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.