3.1B Periodic Trends
Understanding Periodic Trends
Periodic trends are determined by the attraction between the nucleus and electrons, especially the valence electrons.
Effective Nuclear Charge
Each electron in a filled inner shell cancels 1 unit of nuclear charge
Z = atomic number
S = number of shielding electrons
Atomic Radius
The atomic radius increases down a group of elements
As more energy levels are added, the size of an atom increases
The atomic radius decreases across a period
As the Zeff increases across the period, the nuclei of elements pull the electrons in the valence shell increasingly closer to the nucleus
Ionic Radius: Positive Ions
Losing electrons makes a positive ion smaller than the atom that it came from
Comparing positive ions, the size of ions decreases as one moves from left to right across a period
Ionic Radius: Negative Ions
Adding electrons makes a negative ion larger than the atom that it came from
The added electrons are repelled by the other electrons in the valence shell, causing the electrons to spread out
Comparing negative ions, the size of ions decreases as one moves from left to right across a period
Ionization Energy
Ionization energy is the energy required to remove one mole of most loosely held electron from one mole of gaseous atoms, to produce one mole of gaseous ions, each with a charge of +1
X(g) + IE = X+1(g) + e-
Ionization energy increases across a period from left to right
Ionization energy decreases moving down a group
Electron Affinity
Electron affinity is the energy released when one mole of gaseous atoms each acquire electrons to form one mole of gaseous -1 ions
X(g) + e- = X-1(g) + Energy
A high electron affinity means an element has a high tendency to gain electrons
Electron Affinity increases across a period
Electronegativity
The electronegativity of an element is a measure of the ability of its atoms to attract electrons in a covalent bond.
Electronegativity increases from left to right across a period due to the increase in nuclear charge, resulting in increasing attraction to electrons
Electronegativity decreases down a group. The bonding electrons are further from the nucleus, resulting in reduced attraction.
Understanding Periodic Trends
Periodic trends are determined by the attraction between the nucleus and electrons, especially the valence electrons.
Effective Nuclear Charge
Each electron in a filled inner shell cancels 1 unit of nuclear charge
Z = atomic number
S = number of shielding electrons
Atomic Radius
The atomic radius increases down a group of elements
As more energy levels are added, the size of an atom increases
The atomic radius decreases across a period
As the Zeff increases across the period, the nuclei of elements pull the electrons in the valence shell increasingly closer to the nucleus
Ionic Radius: Positive Ions
Losing electrons makes a positive ion smaller than the atom that it came from
Comparing positive ions, the size of ions decreases as one moves from left to right across a period
Ionic Radius: Negative Ions
Adding electrons makes a negative ion larger than the atom that it came from
The added electrons are repelled by the other electrons in the valence shell, causing the electrons to spread out
Comparing negative ions, the size of ions decreases as one moves from left to right across a period
Ionization Energy
Ionization energy is the energy required to remove one mole of most loosely held electron from one mole of gaseous atoms, to produce one mole of gaseous ions, each with a charge of +1
X(g) + IE = X+1(g) + e-
Ionization energy increases across a period from left to right
Ionization energy decreases moving down a group
Electron Affinity
Electron affinity is the energy released when one mole of gaseous atoms each acquire electrons to form one mole of gaseous -1 ions
X(g) + e- = X-1(g) + Energy
A high electron affinity means an element has a high tendency to gain electrons
Electron Affinity increases across a period
Electronegativity
The electronegativity of an element is a measure of the ability of its atoms to attract electrons in a covalent bond.
Electronegativity increases from left to right across a period due to the increase in nuclear charge, resulting in increasing attraction to electrons
Electronegativity decreases down a group. The bonding electrons are further from the nucleus, resulting in reduced attraction.
