The History of the Periodic Table Originally constructed

The History of the Periodic Table

• Originally constructed to represent the patterns observed in the chemical properties of the elements. • First chemist to recognize patterns was Johann Dobereiner (1780 -1849). • Noticed several groups of three elements had similar properties, for example, chlorine, bromine, and iodine. • Tried to expand his model of triads but it was severely limited.

• Next notable attempt was made by John Newlands in 1864. • Suggested that elements should be arranged in octaves. • This was based on the idea that certain properties seemed to repeat for every eighth element. • Model did attempt to group based on properties but not generally successful.

• Present form of periodic table conceived by Julius Lothar Meyer (1830 -1885) and Dmitri Mendeleev (1834 -1907). Meyer Mendeleev

• Mendeleev is given most of the credit because he emphasized the table could be used to predict the existence and properties of unknown elements. • He published his table in 1872.

• Mendeleev predicted the existence and properties of the elements gallium, scandium, and germanium from gaps in his periodic table. • Germanium was discovered in 1886 and his predicted values and those observed are in excellent agreement.

• Mendeleev was also able to predict atomic masses of several elements, including indium, beryllium and uranium. • Mendeleev’s table was almost universally adopted and remains one of the most valuable of a chemist’s tools. • The fundamental difference between Mendeleev’s table and the modern periodic table is the modern table uses atomic number to order the elements rather than atomic mass.

• Valence electrons are the electrons in the outermost principal quantum level (outermost energy level) of an atom. • Electron configuration for nitrogen: 1 s 22 p 3 • The valence electrons for nitrogen are the 2 s and 2 p electrons; therefore, nitrogen has five valence electrons. • Valence electrons are important because they are involved in bonding. • Core electrons are the inner electrons.

• Elements with the same valence configuration show similar chemical behavior. • Groups 1, 2, 13 -18 are often called the main-group or representative elements. • Every member of these groups has the same valence electron configuration. • Predicting the valence electron configurations of the transition metals, the lanthanides, and the actinides is somewhat more difficult because of the many exceptions.

Periodic Trends • There are observed trends in several important atomic properties: ionization energy, electron affinity, and atomic size.

Effective Nuclear Charge Zeff • The effective nuclear charge is the pull that an electron “feels” from the nucleus. • Effective Nuclear Charge (Zeff) = # protons - # core electrons • The closer an electron is to the nucleus, the more pull it feels. • As effective nuclear charge increases, the electron cloud is pulled in tighter.

IONIZATION ENERGY • Ionization energy is the energy required to remove an electron from a gaseous atom or ion: X (g) → X+ (g) + e • Consider the energy required to remove several electrons from aluminum in the gaseous state. Al (g) → Al+ (g) + e. I 1 = 580 k. J/mol Al+ (g) → Al 2+ (g) + e. I 2 = 1815 k. J/mol Al 2+ (g) → Al 3+ (g) + e. I 3 = 2740 k. J/mol Al 3+ (g) → Al 4+ (g) + e. I 4 = 11, 600 k. J/mol

Al (g) → Al+ (g) + e. Al+ (g) → Al 2+ (g) + e- I 1 = 580 k. J/mol I 2 = 1815 k. J/mol • The highest energy electron (the one bound least tightly is removed first. • I 1 is the first ionization energy and for aluminum, this electron comes from the 3 p orbital ([Ne]3 s 23 p 1). • I 2 is the second ionization energy and this electron comes from the 3 s orbital. • Why is I 1 smaller than I 2? • The first electron is removed from a neutral atom and the second is removed from a 1+ ion. • The increase in positive charge binds the electrons more firmly and it takes more energy to remove an electron.

Al (g) → Al+ (g) + e. Al+ (g) → Al 2+ (g) + e. Al 2+ (g) → Al 3+ (g) + e. Al 3+ (g) → Al 4+ (g) + e- I 1 = 580 k. J/mol I 2 = 1815 k. J/mol I 3 = 2740 k. J/mol I 4 = 11, 600 k. J/mol • Why is I 4 so high? • The fourth electron is “core” electron (Al 3+ = 1 s 22 p 6) and core electrons are bound more tightly than valence electrons.

• In general as we go across a period from left to right, the first ionization energy increases. • Reason: increase in effective nuclear charge (more protons in nucleus) felt by the valence electrons across a period. • Causes the valence electrons to be held more tightly, which makes it more difficult to remove them. • Note: there are exceptions in ionization energy trends in going across a period. Due to shielding and electron repulsions.

• First ionization energy decreases in going down a group. • Reason: going down a group the electrons being removed are, on average, farther from the nucleus. • As n increases, the size of the orbital increases, and the electrons are farther from the nucleus, and thus are easier to remove.

Electron Affinity • Electron affinity is the energy change associated with the addition of an electron to a gaseous atom: X (g) + e- → X- (g) • If the addition of the electron is exothermic, the corresponding value for electron affinity will carry a negative sign. • The incoming electron experiences an attraction to the nucleus, which causes the potential energy to be lowered as the electron approaches the atom.

• The trends in electron affinity are similar to those for ionization energy. • Electron affinity becomes more exothermic from left to right across a period. A valence shell that holds its electrons tightly will also tend to bind an additional electron tightly. • Electron affinity becomes less negative down a group. A valence shell that loses electrons easily (low IE) will have little attraction for additional electrons (small EA). • Note: there are exceptions.

Atomic Radius • The radius of an atom (r) is defined as half the distance between the nuclei in a molecule consisting of identical atoms. • For nonmetallic atoms that do not form diatomic molecules, the atomic radii are estimated from their various covalent compounds. • The radii for metal atoms (metallic radii) are obtained from half the distance between metal atoms in solid metal crystals.

• Atomic radii decrease in going from left to right across a period. • Due to increasing effective nuclear charge in going from left to right. Valence electrons are drawn closer to the nucleus, decreasing the size of the atom.

• Atomic radius increases down a group, because of the increases in the orbital sizes in successive principal quantum levels.

Trends in the Sizes of Ions • Negative ions are always larger than the atoms from which they are formed. • When electrons are added to an atom, the mutual repulsions between them increase. • The causes the electrons to push apart and occupy a larger volume.

• Positive ions are always smaller than the atoms from which they are formed. • When electrons are removed from the valence shell, the electron-electron repulsions decrease, which allows the remaining electrons to be pulled closed together around the nucleus.

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. • This uneven concentration of charge has a significant effect on the chemical properties of a compound.

Electronegativity • Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons (the most electronegative element is fluorine). • Electronegativity increases across each period. • Electronegativity decreases or stays the same down a group.