The Periodic Table History of the Periodic Table 1) Doberiner - Doberiner’s triads Grouped together elements in groups of 3’s with similar chemical properties. (ex) Cl, Br, I Ca, Ba, Sr Discovered that the atomic mass of one element from the triad is close to the ave atomic mass of the the other two elements from the triad. Found 4-5 triads First attempt to classify the elements in any way.

2) John Newlands First to lay things out horizontally by increasing at. mass. When the properties of a series (row) began to repeat themselves he started a new row. Noted that every 8th element was similar to the first Law of Octaves: That when the elements were placed in order by increasing atomic mass every 8th element had similar properties to the 1st element. He obtained 7 groups of 7.

3) Sir William Ramsay - discovered the Noble (Inert) gases in the 1890’s. 4) Demitri Mendeleev (1869) Consisted of rows & columns. Elements placed in order by increasing at. Mass Left to right and top to bottom. Placed elements with similar properties in the same column. 8 columns wide. Sometimes the elements properties did not line up properly. As a result he believed that not all the elements had been discovered.

Therefore, left gaps in his P. Table for the undiscovered elements. Predicted the chemical properties of these undiscovered elements and was very accurate in doing so. Mendeleev’s Periodic Law: Properties of elements are a periodic function of their at. mass. (If you order the elements by increasing at. mass the properties of the elements go thru a cycle and repeat themselves).

Also discovered a problem with the P. Table. If he ordered the elements by increasing at. Mass some of the properties of the elements were out of order. (ex) Co and Ni But if he lined the elements up by properties not all of the atomic masses were in order. Discovered the problem, but did not know what to do about it.

5) Mosely 45 years later Mosely solved Mendeleev’s problem with the properties not lining up when you ordered the elements by increasing at. mass. Mosely was performing x-ray experiments that showed the # of protons per nucleus varied progressively from element to element. X-rays are a form of electromagnetic radiation. X-rays have high frequency and therefore short wavelengths. X-rays are produced when high speed electrons hit a metal target in an evacuated tube.

Mosely discovered that the higher the atomic # (Z) = #p, the shorter the wavelength of the x-ray. He found in some cases the wavelength was 2x shorter than he expected - proved Mendeleev correct in that some elements had not yet been discovered. When he listed elements by increasing at. # (Z) instead of increasing at. mass the properties lined up. The Periodic Law - the properties of the elements are a periodic function of their at. #. When you list the elements by increasing at. # the properties of the elements go thru a cycle and repeat themselves.

Group IIIA = 3 valence e- = B family Group IVA = 4 valence e- = C family Group VA = 5 valence e- = N family Group VIA = 6 valence e- = O family (Groups IIIA - VIA are named according to the top element in the family.) Group VIIA = 7 valence e- = halogens Group VIIIA = 8 valence e- = Nobel or Inert gases (octet = stable) No compounds containing He, Ne, and Ar are known.

A few compounds contain Kr, Xe, and Rn. Transition Elements = d sect. of P. Table Transition elements and the bottom of the p sect. = heavy metals. 4f series = lanthanide series (La - Yb). 5f series = actinide series (Ac - No). f sect. = rare earths Metals, nonmetals, and metalloids Metals - left of the P. steps Nonmetals - right of the P. steps. Metalloids - border the P. steps on 2 sides (except Al)

Most active metal is left and bottom. Also the most basic. Most active nonmetal is right and top. Also the most acidic. Metals tend to have 3 or fewer valance e-. Nonmetals tend to have 5 or more valence e-.

Determining the Outer Most Electrons by Position on the Periodic Table (ex) 51 Sb Steps: 1) Locate sect. = sublevel 2) count down = en level 3) count over = # of valence e- (ex) 46 Pd (ex) 74 W (ex) 95 Am

Locating the element from the last subshell to receive electrons Steps: 1) sect. 2) en level 3) # valence e- (ex) 4p 4 (ex) 5d 6 (ex) 5f 4

Periodic Properties An element’s position on the periodic table and its properties are a result of e- configuration. Atoms of elements in the same column have similar outer e- configurations. The change in structure from one column to the next as we scan the p. table varies in a set way. Properties can be predicted by e- configuration as well as the position on the P. table.

Trends in Properties 1) Types of Elements A) Series metals -- metalloids -- nonmetals -- Noble gases (border the steps) B) Family IAVA VIA Metals nonmetalsnonmetals metalloids metalsmetalloids

Generally as you go down a family, each element has one more shell or en level than the element above it. An increase in the number of en levels means an increase in the distance from the nucleus making the atoms larger. Due to an increase in the number of en levels there is also an increase in the screening or shielding effect (when inner en levels block the nuclear pull on outer en level e-). Since the e- are not held as tightly the size of the atom increases. Atoms are becoming more metallic and therefore want to lose e- and do not hold them as tightly making the atom larger.

Metallic atoms lose e- to from smaller ions. Nonmetallic atoms gain e- to form larger ions. 3) First Ionization Energy (I.E.) The amount of energy needed to remove 1 e- from an atom. Atom + en -----> cation + 1e-

Atom + 1 e > en + Anion A) series How? E. A. increases as you go across a series. Why? The atoms are becoming more nonmetallic, therefore they want to gain e-. The atoms are becoming smaller - strong nuclear pull. The atoms are becoming closer to being stable and gaining e- will make them even closer.

B) Family How? E.A. decreases down a family Why? Atoms are becoming more metallic - want to lose e- not gain. Atoms are becoming bigger, decrease in nuclear pull - can’t hold their own e- very tightly so do not want more. More en levels therefore atoms are larger and have a larger screening/shielding effect which decreases the nuclear pull - own e- are not being held tightly - do not want more.

5) Electronegativity An atom’s desire to share e-. A) series How? Electronegativity increases across a series. Why? The same reasons as for E.A. B) Family Electronegativity decreases down a family. Why? The same reasons as for E.A.