Electron Configurations And Periodic Properties Atomic Radii Atomic
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Electron Configurations And Periodic Properties
Atomic Radii • Atomic radius – one-half the distance between the nuclei of identical atoms that are bonded together. • The left hand diagram shows bonded atoms.
• As you move across the period on the periodic table the size of the atoms gets smaller. • The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus. The electrons are pulled closer to the more highly charged nucleus.
• As you read down a group on the periodic table the size of the atoms increase. • Electrons are farther from the nucleus.
Ionization Energy • An ion is an atom or group of bonded atoms that has a positive or negative charge. • Sodium (Na), for example, easily loses an electron to form Na+. • Any process that results in the formation of an ion is referred to as ionization. • The energy required to remove one electron from a neutral atom of an element is the ionization energy (IE). • To compare the ease with which atoms of different elements give up electrons, chemists compare ionization energies.
• The ionization energies of the main-group elements (the s-block and p-block elements) increase across each period. • This is due to the increasing nuclear charge. • A higher charge more strongly attracts electrons in the same energy level.
• Among the main-group elements, ionization energies generally decrease down the groups. • Electrons are farther from the nucleus and removed more easily.
Electron Affinity • The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. • Electron affinity is, essentially the opposite of the ionization energy. Instead of removing an electron from the element we add an electron to create an anion.
• Electron affinities become more negative across a period (easier to gain electrons). • Electrons add with more difficulty going down a group.
Ionic Radii • A positive ion is known as a cation (formed by the loss of one or more electrons). • A negative ion is known as a anion (formed from the addition of one or more electrons). • The metals tend to form cations and the nonmetals tend to form anions. • Cationic radii decrease across a period due to increasing nuclear charge and anionic radii decreases for the elements in groups 15 -18.
• There is a gradual increase of ionic radii down a group. • The outer electrons in both cations and anions are in higher energy levels as one reads down a group and are less affected by the nuclear charge.
Valence Electrons • The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons (in the outer s and p levels). • For main-group elements, the valence electrons are the electrons in the outermost s and p sublevels.
Electronegativity • Valence electrons hold atoms together in chemical compounds. • In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another. • This uneven concentration of charge has a significant effect on the chemical properties of a compound.
Electronegativity • Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons (the most electronegative element is fluorine). • Electronegativity increases across each period. • Electronegativity decreases or stays the same down a group.
- Periodic table radius
- Atomic radius trend
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