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27 Cards in this Set

  • Front
  • Back
Types of Chemical Analysis
• Two types of analysis are performed in an analytical laboratory; qualitative analysis and quantitative analysis.
• Qualitative analysis – finds what substances are present.
• Quantitative analysis – measures the quantity of a substance present.
Deciding on an Analytical Method
• The type of analysis used will be determined by the properties of the substance to be analysed.
• When looking at substances properties things like; melting and boiling temperature as well as solubility and colour are considered.
• For example, atomic absorption spectroscopy can be used for metal ions.
Writing Chemical Equations
• Chemical reactions are merely rearrangements of existing atoms, and chemical equations must therefore be balanced.
• To write a chemical equation, we must know the names of the formulae of the original substances (the reactants) and the new substances (the products).
Writing Chemical Equations
- continued -
• The chemical equation is then derived by using the following steps.
 Write out the word equation (names of the reactants on the left, products on the right and an arrow between them).
 Replace the name of each substance with its correct chemical formula.
 Insert coefficients to balance the equation and thus indicate the relative proportions of the substances involved.
 Add the correct symbol of state for each substance (aq, g, l, s).
Writing Chemical Equations
- continued -
Example:
sulfuric acid + sodium hydroxide sodium sulfate + water
H2SO4 + NaOH Na2SO4 + H2O
H2SO4 + 2NaOH Na2SO4 + 2H2O
H2SO4 (aq) + 2NaOH (aq) Na2SO4 (aq) + 2H2O (l)
Writing Chemical Equations
- continued -
When balancing an equation, the numbers in the formulae must not be altered. Instead, numbers (or coefficients) are inserted in front of the whole formula. These coefficients then tell us the relative proportions of moles (or molecules) of the substances that are involved in the reaction.
Ionic Equations
• To write an ionic equation, we usually begin with the full equation, we then use our knowledge of dissociation and ionisation to write down the ions that are present.
• Dissociation – the process by which ions separate when an ionic compound dissolves in a solvent. The end result is that mobile ions are produced from ions that were originally static.
Ionic Equations
- continued -
• Ionisation – a reaction in which uncharged atoms or molecules gain or lose electrons to form ions. The end result is that mobile ions are produced from a substance which did not originally contain ions.
• Any ions that remain unchanged (called spectator ions) may then be cancelled out, leaving behind the ionic equation.
Ionic Equations
- continued -
Example:
AgNO3 (aq) + KBr (aq) AgBr (s) +KNO3 (aq) - full equation
 As the three aqueous – phase substances will dissociate into their constituent ions upon dissolving, the above equation may be rewritten as:
Ag (aq) + NO3 – (aq) + K+ (aq) + Br- (aq) AgBr (s) + K+ (aq) + NO3- (aq)
 As the potassium and the nitrate ions remain unchanged (they are spectator ions), they may be cancelled out to leave the ionic equation:
Ag+ (aq) + Br- (aq) AgBr (s)
Reaction Patterns
• There are a range of patterns that have been recognised in reactions however remember that exceptions do exist.
Reaction Patterns
- continued -
• The types of patterns are:
1. Acid + base salt + water
2. Acid + reactive metal salt + hydrogen
3. Acid + metal carbonate salt + carbon dioxide + water
4. Metal carbonate metal oxide + carbon dioxide
5. Hydrocarbon + oxygen carbon dioxide + water
Note: when dealing with hydrocarbons this equation is very useful for balancing reactions of this type.
CxHy + (x+y/4) O2 xCO2 + Y/2 H2O
(if this formula leads to any halve, the coefficients can be double)
Redox Reactions
• Redox reactions are also referred to as oxidation and reduction reactions.
• Redox reactions involve the transfer of electrons from one substance to another.
• Reduction – gain of electrons (addition of hydrogen removal of oxygen)
• Oxidation – loss of electrons (addition of oxygen removal of hydrogen)
Note – ‘OIL RIG’ (Oxidation is Loss; Reduction is Gain)
Redox Reactions
- continued -
• For reduction, the electrons will be on the left side of the half-equation.
• For oxidation, the electrons will be on the right side of the half-equation.
• The processes of oxidation and reduction are complementary; one cannot occur without the other.
• An oxidant (oxidising agent) causes another substance to be oxidised and is itself reduced.
• A reductant (reducing agent) causes another substance to be reduced and is itself oxidised.
