Periodic Table

Antoine Lavoisier (1700’s)

• Listed all known elements (33) at the time

• 4 groups: gases, metals, nonmetals, and earths

Dobereiner (early 1800’s)

• Dobereiner arranged the elements into triads (groups of three elements) based on similarities in properties.

John Newlands (1864)

• Arranged elements by increasing atomic mass (70)

• Noticed a repeating pattern of properties

• Created the law of octaves (repeating patterns at every eighth element)

Newlands

Lothar Meyer (1869)

• Identified and proved that there was a connection between atomic mass and the property of the element

• Arranged the elements by increasing atomic mass (added the new ones)

Dmitri Mendeleev (1869)• Proved a connection

between atomic mass and element properties

• Arranged elements by increasing atomic mass

• Predicted the existence and properties of elements yet to be discovered

Henry Moseley (1913)

• Discovered atomic number

• Arranged elements by increasing atomic number

• By doing this a pattern of properties was discovered and fixed previous problems

Periodic Law

• When elements are arranged by increasing atomic number, there is a periodic repetition of physical and chemical properties.

Modern Periodic Table

• Periods (rows)- contain a variety of elements ranging from metals to nonmetals to Noble gases. There are 7.

• Groups or Family (columns)- contain elements that share similar properties. There are 18.

Periods

Representative (Main) Elements

• Marked by “A” on most groups.

• Elements in the ‘s’ and ‘p’ block

• Wide range of characteristics

• This is Newlands octaves

Transition Elements (B)

• Consists of only metals.

• Found in the center of the period table.

• Elements of the “d” block

Metals, Nonmetals, and Metalloids

Metals

• Make up most of the periodic table

• Solid at room temperature (except Mercury)

• Good conductors of heat and electricity

• Ductile and malleable• Have Luster (shiny)

Nonmetals

• Gases or solids at room temperature (except Br, it is a liquid)

• Poor conductors of heat or electricity

• Brittle• Dull

Metalloids

• Combination of characteristics of both metals and nonmetals

• Silicon and Germanium both used in computer chips

Answer the following questions about iron:

1. In what period is iron found?

2. In what group is iron found?

3.How many protons does an atom of iron have?

4.Write the electron configuration for iron.

5.Write the dot notation for iron.

6. Is iron a representative or transition element?

7. Is iron a metal or nonmetal?

8.List three properties of iron.

9.Describe one way in which classification was used in biology.

Elements ColorHydrogen Aqua

Group 1: Alkali Metals Red

Group 2: Alkali Earth Metals Yellow

Groups 3 -12: Transition Elements Blue

Group 13: Boron Group Grey

Group 14: Carbon Group Peach

Group 15: Nitrogen Group Brown

Group 16: Oxygen Group Purple

Group 17: Halogens Orange

Group 18: Noble Gases Green

Rare Earth Elements: Lanthanide Series

Pink

Rare Earth Elements: Actinide Series Magenta

Color Code the Periodic Table

S- Block• Alkali Metals

– 1 valence e-. This makes them highly reactive

– Exist only as compounds

– Silvery white in color– Often bond with

halogens– Used in salts and

batteries– Forms ions with a 1+

charge.

• Alkaline Earth Metals– 2 valence e-. Makes

them highly reactive– Ca and Mg are

important components of living cells

– Silvery in color– Used to make laptop

casings– Forms ions with a 2+

charge.

S- Block

P- Block (Families 13-18)

• Boron Family (13)– 3 valence e-– Tends to give its

valence e- away– Most are metals– Not as reactive as

group one and two– Forms ions with a 3+

charge

• Carbon Family (14)– 4 valence e-– Can either give away

its valence electrons or take additional electrons

– Sn and Pb will form ions with 4+charges

P- Block

• Nitrogen Family (15)– 5 valence e-, but will

form 3- (it prefers to gain 3 e- rather than give away 5)

– N and P are reactive and found in many molecular compounds

• Oxygen Family (16)– 6 valence e-– Forms 2- ions. It

prefers to gain 2 e- rather than give away 6)

– O and S are reactive and found in many compounds

P- Block

• Halogens (17)– 7 valence e-– Form 1- ions (gains 1

e-)– Highly reactive

nonmetals– Will often bond with

metals to make salts

• Noble Gases (18)– 8 valence e-, full p

sublevel– Does not form ions– Inert gases

(unreactive)– They do not bond with

other elements because they do not need anymore e-

P- Block

D- Block (3-12)

• All transition metals

• Most are hard metals

• All can exist as free elements in nature

• Will form a variety of charged ions due to the fact that the s and d sublevels are close in energy amounts

D- Block (3-12)

F- Block (Period 6 and 7)

• Lanthanide Series– Shiny metals– Highly reactive– Fits in period 6

• Actinide Series – Radioactive– The first 4 are

naturally occurring the rest are lab created

– Fits in period 7

F- Block

Identify each of the following elements described below:

1. Nonmetal of the second period and group 4A.

2. The noble gas in period 3.

3. This element has two more protons than phosphorus.

4. The only nonmetal in group 1A.

5. Metal in period 7 with two valence electrons.

6. The element whose electron configuration ends with 3p1.

7. The nonreactive element consisting of 4 energy levels.

8. The metalloid with three valence electrons.

9. The only noble gas that does not have 8 valence electrons.

10. The element in group 2A that has fewer energy levels than magnesium.

Periodic Trends

• Patterns in the periodic table that can be determined by comparing a period or a group

• Atomic Radius, Ionic Radius, Ionization energy, Electronegativity

Atomic Radius

• Size of the atom• Half the distance

between two identical nuclei

increases

incr

ease

s

Atomic Radius Increases:

• Top to Bottom within a group because as you move down a column, the number of energy levels is increasing.

• Right to Left because as you move left across a period the atomic number decreases which results in a larger atomic radius

Ionization Energy

• Energy needed to remove the outermost electron.

• Follows this trend because small atoms have a high ionization energy

increases

incr

ease

s

Formation of Ions

• A positive ion (called a cation) results when an atom loses electrons.

• Metals have low I.E. and therefore, form positive ions.

• Positive ions are smaller than the parent atom.• A negative ion (called an anion) results when an

atom gains electrons.• Nonmetals have high I.E. and therefore, gain

electrons.• Negative ions are larger than the parent atom.

Octet Rule

• Every Element wants 8 valence e-

Ionic Radius• Increases • The

formation of an ion

• When atoms lose e- they become smaller

• When atoms gain e- they become larger

incr

ease

s

Ionization Energy

• Energy required to remove an e- from a gaseous atom. (Pulling an e- off)

Electronegativity

• The ability of an atom to attract an electron while in a chemical bond