Download ppt - Periodic trends

Objectives

• Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity.

• Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations.

• Define valence electrons, and state how many are present in atoms of each main-group element.

Section 2 Periodic Trends

Atomic Radii

• The boundaries of an atom are fuzzy, and an atom’s radius can vary under different conditions.

• To compare different atomic radii, they must be measured under specified conditions.

• Atomic radius - one-half the distance between the nuclei of identical atoms that are bonded together.


Atomic Radii


Atomic Radii

• Atoms tend to be smaller as you go from left to right on the periodic table.

• This is due to increasing positive charge in the nucleus, pulling the electrons closer to the center.

• Atoms tend to be larger down a group.

• This trend is due to the increasing size of the electron cloud as electrons fill up larger energy levels.


Atomic Radii


Ion definition


• Watch the video on entitled “Ion definition” found in this section of the course

Ionic Radii trends


Ionic Radii trends


• Positive ions tend to be smaller than neutral atoms.

• The larger the positive charge, the smaller the ion.

• Again, this is due to increasing positive charge in the nucleus, pulling the electrons closer to the center.

Ionic Radii trends


• Negative ions tend to be larger than neutral atoms.

• The more negative the charge, the larger the ion.

• This is due to the repelling forces between electrons, causing them to occupy a larger space around the atom

Ionization Energy


• The process to form an ion is called ionization.

• The energy required to remove one electron from a neutral atom is called ionization energy.

• Sometimes refered to as IE1 or first ionization energy.

Ionization Energy


• Ionization energies increase across a period.

• Caused by increasing effect nuclear charge

• Higher positive charge more strongly attracts electrons in the same energy level

Ionization Energy


• Ionization energies decrease down a group.

• Electrons removed from larger atoms are at higher energy levels, and are farther away from the nucleus.

• Electrons are removed more easily due to their distance from the nucleus.

Ionization Energy


Ionization Energy


• Watch the “Ionization” in this lesson when you get to this slide before continuing.

Electron affinity


• The energy change that occurs when an electron is acquired by a neutral atom is called electron affinity.

• Electron affinity generally increases across periods.

• Electron affinity generally decreases down groups.

• All these can be explained by effective nuclear charge and by the distance of an electron from the nucleus.

Electron Affinity


• Watch the video on “Electron Affinity” that is included in this lesson.

Valence electrons


• Compounds form because electrons are lost, gained, or shared between atoms.

• Electrons involved in this behavior are called valence electrons.

• Valence electrons are the outermost energy level electrons in an atom.

• Atoms tend to want to have a full set of 8 valence electrons to be stable.

Valence electrons


• Watch the video on “Valence electrons” included in this lesson.

Electronegativity


• Valence electrons hold atoms together in compounds.

• In many compounds, the negative charge of the electrons is concentrated closer to one atom than another.

Electronegativity


• Electronegativity is the ability of an atom in a compound to attract electrons from another atom in the compound.

• Electronegativity generally increases across rows, and decreases down a group

Electronegativity


• Watch the video on “Electronegativity” found in this lesson.

Summary of periodic trends
