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- 1. Periodic Table Metals, Non-Metals, Groups and Periods
- 3. Metals <ul><li>Metals are located left of the black line on the periodic table. </li></ul><ul><li>Metals become cations, they lose electrons. Positive charge. </li></ul><ul><li>Metals are maleable and ductile and they are also conductors of heat and electricity. </li></ul>
- 4. Non-Metals <ul><li>Located right of the black line on the periodic table. </li></ul><ul><li>Non-Metals gain electrons and become negatively charged. </li></ul><ul><li>Not conductors, brittle (if solid), not ductile. </li></ul>
- 5. Metaloids <ul><li>Located along the line on the periodic table. </li></ul><ul><li>Share properties of metals and non-metals. </li></ul><ul><li>Typically used in electronics. </li></ul>
- 6. Groups <ul><li>Group IA has a +1 charge, lose 1 electron. Also known as the Alkali Metals. </li></ul><ul><li>Soft and white and highly reactive. </li></ul><ul><li>Group IIA has a +2 charge, lose 2 electrons. Also known as the Alkaline Earth Metals. React easily with the halogens to form salts. </li></ul>
- 7. More Groups <ul><li>Group VIIA has a -1 charge. They gain one electron. This group is known as the halogens. Highly reactive, fluorine is one of the most reactive elements in existence. </li></ul><ul><li>Group VIIIA are known as the Noble Gases. Full valence electron shell. Non-reactive. Important for use in welding, lighting, and space exploration. </li></ul>
- 8. Oxidation-Reduction <ul><li>Oxidation is the losing of an electron in a reaction. Original meaning was combining with oxygen. </li></ul><ul><li>Reduction is the gaining of an electron in a reaction. Original meaning was removing oxygen. </li></ul><ul><li>LEO says GER or OIL RIG </li></ul>
- 9. Examples of Oxidation
- 10. Examples of Oxidation
- 11. Reduction
- 12. Oxidation Characteristics <ul><li>Complete loss of electrons </li></ul><ul><li>Shift of electrons away from an atom </li></ul><ul><li>Gain of oxygen </li></ul><ul><li>Increase in oxidation number </li></ul>
- 13. Characteristics of Reduction <ul><li>Complete gain of electrons </li></ul><ul><li>Shift of electrons toward an atom </li></ul><ul><li>Loss of oxygen </li></ul><ul><li>Decrease in oxidation number </li></ul>
- 14. Rules for Assigning Oxidation #’s <ul><li>1. Oxidation number of a monatomic ion is equal to its charge. Ex: Br 1- is -1 and Fe 3+ is +3. </li></ul><ul><li>2. Oxidation number of hydrogen in a compound is +1, except in metal hydrides like NaH then it is +1. </li></ul><ul><li>Oxidation number of oxygen in compounds is -2. </li></ul>
- 15. continued <ul><li>4. The oxidation number of an atom in an uncombined elemental form is 0. </li></ul><ul><li>5. For any neutral compound the sum of the oxidation numbers must equal zero. </li></ul><ul><li>For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion. </li></ul>
- 16. Trends in Atomic Radius
- 17. Octet Rule <ul><li>Atoms, gain or lose electrons so they have 8 electrons in their outer shell. </li></ul><ul><li>Think in terms of the Noble Gases. </li></ul><ul><li>Electron configurations will be extremely important to understand here. </li></ul><ul><li>The s and p sublevels must be full!!! </li></ul>
- 18. Octet Rule <ul><li>Na is in Group IA. It becomes Na + . </li></ul><ul><li>Na has 11 electrons, 1 valence electron. Valence electrons are in the outer most shell. </li></ul><ul><li>If Na + has one less electron, it now has 10. Which element has 10 e? Neon </li></ul>
- 19. Octet Rule <ul><li>Magnesium has 12 electrons. It is in group IIA. Its oxidation number is +2. </li></ul><ul><li>Mg becomes Mg 2+ </li></ul><ul><li>It loses 2 e- and now has 10 electrons, it has 8 valence electrons, just like neon. </li></ul><ul><li>Mg 2+ electron configuration is: </li></ul><ul><li>1s 2 2s 2 2p 6 </li></ul><ul><li>Neon’s configuration is 1s 2 2s 2 2p 6 </li></ul>
- 20. Octet Rule <ul><li>Fluorine becomes F - </li></ul><ul><li>Fluorine has 7 electrons in the valence shell. Gaining one electron gives it 8. </li></ul><ul><li>It now has 10 total e-, just like neon. </li></ul><ul><li>What is the electron configuration for this ion? </li></ul>