Writing Half Equations
• To balance redox reactions, the reduction and oxidation are dealt with separately in a six step process.
1) Write down the substance and what it turns into (e.g its conjugate)
2) Balance all atoms, except for oxygen and hydrogen.
3) Balance oxygen atoms by adding water molecules to the other side.
4) Balance hydrogen atoms by adding hydrogen ions.
5) Balance the charge by adding electrons.
6) Finally insert the states in the equation (aq, s, g, l)
Writing Half Equations
- continued -
Note – Ensure both half equations have the same amount of electrons on either side, if they don’t, multiply through the half equations so that they do have the same number of electrons.
Writing Overall Equations
Writing overall equations –
• Ensure that the number of electrons in each equation is equal.
• An overall equation can be produced by adding the two half equations.
• Cancel out the electrons.
Rules of Oxidation Numbers
• Oxidation numbers have been created to help identify whether a reaction is redox.
• If an oxidation number increases, then oxidation has occurred.
• If an oxidation number decreases, then reduction has occurred.
• To calculate an oxidation number, the following rules are used;
 Elements by themselves have an oxidation number of zero, e.g. O2
 Ions have oxidation numbers that match their charges e.g. Ca2+ = 2+
Rules of Oxidation Numbers
- continued -
 Hydrogen is usually assigned an oxidation number of +1
 except if it is part of a metal hydride such as NaH where its oxidation number is -1
 Oxygen is usually assigned the number of -2
 Except if it is a peroxide compound like H2O2 or Na2O2, then the oxidation number for oxygen is -1
 In the compound OF2, oxygen has an oxidation number of +2 because fluorine is so electronegative.
 All compounds have overall oxidation numbers of zero.
Precipitation Reactions
• In precipitation reactions, soluble substances in a solution are mixed together and an insoluble product is formed.
• The product formed is known as a precipitate.
Acid-Base Reactions
• Lowry-Bronsted theory states:
 Acids are substances that donate hydrogen ions (protons) to another substance.
 Bases are proton acceptors.
• For the reaction to occur there must be a proton donor (the acid) and a proton receptor (the base).
Acid - Base Reactions
- continued -
• Conjugate pairs – a conjugate pair is the original acid/base and the consequent base/acid that it turns into.
• A conjugate acid-base pair consists of the acid and the same species with one less proton, from the other side of the equation.
Acid - Base Reactions
- continued -
Example –
HCl(aq) + OH-(aq) Cl- (aq) + H2O(l)
 Here HCl has donated a proton to the OH- so HCl is the acid and OH- is the base.
 The conjugate acid-base pairs are HCl(aq) /Cl- (aq) and H2O(l)/OH- (aq)
 The acid is always written first.
Acid - Base Reactions
- continued -
• Amphiprotic substances – these substances can be both proton donors and receptors, they act as an acid in some reactions and a base in other reaction. An example is water.
• Polyprotic acids – these are acids that can donate more than one proton (per molecule)
• The strength of an acid or base is a measure of how readily it donates or accepts H+ ions.
• The pH scale is used to measure how acidic or how basic a solution is;
 Values less than 7 are acidic
 Values higher than 7 are basic
 A value of 7 is neutral
The Mole Concept
• The mole is defined as ‘the amount of substance that contains the same number of particles as there are atoms in exactly 12 g of pure 12C isotope’.
• Avogadro’s’ number (NA) = 6.023 x 10 ^23
• The molar mass of a substance is obtained by simply adding together the relative atomic masses of each atom that appears in the formula. The unit is ‘g mol -1’
Limiting and Excess Reagents
• In some cases, more than one reactant may be known, before the amount of product can be predicted, it is necessary to work out which reactant is in excess and which reactant is used up.
• Once the number of moles of each reactant has been calculated, the mole ratio (determined from the reaction equation) is used to calculate the number of moles of all reactants used in the reaction.
Limiting and Excess Reagents
- continued -
• The excess reactant no longer takes part in the reaction.
• The amount of product may then be forecast, based upon the substance that is completely consumed (this is called the limiting reagent) the moles of this substance must be used in all future calculations.
• Excess reactants are often used to ensure complete reaction of the other reactants.
Dilution of Solutions
• The equation c1V1 = c2V2 is used to calculate the increased volume or decreased concentration of the diluted solution.