Chapter 6 – Periodic TablePeriodic Law

Section 1 : Development of the Modern Periodic Table?

Late 1700’s Lavoisier (French) grouped 33 elements into four categories; (Table 6., Page 174)

In 1817 Dobereiner (German)grouped elements into

triads (groups of 3)

Ca

Sr *- similar properties *mass of Sr average of Ca & Ba

Ba

ClS

Br *Se*

ITe

1863 John Newlands (English)

  • 49 elements known at the time
  • arranged the elements in order of increasing atomic masses
  • noticed there appeared to be a repetition of similar properties every 8th element
  • law of octaves

1869 Mendeleev (Russian) & [Lothar Meyer (German)]

  • like Newlands
  • Properties of elements were a function of their atomic masshowever, similar properties occurred after periods (horizontal rows) that vary in length
  • Periodic law
  • Left blank spots in some columns – elements existed that were not discovered
  • Predicted properties of unknown element ekasilicon (Germanium)

Table 6.2,Page 175

Error with Mendeleev’s table – Te; I in wrong spots

  • Their properties were different from other elements on same column – switch their positions

SrK NiCo TeI

1913 Moseley – performed x-ray experiments

  • Periodic law-properties of the elements are a periodic function of their atomic number (# of protons)

Modern Periodic Table

Rows – Periods – 7 end with noble gas

Columns – Groups or families

  • All elements in same group have similar chemical and physical properties

Z (atomic #) = # of protons

  • Elements in a same column or group have similar e- configurations
  • Elements in a column in order of increasing principal quantum number.

Section 2: Classification of the Periodic Table

114 named elements; a few more named soon.

Main- Group Elements –Group 1 alkali metals

Group 2alkaline earth metals

Group 13

Group 14carbon

Group 15 nitrogen / pnictogen

Group 16oxygen / chalcogen

Group 17 halogens

Group 18noble gases

Hydrogen

Transition metals (Groups 3 through 12)

Metals, Non-metals

Metalloids or Semiconductors ; B, Si, Ge, As, Sb, Te, Po, At

Lanthanides (elements 58 to 71) & Actinide (elements 89 to 103)

Section 3: Periodic Trends

Elements in the same column have similar e- configurations; therefore they have the same chemical properties.

Radii of atoms

As the principal quantum number increases the size of e- cloud increases.

Down a column the quantum number increases so the size of atom gets bigger.

Atomic radius: radius of an atom without regard to surrounding atoms.

As you go across a row (same period #) but the # of protons increase, the nuclear charge is stronger. Therefore the size becomes smaller. Figure 6.12, Page 188.

Radii of Ions

Atoms unite to form compound to produce a more stable configuration (to reach an octet).

Na : [Ne] 3s1Na+1 : [Ne]

Cl : [Ne] 3s23p5Cl-1: [Ne] 3s23p6

Ions resemble the noble gas configurations.

NaCl crystal will not conduct an electric charge because individual ions are tightly bonded. But, when dissolved or melted ions dissociate into Na+ and Cl- and will conduct an electrical current.

Ions are free to move.

Na+ small the Na atom (lose and e-)

Cl- larger than Cl atom (gains an e-)

Metallic Ions

  1. Left of center of table ; lose e- to form cations
  2. Have a stable outer e- configuration which resembles noble gases at end of preceding row.
  3. atoms are larger than ions.

Non-metallic ions

  1. Right of table ; form anions by gaining e-
  2. have a stable outer configurations which resembles noble gas at end of its row.
  3. atoms are smaller than ions

Prediction of oxidation numbers

Oxidation # comes from e- in outer energy level

Group 1IA → +1

Group 2IIA → +2

Group 3IIIB → s2d1 Y 4s23d1

*For transition metals it is possible to lose not only the outer level e- also lose d level; *d can be lost only after s is lost because it is one level lower

Transitions metals can have +1 to +8

Sc → 4s23d1 → +2 or +3

Ti → 4s23d2 → +2, +3, +4

V → 4s23d3 → +2, +3, +4, +5

Cr → 3d4 → +2 3d3→ +3 3d2→ +4 [Ar]→ +6

Mn → 4s23d5 → +2, → +7

Fe → 4s23d6 → +2, +3 (half filled)

Group 13 (IIIA) – have +3 → Exception Boron (only shares)

Tl 6s24f145d106p1 → +1, +3

Sn 5s24d105p2 → +2 +4

Pb 6s24f145d106p2 → +2 +4

Halogen → s2p5 → -1

Oxygen → s2p4 → -2

First Ionization energy

Ionization energy – energy required to completely remove an e- from an atom.

First Ionization energy – energy required to completely remove the most loosely held e- in an atom.

Unit = kJ/mol

Rules:

  1. Ionization energy tends to increase as the atomic number increases in any row.
  2. Ionization energy decreases as the atomic number increases down a column.

Metals have low ionization energy

Nonmetals have high ionization energy

Ionization energy becomes lowered by:

  1. Increased distance of outer e- from the nucleus
  2. Shielding effect caused by the repulsion between the kernel e- (inner e-) and valence e- (outer e-)

Ionization energy becomes increased by: Increased nuclear charge of elements with a greater atomic #.

Therefore:

  1. Go across a row. Same energy level, atom smaller –takes more energy to remove an e-
  2. Go down a column – size larger, outer e-are farther from nucleus take less energy to remove an e-.

Multiple Ionization energy

Possible to determine ionization energy for 2nd, 3rd, 4th → 8th e- → not only 1# e- removed

Al → 1s22s22p63s23p1

577.5 kJ/mol1st → 1s22s22p63s2

1810 kJ/mol2nd → 1s22s22p63s1

2750 kJ/mol3rd → 1s22s22p6

11, 580 kJ/mol4th → 1s22s22p5

(break up noble gas configuration)

Electron affinity

Electron affinity – attraction of an atom for an e-.

Both ionization and electron affinity are properties of isolated atoms.

Metals – low ionization energy – low electron affinity

Nonmetals – high ionization energy – high electron affinity

Noble gases – highest ionization energy (lowest least) attraction for an e- (electron affinity).

In general, greater ionization energy – greater e- affinity

Down a column – decreasing tendency to gain e- because outer e- are farther from the nucleus little or not strength to gain e-.

Go across a row – increasing tendency to gain e-:size gets smaller – nuclear charge increases pulls in e-.

Electronegativy:

a measure of the ability of an atom in a chemical compound to attract electrons.

Alkali Family IA (s1)– as the atomic # increases

  1. Atom becomes larger
  2. Outer e- are farther from the nucleus
  3. Inner e- shield the effect of a larger nucleus
  4. Outer e- are held less tighter
  5. Atom becomes more active

** Metals more active left and down ** lose e- = + ion (cation)

Halogen Family (VII A) – s2p5 – as the atomic # increases

Non metals tendency is to gain e-

  1. Atom becomes larger
  2. Outer e- are farther from the nucleus
  3. Inner e- shield the effect of a larger nuclei
  4. Nucleus becomes less attracted to gain e-
  5. Atoms become less active

** non-metals – more active to right and up

gain e- => ions => (anions)