chemistry

• Atom


  The smallest particle of a substance that can exist
• Element


  Contains only one type of atom and cannot be split by chemical means
• Compound


  Two or more elements chemically combined, cannot be separated back into constituent elements
• Mixture


  Contains two or more elements not chemically combined, can be separated
• An atom is the smallest particle of a substance that can exist
• Molecule


  Two or more atoms bonded together
• Structure of an atom

  • Nucleus containing protons and neutrons
  • Shells of electrons surrounding the nucleus
• Proton


  Positive charge, mass of 1
• Neutron


  Neutral charge, mass of 1
• Electron


  Negative charge, mass of 1/2000
• Atoms are uncharged because they have equal numbers of protons and electrons
• Atomic number


  Number of protons in an atom
• Mass number


  Total number of protons and neutrons in an atom
• Group number


  Number of electrons in the outer shell of an atom
• Period number


  Number of electron shells in an atom
• Elements in the same group have similar chemical properties because they have the same number of electrons in the outer shell
• Noble gases


  Group 0 elements, highly unreactive because they have full outer shells
• Isotopes


  Atoms of the same element with the same number of protons but different number of neutrons
• Ion


  A charged particle formed by gaining or losing electrons
• Ionic bonding


  1. Metal atom loses electrons

  2. Non-metal atom gains electrons

  3. Resulting ions have opposite charges and are attracted to each other
• Covalent bonding


  Sharing of electrons between two non-metal atoms to achieve full outer shells
• Covalent bonding examples


  • Water (H2O)
  • Methane (CH4)
• Water (H2O) has a central oxygen atom with two hydrogen atoms coming to the side
• Hydrogen


  Has one electron in its outer shell
• Oxygen


  Has six electrons in its outer shell, which becomes full with eight electrons
• Methane (CH4) has a central carbon atom with four hydrogen atoms
• Carbon dioxide (CO2) has double covalent bonds
• Ethene (C2H4) has a central carbon-carbon double bond with four hydrogen atoms
• Giant ionic structures


  • Made up of a metal and a non-metal
  • Held together by strong electrostatic forces of attraction between oppositely charged ions
  • Have high melting and boiling points as a lot of energy is required to break the strong bonds
  • Don't conduct electricity when solid as the ions aren't free to move
  • Conduct electricity when molten/liquid as the ions are free to move
  • Are brittle and easily smash as the layers of ions can slide and positive charges repel
• Giant covalent structures (diamond and graphite)


  • Have many strong covalent bonds which require a lot of energy to break
  • Diamond has a tetrahedral structure with each carbon bonded to four others
  • Graphite has a layered structure with weaker intermolecular forces between the layers
• Covalent bond


  A shared pair of electrons
• Simple molecular substances


  • Have weak intermolecular forces that require little energy to break
  • Have increasing boiling points with increasing relative molecular mass (Mr) due to stronger intermolecular forces
• Giant metallic structures


  • Have strong metallic bonds between positive metal ions and a 'sea' of delocalized electrons
  • Have high melting points due to the strong bonds
  • Are good conductors of heat and electricity due to the delocalized electrons
  • Are malleable and ductile as the layers of ions can slide over each other
• Balancing chemical equations


  1. List the elements on each side

  2. Make a tally chart to count the number of each element

  3. Adjust the coefficients to balance the equation
• Relative atomic mass (Mr)


  The ratio of the average mass of an element compared to one atom of carbon-12
• Calculating moles


  1. Use the formula: mass = number of moles x Mr

  2. Rearrange to: number of moles = mass / Mr
• Empirical formula


  The simplest whole number ratio of atoms of each element in a compound
• Molecular formula


  The actual number of atoms of each element present in a compound
• Determining empirical formula from percentage composition

  1. List the masses and Mr of each element

  2. Calculate the number of moles of each element

  3. Divide the mole values by the smallest to get the simplest ratio
• Calculating moles in a hydrated salt

  1. Set up a table with the anhydrous salt, water, mass, Mr and moles

  2. Use the mass of the anhydrous salt remaining after heating to calculate the original amount of water