AP Chem

• Periodic table


  Gives very basic but essential information about each element

  
• Carbon


  • Symbol (C) tells you the name
  • Atomic number (6) tells you how many protons and neutrons in the element in addition to the amount of electrons surrounding the element when it is neutrally charged
  • Molar mass (12.011) tells you the average atomic mass of a single atom of carbon, measured in atomic mass units (amus) and the average mass of a mole of carbon atoms, measured in grams

  
• Periodic table


  • Horizontal rows are called periods
  • Vertical columns are called groups
• Groups


  • IA/1 - Alkali Metals
  • IIA/2 - Alkaline Earth Metals
  • B/3-12 - Transition Metals
  • VIIA/17 - Halogens
  • VIIIA/18 - Noble Gases
• Periodic table


  • Two rows beneath the periodic table are the lanthanides and actinides, the rare Earth elements, or inner transition metals

  
• Atom identity


  • Determined by the amount of protons in the nucleus, which also contains neutrons
  • Mass number is the sum of its neutrons and protons (valued at 1 each)
  • Electrons have significantly less mass than protons and neutrons and do not contribute to the mass
• Isotopes


  Atoms of an element with different numbers of neutrons, but same amount of protons
• Isotopes of carbon


  • Carbon-12 has 6 protons and 6 neutrons
  • Carbon-14 has 6 protons and 8 neutrons
• Molar mass


  Determined by the average of the mass numbers of all known isotopes of an element weighted by their percent abundance

  
• The most abundant isotope of selenium has a mass of 80, but there are four other naturally occurring isotopes
• Molar mass


  Gives a general idea of the most common isotope, e.g. carbon's molar mass of 12.01 indicates 99% of carbon is carbon-12

  
• Mole


  Connects all the different quantities in chemical equations, with coefficients telling you how many moles there are of that substance
• Avogadro's Number


  The number of atoms in a single mole of any given element, always 6.022x10^23

  
• Moles


  • Particles/(6.022x10^23)
  • Can convert between moles and grams using an element's atomic mass in amu, which also signifies grams
• Ideal gas law


  PV=nRT where P = pressure in atm, V = volume (liters), n = mol, R = .0821 (ideal gas constant), T = temperature (kelvin)
• STP (standard temperature and pressure)

  P =1 atm, T=273 Kelvin, where 1 mol ALWAYS occupies 22.4 liters
• Molarity (M)


  Expresses the concentration of a solution in terms of volume, used in equilibrium calculations, acids and bases, electrochemistry, etc.
• Percent composition


  The percent by mass of each element in a compound, calculated by dividing the mass of each element or component by the total mass of the compound
• Empirical formula

  The simplest ratio of one element in a compound to another
• Molecular formula

  Represents the actual formula for the substrate
• Determining empirical and molecular formulas

  1. Part A: Determine the empirical formula from percent composition

  2. Part B: If the molar mass is known, determine the molecular formula

  
• Electron energy

  • The farther an electron is from the nucleus, the more the potential energy is
  • Electrons can only exist at certain quantized energy levels
• Coulomb's law


  Describes the electrostatic force of attraction between the positively charged nucleus and negatively charged electrons, which decreases with distance

  
• Bohr model


  Correlates energy levels to rows on the periodic table, with the closer levels having fewer electrons
• Photoelectron spectroscopy


  Measures the ionization energies required to eject electrons from atoms, with the kinetic energy of the ejected electrons related to their binding energy
• Photoelectron spectrum


  • Y-axis shows relative number of electrons ejected, X-axis shows binding energy decreasing left to right
  • Peaks represent different energy levels and subshells with different electron configurations

  
• Electron configuration

  Describes the arrangement of electrons in an atom's orbitals, subshells, and shells in order of increasing energy

  
• Electron configuration rules


  • Aufbau principle - electrons fill orbitals in order of increasing energy
  • Pauli exclusion principle - electrons in an orbital must have opposite spin
  • Hund's rule - electrons singly occupy orbitals before pairing up
• Stable electron configuration


  Having a completely full outer shell, which for s and p block elements means 8 valence electrons

  
• Ions


  Atoms that have gained or lost electrons to become stable, with anions being negatively charged and cations being positively charged

  
• Ionic charges


  • Halogens as anions have -1 charge
  • Oxygen group elements as anions have -2 charge
  • Alkali metals as cations have +1 charge
  • Alkaline earth metals as cations have +2 charge
  • Transition metals can have multiple possible cation charges

  
• Periodic table trends


  • Electrons are attracted to protons in the nucleus, with attraction increasing with smaller distance and more protons
  • Electron shielding decreases attraction between outer electrons and nucleus
  • Completed electron shells are stable

  
• Metals, metalloids, nonmetals


  Metals on the left give up electrons, nonmetals on the right gain electrons, metalloids have both metallic and nonmetallic character
• Atomic radius


  Decreases left to right due to more protons and less electron shielding, increases down a group due to more electron shells
• Ionic radius


  Cations are smaller than atoms, anions are larger than atoms
• Ionization energy


  • Energy required to remove an electron, increases left to right and decreases down a group
  • Second ionization energy is greater than first due to stronger proton-electron attraction
• Electronegativity


  Measure of an element's ability to attract electrons, increases left to right and decreases down a group
• Fluorine is the most electronegative element
• Ionization energy


  The energy required to remove an electron from an atom
• Ionization energy decreases when moving down a group


  There are more electron shells that cause electron shielding and repulsion, causing the electrons to be not as attracted to the nucleus and easily removed