- 21. Octet Rule <ul><li>The “A” Group numbers refer to the number of valence electrons. </li></ul><ul><li>Group IA has 1. </li></ul><ul><li>Group IIA has 2. </li></ul><ul><li>Group IIIA has 3. </li></ul><ul><li>All the way to group VIIIA which has 8. </li></ul><ul><li>You cannot go higher than VIIIA. </li></ul>
- 22. Oxidation Numbers <ul><li>For each e- the atom loses, your number is +1. For example, Group IA is +1, Group IIA is +2. </li></ul><ul><li>For each e- the atom gains, your number is -1. For example, Group VIA is -2, Group VIIA is -1. </li></ul>
- 23. Oxidation Numbers <ul><li>The oxidation numbers of a neutral compound must equal 0. </li></ul><ul><li>For example, Na + must combine with something that will have a -1 charge. </li></ul><ul><li>Na + + Cl - NaCl </li></ul><ul><li>(+1) + (-1) =0 </li></ul><ul><li>Mg 2+ + S 2- MgS </li></ul><ul><li>(+2) + (-2) = 0 </li></ul>
- 24. People <ul><li>Dmitiri Mendeleev—developed the modern periodic table. </li></ul><ul><li>John Newlands—first to discover that elements fall into categories by increasing atomic mass. First to assign atomic mass to elements. </li></ul><ul><li>Henry Moseley—discovered atomic mass had a physical significance and helped prove isotopes. </li></ul>
- 25. Terms <ul><li>Organic Chemistry—study of carbon compounds. </li></ul><ul><li>Ore—material in which minerals can be removed—ex: iron-ore. </li></ul><ul><li>Alloy—mixture of two or more elements with one being a metal. </li></ul><ul><li>Inorganic Chemistry—deals with non-organic compunds. </li></ul>
- 26. Terms <ul><li>Actinide Series—group of radioactive elements in Group 3. </li></ul><ul><li>Lanthanide Series—very rare, first row of the inner transition elements. Located in period 7. </li></ul><ul><li>Inner Transition—the “f” grouping, located at the bottom of the periodic chart. </li></ul><ul><li>Diagonal relationships—relationships between elements in neighboring groups. </li></ul>
- 27. Terms <ul><li>Allotrope—elements with the same elements, but different forms. Ex: O 2 and O 3 , oxygen vs. ozone. </li></ul><ul><li>Metallurgy—the ability to extract metal from ore. </li></ul><ul><li>Ferromagnetism—substance whose ions align in the direction of a magnetic field. </li></ul><ul><li>Mineral—something found in nature as solid crystals. </li></ul>
- 28. Types of Bonds <ul><li>Ionic Bonds </li></ul><ul><li>Anions and cations have opposite charges (negative and positive, respectively). </li></ul><ul><li>The positive and negative charges are attracted by electrostatic forces. </li></ul>
- 29. Types of Bonds <ul><li>Covalent Bonds </li></ul><ul><li>Two atoms share electrons in order to complete their octet. </li></ul><ul><li>Only between non-metals. </li></ul>
- 30. Ionic Bonding <ul><li>Ionic bonding occurs between a cation and anion. </li></ul><ul><li>The opposite charges cause the attraction and the bond. </li></ul><ul><li>Understanding how to balance the charges is extremely important. </li></ul>
- 31. Understanding Charges <ul><li>All non metals have a negative charge. When the non-metal gains an electron, it acquires a net negative charge (more electrons than protons). </li></ul><ul><li>Take Cl for example. It is group VIIA or Group 17. It needs one more electron to complete its valence shell. </li></ul>
- 32. Understanding Charges <ul><li>Na is located in IA or Group 1. It can lose 1 electron to achieve the octet rule. If it is 3s 1 then it drops to 2s 2 2p 6 . </li></ul><ul><li>Therefore the positive of Na is attracted to the negative of F. </li></ul>
- 33. The Ionic Bond <ul><li>Na + + F - --> NaF </li></ul><ul><li>Na is +1 F is -1, when you add the charges together you get “0”. </li></ul><ul><li>You will always want a net “0” charge for a neutral compound. Remember, we are trying to achieve stability. </li></ul>
- 34. More Examples <ul><li>Mg 2+ + Cl - ??? </li></ul><ul><li>When writing a chemical formula, you need to cross multiply. </li></ul><ul><li>If you have +2 and -1, what is your net charge? How will you get “0”. </li></ul>
