chemical

5.1 Elements, atoms and symbols

Elements are basic substances consisting of only one type of atom that cannot be simplified or broken down by chemical reactions. Each element is unique and defined by its specific atomic number.
Atoms are the fundamental building blocks of matter. They consist of a dense central nucleus surrounded by a cloud of orbiting electrons.
Symbols are standard abbreviations for elements. Most symbols are based on the English name (e.g., C for Carbon, O for Oxygen), while others come from Latin or Greek (e.g., Fe for Ferrum/Iron, K for Kalium/Potassium).

5.2 Structure of an atom

Atoms are electrically neutral because the number of positive protons equals the number of negative electrons.
The Nucleus: Located at the center, containing protons and neutrons. It holds almost all the atom's mass but occupies a tiny fraction of its volume.
Subatomic Particles:
- Protons: Positive charge (+1), relative mass of 1 unit. The number of protons determines the element's identity.
- Neutrons: Neutral charge (0), relative mass of 1 unit. They act as 'glue' to stabilize the nucleus.
- Electrons: Negative charge (-1), negligible mass (1/1840). They move rapidly in regions called shells or energy levels.

5.3 Atomic number and mass number

Atomic Number (Z): The number of protons in an atom. For neutral atoms, Z = number of electrons.
Mass Number (A): The sum of protons and neutrons in the nucleus (A = p + n).
Nuclide Notation: Represented as {}_Z^A X, where X is the chemical symbol.
Calculation: To find neutrons, use n = A - Z.

5.4 Isotopes

Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A).
Examples: Chlorine has two main isotopes, {}{17}^{35}Cl and {}{17}^{37}Cl.
Since chemical reactions involve electrons, isotopes have identical chemical properties but differ in physical properties like density and boiling point.

5.5 Relative isotopic mass and relative atomic mass (A_r)

Relative Isotopic Mass: The mass of one atom of an isotope compared to 1/12th the mass of a Carbon-12 atom.
Relative Atomic Mass (A_r): Calculated as the weighted average:
- A_r = \frac{\sum (isotope\ mass \times percentage\ abundance)}{100}

5.6 Electronic arrangement of atoms

Electrons occupy shells (K, L, M…).
Maximum capacity: 1st shell (2), 2nd shell (8), 3rd shell (8 for the first 20 elements).
Valence Electrons: Electrons in the outermost shell that determine the chemical reactivity and bonding behavior of the element.

6.1 Classifying elements

Metals: Located on the left side. Characteristics include high melting points, ductility, and tendency to lose electrons to form positive ions (cations).
Non-metals: Located on the right side. They tend to gain electrons to form negative ions (anions) and are usually insulators.
Metalloids: Elements like Silicon (Si) and Germanium (Ge) that have intermediate properties.

6.2 Development of the Periodic Table

Dmitri Mendeleev: Arranged elements by atomic mass but left gaps for undiscovered elements.
Modern Periodic Table: Elements are arranged by increasing atomic number, which resolved inconsistencies in Mendeleev's version.

6.3 Patterns in the Periodic Table

Periods (Rows): Indicate the number of occupied electron shells. Moving left to right, metallic character decreases.
Groups (Columns): Indicate the number of valence electrons. Elements in a group share similar chemical properties.

6.4 Group I: The alkali metals

Properties: Soft enough to be cut with a knife, low density (float on water), and low melting points.
Reactivity: They are stored under oil to prevent reaction with air and water. Reactivity increases down the group (Li < Na < K).

6.5 Group II: The alkaline earth metals

6.6 Group VII: The halogens

Physical state: F2 and Cl2 (gases), Br2 (liquid), I2 (solid).
Trends: Color gets darker down the group. Reactivity decreases down the group because it becomes harder to attract an electron into the outer shell.
Displacement: A more reactive halogen will displace a less reactive one from its halide solution (e.g., Cl2 + 2KI \rightarrow 2KCl + I2).

6.7 Group 0: The noble gases

Monoatomic gases. They are chemically stable because they possess a duplet (He) or octet (Ne, Ar) electronic configuration.

6.8 Predicting properties

Periodic trends allow for the prediction of properties for elements like Francium (Fr) or Astatine (At). For example, since reactivity increases down Group I, Fr is the most reactive metal.

7.1 Metallic bonding

Definition: The electrostatic force of attraction between the lattice of positive metal ions (cations) and the 'sea' of delocalized electrons.
Structure: Metals exist as a giant metallic lattice.
Properties:
- Electrical conductivity: Delocalized electrons are free to move through the structure.
- Malleability: Layers of atoms can slide over each other without breaking the bonds.

7.2 Formation of ions from atoms

Atoms lose or gain electrons to achieve a stable octet (full outer shell) configuration.
Cations: Positive ions formed when metal atoms lose valence electrons (e.g., Na \rightarrow Na^+ + e^-).
Anions: Negative ions formed when non-metal atoms gain electrons (e.g., Cl + e^- \rightarrow Cl^-).

7.3 Ionic bonding and ionic compounds

Ionic Bond: The strong electrostatic attraction between oppositely charged ions.
Ionic Compounds: Typically formed between metals and non-metals. They form a giant ionic lattice structure.
Properties:
- High melting and boiling points due to strong electrostatic forces.
- Conduct electricity when molten or aqueous (ions are mobile) but not when solid.

7.4 Names and formulae of ions

Common Ions:
- Metals (Cations): Na^+, Mg^{2+}, Al^{3+}.
- Non-metals (Anions): Cl^-, O^{2-}, N^{3-}.
Polyatomic Ions:
- Hydroxide (OH^-), Sulfate (SO4^{2-}), Nitrate (NO3^-), and Carbonate (CO_3^{2-}).
Formula Writing: The total charge in an ionic compound must be zero.