Atom structures and periodic table

• Atoms


  The smallest part of an element that can exist
• Chemical symbols


  Represent an atom of an element e.g. Na represents an atom of sodium
• Compounds


  Formed from elements by chemical reactions, contain two or more elements chemically combined in fixed proportions, can be represented by formulae
• Mixtures


  Consist of two or more elements or compounds not chemically combined together, can be separated by physical processes
• Development of the model of the atom

  1. First thought to be tiny spheres

  2. Discovery of electron led to plum pudding model

  3. Alpha particle scattering experiment led to conclusion mass concentrated in nucleus

  4. Bohr suggested electrons orbit nucleus at specific distances

  5. Later experiments showed positive charge subdivided into protons

  6. Chadwick's work provided evidence for existence of neutrons
• Atomic number


  The number of protons in an atom of an element
• Relative charge of subatomic particles

  • Proton: +1
  • Neutron: 0
  • Electron: -1
• Atom


  Has an overall charge of 0, number of protons = number of electrons
• Size and mass of atoms

  • Atoms are very small, radius about 0.1 nm, nucleus radius less than 1/10,000 of atom but holds almost all mass
  • Relative mass: Proton 1, Neutron 1, Electron very small
• Mass number


  Sum of protons and neutrons in an atom
• Isotopes


  Atoms of the same element with different numbers of neutrons
• Relative atomic mass


  An average value that takes account of the abundance of the isotopes of the element
• Calculating relative atomic mass

  • Carbon has 2 isotopes: carbon-14 with 20% abundance, carbon-12 with 80% abundance. Relative atomic mass = ((14 x 20) + (12 x 80)) / 100 = 12.4
• Electronic structure


  Electrons occupy the lowest available energy levels (shells closest to nucleus), can be represented as 2,8,1 for sodium

  
• Transition metals compared to group 1

  • Harder and stronger
  • Higher melting points (except for mercury) and higher densities
  • Much less reactive and don't react as vigorously with oxygen or water
• Transition metals


  • chromium
  • manganese
  • iron
  • cobalt
  • nickel
  • copper
• chromium

  lustrous, brittle, hard metal
• manganese

  hard and very brittle, difficult to fuse, but easy to oxidise
• iron

  good conductor, rusts easily in air, strong, ductile
• cobalt

  malleable, brittle, hard, high melting point
• nickel

  hard, malleable, and ductile, fairly good conductor of heat and electricity
• copper

  highly ductile and conductive, malleable and soft
• Typical properties of transition metals

  • They have ions with many different charges
  • Form coloured compounds
  • Are useful as catalysts
• Uses of transition metals as catalysts


  • Decomposition of hydrogen peroxide (2H2O2 -> 2H2O + O2)
  • Haber process (N2 + 3H2 ↔ 2NH3)
  • Manufacture of margarine (adding H2 to double bonds)
• Periodic table


  Elements are arranged in order of atomic (proton) number (smaller number) and so that elements with similar properties are in columns, known as groups

  
• Elements in the same periodic group

  Have the same amount of electrons in their outer shell, which gives them similar chemical properties
• John Newlands


  • Ordered his table in order of atomic weight
  • Realised similar properties occurred every eighth element – 'law of octaves' but broke down after calcium
• Dmitri Mendeleev


  • Ordered his table in order of atomic mass, but not always strictly – i.e. in some places he changed the order based on atomic weights
  • Left gaps for elements that he thought had not been discovered yet
• The table is called a periodic table because similar properties occur at regular intervals

  
• Elements with similar properties are found in the same column (groups)
• Elements with properties predicted by Mendeleev were discovered and filled the gaps
• Knowledge of isotopes made it possible to explain why the order based on atomic weights was not always correct
• When electrons, protons and neutrons were discovered in the early 20th century, elements were ordered in atomic (proton) number
• When this was done, all elements were placed in appropriate groups
• Metals


  Elements that react to form positive ions
• Non-metals


  Elements that do not form positive ions
• Group 1 - Alkali metals


  • They have characteristic properties due to the single electron in their outer shell
  • Metals in group one react vigorously with water to create an alkaline solution and hydrogen
  • They all react with oxygen to create an oxide
  • They all react with chlorine to form a white precipitate
  • The reactivity of the elements increases going down the group
• Group 0 - Noble gases

  • They have 8 electrons in their outer shell (except helium, which has 2)
  • All of them (including helium) have full outer shells
  • They are unreactive and do not easily form molecules, because they have a stable arrangement of electrons (full outer shell)
  • The boiling points of the noble gases increase with increasing relative atomic mass (going down the group)
• Group 7 - The halogens


  • Similar reactions due to their seven electrons in their outer shell
  • Non-metals and exist as molecules made of pairs of atoms (e.g. Cl2)
  • They react with metals to form ionic compounds in which the halide ion carries a -1 charge
  • They react with nonmetals to form covalent compounds, where there is a shared pair of electrons
  • As you go down the group, relative molecular mass, melting point and boiling point all increase
  • Reactivity decreases down the group because halogens react by gaining an electron (to increase their number of outer shell electrons from 7 to 8) and the number of shells of electrons increases down the group, so down the group the element attracts electrons from other atoms less, so can't react as easily
  • A more reactive halogen (one from higher up group 7) can displace a less reactive one in an aqueous solution of its salt