Chemistry Edexcel GCSE

• Dalton model of an atom

  Has changed over time because of the discovery of subatomic particles
• John Dalton (1803)
• Relative charge and relative mass of subatomic particles


  • Proton: +1, 1
  • Neutron: 0, 1
  • Electron: -1, 1/1836
• Atoms contain equal number of protons and electrons
• Atoms are neutral


  The charges on a proton are +1 and on an electron are -1, so the amount of protons = amount of electrons, to cancel out the charges
• The nucleus of an atom is very small compared to the overall size of the atom
• Most of the mass of an atom is concentrated in the nucleus
• Mass (nucleon) number


  Number of protons + neutrons
• Atoms of a given element have the same number of protons in the nucleus, which is unique to that element
• Isotopes


  Different atoms of the same element containing the same number of protons but different numbers of neutrons in their nuclei
• Calculating the numbers of protons, neutrons and electrons in atoms


  1. Atomic (proton) Number = number of protons (= number of electrons if it's an atom, because atoms are neutral)

  2. Number of neutrons = mass number - atomic number
• Calculating the relative atomic mass of an element from the relative masses and abundances of its isotopes


  1. R.A.M. = (mass of isotope-A x % of isotope-A) + (mass of isotope-B x % of isotope-B) / 100

  2. Example calculation for chlorine: R.A.M. = (35 x 75) + (37 x 25) / 100 = 35.5
• Ionic bonds


  Formed by the transfer of electrons between atoms to produce cations and anions

  
• Metals + nonmetals: electrons in the outer shell of the metal atom are transferred

  
• Cation


  Positively charged ion formed when a metal atom loses electrons
• Anion


  Negatively charged ion formed when a nonmetal atom gains electrons
• Electron transfer during the formation of an ionic compound can be represented by a dot and cross diagram

  
• Ion


  An atom or group of atoms with a positive or negative charge
• Calculating the numbers of protons, neutrons and electrons in simple ions

  1. Atomic number = proton number = number of protons

  2. Mass number = nucleon number = number of protons + neutrons

  3. In an ion, the number of electrons is different to the number of protons, so to work out electrons in an ion: 1) work out how many electrons an atom of the element would have (same as proton number), 2) work out how many electrons have been lost or gained (using charge - -ve means electrons gained, +ve means electrons lost), 3) calculate number of electrons in atom plus electrons gained or minus electrons lost
• Deducing the formulae of ionic compounds


  1. Oxide -> involves O2- ion

  2. Hydroxide -> involves OH1- ion

  3. Halide -> involves a -1 halide ion

  4. Nitrate -> involves NO31- ion

  5. Carbonate -> involves CO32- ion

  6. Sulfate -> involves SO42- ion

  7. To deduce the formula, balance the +ve and -ve charges to make the overall charge 0, by writing a little number below the element or drawing a bracket around ions with more than one element

  
• Types of ionic compounds


  • Oxides
  • Hydroxides
  • Halides
  • Nitrates
  • Carbonates
  • Sulfates
• Oxide


  Involves O2- ion (e.g. sodium oxide: Na2O)
• Hydroxide


  Involves OH1- ion (e.g. sodium hydroxide: NaOH)
• Halide


  Involves a -1 halide ion (e.g. sodium chloride NaCl)
• Nitrate


  Involves NO31- ion (e.g. sodium nitrate: NaNO3)
• Carbonate


  Involves CO32- ion (e.g. sodium carbonate: NaCO3)
• Sulfate


  Involves SO42- ion (e.g. sodium sulfate: NaSO4)
• Deducing formula of ionic compounds


  1. Balance out the + and - charges to make the overall charge 0

  2. Write a little number below the element (e.g. Cl3)

  3. For ions with more than one element, draw a bracket round first (e.g. (SO4)2)
• Ionic compound


  • Giant structure of ions
  • Held together by strong electrostatic forces of attraction between oppositely charged ions
  • Forces act in all directions in the lattice, called ionic bonding
  • Lattice has a regular arrangement of ions

  
• Ionic compound


  • Sodium chloride (salt): Na+ (small blue particles) and Cl- (larger green ones)

  
• Covalent bonding


  • Occurs in most non-metallic elements and in compounds of nonmetals
  • When atoms share pairs of electrons, they form covalent bonds
• Covalent bonding results in the formation of molecules
• Types of covalently bonded substances


  • Small molecules (e.g. HCl, H2, O2, Cl2, NH3, CH4)
  • Very large molecules, such as polymers
  • Giant covalent structures (macromolecules) (e.g. diamond, silicon dioxide)
• Atoms are smaller than small molecules
• Formation of simple molecular, covalent substances


  Using dot and cross diagrams, including: hydrogen, hydrogen chloride, water, methane, oxygen, and carbon dioxide

  
• Types of substances


  • Ionic
  • Simple molecular (covalent)
  • Giant covalent
  • Metallic
• Ionic compounds


  • Made up of a metal and a nonmetal
  • Have regular structures (giant ionic lattices) with strong electrostatic forces of attraction in all directions between oppositely charged ions
  • Have high melting and boiling points
  • Conduct electricity when melted or dissolved in water (ions are free to move)
  • Do not conduct electricity as solids (ions are fixed and cannot move)
  • Often dissolve in water to form an aqueous solution
• Simple molecular compounds


  • Usually gases or liquids with low boiling and melting points
  • Made up of nonmetal elements
  • Have weak intermolecular forces between the molecules (not the covalent bonds)
  • Do not conduct electricity (small molecules do not have an overall electric charge)
  • Many are insoluble in water, but some are soluble
• Giant covalent structures


  • Made up of nonmetal elements
  • Solids with very high melting points (all atoms linked by strong covalent bonds)
  • Some can conduct electricity, others cannot
• Metals


  • Consist of giant structures of atoms arranged in a regular pattern
  • Always made up of just metallic elements
  • Electrons in the outer shell are delocalised and free to move through the whole structure
  • Sharing of delocalised electrons gives rise to strong metallic bonds
  • Have high melting and boiling points
  • Can conduct heat and electricity
  • Layers of atoms can slide over each other, so metals can be bent and shaped
  • Insoluble in water, but some will react with it