Atoms are composed of protons, which have a positive charge, and electrons, which have a negative charge, as well as neutrons, which have no charge. The reason atoms have no overall electrical charge is due to the balance between the number of protons in the nucleus and the number of electrons orbiting the nucleus. Each proton's positive charge is exactly balanced by the negative charge of an electron. In a neutral atom, the number of protons equals the number of electrons, leading to an overall charge of zero. This balance is crucial for the stability of atoms and their ability to form molecules through chemical bonds. When an atom loses or gains electrons, it becomes an ion with a net charge, positively charged if it loses electrons (more protons than electrons) and negatively charged if it gains electrons (more electrons than protons). The precise balance of electrical charges in neutral atoms underpins the structure of matter and the interactions that lead to the formation of complex chemical compounds.

The quantum mechanical model of the atom is a fundamental advancement over the Bohr model, incorporating principles of quantum mechanics to provide a more accurate and comprehensive description of electron behavior. The Bohr model, introduced by Niels Bohr in 1913, proposed that electrons orbit the nucleus in fixed, circular paths called orbits, each with a specific energy level. This model was revolutionary because it introduced the concept of quantized energy levels but failed to accurately describe the electron's behavior in atoms other than hydrogen.

The quantum mechanical model, developed through the work of scientists like Schrödinger and Heisenberg, replaces the notion of fixed orbits with orbitals, regions of space around the nucleus where there is a high probability of finding an electron. This model accounts for the wave-like nature of electrons and introduces the concept of quantum numbers to describe their properties. It provides a more accurate representation of electron distribution and energy levels, explaining the chemical behavior of atoms more effectively than the Bohr model. The quantum mechanical model's introduction of orbitals, probability density, and quantum numbers allows for a more nuanced understanding of atomic and molecular structure, bonding, and electron transitions, aligning with experimental observations across all elements, not just hydrogen.

The electron configuration of an atom describes the distribution of electrons in the atom's orbitals and is fundamental to understanding the chemical behavior of the element. It determines how an atom interacts with other atoms, influencing its ability to form bonds, its reactivity, and its placement in the periodic table. The arrangement of electrons across different energy levels and sublevels follows specific rules, such as the Aufbau principle, which states that electrons fill orbitals starting from the lowest available energy levels before moving to higher ones.

This configuration is crucial because electrons in the outermost shell, or valence electrons, are the ones involved in chemical reactions. Elements with similar valence electron configurations tend to exhibit similar chemical properties, which is the basis for the organization of the periodic table into groups. For instance, the noble gases have full valence shells, making them largely unreactive, whereas alkali metals have a single valence electron, making them highly reactive. Understanding electron configuration enables chemists to predict and explain the chemical behavior of elements, from bonding patterns to reactivity trends, making it a cornerstone of chemical science.

Isotopes of the same element have the same number of protons and, therefore, the same electron configuration, which dictates their chemical properties. This means they will form the same types of bonds and participate in chemical reactions in the same way. However, isotopes differ in their number of neutrons, which affects the atom's mass and can influence its physical properties. Physical properties such as melting point, boiling point, and density are influenced by the mass of the atoms. For example, heavier isotopes will have slightly higher densities compared to lighter isotopes of the same element.

Additionally, the difference in neutron number can affect the nucleus's stability, leading to variations in radioactive behavior among isotopes. Radioactive isotopes, or radioisotopes, may emit radiation and decay into other elements, a physical property not shared by their more stable counterparts. This variation does not affect how isotopes chemically react with other elements but can significantly impact their application in scientific research, medicine (such as in diagnostic imaging or cancer treatment), and dating ancient objects through radiometric dating techniques.

The nucleus of an atom remains intact due to the strong nuclear force, one of the four fundamental forces in nature, which acts between all nucleons (protons and neutrons). This force is much stronger than the electromagnetic force but operates over very short distances, effectively within the nucleus itself. While protons do repel each other due to their positive charge, the strong nuclear force overcomes this repulsion and binds protons and neutrons together in the nucleus.

The strong nuclear force is attractive and becomes stronger as nucleons come closer, reaching a maximum at very short distances (on the order of 1 femtometer, which is 10⁻¹⁵ meters). This force not only counteracts the electromagnetic repulsion between protons but also contributes to the overall stability of the nucleus. However, as the nucleus gets larger (with more protons and neutrons), the strong nuclear force must work harder to maintain stability, which is why larger atoms tend to have more neutrons than protons. This excess of neutrons increases the range of the strong nuclear force to cover the repulsion among the greater number of protons, maintaining nuclear stability.