluoride IF 7  Iodine heptafluoride Si 2 Cl 6  disilicon hexachloride