Environmental Science 101

Chemistry: An Overview

The scientific discipline involved with compounds composed of atoms: their composition, structure, properties, behavior and the changes they undergo during chemical reactions.
Term is derived from Alchemy: a philosophical and protoscientific tradition practiced throughout the world. It aimed to purify, mature, and perfect certain objects.


Anything that has mass and volume and is made up of particles.
Atom: smallest particle of an element which can enter into a chemical combination.
Element: a single type of atom, characterized by its particular number of protons in the nuclei of its atoms, known as the atomic number. (Periodic Table of Elements)
Compound: a chemical substance containing more than one kind of atom
Molecule: an aggregate of atoms, bonded together by covalent bonds (shared electrons), and has relatively little interaction with other molecules (electrically neutral)
Substance: A chemical substance is a kind of matter with a definite composition and set of properties.
Mixture: A collection of substances is called a mixture. Examples: air and alloys.
Mole: Unit of measurement that denotes an amount of substance
Avogadro constant: number of entities per mole; ̶ approximately 6.022×1023 per mole

Phases of Matter

A set of states of a chemical system that have similar bulk structural properties, over a range of conditions, such as pressure or temperature.
Common Phases and Transition Processes
     Solid to Liquid –
     Liquid to Solid –
     Liquid to Gas –
     Gas to Liquid –
     Gas to Solid –
     Solid to Gas –
Other Phases:
     Bose–Einstein Condensates
     Fermionic condensates
     Paramagnetic and ferromagnetic phases and transition processes


Bonding of atoms occurs because of interactions between electrons of neighboring atoms Most free atoms have too many or too few electrons
This results in an atom that is either negatively or positively charged - an ion
Types of Bonds:
Ionic Bonds result when one atom gives up an electron to another – weak bonds
Covalent Bonds result when two atoms share an electron – very strong bonds
Metallic Bonds result when electrons are shared by a large number of atoms – electrons are free to move
Van der Waals Bonds develop because the zone around an atom in which the electrons are found is not spherical – electrical charges develop


A chemical substance is transformed as a result of its interaction with another substance or with energy
Can result in a phase change, or the formation or dissociation of molecules
Usually involve making or breaking of chemical bonds


Chemical reactions involve energy changes
Enthalpy: the system's internal energy plus the product of its pressure and volume
Exothermic reactions: release energy during the reaction, increasing the temperature of the surroundings (combustion reactions)
Endothermic reactions: absorb energy during the reaction, decreasing the temperature of the surrounding (phase change of melting ice)

Acids and Bases

     Produce H+ ions in aqueous solutions (Arrhenius definition)
     Is a proton (H+ ion) donor (Bronsted-Lowry definition)
Produce OH- ions in aqueous solutions (Arrhenius definition)
Is a proton (H+ ion) acceptor (Bronsted-Lowry definition)
Acid strength is measured by pH.
Scale runs from 0 to 14.
pH = 7 is neutral (neither acidic or basic)

Redox: Reduction and Oxidation Reactions

Named for reactions involving elements reacting with oxygen (oxidation) or losing an oxygen (reduction).
Oxidation: is the loss of electrons
Reduction: is the gain of electrons
Oxidation and reduction reactions must occur together.

Dynamic Equilibrium

In a chemical reaction, the condition in which the rate of forward reaction equals the rate of the reverse reaction.

Chemical Laws

Avogadro's Law: Equal volumes of all gases, at the same temperature and pressure, have the same number of molecules

Beer–Lambert Law: The transmittance of material sample is related to its optical depth and to its absorbance.

Boyle's Law: Describes how the pressure of a gas tends to increase as the volume of the container decreases, if the temperature and amount of gas remain unchanged.

Charles's Law: Describes how gases tend to expand when heated.

Fick's Laws of Diffusion: Describes how concentrated solute in one portion of a container diffuses to fill the whole container.

Gay-Lussac's Law: Describes how as a gas's temperature increases, then so does its pressure, if the mass and volume of the gas are held constant.

Le Chatelier's Principle: Describes and used to predict the effect of a change in conditions on some chemical equilibria.

Henry's Law: States that the amount of dissolved gas is proportional to its partial pressure in the gas phase.

Hess’s Law: States that the total enthalpy change during the complete course of a chemical reaction is the same whether the reaction is made in one step or in several steps.

Ideal Gas Law: PV = nRT Combination of Gay-Lussac's Law, Charles's Law, Boyle's Law and Avogadro’s Law P = pressure, V = volume, n = number of moles, R = ideal gas constant, T = absolute temperature

Law of Conservation of Energy: States that the total energy of an isolated system remains constant.

Law of Conservation of Mass: States that for any system closed to all transfers of matter and energy, the mass of the system must remain constant over time. The law implies that during chemical reactions, mass can neither be created nor destroyed.

Law of Definite Composition: States that a given chemical compound always contains its component elements in fixed ratio (by mass) and does not depend on its source and method of preparation.

Law of Multiple Proportions: If two elements form more than one compound between them, then the ratios of the masses of the second element which combine with a fixed mass of the first element will be ratios of small whole numbers.

Raoult's Law: States that the partial vapor pressure of each component of an ideal mixture of liquids is equal to the vapor pressure of the pure component multiplied by its mole fraction in the mixture.