protons, neutrons and electrons textbook reading notes

I. Introduction

  • Explores atoms as the building blocks of matter.

  • Introduces fundamental aspects of modern atomic theory.

  • Examines three types of subatomic particles: protons, neutrons, and electrons.

  • Interprets atomic symbols.

II. Learning Objectives for Subatomic Particles

  • Recognize that theatom is comprised of protons, neutrons, and electrons; distinguish among these building blocks.

  • Know the names and symbols of certain elements.

III. The Structure of the Atom

  • Early Atomic Understanding (early 20th century)

    • Thomson, Millikan: revealed the charge and mass of electrons (e^-).

    • Rutherford: indicated that positively charged subatomic particles—protons (p^+)—are much more massive than electrons (by a factor of about 2,000) and highly concentrated in the atom’s nucleus.

    • Neutrons (n^0): bear no electrical charge, have masses similar to protons, and co-locate with protons in the nucleus.

  • Relative Sizes

    • If the nucleus were the size of a blueberry, the atom would be about the size of a football stadium.

    • Their diameters differ by about 20,000-60,000 times.

  • Electron Cloud

    • Most of the volume of an atom is occupied by the “cloud” of electrons.

    • The exact positions of electrons cannot be known; they are distributed about the volume.

    • Example: For a helium atom, the electron cloud spans about 1 angstrom (10^{-10} m) in diameter, while the nucleus is about 1 fm (10^{-15} m).

  • Units of Measure for Atoms

    • Atoms are extremely small (e.g., a carbon atom weighs less than 2 \times 10^{-23} g, and an electron has a charge of less than 2 \times 10^{-19} C).

    • Atomic Mass Unit (amu):

      • Originally defined based on hydrogen, then oxygen; since 1961, based on carbon-12.

      • Exactly \frac{1}{12} of the mass of one carbon-12 atom.

      • 1 \text{ amu} = 1.6605 \times 10^{-24} g.

      • Alternative units: Dalton (Da) and unified atomic mass unit (u).

    • Fundamental Unit of Charge (e):

      • Equals the magnitude of the charge of an electron.

      • e = 1.602 \times 10^{-19} C.

  • Properties of Subatomic Particles

    • Electron (e^-):

      • Location: outside nucleus

      • Unit Charge: 1-

      • Mass: 0.0005486 amu (it takes 1836 electrons to equal the mass of one proton).

    • Proton (p^+):

      • Location: nucleus

      • Unit Charge: 1+

      • Mass: 1.0073 amu.

    • Neutron (n^0):

      • Locat ion: nucleus

      • Unit Charge: 0

      • Mass: 1.0087 amu.

  • Atomic Number (Z) and Mass Number (A)

    • Atomic Number (Z):

      • Number of protons in the nucleus of an atom.

      • The defining trait of an element; determines the identity of the atom.

      • In a neutral atom, Z = number of electrons.

    • Mass Number (A):

      • The total number of protons and neutrons in an atom (A = \text{protons} + \text{neutrons}).

      • Number of neutrons = A - Z.

  • Isotopes

    • Atoms with the same atomic number but different mass numbers.

    • Differ only in the number of neutrons within the nucleus.

  • Ions

    • Electrically charged atoms formed when the numbers of protons and electrons are not equal.

    • Atomic charge = \text{number of protons} - \text{number of electrons}.

    • Anion: A negatively charged atom or molecule (gains one or more electrons, contains more electrons than protons).

    • Cation: A positively charged atom or molecule (loses one or more electrons, contains fewer electrons than protons).

    • Example: A neutral sodium atom (Z = 11, 11 electrons) that loses one electron becomes a 1+ cation (11 - 10 = 1+).

    • Example: A neutral oxygen atom (Z = 8, 8 electrons) that gains two electrons becomes a 2- anion (8 - 10 = 2-).

IV. Chemical Symbols

  • Definition: An abbreviation used to indicate an element or an atom of an element (e.g., Hg for mercury).

  • Rules for Symbols:

    • Only the first letter of a symbol is capitalized (e.g., Co for cobalt, not CO for carbon monoxide).

    • Most symbols have one or two letters; three-letter symbols are used for some elements with atomic numbers greater than 112.

  • Naming Conventions (IUPAC):

    • Traditionally, the discoverer names the element.

    • Until name recognition by the International Union of Pure and Applied Chemistry (IUPAC), a recommended name is based on Latin words for its atomic number (e.g., unnilhexium for element 106).

    • Elements are now often named after scientists or locations (e.g., seaborgium (Sg) for element 106, honoring Glenn Seaborg).

V. Key Concepts and Summary

  • An atom consists of a small, positively charged nucleus (protons and neutrons) surrounded by electrons.

  • The nucleus diameter is about 20,000-60,000 times smaller than that of the atom.

  • Units:

    • Atomic mass expressed in atomic mass units (amu), defined as exactly \frac{1}{12} of the mass of a carbon-12 atom (equal to 1.6605 \times 10^{-24} g).

    • The fundamental unit of charge (e) equals the magnitude of the charge of an electron (1.602 \times 10^{-19} C).

  • Subatomic particle properties:

    • Protons: relatively heavy particles with a charge of 1+ and a mass of 1.0073 amu.

    • Neutrons: relatively heavy particles with no charge and a mass of 1.0087 amu.

    • Electrons: light particles with a charge of 1- and a mass of 0.0005486 amu.

  • Atomic Number (Z): The number of protons in the nucleus, defining an atom’s elemental identity.

  • Mass Number (A): The sum of the numbers of protons and neutrons in the nucleus.

  • A neutral atom contains equal numbers of electrons and protons.

  • A chemical symbol identifies the atoms in a substance using one-, two-, or three-letter abbreviations.

VI. Glossary

  • Contains definitions for important terms related to atomic structure.