Periodic Table: Organization & Trends

Overview of the Periodic Table

• The periodic table organizes elements and reveals patterns about nature's operations.
• Dmitri Mendeleev arranged elements in rows (periods) and columns (groups).
• Dmitri Mendeleev's table gained acceptance because it correlated data and predicted the existence/properties of undiscovered elements.
• Elements are arranged in rows (periods) and columns (groups) based on similar behavior.
• The arrangement of elements explains metals, metalloids, and nonmetals.
• Rows represent periods.
• Columns represent groups.
• The properties of newly discovered elements matched Mendeleev's predictions.
• Elements in groups exhibit similar behavior.
• Elements in the same group behave similarly, due to the same number of valence electrons.
• Similar behavior among elements helped correlate existing data.
• Predicted elements were later found.
• Predicted elements were later discovered with expected properties, filling gaps in the table.
• The number of valence electrons significantly influences each element's characteristics.

Valence Electrons and Element Behavior

• Elements in the same group behave similarly, possessing the same number of valence electrons.
• Group 1 elements have one valence electron in their outermost shell.
• The number of valence electrons determines many characteristics of each element.
• As you move down the table (increase n value), you gain a shell, but the outermost shell contains a solitary valence electron in group 1.
• Elements in group 2 have two electrons in their outermost shell.
• The number of valence electrons dictates element characteristics.
• Elements in group 1 have one valence electron, those in group 2 have two, and so on.

Trends on the Periodic Table

• Periodic trends help understand element properties based on their position in the table.

Atomic Radius

• Atomic radius refers to the size of an atom.
• Atomic size increases down the table due to the addition of electron shells.
• Atomic size increases as you proceed downward on the table due to adding shells.
• Atomic radius decreases as you move to the right due to increased protons in the nucleus, leading to stronger electromagnetic attraction and radius shrinkage.
• Atomic radius decreases as you go to the right due to stronger electromagnetic attraction.
• Increasing protons in the nucleus causes the electromagnetic attraction felt by the electrons to increases, causing the radius to shrink.
• Overall atomic radius increases from right to left and top to bottom on the periodic table.
• Overall atomic radius increases from right to left and from top to bottom on the periodic table.
• Atomic radius increases from right to left and from top to bottom on the periodic table.
• Atomic radius increases diagonally from the top right to the bottom left of the periodic table.
• Atomic radius decreases as you move right across the periodic table within a shell.

Ionic Radius

• Ionic radius differs, with added electrons making an atom bigger and removed electrons making it smaller.
• Ionic radius is affected by electron repulsion; adding an electron makes an atom bigger, and taking one away makes it smaller.
• Adding an electron increases the size of an atom due to electron repulsion.
• Removing an electron decreases the size of an atom.
• Ions sharing the same electron configuration exhibit decreasing radii with increasing atomic number.
• Increasing nuclear charge within the same shell also contributes to decreasing atomic radius as you move to the right.
• Ions with the same electron configuration have radii that decrease as atomic number increases.
• For ions isoelectronic, radii lessen with atomic number increase
• Ionic radius is the radius of an atom's ion.

Ionization Energy

• Ionization energy is the energy required to remove an electron from an atom's outermost shell.
• Ionization energy is the energy required to remove an electron from an atom's outermost shell.
• Electromagnetic force decreases with distance, facilitating easier electron removal farther from the nucleus.
• Electromagnetic force weakens with distance, so electrons farther from the nucleus are easier to remove.
• The ionization energy trend is opposite to the atomic radius trend.
• Ionization energy trend is opposite to atomic radius trend.
• Francium, a large atom with one valence electron, is easily ionized because this electron is distant from the nucleus.
• Francium (large atom with one valence electron) is easy to ionize.
• Francium (large atom, one valence electron) is easy to ionize.
• Atoms tend to prefer a full outermost shell.
• Atoms like to have their outermost shell completely full.
• Atoms prefer full outermost shells; elements in group 1 easily lose one.
• Elements in group 1 readily lose one electron.
• Elements in group 1 easily lose one electron.
• Helium, featuring only one full shell, demonstrates high stability and necessitates substantial energy for ionization.
• Helium (small atom with full shell) requires much more energy to ionize.
• Helium (one shell, full) is stable and requires much energy to ionize.
• Ionization energy increases from left to right and from bottom to top on the periodic table.
• Ionization energy increases from left to right and bottom to top on the periodic table.
• Ionization energy trend is opposite to atomic radius, increasing up and to the right on the periodic table.
• Removing the last electron in a shell causes a significant jump in ionization energy, achieving a noble gas configuration.
• Successive ionization energies increase because removing more electrons makes the atom less stable.
• There's a huge jump in ionization energy after removing the last electron in a shell.
• Removing the last electron gives noble gas electron configuration.
• Ionization energy increases diagonally from the bottom left to the top right

Successive Ionization Energies

• Removing multiple electrons requires successive ionization energies.
• Elements possess successive ionization energies for removing multiple electrons.
• Each subsequent ionization energy exceeds the previous one because the atom becomes less stable.
• Second ionization energy > first, increasing with each removal, because the atom becomes less stable as electrons are removed.
• There is a large jump in ionization energy when removing the last electron in a shell, as it disrupts the noble gas configuration of the previous shell.
• A significant jump in ionization energy occurs after removing the last electron from a shell, because you jump to the noble gas electron configuration.
• Removing the last electron in a shell results in a significant jump in ionization energy, achieving a noble gas configuration.

