TOPIC 3

• Law of Conservation of Mass


  No matter is lost or gained during a chemical reaction. Mass is always conserved, therefore the total mass of the reactants is equal to the total mass of the products, which is why all chemical equations must be balanced
• If the reaction flask is closed and no other substance can enter or leave the system, then the total mass of the reaction flask will remain constant
• Precipitation reaction

  Two solutions react to form an insoluble solid called a precipitate
• If the reaction flask is open and a gaseous product is allowed to escape, then the total mass of the reaction flask will change as product mass is lost when the gas leaves the system
• Mass will be lost from the reaction flask as carbon dioxide gas escapes to the atmosphere
• If the mass of a reaction flask is found to increase then it may be due to one of the reactants being a gas found in the air and all of the products are either solids or liquids
• Chemical formula


  The numbers involved give information about the chemicals involved

  
• Information given by chemical formulae

  • If no subscript number after an element, there must be one of that element
  • If subscript number after an element, that number belongs to the element just before it
  • If subscript number after brackets, that number belongs to all elements inside the bracket
  • Most complicated examples contain a subscript number inside the bracket as well as outside
• Chemical equations use the chemical symbols of each reactant and product
• When balancing equations, there has to be the same number of atoms of each element on either side of the equation in accordance with the Law of Conservation of Mass
• Non-metals that must be written as molecules

  • H2
  • N2
  • O2
  • F2
  • Cl2
  • Br2
  • I2

  
• Balancing an equation

  1. Work across the equation from left to right, checking one element after another

  2. If a group of atoms has not changed from one side to the other, count the whole group as one entity rather than counting the individual atoms
• Chemical equations


  • Acid-base neutralisation reaction: NaOH (aq) + HCl (aq) ⟶ NaCl (aq) + H2O (l)
  • Redox reaction: 2Fe2O3 (aq) + 3C (s) ⟶ 4Fe (s) + 3CO2 (g)
• A large number before any chemical applies to that entire chemical
• Balancing equations

  1. Best approach is to practice lots of examples

  2. Change the coefficients (multipliers) in front of the formulae, one by one checking the result on the other side

  3. Balance elements that appear on their own, last in the process
• Relative atomic mass (Ar)

  Calculated from the mass number and relative abundances of all the isotopes of a particular element

  
• Relative formula mass (Mr)

  The total mass of the molecule, calculated by adding up the relative atomic masses of all the atoms present in the formula

  
• The sum of the relative molecular masses of the reactants will be the same as the sum of the relative molecular masses of the products
• Calculating percentage by mass of an element in a compound


  Use the equation: Percentage by mass = (Mass of element / Total mass of compound) x 100

  
• Particles in the gas state still have mass! Remember mass is always conserved in chemical reactions
• If the reaction flask is completely closed, then no gases should be able to escape
• Error

  The difference between a value or quantity obtained in an experiment and an accepted or literature value
• Uncertainty

  The same as random errors, expressing the confidence to which the measurement can be taken
• Random errors

  • Can be affected by: how easily the instrument or scale is to read, the person reading the scale poorly, changes in the environment like temperature fluctuations or air currents
  • Pull a result away from an accepted value in either direction (too high or too low)
• Systematic errors


  • Occur as a result of a faulty or poorly designed experimental procedure
  • Always pull the result away from the accepted value in the same direction (always too high or always too low)
• Calculating uncertainty

  1. For analogue instruments, the uncertainty is half the smallest division on the scale

  2. For digital instruments, the uncertainty is the smallest division on the scale

  3. For results from repeated experiments, the uncertainty is half the range of results
• Other sources of uncertainty can arise where the judgement of the experimenter is needed to determine a changing property
• Chemical amounts

  (Higher Tier Only)
• Uncertainty

  The smallest division on the digital scale
• Uncertainty example

  • An electronic balance that reads to 0.01 g, the uncertainty would be +0.01 g
• Uncertainty in results

  For results that are obtained from a series of repeated experiments, the uncertainty is ± half of the range of results
• Calculating uncertainty in results

  1. Calculating the mean average and then determining the deviation of the highest and lowest results from the mean value

  2. Calculating the range of the results and then dividing this value by 2
• Other uncertainties

  Arise where the judgement of the experimenter is needed to determine a changing property
• Examples of other uncertainties

  • Judging the end point of a titration by looking at the colour of the indicator
  • Controlling a stopwatch in a rate of reaction experiment
  • Deciding when to extinguish the flame in an experiment
• Mole

  Chemical amounts are measured in moles
• One mole of a substance contains the same number of the stated particles, atoms, molecules, or ions as one mole of any other substance
• Avogadro constant

  The number of atoms, molecules or ions in a mole (1 mol) of a given substance

  
• Examples of one mole

  • One mole of sodium (Na) contains 6.02 x 10^23 atoms of sodium
  • One mole of hydrogen (H2) contains 6.02 x 10^23 molecules of hydrogen
  • One mole of sodium chloride (NaCl) contains 6.02 x 10^23 formula units of sodium chloride
• The measurement of amounts in moles can apply to atoms, molecules, ions, electrons, formulae and equations

  
• Molar mass


  One mole of any element is equal to the relative atomic mass of that element in grams