Ch 2

• Atom


  The basic structural unit of an element. The smallest unit of an element that retains the chemical properties of that element.
• Atomic particles

  • Electrons
  • Protons
  • Neutrons
• Electrons


  • Negatively charged particles located outside of the nucleus of an atom
  • Move very rapidly in a relatively large volume of space while the nucleus is small and dense
  • Protons and electrons have charges that are equal in magnitude but opposite in sign
• Neutral atom


  Has the same number of protons and electrons (no electrical charge)
• Atomic number (Z)


  The number of protons in the atom
• Mass number (A)


  Sum of the number of protons and neutrons
• Atomic Calculations

  1. mass number = number of protons + number of neutrons

  2. number of neutrons = mass number − number of protons

  3. number of neutrons = mass number − atomic number

  4. number of neutrons = A − Z
• Isotopes


  • Atoms of the same element having different masses
  • Contain same number of protons
  • Contain different numbers of neutrons
• Isotopes of the same element have identical chemical properties
• Some isotopes are radioactive
• Atomic mass


  The weighted average of the masses of all the isotopes that make up an element
• Determining Atomic Mass

  1. Step 1: Convert the percentage to a decimal fraction

  2. Step 2: Multiply the decimal fraction by the mass of that isotope to determine the contribution of each isotope

  3. Step 3: Add the mass contributed by each isotope
• Periodic Law


  The physical and chemical properties of the elements are periodic functions of their atomic numbers
• Classification of the Elements

  • Metals
  • Nonmetals
  • Metalloids
• Metals

  • Tend to lose electrons during a chemical change
  • Found primarily in the left of the periodic table
  • High thermal and electrical conductivities
  • High malleability and ductility
  • Metallic luster
  • Solid at room temperature
• Nonmetals

  • May gain electrons, forming negative ions
  • Found in the right of the periodic table
  • Brittle
  • Powdery solids or gases
  • Opposite of metal properties
• Electron configuration

  Describes the arrangement of electrons in atoms
• Bohr's model of the hydrogen atom


  • Did not clearly explain the electron structure of other atoms
  • Electrons in very specific locations, principal energy levels
  • Wave properties of electrons conflict with specific location
• Schröedinger's equations


  Determine the probability of finding an electron in specific region in space, quantum mechanics
• Principal energy levels


  • Regions where electrons may be found
  • Have values designated as n
  • The larger the value of n, the higher the energy level and the farther away from the nucleus the electrons are
  • The number of sublevels in a principal energy level is equal to n
• Electron capacity of a principal energy level

  • n = 1 can hold 2 electrons
  • n = 2 can hold 8 electrons
  • n = 3 can hold 18 electrons
• Sublevel


  A set of energy-equal orbitals within a principal energy level
• Subshells increase in energy: s < p < d < f
• Electrons in 3d subshell have more energy than electrons in the 3p subshell
• Possible subshells in each principal energy level

  • 1: 1s
  • 2: 2s, 2p
  • 3: 3s, 3p, 3d
  • 4: 4s, 4p, 4d, 4f
• Atomic Orbital


  A specific region of a sublevel containing a maximum of two electrons
• Orbitals are named by their sublevel and principal energy level
• s orbital is spherically symmetrical
• p orbital has a shape much like a dumbbell
• Number of orbitals in each subshell


  • s: 1
  • p: 3
  • d: 5
• Electron Configuration


  The arrangement of electrons in atomic orbitals
• Aufbau Principle


  Electrons fill the lowest-energy orbital that is available first
• Pauli Exclusion Principle: each orbital can hold up to two electrons with their spins in opposite directions (paired)
• Hund's Rule: each orbital in a subshell is half-filled (with one electron) before any orbital becomes completely filled (with two electrons)
• Orbital diagrams show orbitals as boxes and electrons as arrows
• Rules for Writing Electron Configurations

  1. Obtain the total number of electrons in the atom from the atomic number

  2. Electrons in atoms occupy the lowest energy orbitals that are available beginning with 1s

  3. Fill subshells according to the order depicted
• The s sublevel has one orbital
• Aufbau Principle


  Helps determine the electron configuration: Electrons fill the lowest-energy orbital that is available first
• Remember s<p<d<f in energy
• Pauli Exclusion Principle


  Each orbital can hold up to two electrons with their spins in opposite directions (paired)