- 35. Writing the formula <ul><li>Mg 2+ + Cl - MgCl 2 </li></ul><ul><li>Cross multiply and drop the charges. </li></ul><ul><li>You have 1(+2) and 2(-1) the net charge “0”. </li></ul>
- 36. Writing a formula <ul><li>Polyatomic ions are a group of atoms with a charge. Ex: (SO 4 ) 2- </li></ul><ul><li>Al 3+ + (SO 4 ) 2- </li></ul><ul><li>Cross multiply the charges: </li></ul><ul><li>Al 2 (SO 4 ) 3 </li></ul><ul><li>Al (+3) and Sulfate (-2) the LCF is 6, cross multiplying charges will achieve “0”. 2(+3) and 3(-2) = 0 </li></ul>
- 37. Review <ul><li>Ionic Compounds are a metal and non-metal (cation and anion). </li></ul><ul><li>Covalent Compounds are 2 or more non-metals that share electrons. </li></ul><ul><li>Oxidation numbers are the charges of the ions. </li></ul><ul><li>Remember to find the LCF of the charges and cross multiply when creating an ionic compound. </li></ul>
- 38. Review <ul><li>The electron dots only represent the valence electrons. The electrons go around the symbol for the element and then after you have 4 lone electrons, begin pairing. </li></ul>
- 39. Review e- dots <ul><li>Li </li></ul><ul><li>Mg </li></ul><ul><li>Al </li></ul><ul><li>Ge </li></ul><ul><li>N </li></ul><ul><li>S </li></ul><ul><li>Cl </li></ul><ul><li>Ar </li></ul>
- 40. Naming Compounds <ul><li>The first word is the cation, the second word is the anion with –ide as the ending. </li></ul><ul><li>Take NaCl for example. </li></ul><ul><li>Na is Sodium and Cl is chlorine. </li></ul><ul><li>It is called Sodium Chloride. </li></ul>
- 41. Naming Ionic Compounds <ul><li>Here is another; Li 3 P </li></ul><ul><li>The number of atoms of each element does not change any part of the name. </li></ul><ul><li>This compound is now called Lithium Phosphide. </li></ul>
- 42. Naming Covalent Compounds <ul><li>Like ionics, use the name of the first element and drop the ending of the name of the second element. </li></ul><ul><li>HF has hydrogen and fluorine. </li></ul><ul><li>HF is called hydrogen fluoride. </li></ul>
- 43. Prefixes <ul><li>Covalent compounds with multiple atoms use one of the following prefixes: </li></ul><ul><li>1=mono 7=hepta </li></ul><ul><li>2=di 8=octa </li></ul><ul><li>3=tri </li></ul><ul><li>4=tetro </li></ul><ul><li>5=penta </li></ul><ul><li>6=hepta </li></ul>
- 44. Naming with a prefix <ul><li>CO 2 </li></ul><ul><li>One carbon, 2 oxygens </li></ul><ul><li>Carbon Dioxide </li></ul><ul><li>Do not use a prefix with an ionic compound: </li></ul><ul><li>MgCl 2 </li></ul><ul><li>Magnesium Chloride </li></ul>
- 45. Common Polyatomic Ions <ul><li>CN - Cyanide </li></ul><ul><li>OH - Hydroxide </li></ul><ul><li>NO 3 - Nitrate </li></ul><ul><li>NO 2 - Nitrite </li></ul><ul><li>CO 3 2- Carbonate </li></ul><ul><li>To name something with a polyatomic ion, use the first element then the name of the polyatomic. </li></ul>
- 46. Covalent Bonding <ul><li>Covalent bonds occur when atoms share electrons in order to complete their octet. </li></ul><ul><li>Covalent bonds are much weaker when compared to an ionic bond. </li></ul>
- 47. Examples <ul><li>Fluorine has 7 valence electrons and needs 1 more to complete it’s octet. </li></ul><ul><li>Hydrogen has 1 valence electron and needs 1 more to complete its “s” sublevel. </li></ul>
- 48. Carbon Tetra Chloride <ul><li>Carbon has 4 valence electrons and needs 4 more. </li></ul><ul><li>Chlorine has 7 valence and needs 1 more. </li></ul>
- 49. Diatomic Molecules <ul><li>Some of the non-metals form what are called diatomic molecules. </li></ul><ul><li>A diatomic molecule is two atoms of the same element bonding together. </li></ul><ul><li>All of the Halogens are diatomic, as well as nitrogen, and oxygen. </li></ul>
- 50. Halogens <ul><li>Each halogen forms a single bond, sharing one electron. </li></ul><ul><li>Let’s take a look at fluorine. </li></ul>
- 51. Polar Molecules <ul><li>In a polar molecule, one end is slightly more negative than the other end. </li></ul><ul><li>Hydrogen Chloride is polar. The Chlorine is more negative than the hydrogen. </li></ul><ul><li>Diatomic Fluorine is not polar. Each fluorine pulls equally. </li></ul>

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