Atomic Structure and Quantum Model

Protons, neutrons and electrons are the subatomic particles that make up an atom. Protons have a positive charge ( +1 ), a relative mass of 1, and are located in the nucleus. Neutrons are neutral, have a relative mass of 1, and reside in the nucleus. Electrons are negatively charged (-1), have a negligible mass relative to protons and neutrons (1/1840), and orbit the nucleus.

The Bohr model, while a simplified representation, was insufficient to fully understand the behavior of electrons in atoms. The Quantum Model emerged as a more accurate and complex model that explains the distribution and energy levels of electrons within an atom.

In the Quantum Model, electrons occupy shells, or energy levels, around the nucleus. These shells are numbered using the principal quantum number (n), starting with n=1, 2, 3, etc. Each shell contains subshells, and the number of subshells corresponds to the value of n.

Each subshell is designated using the letters s, p, d, and f. The number of orbitals within a subshell increases in odd numbers: s has 1 orbital, p has 3, d has 5, and f has 7. Each orbital can hold a maximum of 2 electrons, with opposite spins.

Electrons fill orbitals in order of increasing energy, with the lowest energy levels being filled first. If multiple orbitals have the same energy level, Hund's rule states that electrons fill each orbital with one electron before pairing up in any orbital.

The electronic configuration of an element describes the arrangement of electrons in its orbitals. It is written in a shorthand way, using the noble gas core and filling the orbitals in order of increasing energy.

The electronic configuration of chromium (Cr) and copper (Cu) are anomalous because they deviate from the expected filling of the d orbitals. In both cases, an electron from the 4s orbital is moved to the 3d orbital to achieve a more stable half-filled or completely filled d orbital.

The outermost electron (in the highest energy level) is responsible for an element's chemical properties. It is called the 'valence electron', and it is involved in forming chemical bonds.

The electronic configuration of an ion is determined by adding or subtracting electrons from the neutral atom's electronic configuration. For example, a +1 ion loses one electron, and a -2 ion gains two electrons. When removing electrons from a transition metals, the electrons are first taken from the 4s orbital then the 3d orbital.

Electronic configuration can be visualized on a graph to represent the energy level and orbital of electrons in an atom.

Electrons occupy the lowest energy levels first, filling each subshell in ascending order.

The maximum number of electrons in a subshell is 2(2l + 1).

When filling orbitals in a subshell, electrons occupy each orbital individually before pairing up.

The arrangement of electrons in an atom's energy levels and subshells.

The arrangement of electrons in the energy levels and subshells can be visualized using arrows representing electrons in boxes representing orbitals.

A shorthand notation using the previous noble gas element and the remaining electronic configuration.

Elements in the same group have similar configurations.

The minimum energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of 1+ charged gaseous ions.

A full d orbital is more stable than a half-filled d orbital, resulting in an exception to the Aufbau principle.

Elements in the same period have the same number of electron shells.

The minimum energy required to remove one mole of electrons from one mole of 1+ charged gaseous ions to form one mole of 2+ charged gaseous ions.

A configuration with more than one electron in the outer layer is considered reactive; one electron means less reactive.

The energy needed to remove an electron from an already positively charged ion is greater than the first ionization energy because the remaining electrons are more strongly attracted to the nucleus.

The sudden jump in ionization energy indicates the removal of an electron from a new electron shell that is closer to the nucleus.

The shape of the 1s orbital describes the region of space where there is a high probability of finding an electron.

The shape of the 2s orbital is also a sphere, but it is larger than the 1s orbital because the electron is further from the nucleus.

The shape of the 2p orbitals are dumbbell-shaped and oriented along the x, y, and z axes. The 2px, 2py, and 2pz orbitals are all degenerate.

The shape of the 3p orbitals are also dumbbell-shaped and oriented along the x, y, and z axes. However, 3px, 3py, and 3pz orbitals are larger than the 2px, 2py, and 2pz orbitals.

The first ionization energy generally increases across a period of the periodic table due to increasing nuclear charge, while shielding remains constant. This leads to stronger attraction between the nucleus and valence electrons, requiring more energy to remove one electron.

The first ionization energy generally decreases down a group of the periodic table because nuclear charge increases, but shielding effect also increases, leading to larger atomic size and weaker attraction between the nucleus and valence electrons.

Across a period, the atomic radius decreases due to increasing nuclear charge and constant shielding. This means that the valence electrons are closer to the nucleus, making it harder to remove them.

The shielding effect describes the ability of inner shell electrons to reduce the electrostatic attraction between the nucleus and the valence electrons. Inner shell electrons effectively 'shield' the valence electrons from the full positive pull of the nucleus.

The removal of an electron from an atom or ion is represented as a chemical equation, where the atom or ion in its gaseous state transitions to a higher charged state by releasing an electron.

The variation in ionization energies across the periodic table, with general trends of increasing energy across periods and decreasing energy down groups, provides strong evidence for the validity of the Quantum Model of atomic structure. The model successfully explains the distribution and energy levels of electrons within atoms, leading to these observed trends.

The element from which an electron is removed from a higher energy level is easier to ionize, so it requires less energy.

As we move across a period, the nuclear charge increases while the shielding by inner electrons remains constant. This results in a stronger attraction between the nucleus and valence electrons, leading to a smaller atomic radius.

The electron in magnesium is removed from a 3s orbital, which is lower in energy and closer to the nucleus than a 3p orbital from which the electron of aluminum is removed. Removing an electron from a lower energy level requires more energy.

Na⁺ (g) → Na²⁺ (g) + e^-

As we move across Period 3 from sodium to argon, the nuclear charge increases, but the shielding by inner electrons remains constant. This leads to a stronger attraction between the nucleus and valence electrons, resulting in a smaller atomic radius and higher ionization energy.

The element with a half-filled p orbital is oxygen (O).

The alkali metals are located in Group 1 of the periodic table. From the graph, Element S is located in Group 1, making it the alkali metal.