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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: Br1- is -1 and Fe3+ 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 Mg2+
• It loses 2 e- and now has 10 electrons, it has 8 valence electrons, just like neon.
• Mg2+ electron configuration is:
• 1s2 2s2 2p6
• Neon’s configuration is 1s2 2s2 2p6
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
• Mg2+ + S2- 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: O2 and O3, 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 3s1 then it drops to 2s22p6.
• 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
• Mg2+ + 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
• Mg2+ + Cl- MgCl2
• 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: (SO4)2-
• Al3+ + (SO4)2-
• Cross multiply the charges:• Al2(SO4)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; Li3P
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
• CO2
• One carbon, 2 oxygens• Carbon Dioxide
• Do not use a prefix with an ionic compound:
• MgCl2• Magnesium Chloride
Common Polyatomic Ions
• CN- Cyanide
• OH- Hydroxide
• NO3- Nitrate
• NO2- Nitrite
• CO32- 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.