Medical Chemistry L2: Periodic Table of Elements

Transition metals located in the central part of the periodic table, filling 10 electrons into the five d orbitals.

Elements at the bottom of the periodic table, filling 14 electrons into the seven f orbitals.

Half the distance between the nuclei of two identical atoms bonded together.

Increases down a group and decreases across a period due to changes in nuclear charge and electron shielding.

The size of an ion compared to its neutral atom.

Addition of electrons increases ionic radius due to increased electron-electron repulsion.

Loss of electrons decreases ionic radius due to increased effective nuclear charge.

The attraction between the nucleus and outermost electrons, affected by the number of protons and shielding by inner electrons.

A chart that organizes all known chemical elements based on their properties.

Horizontal rows in the periodic table. Elements in the same period have different properties but share the same number of electron shells.

Vertical columns in the periodic table. Elements in the same group have similar chemical properties because they have the same number of valence electrons.

Elements that are typically shiny, malleable, and good conductors of electricity. They form positive ions when reacting.

Elements that are usually gases or brittle solids, poor conductors of electricity, and form negative ions when reacting.

Elements that have properties of both metals and non-metals. They are semiconductors used in technology.

The standard unit used to measure the relative atomic mass of elements. It's defined as 1/12 the mass of a carbon-12 atom.

The ratio of the mass of an atom to the atomic mass unit. It tells us how many times the mass of an atom is greater than one twelfth of the mass of a carbon-12 atom.

The ionic radius of a cation is smaller than its atomic radius because the loss of electrons leads to a stronger attraction between the remaining electrons and the nucleus, pulling them closer.

The ionic radius of an anion is larger than its atomic radius because the gain of electrons leads to a weaker attraction between the electrons and the nucleus, pushing them further out.

The energy required to remove one electron from a neutral atom in its gaseous state.

First ionization energy generally increases across a period because effective nuclear charge increases, making it harder to remove an electron.

First ionization energy generally decreases down a group because the atomic radius increases, making it easier to remove an electron from a larger atom.

The change in energy that occurs when an atom gains an electron in its gaseous state.

The tendency of an atom to attract electrons in a chemical bond.

Measured on the Pauling scale, with fluorine (F) being the most electronegative element and cesium (Cs) being the least.

Across a period (left to right), electron affinity generally increases due to the increasing effective nuclear charge, making the atom more attractive to electrons.

Down a group, electron affinity generally decreases due to the increasing distance between the nucleus and the outer electron shell.

A set of elements with similar chemical properties due to the same outermost electron configuration, arranged in columns of the periodic table.

Elements in a group, also known as a family, have the same number of valence electrons, resulting in similar chemical behavior.

Periods are horizontal rows of the periodic table.They represent elements with increasing numbers of electron shells.

Represents the outermost energy level where electrons are present in an atom.

The four main block structures on the periodic table (s, p, d, and f) reflect different types of atomic orbitals and their capacity to hold electrons.

The systematic filling of atomic orbitals in an atom, as determined by the order of increasing energy levels.