Definitions paper 1

• Isotope

  atoms with the same number of protons but a different number of neutrons
• S block element

  Atom whoses outer electron is filling an s shell
• First ionisation energy

  The enthalpy change when one mole of gaseous atoms each have an electron removed, forming one mole of gaseous ions with a single positive charge
• Second ionisation energy

  The enthalpy change when one mole of gaseous ions with a single positive charge each have an electron removed, forming one mole of gaseous ions with a double positive charge
• Mole

  The amount of substance in grams that has the same number of particles as there are atoms in 12 grams of carbon 12
• Relative atomic mass

  The average mass of one atom in ratio to one twelfth of the mass one atom of carbon 12
• Relative molecular mass

  The average mass of a molecule in ratio to one twelfth of the mass of one carbon 12
• Avogadros constant

  6.02x10^23
• Empirical formula

  The simplest ratio of atoms of each element in the compound
• Molecular formula

  The actual number of atoms of each element in the compound
• Ionic bond

  The strong electrostatic force of attraction between oppositely charged ions formed by electrons tranfer
• Covalent bond

  Shared pair of electrons
• Metallic bond

  The electrostatic force of attraction between positive metal ions and delocalised electrons
• Electronegativity

  the power of an atom to attract electrons in a covalent bond
• Polar covalent bond

  Form when the elements in the bond have different electronegativity
• Symmetric molecule

  All bonds are identical with no lone pairs. There is no net movement
• Van der waal forces ( temporary dipoles )

  The weakest type of intermolecular force and acts as an induced dipole between molecules. They are formed by a random asymmetry of electro clouds which forms a temporary dipole. This induces a dipole on a neighbouring molecule and this attracts an induced dipole.
• Permenant dipole dipole bonding

  Stronger than van der waal forces. Occurs between polar molecules.
• Hydrogen bonding

  Strongest type of intermolecular force. It occurs in compounds containing a hydrogen atom and either nitrogen, oxygen and fluorine which must have an avaidable lone pair. They have a large electronegativity difference.
• Enthalpy change

  The amount of heat energy taken in or given out during any chsnge in a system provided the pressure is constant
• Exothermic change

  Energy is transferred from the system to the surrounding. The products have less energy than the reactants
• Endothermic change

  Energy is transferred from the surroundings to the system. They require an input of heat energy. The products have more energy than the reactants
• Standard enthalpy change of formation

  Enthalpy change when one mole of the compound is formed from its elements under standard condition. All reactants and products are in there standard state
• Standard enthalpy change of combsustion

  Enthalpy change that occurs when one mole of substance is combusted completely in oxygen under standard conditions. All reactants and products are in their standard states
• Hess law

  The Total enthalpy change for a reaction is independent of the route by which the chemical change takes place
• Mean bond energy

  The enthalpy needed to break the covalent bond into gaseous atoms averaged over different molecules
• Dynamic equilibrium

  The forward and bacjward reactions are occuring at equal rates and the concentrations of reactants and products stay constant
• Le chatelier principle

  If external conditions are changed the equilibrium will shift to oppose the change. If temperature is increased the equilibrium will shift to oppose this and move in the endothermic direction to reduce the temperature by absorbing heat. Increasing pressure will cause the equillibrium to shift towards the side with fewer moles of gas to oppose the change and reduce the pressure. A catalyst has no effect om the position of equilibriu, but it will spee up the rate at which the equilibrium is achieved.
• Oxidation

  The process of electron loss, or an increase in oxidation number
• Reduction

  The process of electron gain or decrease in oxidation number
• Reducing agent

  Electron donors
• Oxidising agents

  Electron acceptors
• alkali

  water soluable base