In Group 7A(17), the first two elements are gases, the next one is a liquid, and the fourth one is a solid. What do they have in common that causes them to be in the same group?

When discussing the Group 7A(17) elements, often referred to as the halogens, their chemical properties stand at the forefront of their group classification. All halogens share a distinctive feature: they have seven electrons in their outermost shell, known as valence electrons. This configuration leads to a strong desire to gain an additional electron to achieve a stable noble gas electron configuration. As a result, halogens are highly reactive, particularly with alkali and alkaline earth metals, with which they can form ionic compounds known as salts.

Additionally, the halogens are known for their electronegativity, with fluorine being the most electronegative element on the periodic table. This high electronegativity allows them to form strong polar covalent bonds as they attract electrons from less electronegative atoms. This property is key in many organic and inorganic reactions, making halogens widely applicable in various chemical processes.

Understanding electron configuration is crucial in explaining the behaviors of Group 7A(17) elements. The general electron configuration for halogens is expressed as ns^2np^5. For instance, fluorine has the configuration 2s^2 2p^5, and chlorine has 3s^2 3p^5. This configuration indicates that each halogen element is one electron short of a complete octet, which would provide optimal stability according to the octet rule. This drives the elements’ significant reactivity.

The consistent presence of seven valence electrons among the halogens is the underlying reason why they are placed in the same group on the periodic table. It's also the basis for their similarity in chemical properties, despite differing physical states. Their electron configuration leads to a predictable addition of one electron during reactions, often forming a negative ion with a charge of -1, known as a halide ion.

The halogens exhibit a trend in states of matter that is dependent on their position in the periodic table. At room temperature, the first two halogens, fluorine and chlorine, are found as diatomic gases, bromine is in a liquid state while iodine exists as a diatomic solid. This variation in physical state is primarily related to the increasing molecular weight and intermolecular forces as one moves down the group in the periodic table.

Intermolecular forces such as London dispersion forces (a type of van der Waals force) become stronger with an increase in the number of electrons. As a result, the halogens require more thermal energy to change states as the molecular weight increases. Thus, with each successive element, from fluorine to iodine, the boiling and melting points increase, changing from gaseous to liquid to solid under standard conditions. Despite these differences, the halogens' placement in the same group signals their resemblant behavior in a chemical context rather than their physical state.