The effective nuclear charge, Zeff, is the charge of the nucleus that an electron experiences.
For the hydrogen atom, the electron experiences the full nuclear charge. For a many electron atom, the other electrons shield part of the nuclear charge for any particular electron.
Obviously, inner electrons (lower principal quantum number) provide a better shielding for outer electrons, than the opposite direction.
The effective nuclear charge is lower than the actual
nuclear charge, Z.
Zeff < Z
It is described as:
Zeff = Z – S
where S is the screening constant. For valence electrons, the screening
constant is close to the number of core electrons. Electrons in the same
valence shell have a small impact on the screening constant.
The greater the effective nuclear charge, the more strongly an electron is held (lower energy).
The lower the effective nuclear charge, the more weakly an electron is held (higher energy).
The effective nuclear charge varies with different orbitals within each atom.
Valence electrons are shielded by core electrons, and experience a small effective nuclear charge.
For principal energy level 2; the 2s orbital has an inner lobe that is closer to the nucleus than the 2p orbitals. So the 2s orbital experiences a greater effective nuclear charge and is lower energy than the 2p orbitals.
This property continues with the d and f orbitals. That is why the energy levels of the sublevels is ns < np < nd < nf.
Going left to right across a row (period), the nuclear charge increases, but the number of core electrons does not increase, so the effective nuclear charge increases for each valence electron. This makes the valence electrons become lower energy, they are held more strongly by the atom, and contract in diameter.
Going down a column, the effective nuclear charge would appear to not change. Except that the higher orbitals have inner lobes that penetrate closer to the nucleus, so that the effective nuclear charge actually increases slightly as we go down a column.
Atomic Radius is an estimate and there are different methods of estimation.
The nonbonding radius (van der Waals radii) is a measure of the radius during collisions. This is commonly used in space filling models of atoms.
The bonding atomic radius of elements is estimated from experimental measurements of the distance between atoms in elements and compounds. For identical atoms bonded together, the bonding atomic radius is 1/2 of the bond distance.
The bonding atomic radii can be used to estimate bond distances.
These bonding atomic radii are smaller than the nonbonding radii of isolated atoms, because when atoms bond, their electron orbitals interpenetrate.
The atomic radii decreases from left to right across a period. Because of increasing nuclear charge.
The atomic radii increases down a group. Because of increases in orbital sized with successive principal quantum numbers
The atomic radius will increase for an element when an electron is added to form an anion. Anions are larger than the parent atoms.
The atomic radius will decrease for an element when an electron is removed to form a cation. Cations are smaller than the parent atoms.
An isoelectronic series is a group of ions that have the same number of electrons. ExampleS2-, Cl-, K+, Ca2+.
Listed in order of increasing atomic number these ions have an increasing effective charge and a decreasing radii.
Ionization Energy is the energy required to remove a ground state electron from a gaseous atom or ion. Ionization energy is positive showing that the energy must be added to ionize the atom.
The greater the Ionization energy, the more difficult to remove the electron.
The First Ionization Energy, I1, is the energy required to remove the highest energy electron from an atom.
Example: Mg à Mg+ + e-
The first ionization energy is lowest in the lower left of the periodic table and highest in the upper right.
The first ionization energy increases from left to right across a period as effective nuclear charge increases and radius decreases.
The first ionization energy decreases going down a group as radius increases with little change in effective nuclear charge.
Discontinuities occur in these trends. These are seen when a sublevel is filled followed by adding the next electron into the next sublevel (shielding). Also when a sublevel is half filled with a single electron in each orbital followed by putting a second electron into one of the orbitals (electron-electron repulsion).
The Second Ionization Energy, I2, is the energy required to remove the highest energy electron from a plus-one ion.
Example: Mg+ à Mg2+ + e-
The second ionization energy will be larger than the first ionization energy.
Subsequent ionization energies continue to increase. The ionization energy increases dramatically when it involves an inner shell electron.
I1 < I2 < I3 …
Electron Affinity is the energy associated with the addition of an electron to a gaseous atom.
An exothermic reaction for the addition of an electron will have a negative electron affinity, resulting in a stable anion.
The greater the attraction of the atom and electron, the more negative the electron affinity.
A positive electron energy means the anion is not stable.
Electron Affinities generally become more negative from left to right across a period (excluding the noble gases). The halogens have the most negative electron affinities.
Electron affinities are smaller or positive for Group 2A elements because the new electron would occupy a previously empty sublevel.
Also Electron affinities are smaller or positive for Group 5A elements because the new electron would be the second electron in an orbital that only held 1 electron.
Electron Affinities do not have a general trend in a group. This difference in values is quite small.
Exceptions to these trends are caused by electron repulsion between electrons in doubly occupied orbitals
Metals tent to have a shiny luster, conduct heat and electricity, are malleable and ductile, and have high melting points.
Metals tend to have low ionization energies and tend to form positive ions easily.
Metals oxidize (lose electrons) easy, often by O2 or acid.
Compounds of metals with nonmetals form ionic compounds.
Most metal oxides are basic. Those that dissolve in water form the metal hydroxide.
MgO(s) + H2O(l) à Mg(OH)2(aq)
The basicity of metal oxides is due to the oxide ions which reacts with water according to the equation.
O2-(aq) + H2O(l) à 2OH-(aq)
Metal oxides also demonstrate their basicity by reacting with acids to form a salt plus water.
CaO(s) + 2HCl(aq) à CaCl2(aq) + H2O(l)
Nonmetals vary greatly in appearance. Most nonmetals have low melting points (except carbon). Seven nonmetals form diatomic molecules under normal conditions.
Nonmetals have higher electron affinities and form anions (and ionic compounds) in reactions with metals.
Compounds composed of only nonmetals are covalent/molecular substances.
Most nonmetal oxides are acidic. Those that dissolve in water form acids.
SO3(g) + H2O(l) à H2SO4(aq)
Showing their acidity, nonmetal oxides dissolve in basic solutions to form salt and water.
SO2(g) + 2NaOH(aq) à Na2SO3(aq) + H2O(l)
The alkali metals are in group 1A. Although hydrogen is usually listed in this group, it behaves as a nonmetal.
In addition to the general trends listed above: the density of the alkali metals increases down the column. This is typical of all groups because atomic mass increases faster than atomic size.
Melting point and freezing point decrease going down the group.
Alkali metals are very reactive and react with nonmetals to form ionic solids.
Alkali metals react vigorously with water forming hydrogen gas and the metal hydroxide. These reactions are very exothermic and become more vigorous with the heavier elements.
They have a low ionization energy.
Alkali metals react with oxygen in variable ways.
Lithium reacts with oxygen to form lithium oxide (O2-).
Sodium reacts with oxygen to form sodium peroxide (O22-)
Potassium, rubidium, and cesium react with oxygen to form the metal peroxides and also metal superoxides (O2-)
Compared to the Alkali metals, the Alkaline Earth metals are harder, more dense, have higher melting points, have higher ionization energies, and are less reactive.
Beryllium does not react with water or steam.
Magnesium does not react with water but does react with steam to form magnesium oxide plus hydrogen.
Calcium and the heavier alkaline earth metals reacts with water to form metal hydroxide plus hydrogen.
The heavier elements increase in reactivity.