1. The Periodic Law Modern Chemistry © 2009 Holt, Rinehart, & Winston Chapter 5, pp 133 - 165

2. Periodic Law Chemical & physical properties regularly repeat when elements are listed according to their atomic numbers  Atomic number is equal to # of protons  # of protons = # of electrons  # of valence electrons = Group # for s block = Group # - 10 for p block o “Main-group elements” are those in the s & p blocks o Properties are determined mostly by valence electrons

3. Contributors Stanislao Cannizzaro – Reliable method of measuring atomic mass Dmitri Mendeleev – Group elements of similar properties together – Arrange groups according to their atomic masses Henry Moseley – Rearrange elements according to atomic numbers – Maintain groupings according to properties

4. Periodic Properties’ Trends Atomic & ionic radii: as go left & down – Across the period, nuclear charge increases. As the positive charge increases, electrons are pulled in more tightly, thereby the radius. – Down the group, the # of energy levels increase. Cations have smaller ionic radii than the atom. – When the valence electron shell is lost, the ion is smaller Anions have larger ionic radii than the atom. – Extra electrons do not compact as readily, mostly because of their electrostatic repulsion for each other.

5. Periodic Properties’ Trends Ionization energy = energy required to remove electron from neutral atom  as you go right across the period o As nuclear charge increases, electrons are held more tightly.  as you go down the group o As energy levels are added, electrons are held more loosely. o Additional layers of electrons shielding of protons and further repel valence electrons.

6. Periodic Properties’ Trends Electron affinity = energy change accompanying acquisition of electrons by neutral atoms.  Most atoms release energy when get electron. o Released energy is noted as a negative value.  A + e -  A - + energyF has the highest value of -339.9. o Positive or less negative values indicate atom was “forced” to get electron, i.e., energy was absorbed. o Most atoms so forced will spontaneously lose electron.  A + e - + energy  A - Values are often listed as “(0)”.

7. Periodic Properties’ Trends Electronegativity = ability to attract electron in a compound F has highest value.  Fluorine becomes most like a noble gas (filled valence shell = greater stability) when is F -. Fr/Cs has lowest.  Alkali metals are energetically most stable when they lose their valence electron, not gain one.

8. Additions to Periodic Table Noble gases – 1894: Ar; 1895: He; 1898: Kr & Xe; 1900 Rn Lanthanides – Early 1900s – Very similar in chemical & physical properties – Soft and shiny; reactivity like alkaline-earth metals Actinides – All are radioactive; 1 st 4 are natural, rest synthetic – Those after uranium are “transuranium elements”.

9. The s-Block Elements Group 1 = alkali metals (but H is a nonmetal) – Electron configuration for all end in ns 1 – The most reactive of all elements (with air or H 2 O) Never found free in nature; stored in kerosene – Are silvery and soft Group 2 = alkaline earth metals – Electron configuration for all end in ns 2 – Too reactive to be found free in nature Used in fireworks (Mg – white; Sr – red) – Harder, denser, stronger than alkali metals

10. d-Block, or Transition, Elements Less reactive than elements of the s-block May have different # valence electrons/ group – Sum of outer s & inner d electrons = group # Have typical metal characteristics – Good conductors of electricity – High luster Some are don’t react, staying free in nature – Pt, Pd, & Au are least reactive metals

11. The p-Block Elements All have filled s-orbitals Includes metals, all metalloids, & nonmetals – Metals: Al, Ga, In, Tl, Sn, Pb, Bi, Po, & Uuq & Uuh – Metalloids: B, Si, Ge, As, Sb, Te, & (At) Properties of both metals and nonmetals Brittle solids; have luster Semi-conductors – Nonmetals: noble gases, halogens, O, S, N, P, C, Se Halogens are the most reactive of all nonmetals Halogens have all 3 states of matter in their group

12. Trends for p- & d-blocks Atomic radii: going right across period Ionization energy (IE): going right across period & going down the group for the 1 st IE – Outer s electrons are less shielded by d electrons Ion formation & radii: all lose ns 2 electrons 1 st – Ions of a 2 + charge in size across period Electronegativity: {F = 4.0; others are relative} – d block: all are 1.1 – 2.54 – f block: all are 1.1 – 1.5 – Inversely proportional to atomic radii for both.