Fundamental Chemistry Concepts: Atoms and Bonds
Subatomic Particles Discovery
Electron
Electrons allow electric current conduction in discharge tubes. Discovered by J.J. Thomson (1897), cathode ray particles were identified as negatively charged and possessing very little mass.
Proton
Observed in discharge tubes as rays traveling opposite to cathode rays when using a perforated cathode (canal rays). Discovered by E. Goldstein (1886), canal rays were found to consist of particles with a positive electric charge and a mass approximately 1837 times that of an electron.
Neutron
Identified when unknown radiation striking paraffin ejected protons (or particles of similar mass). These ejected particles themselves carried no electric charge. Discovered by J. Chadwick (1932), neutrons were identified as neutral particles residing in the nucleus, with a mass similar to that of protons.
Atomic Structure Basics
Atomic Number (Z)
The number of protons in the nucleus of an atom. This number uniquely identifies an element, as each element has a distinct number of protons.
Mass Number (A)
The total number of particles with significant mass in the nucleus, meaning the sum of protons and neutrons in an atom.
Isotopes
Different forms of the same element. Isotopes have the same atomic number (number of protons) but differ in their mass number due to having different numbers of neutrons.
Periodic Table Organization and Trends
Groups
Vertical columns in the periodic table. Elements within the same group typically have the same number of electrons in their outermost energy level (valence electrons), leading to similar chemical properties.
Periods
Horizontal rows in the periodic table. Elements are arranged in order of increasing atomic number. The period number corresponds to the principal energy level of the valence electrons.
Atomic Size Trend
Within the same period (across), atomic size generally decreases from left to right. This is because the increasing positive charge of the nucleus attracts the electrons more strongly, pulling them closer. Down a group, atomic size generally increases as electrons occupy higher energy levels, further from the nucleus.
Chemical Bonding Types
Ionic Bond
Typically formed between metallic and non-metallic elements. It involves the electrostatic attraction between positively charged ions (cations) and negatively charged ions (anions) formed by the transfer of electrons.
Properties of Ionic Compounds:
- Usually solid at room temperature with high melting points (strong electrostatic forces require significant energy to overcome).
- Often brittle; applying force can shift ions, causing repulsion and fracture.
- Many are soluble in water (polar water molecules interact with and surround the ions, breaking the lattice).
- Do not conduct electricity in the solid state (ions are fixed in the crystal lattice).
- Conduct electricity when dissolved in water or molten (ions become mobile and can carry charge).
Covalent Bond
Formed by the sharing of one or more pairs of electrons between two atoms, typically non-metals.
Molecular Covalent Substances:
Consist of discrete molecules.
Properties:
- Often gases, liquids, or solids with low melting and boiling points (weak intermolecular forces exist between molecules).
- Generally insoluble in water but soluble in organic solvents (often nonpolar, thus interacting poorly with polar water).
- Do not conduct electricity in any state (no free-moving ions or electrons).
Atomic Covalent Substances (Network Solids):
Atoms are linked in a continuous network by covalent bonds.
Properties:
- Typically very hard solids with very high melting points (strong covalent bonds throughout the structure).
- Generally insoluble in water and other solvents (strong bonds and lack of polarity).
- Usually do not conduct electric current (electrons are localized in bonds, no free ions).
Metallic Bond
The electrostatic attractive force between a lattice of positive metal ions and a surrounding ‘sea’ of delocalized valence electrons.
Properties of Metals:
- Generally high melting points (strong metallic bonding leads to stable structures).
- High density (atoms are usually packed closely).
- Metals can dissolve in each other to form alloys.
- Good conductors of heat (energy transferred via electron movement and lattice vibrations).
- Good conductors of electricity (delocalized electrons are free to move and carry charge).
- Malleable (can be hammered into sheets) and ductile (can be drawn into wires) because layers of ions can slide past each other without breaking the metallic bonds.