Inorganic Chemistry

• Periodicity

  Repeating pattern/trend of physical or chemical properties/reactions
• Classification of elements

  • s block
  • p block
  • d block
  • f block
• Periods

  • Rows in the periodic table
  • Number of the period corresponds to number of electron shells
• Groups

  • Vertical columns in the periodic table
  • Correspond to the number and the subshell of the outermost electrons
• Atomic radius (period)

  • Along a period, atomic radius decreases
  • Number of protons is increasing, but the amount of shielding (core electrons repelling outer electrons) remains the same
  • This increases the force of attraction for the electrons and causes them to be attracted more closely to the nucleus.
  • This leads to decreasing atomic radius
• Atomic radius (group)

  • Atomic radius increases down a group
  • There is an increasing number of full shells between the nucleus and the outermost electrons
• First ionisation energy

  The energy required to remove one electron from a mole of gaseous atoms to form a mole of gaseous +1 ions

  
• The first ionisation energy increases along a period due to the increasing nuclear force of attraction arising from the increasing number of protons and the same amount of shielding
• Aluminium and Sulfur have ionisation energies that are lower than would be expected based on the overall trend
• Aluminium

  The electron removed when aluminium is ionised is in a 3p sub-level, which is higher in energy than the 3s electron removed when magnesium is ionised. Removing an electron from a higher energy orbital requires less energy.
• Sulfur

  The highest energy electron in both phosphorus and sulfur is in the 3p sub-level. However, in sulfur this electron is paired, while in phosphorus each 3p orbital is singly occupied. Mutual repulsion between paired electrons means less energy is required to remove one of them.
• Melting point

  The strength of the metallic bonds increase due to increasing charge density and number of free electrons
• Silicon

  • 4th element of the period
  • Highest melting point of the period
  • Giant covalent structure similar to diamond
  • All four bonds must be broken to melt silicon and this takes a very large amount of energy
• Phosphorus, sulfur, and chlorine(5,6,7)


  • Much lower than that of silicon (4)
  • Have a simple molecular structure with weak van der Waal's forces holding the molecules together
  • m.p.s depending on the strength of their van der wals forces.
• Sulfur

  • S8 has the highest van der waals forces out of Cl2, P4 and S8.
  • This is because its molecules have the formula S8 with the highest size and Mr
  • This means it also has the highest m.p.
• Chlorine

  • Cl2 is a small diatomic molecule with the lowest Mr of simple molecules (Argon is not a molecule)
  • Therefore has the lowest boiling point of the simple molecules
• Argon

  • Monatomic gas with very low van der Waal's forces between atoms
  • Argon has the lowest boiling point of the period 3 elements
• What does the number of the period correspond to?

  Number of electron shells
• What do groups correspond?

  The number and shell of valence electrons
• What is shielding?

  Electrons closer to nucleus "defend" electrons further away from +ve charge of nucleus
• What is the the first ionisation energy?

  Measure of the strength of the attracting force between the outermost electrons and the nucleus
• Why does the melting point increase across the first 3 metallic elements of each period?

  • Charge density increases (ratio of charge to size) - Al3+ ions are small and have a higher charge than Mg2+/Na2+ ions. As they have the highest charge density and are attracted more strongly to the delocalised electrons
  • Number of free electrons increases - Na has one free electron per metal ion, Mg has 2 and Al has 3. This means that more energy must supplied to break the individual attractions in Mg and Al compared to Na
• Melting points depend on bonding and structure

  
  

  • Metallic bonding: Strong electrostatic force of attraction between +ve ions and -ve electrons. Lots of energy needed to overcome the attraction
  • Giant covalent structure: Many covalent bonds which are strong. Lots of E needed to break bonds
  • Simple molecular structure: Weak IMF b/w molecules. Little E needed to break forces
• What are groups?

  Vertical columns in the periodic table
• EQ: Explain why the atomic radius across a period.

  1. The number of protons increases

  2. But the shielding stays the same

  3. Since the attraction is greater, the electron shells get attracted more to the nucleus
• EQ: Identify the element with the largest atomic radius in a given period.

  1. The G1 element

  2. Least number of protons for a given amount of shielding in the period

  3. Thus has the least attraction out of the other elements in the period, leading to a smaller radius
• EQ: Explain why the melting point of S8 is greater than that of P4

  The van der waals forces in S8 are greater than those in P4 due to greater Mr and size
• What are redox reactions?

  Reactions that involve a transfer of electrons from the reducing agent to the oxidising agent
• The change in oxidation state of of an element in a compoud or an ion is used to identify the element that has been oxidised/reduced in a given reaction.

  
  

  Separate half equations are written for the oxidation/reduction processes. These half equations can be combined to give an overall equation for any redox reaction.
• Formal oxidation states

  • Hypothetical
  • Useful tool to help us determine whether elements have been oxidised or reduced when they form compounds or when substances react with each other.
  • Technique, not real - calculated as if a substance is completely ionic.
  • We define any (uncombined) element to be in it's 0 oxidation state when it is formed from its element
  • Whenever a chemical compound is formed from its elements, a change in oxidation state occurs, meaning that oxidation and reduction (redox) has happened.
• Oxidation of metals is also called corrosion
• Reducing agent: Oxidised, loses electrons (electron donor)

  Oxidising agent: Reduces, gaines electrons (electron acceptor)
• Unusual reactions

  If there is no overall change in oxidation state, it is NOT redox, it is an acid/base reaction - we call it a 'neutralisation' reaction

  If a species is both oxidised and reduced in a reaction, it is called a 'disproportionation' reaction.
• Oxidation state can also be given in roman numerals ie (II)
• Rules for assigning oxidation states

  • Uncombined element = 0
  • Applied regardless of structure, ie S8 or P4
  • Sum of oxidation states in a neutral compound is 0
  • Sum of oxidation states in an ion is the magnitude of the charge of the ion
  • Oxidation state of a simple metal ion is the charge on the ion
  • The more electronegative element in a compound is assigned a negative oxidation state
• Elements with the same oxidation state

  • G1 : always +1
  • G2 : always +2
  • Oxygen : usually -2 (exception : peroxides (anything with 2 Os) and F2O)
  • Hydrogen : usually +1 (exception : metal hydrides)
  • Fluorine : always -1
  • Chlorine : usually -1 (exception : compounds with O or F)
• Remember:

  Oxidation states do NOT have to be the same on both sides this is NOT a chemical reaction to balance

  If a species loses lots of electrons it is a powerful oxidising agent
• In acidic conditions:

  • Balance the equations by adding the correct number of electrons AND...
  • (if required) either H20 or H+ to one side or the other to balance any oxygen atoms that are present as part of the ion
• Oxidation:

  ... + H20 -> ... + H+

  Reduction:

  ... + H+ -> ... + H20
• Combining half equations

  • To combine half equations there must be the same number of electrons on each side in order for them to be able to cancel out
  • Multiply out just like in a simultaneous equation