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Metals, Non Metals And Oxidation

Metals, Non Metals And Oxidation






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    Metals, Non Metals And Oxidation Metals, Non Metals And Oxidation Presentation Transcript

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