AH Chem Notes - Unit 1 Inorganic

• Electromagnetic radiation

  Can be described as a wave (has a wavelength and frequency), and as a particle, and is said to have a dual nature
• Wavelength (λ)

  Horizontal distance from one point on a wave to the same point on the next wave, measured in metres (m). When radiation in the visible part of the electromagnetic spectrum is described, it is common to specify the wavelength in nanometres (nm, 1 nm = 1 x 10-9 m). Chemistry usually measures in nm
• Frequency (f)

  Number of waves that pass a point in one second, measured in s-1 or Hertz (Hz)
• Velocity (c)

  The velocity of electromagnetic radiation is constant and has a value of approximately 3 x 108 ms-1 (speed of light in a vacuum)
• Velocity, frequency and wavelength

  Linked by: c=fλ
• Types of electromagnetic radiation arranged in order of wavelength

  • Gamma
  • X-rays
  • Ultraviolet
  • Visible
  • Infrared
  • Microwaves
  • Radiowaves
• Photons


  Particles that electromagnetic radiation appears to behave as when absorbed or emitted by matter
• Photon energy

  Proportional to the frequency of the radiation
• Atomic Spectroscopy

  When energy is transferred to atoms, electrons within the atoms may be promoted to higher energy levels. To allow electrons to return to their original levels, energy must be lost from the atom in the form of a photon
• Atomic absorption spectroscopy

  Electromagnetic radiation is directed at an atomised sample. Radiation is absorbed as electrons are promoted to higher energy levels
• Atomic emission spectroscopy

  High temperatures are used to excite the electrons within atoms. As the electrons drop to lower energy levels, photons are emitted
• Atomic orbitals

  Electrons behave as standing (stationary) waves in an atom. These are waves that vibrate in time but do not move in space
• Quantum numbers

  Electrons within atoms have fixed amounts of energy called quanta. As a result, electrons can be defined in terms of quantum numbers
• Quantum numbers

  • Principal quantum number n
  • Angular momentum quantum number l
  • Magnetic quantum number ml
  • Spin magnetic quantum number ms
• Electron shells

  The shell nearest the nucleus is the first shell and any electron in the first shell has the principal quantum number n=1. Electrons in the second shell have n=2. The higher the value of n, the further the electrons are from the nucleus
• Electron subshells


  • s
  • p
  • d
  • f
• Orbital shapes

  Each type of subshell (s, p, d and f) contains one or more orbitals. These are defined by the angular momentum quantum number, symbol l
• Orbital types

  • s
  • p
  • d
• Pauli exclusion principle

  No two electrons in any one atom can have the same set of four quantum numbers
• Ground state configuration

  Electrons within atoms in the ground state configuration (lowest possible electronic configuration) are arranged according to the aufbau principle, Hund's rule, and the Pauli exclusion principle
• Periodic table blocks

  • s-block
  • p-block
  • d-block
  • f-block
• Helium is from the s-block, with its outer (and only) electrons in the 1s atomic orbital, although its chemical properties are more similar to the p-block noble gases due to its full outer shell
• Electron configurations for the first 20 elements

  • hydrogen 1s1
  • helium 1s2
  • lithium 1s22s1
  • beryllium 1s22s2
  • boron 1s22s22p1
  • carbon 1s22s22p2
  • nitrogen 1s22s22p3
  • oxygen
  • fluorine
  • neon
  • sodium
  • magnesium
  • aluminium
  • silicon
  • phosphorus
  • sulfur
  • chlorine
  • argon
  • potassium
  • calcium
• Electron configurations for transition metal elements

  • scandium 1s22s22p63s23p63d14s2
  • titanium
  • vanadium
  • chromium (Ar) 4s1 3d5
  • manganese
  • iron
  • cobalt
  • nickel
  • copper (Ar) 4s1 3d10
  • zinc
• Electron configurations for transition metal ions

  • Mn2+
  • Ti3+
  • Co2+
  • Co3+
  • Ni2+
  • Cu+
  • Fe3+
  • Fe2+
  • V4+
  • Mn4+
• Transition metals are metals with an incomplete d subshell in at least one of their ions
• Ions of zinc and scandium do not have an incomplete d subshell, so they are often not regarded as transition metals
• Fe(III) compounds are more stable than Fe(II) compounds

  Because Fe(III) has a higher oxidation state and more stable electron configuration
• Oxidation state

  When an element has a specific oxidation number
• Rules for determining oxidation number

  1. All uncombined elements are given the oxidation number zero

  2. For ions containing single atoms, the oxidation number is the same as the charge on the ion

  3. In nearly all of its compounds, oxygen has an oxidation number of -2

  4. In nearly all of its compounds, hydrogen has an oxidation number of +1

  5. In molecules and neutral compounds, the sum of all the oxidation numbers is equal to zero

  6. In polyatomic ions, the sum of all the oxidation numbers is equal to the overall charge on the ion
• Common oxidation states for elements in the first transition series

  • (-2), -1, 0, +1, +2, +3, +4, (+5), (+6), (+7)
• Compounds of the same transition metal in different oxidation states


  May have different colours
• Oxidation states and colours of vanadium compounds

  • VO2+
  • VO2+
  • V3+
  • V2+
• Oxidation

  Increase in oxidation number
• Reduction

  Decrease in oxidation number
• Conversion of VO3- to V3+ is an example of reduction
• Compounds containing metals in high oxidation states are often oxidising agents, whereas compounds with metals in low oxidation states are often reducing agents
• Transition metal complex

  A central metal ion surrounded by ligands
• Ligand

  An electron donor that forms dative covalent bonds with the central metal atom or ion
• Types of ligands

  • Monodentate (one donor atom)
  • Bidentate (two donor atoms)
  • Hexadentate (six donor atoms)