The Periodic Table

( very ) Brief History 1869 Mendeleev* & Meyer published similar tables * First to be recognized at international convention – Elements were ordered based on atomic mass Henry Moseley developed the atomic # concept – Proved more accurate than Mendeleev ’ s atomic mass method

The Periodic Law When elements are arranged by increasing atomic # their chemical & physical properties show a periodic pattern.

Review of Element Symbol Atomic # – #protons Element Symbol Atomic Mass – Weighted average of all isotopes ’ mass # – Listed in AMUs – Equal to MM (g/mol) Element Name

Periodic Table Layout Groups or Families – Columns, Vertical – 18 groups Periods – Rows, Horizontal – 7 periods Kinds of Elements – Metals, Nonmetals, Semi-metals – Varying Properties Groups to know – Group 1 - Alkali Metals – Group 2 - Alkaline Earth Metals – Group 17 Halogens – Group 18 - Noble Gases

Electron Configurations Using the Periodic Table s-block elements: Group 1: Alkali H - 1s 1 Li - 1s 2 2s 1 Na- 1s 2 2s 2 2p 6 3s 1 K - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 s-block elements: Group 2: Alkaline Earth Be - 1s 2 2s 2 Mg- 1s 2 2s 2 2p 6 3s 2 Ca - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2

Electron Configurations (con ’ t) p-block elements: Group 13: B -1s 2 2s 2 2p 1 Group 14: Si -1s 2 2s 2 2p 6 3s 2 3p 2 Group 15: As-1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 Group 16: S - 1s 2 2s 2 2p 6 3s 2 3p 4 Group 17: Br- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5

Noble Gas Configuration Choose Noble Gas Prior to Element FOCUS ON THE VALENCE SHELL Put Noble gas in brackets, ex. [Ne] This represents the “ inner shell ” of the element Add the outer level electrons by row # (2-7) s & p **d level starts on 4th row** 4f & 5f start on 6th & 7th row Examples: Ni: [Ar] 4s 2 3d 8 **remember that 3d fills after 4s and before 4p Sb: [Kr] 5s 2 4d 10 5p 3

Trends in the Periodic Table Definition : Predictable changes in properties of the elements as you move through the table. (Realize there are exceptions) Ionic Size +Ions are Smaller - Ions are larger Atomic Radii Decrease Increas e Ionization Energy Increase Decrease Atomic Radii - distance from the nucleus to outermost electron. Ionization Energy - energy required to remove an electron (kJ/mol) Electron Affinity – energy change when neutral atom gains electron Electron Affinity Increase Decrease

Ionization Energy In general, ionization energies of the main-group elements increase across each period. This increase is caused by increasing nuclear charge. A higher charge more strongly attracts electrons in the same energy level. Among the main-group elements, ionization energies generally decrease down the groups. Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the nucleus.(electron shielding) The electrons are removed more easily.

Additional Trend Information Electron Affinity The energy associated with an atom gaining or losing an electron. (kJ/mol) + Energy means it requires energy …not favorable - Energy means it gives up energy…favorable Electronegativity The ability to attract an electron during bonding Increases up a group and across a period

Electron Affinity The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. Electron affinity generally increases across periods. Increasing nuclear charge along the same sublevel attracts electrons more strongly Electron affinity generally decreases down groups. The larger an atom’s electron cloud is, the farther away its outer electrons are from its nucleus.

Ionic Radii A positive ion is known as a cation. The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius. The electron cloud becomes smaller. The remaining electrons are drawn closer to the nucleus by its unbalanced positive charge. A negative ion is known as an anion. The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius.

Ionic Radii, continued Cationic and anionic radii decrease across a period. The electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level. The outer electrons in both cations and anions are in higher energy levels as one reads down a group. There is a gradual increase of ionic radii down a group.

Electronegativity Valence electrons hold atoms together in chemical compounds. In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another. Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. Electronegativities tend to increase across periods, and decrease or remain about the same down a group.

The Periodic Table

( very ) Brief History 1800 ’ s - Dobereiner introduced “ Triads ” –3 elements with similar properties Newlands introduced “ Law of Octaves ” –At this time 62 known elements –1st behaved like 8th, 2nd like 9th… 1869 Mendeleev* & Meyer published similar tables * First to be recognized at international convention –Elements were ordered based on atomic mass

The Periodic Law When elements are arranged by increasing atomic # their chemical & physical properties show a periodic pattern Henry Moseley developed the atomic # concept. –Proved more accurate than Mendeleev ’ s atomic mass method

Review of Element Symbol Atomic # –#protons Element Symbol Atomic Mass –Weighted average of all isotopes ’ mass # –Listed in AMUs –Equal to MM (g/mol) Element Name

Periodic Table Layout Groups or Families –Columns, Vertical –18 groups Periods –Rows, Horizontal –7 periods Kinds of Elements –Metals, Nonmetals, Semi-metals –Varying Properties Groups to know –Group 1 - Alkali Metals –Group 2 - Alkaline Earth Metals –Group 17 Halogens –Group 18 - Noble Gases

Electron Configurations Using the Periodic Table s-block elements: Group 1: Alkali H - 1s 1 Li - 1s 2 2s 1 Na- 1s 2 2s 2 2p 6 3s 1 K - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 s-block elements: Group 2: Alkaline Earth Be - 1s 2 2s 2 Mg- 1s 2 2s 2 2p 6 3s 2 Ca - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2

Electron Configurations (con ’ t) p-block elements: Group 13: B -1s 2 2s 2 2p 1 Group 14: Si -1s 2 2s 2 2p 6 3s 2 3p 2 Group 15: As-1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 Group 16: S - 1s 2 2s 2 2p 6 3s 2 3p 4 Group 17: Br- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5

Noble Gas Configuration Choose Noble Gas Prior to Element FOCUS ON THE VALENCE SHELL Put Noble gas in brackets, ex. [Ne] This represents the “ inner shell ” of the element Add the outer level electrons by row # (2-7) s & p **d level starts on 4th row** 4f & 5f start on 6th & 7th row Examples: Ni: [Ar] 4s 2 3d 8 **remember that 3d fills after 4s and before 4p Sb: [Kr] 5s 2 4d 10 5p 3

Trends in the Periodic Table Definition: Predictable changes in properties of the elements as you move through the table. Ionic Size +Ions are Smaller - Ions are larger Atomic Radii Decrease Increas e Ionization Energy Increase Decrease Atomic Radii - distance from the nucleus to outermost electron. Ionization Energy - energy required to remove an electron (kJ/mol)

Additional Trend Information Electron Affinity The energy associated with an atom gaining or losing an electron. (kJ/mol) + Energy means it requires energy …not favorable - Energy means it gives up energy…favorable Electronegativity The ability to attract an electron during bonding Increases up a group and across a period