All questions of Properties of Materials for Class 9 Exam

Ionic compounds conduct electricity when dissolved in water because the ions are free to move and carry an electrical charge. This property contrasts with covalent compounds, which typically do not conduct electricity due to the absence of free-moving charged particles. Additionally, ionic compounds like sodium chloride form a crystal lattice structure, characterized by high melting and boiling points due to the strong electrostatic forces between the ions.

The melting points of alkali metals decrease as you move down Group 1. For example, lithium has a melting point of 180°C, while potassium has a melting point of 63°C. This trend can be attributed to the increasing atomic size and the resulting decrease in the strength of metallic bonds as the outermost electron is farther from the nucleus, making it easier to break bonds.

A covalent bond is formed when electrons are shared between two non-metal atoms. This type of bond allows each atom to achieve a full outer electron shell, stabilizing the molecule. Covalent compounds, like water (H₂O) and ammonia (NH₃), exhibit unique properties based on these shared electron interactions, which differ significantly from ionic compounds formed through electron transfer.

Atoms react to form compounds primarily to achieve greater stability by filling their outermost electron shells. Elements with incomplete outer shells will react with others to either share, gain, or lose electrons, thus reaching a stable electronic configuration. For example, sodium loses an electron to stabilize its outer shell, while chlorine gains an electron to fill its outer shell. This drive for stability explains the reactivity of many elements, particularly those in Groups 1 and 7.

Potassium (K) has the highest atomic number among the options listed, with an atomic number of 19. Lithium has an atomic number of 3, sodium has 11, and hydrogen has 1. The atomic number signifies the number of protons in the nucleus, and as you move down Group 1 in the periodic table, the atomic numbers increase, indicating larger atoms.

A distinguishing property of ionic compounds is their high melting and boiling points, which arise from the strong electrostatic forces between the oppositely charged ions in the lattice structure. In contrast, covalent compounds, especially those with simple molecular structures, generally have lower melting and boiling points due to weaker intermolecular forces. This difference plays a significant role in the practical applications and behaviors of these substances in various conditions.

Neon is a noble gas, characterized by having a complete outer electron shell, which makes it chemically inert and unlikely to form compounds under standard conditions. Noble gases are located in Group 8 of the Periodic Table and include helium, neon, argon, krypton, xenon, and radon. Their lack of reactivity is a vital aspect of their properties, making them useful in applications such as lighting and as inert environments for chemical reactions.

Halogens are highly reactive nonmetals, characterized by having seven electrons in their outermost shell. This incomplete outer shell drives their reactivity, as they seek to gain an electron to achieve a stable configuration. Elements such as fluorine and chlorine readily form compounds, especially with alkali metals, due to this reactivity.

The electronic structure of sodium is 2,8,1, indicating that it has two electrons in the first shell, eight in the second, and one in the third. When sodium forms an ion by losing its outermost electron, its structure changes to 2,8, becoming more stable with a full outer shell of eight electrons in the second shell. This transformation results in the sodium ion (Na⁺), which is positively charged due to the loss of one electron.

The mass number of an atom is calculated by adding the number of protons and neutrons in its nucleus. In this case, 6 protons + 8 neutrons equals a mass number of 14. This number is crucial for determining the identity of the element and its isotopes, influencing its stability and radioactive properties.

The second shell of an atom can hold a maximum of 8 electrons. This is due to the principles of electron configuration, where the energy levels of electrons determine their arrangement around the nucleus. The formula 2n², where n is the shell number, helps calculate the maximum number of electrons that each shell can accommodate, leading to 8 for the second shell (2(2)² = 8). Understanding this is fundamental in grasping how atoms bond and interact with one another.

Group 1, known as alkali metals, is characterized by having one electron in their outermost shell. This single valence electron is responsible for their high reactivity, as they readily lose this electron to form positive ions. Elements such as lithium, sodium, and potassium exemplify this group, demonstrating similar chemical properties due to their electron configuration.

The primary factor that determines the reactivity of alkali metals as you move down the group is the increase in distance of the outermost electron from the nucleus. As more electron shells are added, the outermost electron experiences weaker electrostatic attraction to the nucleus, making it easier to lose. This trend explains why potassium is more reactive than sodium or lithium.

The electronic structure of chlorine is 2,8,7, indicating that chlorine has 2 electrons in the first shell, 8 in the second, and 7 in the third. This configuration is significant because it shows that chlorine is one electron short of having a full outer shell, which contributes to its high reactivity as it tends to gain one electron to achieve stability.

An atom becomes more stable when it achieves a full outer electron shell. This stability is a driving force behind chemical reactions, as atoms tend to react with each other to attain this stable configuration. For example, noble gases have full outer shells and are largely unreactive, while other elements will gain, lose, or share electrons to achieve similar stability, often resulting in the formation of compounds.