Principles of Chemistry: From Stoichiometry to Molecular Orbital Theory

The Mole Concept

The mole concept is an International System unit of measurement for a substance. It is also represented by carbon-12, and 1 mole is equal to Avogadro's number (6.022 x 10^23).

The Periodic Table's Role in Stoichiometric Calculations

We use the periodic table to get values for given elements in a reaction to obtain their atomic values. We can then use these values for conversion from grams to moles or moles to grams in stoichiometric-based problems (gram/mole, mole-gram conversions).

Heat vs. Enthalpy

Heat is the energy that flows between a system and its surroundings due to a difference in temperature.

Enthalpy is equivalent to the total heat content of a system.

Why Only Enthalpy Explains Energy Associated with Chemical Bonds

Because enthalpy is a state function, the numerical process doesn't matter. What matters is the beginning function and the result (reactants and products).

How Enthalpy's State Function and Extensive Nature Facilitate Calculations

Enthalpy being a state function and extensive means that the numerical calculation present in the equation can be multiplied up to a certain point.

Stoichiometry and Energy Content

The link between the stoichiometry of a balanced chemical equation and the energy content of the process is that the given energy amount is used to multiply by the chemical equation to get the quantized amount.

Identifying Exothermic and Endothermic Reactions

A reaction is exothermic (-) when it releases energy and heats the external environment.

A reaction is endothermic (+) when it gains energy from the environment, causing it to become colder.

Evolution of Atomic Theories

Bohr's Model

Bohr made a model that represents electrons in orbits around a small, central nucleus.

de Broglie's Contribution

De Broglie was the first to think of electrons as waves and that all particles could be treated as waves.

The Basis of Quantum Mechanical Theory

Quantum mechanical theory proposes that electrons were not particles or waves but had properties of both and neither. It is also based on probabilities.

Periodic Trends of Elements

• All elements in the same group have the same amount of valence electrons.
• Atomic Radius: As you go down the periodic table, the atomic radius increases. As you go to the right, the radius decreases.
• Ionic Radius: Electrons repel each other, so adding an electron makes the atom bigger, and taking one makes the atom smaller.
• Ionization Energy: The energy required to take an electron from an atom. It will always be an atom of the outermost shell because they are farther from the nucleus. Elements on the upper side and to the right tend to require more ionization energy to take an electron.
• Electron Affinity: How much an atom wants to gain an electron to be stable. Electron affinity increases as you go up and right on the periodic table.
• Electronegativity: The tendency to hold electrons tightly.

The Theory of Chemical Bonding

Lewis Dot Structure

A Lewis structure is a drawing of a molecule that illustrates the connectivity of the compound, as well as identifying the lone pairs of electrons available for bonding with other molecules. Lines between atoms in a Lewis structure represent covalent bonds, and lone pair electrons are drawn as a pair of dots.

VSEPR

The premise of VSEPR is that the valence electron pairs surrounding an atom tend to repel each other and will, therefore, adopt an arrangement that minimizes this repulsion, thus determining the molecule's geometry.

Valence Bond Theory

Valence bond theory is a chemical bonding theory that explains that the bonding between two atoms is caused by the overlap of half-filled atomic orbitals. The two atoms share each other's unpaired electron to form a filled orbital, creating a hybrid orbital and bonding together. Examples: Sigma and pi bonds are part of valence bond theory.

Molecular Orbital (MO) Theory

Molecular orbital (MO) theory is a method for describing the electronic structure of molecules using quantum mechanics. Electrons are not assigned to individual bonds between atoms but are treated as moving under the influence of the nuclei in the whole molecule.