General Chemistry
The major trends of the periodic table can be explained at least in part by effective nuclear charge, which is a measure of how much positive charge an electron "feels" from the nucleus. This influences atomic radius, a measure of how large an atom is, and generally increases as you move down and to the left on the periodic table.
Ionization energy is a measure of how much energy it takes to remove an electron from a neutral atom and generally increases as you move towards the top right of the periodic table. Inversely, electron affinity measures how much energy is released when you add an electron to a neutral atom and increases as you move towards the top right of the periodic table. Lastly, electronegativity is a measure of how strongly an atom attracts electrons from other atoms to form bonds and follows the same pattern as ionization energy and electron affinity. Electronegativity is measured using the Pauling scale.
Lesson Outline
<ul> </ul> <li>Effective nuclear charge (Z<sub>eff</sub>)</li> <ul> <li>Measure of positive charge an electron "feels" from nucleus</li> <li>Decreases down a group of the periodic table (as electrons are in higher energy levels) due to shielding</li> <li>Increases across a row of the periodic table due to more protons in the nucleus</li> </ul> <li>Atomic radius trend</li> <ul> <li>Atoms get larger as you move down and to the left</li> <li>Inversely proportional to effective nuclear charge</li> <li>Radius changes when an electron is added or removed</li> </ul> <li>Opposite trends to Atomic radius</li> <ul> <li>Directly proportional to effective nuclear charge: values increase as you move up and to the right of periodic table</li> <li>Ionization energy</li> <ul> <li>Energy required to remove an electron from an atom</li> <li>First and second ionization energies are distinct measures</li> </ul> <li>Electron affinity</li> <ul> <li>Energy released when an electron is added to a neutral atom</li> </ul> <li>Electronegativity</li> <ul> <li>Measure of an atom's tendency to attract electrons when forming bonds</li> <li>Trend arises for similar reasons as electron affinity and ionization energy</li> <li>Measured on the Pauling scale</li> </ul> </ul> <li>Notes on the trends</li> <ul> <li>Patterns aren't fully identical to one another or fully smooth across the periodic table</li> <li>Example: half-full subshells are more stable than expected by the trends</li> <li>Noble gases have lower values for metrics due to full valence shells</li> </ul> </ul>
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FAQs
As one moves across a period (row) on the periodic table from left to right, the atomic radius generally decreases. This is because as the atomic number increases, more protons are added to the nucleus, which results in a greater effective nuclear charge and stronger attraction between the nucleus and electrons. As you move down a group (column) on the periodic table, the atomic radius generally increases due to the addition of energy levels (electron shells) which are further away from the nucleus and, therefore, experience weaker attraction to the nucleus, despite shielding effects remaining constant.
Effective nuclear charge refers to the net positive charge experienced by an electron in an atom and is influenced by shielding and the actual nuclear charge (number of protons in the nucleus). Shielding occurs when electrons in inner energy levels block the attraction between the nucleus and the outer electrons, which leads to a decrease in effective nuclear charge. On the other hand, ionization energy is the energy required to remove an electron from an atom. As the effective nuclear charge increases, the ionization energy generally increases because it takes more energy to remove an electron from an atom with a stronger positive pull.
Electron affinity is the energy change that occurs when an electron is added to a neutral atom, forming a negative ion. It's a measure of an atom's tendency to attract and capture an additional electron. Electronegativity, on the other hand, is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is a commonly used numerical scale to describe electronegativity. Both electron affinity and electronegativity are related in that they describe an atom's attraction for electrons and they both increasing when moving up and to the right on the period table, but while electron affinity pertains to the energy change within an isolated atom, electronegativity deals with the context of atoms within a chemical bond.
Ionization energy generally increases as you move across a period from left to right on the periodic table due to an increasing effective nuclear charge. As the number of protons in the nucleus increases, more energy is required to remove an electron because the electron is more tightly held by the increased positive charge. Conversely, ionization energy usually decreases as you move down a group on the periodic table. This is due to the addition of energy levels that are further from the nucleus, resulting in weaker attraction between the nucleus and the outer electrons, despite shielding effects remaining constant.
Shielding effects occur when electrons in inner energy levels "block" the attraction between the nucleus and the outer electrons, reducing the effective nuclear charge experienced by outer electrons. This can result in an increase in atomic radius, as the outer electrons experience a weaker pull from the nucleus. Shielding effects can also lead to a decrease in ionization energy for an atom, as the outer electrons are less tightly bound and require less energy to be removed. Finally, shielding effects can affect electron affinity by reducing the attractive pull of the nucleus on incoming electrons, making it less likely that an atom will capture an additional electron.
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