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Chemistry LibreTexts

4: Atoms and Elements

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    • 4.1: Cutting Alumimun until you get Atoms
      Take some aluminum foil. Cut it in half. Now you have two smaller pieces of aluminum foil. Cut one of the pieces in half again. Cut one of those smaller pieces in half again. Continue cutting, making smaller and smaller pieces of aluminum foil.
    • 4.2: Indivisible: The Atomic Theory
      You learned earlier how all matter in the universe is made out of tiny building blocks called atoms. All modern scientists accept the concept of the atom, but when the concept of the atom was first proposed about 2,500 years ago, ancient philosophers laughed at the idea. It has always been difficult to convince people of the existence of things that are too small to see. We will spend some time considering the evidence (observations) that convince scientists of the existence of atoms.
    • 4.3: The Nuclear Atom
      While Dalton's Atomic Theory held up well, J. J. Thomson demonstrate that his theory was not the entire story. He suggested that the small, negatively charged particles making up the cathode ray were actually pieces of atoms. He called these pieces "corpuscles," although today we know them as electrons. Thanks to his clever experiments and careful reasoning, J. J. Thomson is credited with the discovery of the electron.
    • 4.4: The Properties of Protons, Neutrons, and Electrons
      Electrons are extremely small. The mass of an electron is only about 1/2000 the mass of a proton or neutron, so electrons contribute virtually nothing to the total mass of an atom. Electrons have an electric charge of −1, which is equal but opposite to the charge of a proton, which is +1. All atoms have the same number of electrons as protons, so the positive and negative charges "cancel out", making atoms electrically neutral.
    • 4.5: Elements: Defined by Their Number of Protons
      Scientists distinguish between different elements by counting the number of protons in the nucleus. Since an atom of one element can be distinguished from an atom of another element by the number of protons in its nucleus, scientists are always interested in this number, and how this number differs between different elements. The number of protons in an atom is called its atomic number (Z). This number is very important because it is unique for atoms of a given element.
    • 4.6: Looking for Patterns: The Periodic Table
      Certain elemental properties become apparent in a survey of the periodic table as a whole. Every element can be classified as either a metal, a nonmetal, or a metalloid (or semi metal). A metal is a substance that is shiny, typically (but not always) silvery in color, and an excellent conductor of electricity and heat. Metals are also malleable (they can be beaten into thin sheets) and ductile (they can be drawn into thin wires).
    • 4.7: Ions - Losing and Gaining Electrons
      Atom may lose valence electrons quite to obtain a lower shell that contains an octet. Atoms that lose electrons acquire a positive charge as a result because they are left with fewer negatively charged electrons to balance the positive charges of the protons in the nucleus. Positively charged ions are called cations. Most metals become cations when they make ionic compounds.
    • 4.8: Isotopes - When the Number of Neutrons Varies
      All atoms of the same element have the same number of protons, but some may have different numbers of neutrons. For example, all carbon atoms have six protons, and most have six neutrons as well. But some carbon atoms have seven or eight neutrons instead of the usual six. Atoms of the same element that differ in their numbers of neutrons are called isotopes. Many isotopes occur naturally.
    • 4.9: Atomic Mass: The Average Mass of an Element’s Atoms
      In chemistry we very rarely deal with only one isotope of an element. We use a mixture of the isotopes of an element in chemical reactions and other aspects of chemistry, because all of the isotopes of an element react in the same manner. That means that we rarely need to worry about the mass of a specific isotope, but instead we need to know the average mass of the atoms of an element.