1.7 Trends in the Periodic Table

Trends in the Periodic Table There are three factors that contribute to the trends we observe in the periodic table. Quantum theory Nuclear charge Electron shielding

Trends in the Periodic Table Quantum theory Bohr’s model of the atom indicates that electrons are held in energy levels at specific distances from the nucleus. This results in some electrons being further from the nucleus resulting in different sizes and electrostatic forces. Places electrons at a distance from the nucleus.

Trends in the Periodic Table Nuclear Charge As protons are added to the nucleus there is a greater electrostatic attraction for the electrons orbiting the nucleus. The increased nuclear charge results in greater attractions on the electrons within the same energy level. Pulls electrons toward the nucleus.

Trends in the Periodic Table Electron Shielding The negative charges on electrons results in electrostatic forces of repulsion between electrons. Electrons in lower energy levels repel those in upper levels pushing them further from the nucleus. As energy levels become more populated with electrons they repel each other within the energy level. Pushes electrons away from the nucleus.

Reactivity Reactivity is a function of atomic structure H   He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Te Ru Rh Pd Ag Cd In Sn Sb I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg Uub Uut Uuq Uuh Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Increasing reactivity in non-metals Increasing reactivity in metals Reactivity is a function of atomic structure or Quantum theory Note: The noble gases are essentially non-reactive.

Atomic Radius – Across a Period Difficult to predict the radius of an atom Generally, atomic radii decrease from left to right across each period Because… the strength of the attraction between the nucleus and the valence electrons is stronger as protons are added

Why? the number of filled inner energy levels stays the same but… As protons are added to the nucleus and electrons are added to the valence shell, the attraction between the nucleus and valence becomes greater with each increase

Sodium atom Chlorine atom Which would have the smaller atomic radius? Chlorine Why? There is a much greater pull towards the nucleus by the valence electrons in chlorine atom compared to the sodium atom.

Atomic Radius – Down a Group the atomic radius increases down a given group because of the increasing number of energy levels Why? The more inner energy levels there are the more shielded the electrons in the outermost shell become This allows the outer shell to experience less nuclear pull thereby increasing the radius

Cesium atom Sodium atom 2 rings between nucleus and valence shell 4 rings between nucleus and valence shell The valence shell in cesium is much more shielded than in sodium, allowing the valence shell to be further out

Atomic Radius

Ionic Radius- Positive Ions removal of an electron from an atom results in the formation of a positive ion “cation” Ex. Li+ , Au3+ Positive ions always have a smaller radius than the neutral atom from which they are formed Ionic radius increases down a group Why? the outermost occupied energy level is farther from the nucleus

As we move across a Period (from left to right), ionic radii for the positive ions decreases Why? attraction of the nucleus to the electrons in positive ions increases across the period, and the ionic radius gets smaller as a result

Ionic Radius- Negative Ion addition of an electron from an atom results in the formation of a negative ion “anion” Ex. Cl-, O2- When an atom gains an electron to form a negative ion, its radius increases Why? It has a greater electron cloud around it which increases repulsion amongst electrons

Ionization Energy amount of energy required to remove an electron from an atom or ion in the gaseous state amount of energy is not constant, depending on which electron is being removed first ionization energy is defined as the amount of energy required to remove the most weakly held electron from a neutral atom

The second ionization energy is similarly defined as the energy required to remove a second electron, this one from the gaseous positive ion

Ionization energies tend to: increase as one progresses across a period because the greater number of protons (higher nuclear charge) attract orbiting electrons more strongly, which increases the energy needed to remove an electron. generally decrease as one progresses down a group, since the valence electrons are further away from the nucleus and experience a weaker attraction to the positive nucleus.

First Ionization Energy

Electron Affinity Electron affinity might be summarized as the energy required to add an electron to an atom. It might be considered as a measure of how well an atom attracts electrons. This may result in the formation of an anion if the added electron produces a stable atom.

Electron affinity tends to: Increase as one progresses from left to right across a period, l, due to the larger attraction from the nucleus, (nuclear charge) and the atom "wanting" the electron more as it reaches maximum stability. Decrease as one moves down a group, due to the large increase in the atomic radius, electron-electron repulsion and the shielding effect of inner electrons against the valence electrons of the atom.

Electronegativity Electronegativity is a measure of the ability of an atom or molecule to attract pairs of electrons in a chemical bond. The type of bond formed is largely determined by the difference in electronegativity between the atoms involved.

Electronegativity tends to: Increase as one moves from left to right across a period, due to the stronger attraction that the atoms obtain as the nuclear charge increases. Decrease as one moves down a group due to the longer distance between the nucleus and the valence electron shell

Metallic properties Metallic properties refer to the chemical and physical properties associated with elements classified as metals. These properties, which arise from the element's ability to lose electrons, are: the formation of cations; the displacement of hydrogen in acids; the formation of basic oxides; the formation of ionic chlorides; etc

Summary