Atomic Size and Ionization Energy

• The increasing trend in atomic size is due to the shielding effect.
• Electrons shield other electrons from the nucleus, increasing with electron number.
• More electrons increase the shielding effect, moving electrons farther from the nucleus.
• Moving down the periodic table (increasing energy level) increases atomic size due to increased shielding.
• The number of electron shells remains constant, so the shielding effect does not change significantly, resulting in a higher effective nuclear charge.
• Moving left to right decreases atomic size due to a decreased shielding effect and increased effective nuclear charge.
• Fewer electrons result in less shielding.
• Effective nuclear charge is the net positive charge experienced by an electron.
• As atomic number increases (left to right), the nucleus's pull increases, drawing electrons closer.
• Effective nuclear charge pulls electrons towards the nucleus.

Ionization Energy

• Ionization energy (IE) is the energy needed to remove the outermost electron from a gaseous atom or ion.
• Energy is needed to remove an electron because it's attracted to the nucleus.
• Smaller atoms require more energy to remove electrons due to the stronger pull of the nucleus.
• Larger atoms require less energy because outer electrons are shielded from the nucleus's pull.
• Ionization energy increases from left to right on the periodic table, as atomic size decreases.
• Ionization energy decreases from top to bottom, as atomic size increases.
• The removal of the last outermost electron is called the first ionization energy

Ionization Energy and Noble Gases

• Noble gases (helium, neon, argon, krypton, xenon) have the highest ionization energies in their respective periods.
• High ionization energy indicates stability, requiring high energy to remove electrons from noble gases due to their stable electron configurations.
• Atoms (excluding noble gases) with high ionization energy tend to gain electrons, forming negative ions (anions).
• Atoms with low ionization energy tend to lose electrons, forming positive ions (cations).

Electronegativity

• Electronegativity is the ability of an atom to attract electrons to itself.
• Higher electronegativity means an atom can more easily receive electrons.
• A weak shielding effect allows an atom to easily add an electron, increasing electronegativity.
• Smaller atoms have higher electronegativity.
• Larger atoms have smaller electronegativity.
• Noble gases are exceptions, having very low electronegativity due to their complete, stable orbitals making them inert.

Electron Affinity

• Electron affinity is the energy change when an electron is added to a gaseous atom or ion.
• If an atom readily gains an electron, energy is released, indicated by a negative electron affinity value.
• Irregularities exist, but generally, electron affinity becomes less negative down a group due to increasing atomic size.
• Increased size makes it harder to add an electron.
• Across a period, electron affinity becomes more negative as atomic size decreases, making it easier for the atom to accept an electron.
• Noble gases are exceptions to this trend.
• Negative electron affinity means energy is released when an electron is added.
• Positive electron affinity means energy is required to add an electron.

Metallic Property

• Metals are good conductors of heat and electricity due to their outermost electrons.
• Metallic property is defined as the looseness of the outermost electrons.
• Metallic property increases down a group because larger atoms loosely hold electrons.
• Metallic behavior decreases from left to right due to the decrease in size.
• Increased effective nuclear charge holds electrons more strongly in smaller atoms.

Ionic Size

• Anion size increases when compared to its neutral atom counterpart.
• The size increases is because of the increase in the shielding effect and repulsion among the electrons.
• Cation size decreases with charge.
• Removing an electron increases the effective nuclear charge.
• Decreasing the number of electrons compared with the number of protons strengthens the attractive force that pulls the electrons into the atom, thus shrinking its size.
• Picometer (pm) is a unit of length in the metric system and it's equal to one trillionth of a meter.

Reactivity of Elements

• Elements typically combine with other elements to form compounds.
• Valence electrons determine how an atom behaves during compound formation.
• Valence electrons can be lost, gained, or shared during chemical changes.
• Metals and nonmetals react and form ionic compounds.
• Ionic compounds are formed between a cation and an anion.
• Metals lose electrons and become cations.
• Nonmetals gain the electrons lost by metals and become anions.
• The number of electrons lost must equal the number gained during ionic compound formation.
• Metals and nonmetals should become stable (like noble gases) when forming ions.