Exceptions to Ionization Energy Trend

• Exceptions exist; in the second row, oxygen's ionization energy dips below nitrogen's due to orbital symmetry.
• Deviations from the ionization energy trend have explanations.
• Oxygen dips downwards from nitrogen's ionization energy because of orbital symmetry.
• Ionization energy should generally increase as protons are added to the nucleus and the radius contracts.
• Oxygen has a lower ionization energy than nitrogen because of orbital symmetry.
• Nitrogen's half-full 2p orbitals offer special stability.
• Nitrogen's 2p orbitals are half full, Nitrogen gains special stability, analogous elements possessing a full outermost shell.
• Nitrogen's half-full 2p orbitals provide special stability.
• When nitrogen loses an electron, it loses stability.
• Nitrogen loses stability when losing an electron, however oxygen gains it, explaining the lower ionization energy.
• If nitrogen loses an electron, it loses stability, while oxygen gains stability upon losing an electron.
• Oxygen gains stability by losing an electron, explaining its lower ionization energy compared to nitrogen.
• When oxygen loses an electron, it gains stability.
• All deviations from the ionization energy trend can be explained by discrepancies in orbital symmetry.
• Deviations are attributable to orbital symmetry discrepancies.
• Discrepancies in orbital symmetry explain all deviations from the ionization energy trend.
• Exceptions to the ionization energy trend can be rationalized by orbital symmetry.

Electron Affinity

• Electron affinity measures an atom's desire to gain an electron.
• Electron affinity measures how much an atom wants to gain an electron.
• Electron affinity is the energy released or absorbed when an electron is added to a neutral atom or ion in the gaseous phase.
• Electron affinity is the opposite of ionization energy.
• Apart from noble gasses, electron affinity increases from left to right and from bottom to top on the periodic table.
• Electron affinity, the opposite of ionization energy, indicates an atom's desire to gain an electron (disregarding noble gasses, shells are full).
• Disregarding noble gases, electron affinity increases from left to right and bottom to top.
• Electron affinity increases up and to the right.
• Fluorine has the highest electron affinity, because it needs one to have a full shell.
• Fluorine's electron affinity is the highest, with the addition of one more electron resulting in a full shell (noble gas configuration).
• Fluorine has the highest electron affinity, because it needs one electron for a full shell.
• Elements on the left side do not want to gain electrons.
• The elements in the opposite corner are more likely to lose electrons than gain them.
• Exceptions to the trend relate to discrepancies in orbital symmetry.
• Exceptions to the electron affinity trend occur for the same reasons as with ionization energy.
• Exceptions arise due to similar reasons as ionization energy exceptions.
• Exceptions to electron affinity trend for the same reasons as ionization energy trend.
• Disregarding noble gases, electron affinity increases diagonally from the bottom left to the top right.

Electronegativity

• Electronegativity signifies an atom's ability to hold electrons tightly.
• Electronegativity is ability of an atom to hold electrons tightly.
• Electronegativity increases from left to right and from bottom to top on the periodic table.
• Electronegativity, an atom's ability to hold electrons tightly, increases up and to the right.
• It increases the closer an atom is to having a full shell of electrons.
• Smaller atoms, possessing more protons for their energy level (higher effective nuclear charge), can hold electrons best.
• Smaller atoms (e.g., fluorine) with more protons for their energy level or a higher effective nuclear charge hold electrons best.
• Noble gases are disregarded in this trend.
• This excludes noble gases from this trend.
• Electronegativity increases from left to right and bottom to top disregarding noble gases.
• Electronegativity increases from left to right and from bottom to top on the periodic table.
• Electronegativity is important for understanding chemical bonds.
• Ignoring Noble Gases, electronegativity increases this way

Summary of Trends

• Atomic radius increases from right to left and from top to bottom on the periodic table.
• Atomic radius increases from right to left and top to bottom.
• Atomic radius increases down and to the left.
• Ionization energy, electron affinity, and electronegativity increases from left to right and from bottom to top on the periodic table.
• Ionization energy, electron affinity, and electronegativity increase from left to right and bottom to top.
• Ionization energy, electron affinity, and electronegativity all increase up and to the right.
• Atomic radius increases diagonally from the top right to the bottom left
• Ionization energy, electron affinity, and electronegativity increase diagonally from the bottom left to the top right