Chemistry Chapter on Ionization and Electron Enthalpy

Study Notes

Ionization enthalpy is the energy required to remove an electron from an atom.
The second ionization enthalpy is always greater than the first because removing an electron from a positively charged ion is more difficult.
Ionization enthalpy decreases as you move down a group because electrons are further from the nucleus.

Shielding effect is the reduction of the effective nuclear charge on the electron cloud due to the repulsion forces between electrons.

Electron gain enthalpy is the energy change when an electron is added to a gaseous atom.
It can be either negative (energy is released) or positive (energy is absorbed).
Halogens have highly negative electron gain enthalpies because adding an electron gives them a stable noble gas configuration.
Noble gases have large positive electron gain enthalpies because the extra electron must be placed in the next higher energy level, making the configuration unstable.
The second electron gain enthalpy of an element is positive because adding an electron to a negative ion is opposed by electrostatic repulsion.
Electron gain enthalpy decreases down a group because the increased atomic size weakens the attraction between the nucleus and the incoming electron.
Electron gain enthalpy increases across a period because of the increasing nuclear charge and decreasing atomic radius.

Electronegativity is the tendency of atom in a molecule to attract shared electrons towards itself.
It generally increases across a period and decreases down a group.
Alkali metals have the lowest electronegativities and halogens the highest.
Fluorine has the highest electronegativity (EN = 4.0) and Caesium the lowest (EN = 0.79).

As you move across a period, the number of valence electrons increases from 1 to 8.
The valency of elements, when combined with H or O, increases from 1 to 4 and then decreases to zero.

Chemical bonds are attractive forces that hold atoms together in a stable molecule.
The five types of chemical bonds are: ionic, covalent, coordinate covalent, metallic, and hydrogen.

An ionic bond is formed by the transfer of an electron from one atom to another.
The electrostatic attraction between the cation and anion forms the ionic bond.
Ionic compounds are solids at room temperature, have high melting points, are hard and brittle, soluble in water, conductors of electricity, and do not exhibit isomerism.
Ionic reactions are fast.

A covalent bond is formed by sharing electrons between two atoms.
The shared pair of electrons forms the covalent bond.
Covalent compounds can be gases, liquids, or solids at room temperature, have low melting and boiling points, are not hard or brittle, are soluble in organic solvents, are non-conductors of electricity, exhibit isomerism and participate in molecular reactions.

A coordinate covalent bond is formed when both electrons are supplied by one atom.
The atom that donates the lone pair is called the donor, and the atom that accepts the lone pair is called the acceptor.
The molecule that contains the donor atom is called the ligand.

A metallic bond is formed by the collective sharing of valence electrons between several positively charged metal ions.
Metallic compounds are good conductors of electricity and heat, highly ductile, have a shiny metallic luster, and have high melting and boiling points.

A hydrogen bond is an electrostatic force of attraction between a hydrogen atom covalently bound to a more electronegative atom (N, O, or F) and a lone pair of electrons in a nearby electronegative atom.