Elements, compounds and mixtures Revision Notes

Ionic Compounds

Ionic compounds are formed between metals and non-metals. They consist of positive and negative ions arranged in a giant ionic lattice. The electrostatic forces of attraction between these oppositely charged ions are very strong. This strong bonding requires a large amount of energy to break, leading to high melting and boiling points. In the solid state, ions are fixed in position, so ionic compounds do not conduct electricity. However, when molten or dissolved in water, the ions become mobile and can carry charge, allowing for electrical conductivity.

Simple Molecular Substances

Simple molecular substances, like water (H2O) or carbon dioxide (CO2), are made up of individual molecules. Within each molecule, atoms are held together by strong covalent bonds. However, between molecules, there are only weak intermolecular forces of attraction. These weak forces require very little energy to overcome, resulting in low melting and boiling points. Most simple molecular substances are gases or liquids at room temperature. They do not conduct electricity because they do not have overall charged particles or delocalised electrons to carry charge.

Giant Covalent Structures

Giant covalent structures, such as diamond and silicon dioxide, are macromolecules where all atoms are held together by strong covalent bonds in a vast, continuous network. This extensive network of strong bonds means a very large amount of energy is needed to break them, leading to extremely high melting and boiling points. These substances are typically very hard. Most giant covalent structures, like diamond, do not conduct electricity as they do not have delocalised electrons. Graphite is an exception due to its layered structure and delocalised electrons. Polymers also have strong covalent bonds within chains and strong intermolecular forces between tangled chains giving them varied properties.

Metals

Metals consist of a giant structure of positive metal ions held in place by strong electrostatic forces of attraction to a 'sea' of delocalised (free-moving) electrons. This strong metallic bonding requires significant energy to overcome, leading to high melting and boiling points. The delocalised electrons are also responsible for metals being good conductors of electricity and heat, as they can move throughout the structure carrying charge and energy. The layers of ions in a metal can slide over each other without breaking the metallic bond, making metals malleable and ductile.

Electron Configuration of Carbon

Carbon is an element with an atomic number of 6. This means a neutral carbon atom has 6 electrons. These electrons are arranged in shells: 2 electrons in the innermost shell and 4 electrons in its outermost shell, also known as the valence shell. The number of electrons in the valence shell dictates how an atom will typically bond with other atoms.

Achieving Stability

Atoms strive to achieve a stable electron configuration, which usually means having a full outermost electron shell. For carbon, with 4 electrons in its valence shell, it needs to acquire 4 more electrons to complete its octet (a total of 8 electrons in the outer shell). Gaining or losing 4 electrons outright requires a lot of energy and is less common for carbon. Instead, carbon achieves stability by sharing electrons.

Covalent Bonding and Carbon

Carbon forms covalent bonds, which involve the sharing of electron pairs between atoms. Because carbon has 4 valence electrons and needs 4 more to complete its outer shell, it characteristically forms four covalent bonds. Each covalent bond allows carbon to share one of its valence electrons with an electron from another atom. This sharing mechanism ensures that carbon attains a stable electron arrangement with eight electrons effectively in its outer shell.

Significance of Four Covalent Bonds

The consistent ability of carbon to form four strong covalent bonds (quadrivalency) with other carbon atoms and with different elements (such as hydrogen, oxygen, and nitrogen) is fundamental to its role in chemistry. This property allows carbon to form long chains, branched structures, and rings, leading to the vast diversity of organic compounds and life on Earth.

What are Nanoparticles?

Nanoparticles are extremely small particles, typically between 1 and 100 nanometers (nm) in size. To put this into perspective, a nanometer is one billionth of a meter. Materials at this scale are known as nanoparticulate materials. Their incredibly tiny size gives them unique properties that are often different from the same materials in their larger, 'bulk' forms. This difference in properties is due to a very high surface area to volume ratio.

Importance of Surface Area to Volume Ratio

The key to understanding the unique properties of nanoparticles lies in their surface area to volume ratio. As a material is broken down into smaller and smaller particles, its total surface area increases significantly relative to its total volume. For example, a sugar cube has a relatively small surface area, but if you powder the same amount of sugar, its total surface area becomes much larger. This large surface area means more of the material is exposed, leading to enhanced reactivity, changes in optical properties (how they interact with light), and different electrical or thermal conductivity.

Applications of Nanoparticles

The unique properties of nanoparticles make them useful in many applications:

• Catalysts: Their high surface area provides many sites for chemical reactions to occur, making them efficient catalysts in industrial processes.
• Sunscreens: Nanoparticles of titanium dioxide and zinc oxide are transparent to visible light but block harmful UV rays effectively, avoiding the white residue left by older sunscreens.
• Self-cleaning windows and fabrics: Titanium dioxide nanoparticles can break down dirt when exposed to sunlight, while silver nanoparticles in fabrics have antibacterial properties.
• Cosmetics: Their tiny size allows them to be absorbed more easily or provide specific textures.
• Medicine: Developing uses include targeted drug delivery systems and medical imaging agents.

Potential Risks and Concerns

While nanoparticles offer exciting possibilities, their unique properties also raise concerns. Due to their extremely small size, some nanoparticles can potentially be inhaled, absorbed through the skin, or ingested, and their effects on human health are not fully understood. There are also environmental concerns regarding their release into water and soil ecosystems. Research is ongoing to assess the long-term impact and ensure safe development and use of these materials.

Mendeleev's Original Arrangement

Dmitri Mendeleev created an early version of the Periodic Table in 1869. He arranged the known elements primarily by increasing atomic mass. A key feature of his work was his focus on the recurring patterns of chemical properties. He placed elements with similar properties in the same vertical columns, now known as groups. This emphasis on properties sometimes led him to leave gaps for undiscovered elements and even to swap the order of a few elements if their chemical properties strongly suggested they belonged to a different group, overriding the atomic mass ordering.

Anomalies in Mendeleev's Table

Despite its brilliance, Mendeleev's table had minor anomalies when ordered purely by atomic mass. For example, tellurium (atomic mass ~127.6) was placed before iodine (atomic mass ~126.9). According to atomic mass, iodine should have come first. However, Mendeleev recognised that iodine's chemical properties were more like those of fluorine and chlorine (halogens), and tellurium's were more like oxygen and sulfur. He correctly prioritised chemical properties over strict atomic mass ordering, foreshadowing a deeper underlying principle.

The Discovery of Atomic Number

The breakthrough that explained these anomalies and refined the Periodic Table came in the early 20th century with Henry Moseley's work. Moseley discovered that the critical identifier for an element is not its atomic mass, but its atomic number (Z), which is the number of protons in the nucleus. Each element has a unique atomic number. When elements are arranged strictly by increasing atomic number, the periodic recurrence of chemical properties becomes perfectly consistent, and the previously inverted pairs (like tellurium and iodine) fall into their correct positions based on their fundamental atomic structure. This atomic number is also directly linked to the electron configuration, which dictates chemical reactivity.

The Modern Periodic Table

The modern Periodic Table arranges elements in order of increasing atomic number. This arrangement ensures that elements with similar chemical properties are grouped together consistently. The atomic number directly corresponds to the number of protons, which defines the element. Furthermore, in a neutral atom, it also equals the number of electrons. The electron arrangement, particularly the number of outer shell electrons, is what primarily determines an element's chemical behaviour. Therefore, arranging by atomic number directly reflects the fundamental reason for periodic trends in properties.