Key Concepts in Chemistry: Atoms, Reactions, Acids
Periodic Trends
- Ionization Energy: Energy required to remove an electron.
- Atomic Radius: The total distance from an atom’s nucleus to the outermost orbital of an electron.
- Electron Affinity: Energy change associated with the addition of an electron.
- Electronegativity: Ability to attract an electron. Fluorine is the most electronegative element.
Electron Configuration
- Hund’s Rule: When there are different orbitals of the same energy, they first get occupied with a single electron.
- Pauli Exclusion Principle: Electrons differ in at least one quantum number. In one orbital, the two electrons have to differ in their spin.
- Quantum Numbers: Used to describe every electron of the atom.
- VSEPR Rule: Valence shell electron pair repulsion: every electron pair wants as much space as possible.
- Isotope: Variants of an element that have the same number of protons but a different number of neutrons.
Atomic Theories
- Dalton’s Theory:
- Matter consists of tiny particles called atoms.
- Atoms are indestructible. During reactions, they rearrange themselves but do not break apart.
- In a sample of a pure element, all atoms are identical in mass and other properties.
- Atoms of different elements differ in mass and other properties.
- When different elements combine to form compounds, new and more complex particles are built, but in the compound, they are always present in the same fixed numerical ratio.
- Thomson’s Plum-Pudding Model: He discovered the electron in 1897. The model is composed of a “positive cloud” (pudding) that surrounds the negative charges (plums), that were still called corpuscles.
- Rutherford’s Experiment: “The Gold-foil experiment”: This experiment tested Thomson’s model by placing a thin gold metal foil in an evacuated chamber and bombarding it with alpha particles. He expected everything to go through, but since that did not happen, he proved that Thomson’s model was wrong.
- Rutherford’s Theory: There are two parts of the atom: the core (nucleus) and the shell. The core is positively charged, contains protons and neutrons (massive particles), and the shell is negatively charged and occupies almost the whole atom with nearly mass-less particles.
The main problem with Rutherford’s model was that he couldn’t explain why negatively charged electrons remain in orbit when they should instantly fall into the positively charged nucleus.
- Bohr’s Theory: He stated that electrons travel in discrete orbits around the nucleus and that the energy is quantized (restricted to some energy levels).
Problems: Starts to fail mathematically with more than one electron (only H can be treated with this model).
- Heisenberg Uncertainty Principle: Certain pairs of physical properties, such as position and momentum, cannot be simultaneously known to arbitrarily high precision. So, the model violates this since it considers the electrons to have known orbits and a definite radius.
Solutions: Electrons may not be seen as just a particle but also as a wave (wave-particle duality).
Reactions and Energy
- Entropy (S): A measure of the energy not available for work in a thermodynamic process. It expresses the disorder or randomness of the constituents of a thermodynamic system. A closed system always tends towards achieving a state with a maximum of entropy.
- Heat Capacity (C): The amount of heat an object must gain to raise its temperature by 1 degree.
- Endothermic Reaction: Absorbs energy from the surrounding that is in the form of heat (consumes energy). H>0
- Exothermic Reaction: Gives off energy to the surrounding system (heat is a product). H<0
- Catalyst: A substance that causes or accelerates a chemical reaction without itself being affected (does not affect the position of equilibrium in a system).
Acids and Bases
- Ampholyte: A molecule containing both acid and base functionality. Can donate/accept a proton.
- Lewis Acid: Any ionic or molecular species that can accept a pair of electrons to form a coordinate bond.
- Lewis Base: Any ionic or molecular species that can donate a pair of electrons to form a coordinate bond.
- Lewis Neutralization: Formation of the coordinate bond between donator and acceptor.
- Buffer: Watery solution where the pH value stays constant upon the addition of small amounts of concentrated acid or base. It is usually composed of an intermediate or weak strength acid and its corresponding base pair.
- Buffer Capacity: Effectiveness of the buffer to resist changes in the pH, and it depends on the molar concentration of the acid and its conjugate base. The more concentrated, the more effective.
- Amphoteric Line: The amphoteric line is a line on the periodic table of the elements separating metals from nonmetals.
Intermolecular Attraction Forces
- Ion-Ion Interaction: Formation of a lattice; each anion is surrounded by cations and vice versa.
- Van der Waals Interactions: Describe the electrostatic interactions between uncharged molecules which are caused by either the formation of a polar bond (permanent dipoles) or by a temporary disequilibrium in electron density across the molecule (induced dipole).
- Permanent Dipole-Dipole Interaction
- Dipole/Dipole: Permanent dipoles lead to an attraction between the molecules in a solution (surface tension) or solid. The permanent dipole interactions try to establish an ideal spatial arrangement and therefore lead to an ordered arrangement of molecules in a solution.
- Hydrogen Bonds: A special class of permanent dipole interactions. They form in compounds involving hydrogen bound in a highly polar bond. The unshielded H nucleus attracts the electron density of free electron pairs. The energy released is higher than in normal permanent dipole interactions but not as high as in atomic bonds.
- London Dispersion Interactions: Induced dipole interactions are predominating in liquids with nonpolar compounds. London dispersion forces increase with the size of molecules involved. The larger the molecule, the easier polarization effects occur and can be stabilized. London forces generate the least amount of energy in intermolecular interactions.