IGCSE CHEMISTRY 0620 Cambridge LESSON 7 EXPLANATION

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 Lesson 7

4 Stoichiometry
4.1 Stoichiometry

• Use the symbols of the elements and write the formulae of simple compounds

 

Element symbols

• Each element is represented by its own unique symbol as seen on the Periodic Table e.g. H is hydrogen.
• Where a symbol contains two letters, the first one is always in capital letters and the other is small e.g. sodium is Na, not NA.
• Atoms combine together in fixed ratios that will give them full outer shells of electrons.
• The chemical formula is what tells you the ratio of atoms.
• E.g. H₂O is a compound containing 2 hydrogen atoms which combine with 1 oxygen atom.
• The chemical formula can be deduced from the relative number of atoms present.
• E.g. if a molecule contains 3 atoms of hydrogen and 1 atom of nitrogen then the formula would be NH₃.
• Diagrams or models can also be used to represent the chemical formula.

• Deduce the formula of a simple compound from the relative numbers of atoms present

 

Chemical formulae

• The structural formula tells you the way in which the atoms in a particular molecule are bonded. This can be done by either a diagram (displayed formula) or written (simplified structural formula).
• The empirical formula tells you the simplest whole number ratio of atoms in a compound.
• The molecular formula tells you the actual number of atoms of each element in one molecule of the compound or element e.g. H₂ has 2 hydrogen atoms, HCl has 1 hydrogen atom and 1 chlorine atom.

Example: Butane

• Structural formula (displayed)

 

• Structural formula (simplified)

          CH₃CH₂CH₂CH₃

• Molecular formula

          C₄H₁₀

• Empirical formula

          C₂H₅

Deducing formulae by combining power

• The concept of valency is used to deduce the formulae of compounds.
• Valency or combing power tells you how many bonds an atom can make with another atom.
• E.g. carbon is in Group IV so a single carbon atom can make 4 single bonds or 2 double bonds.
• The following valencies apply to elements in each group:

• We can use the combining power of each atom to work out a formula.
• Example: what is the formula of aluminium sulfide?

          Write out the symbols of each element and write their combining powers underneath:
Al   S
3    2

• The formula is then calculated by cross multiplying each atom with the number opposite, hence the formula for aluminium sulfide is Al₃S₂.

• Deduce the formula of a simple compound from a model or a diagrammatic representation

 
 
 

• Construct word equations and simple balanced chemical equations

 
 
 

• Define relative atomic mass, Ar, as the average mass of naturally occurring atoms of an element on a scale where the 12C atom has a mass of exactly 12 units

 

Relative Atomic mass

• The symbol for the relative atomic mass is A_r.
• This is calculated from the mass number and relative abundances of all the isotopes of a particular element.

 

Equation:

• The top line of the equation can be extended to include the number of different isotopes of a particular element present.

Example for Isotopes:

The Table shows information about the Isotopes in a sample of Rubidium

Example

 Use information from the table to calculate the relative atomic mass of this sample of Rubidium.
Give your answer to one decimal place:

Define relative molecular mass, Mr, as the sum of the relative atomic masses. (Relative formula mass or Mr will be used for ionic compounds.)

Relative formula (molecular) mass

• The symbol for the relative molecular mass is Mr and it refers to the total mass of the molecule.
• To calculate the Mr of a substance, you have to add up the Relative Atomic Masses of all the atoms present in the formula

 Example:

(Calculations involving reacting masses in simple proportions may be set. Calculations will not involve the mole concept.)

• Determine the formula of an ionic compound from the charges on the ions present

• The formulae of these compounds can be calculated if you know the charge on the ions.
• Below are some common ions and their charges:

• For ionic compounds you have to balance the charge of each part by multiplying each ion until the sum of the charges = 0.
• Example: what is the formula of aluminium sulfate?
  • Write out the formulae of each ion, including their charges.
  • Al³⁺       SO₄^2-

• Balance the charges by multiplying them out: Al³⁺ x 2 = +6 and SO₄^2- x 3 = -6; so +6 – 6 = 0.
• So the formula is Al₂(SO₄)₃

• Construct equations with state symbols, including ionic equations

Word equations

• These show the reactants and products of a chemical reaction using their full chemical names.
• The arrow (which is spoken as  “goes to” or “produces”) implies the conversion of reactants into products.
• Reaction conditions or the name of a catalyst can be written above the arrow.

Names of compounds
For compounds consisting of 2 atoms:

• If one is a metal and the other a nonmetal, then the name of the metal atom comes first and the ending of the second atom is replaced by adding –ide
  • E.g. NaCl which contains sodium and chlorine thus becomes sodium chloride.
• If both atoms are nonmetals and one of those is hydrogen, then hydrogen comes first.
  • E.g. hydrogen and chlorine combined is called hydrogen chloride.
• For other combinations of nonmetals as a general rule, the element that has a lower Group number comes first in the name.
  • E.g. carbon and oxygen combine to form CO₂ which is carbon dioxide since carbon is in Group 4 and oxygen in Group 6.

For compounds that contain certain groups of atoms:

• There are common groups of atoms which occur regularly in chemistry.
• Examples include the carbonate ion(CO₃^2-), sulfate ion (SO₄^2-), hydroxide ion (OH^–) and the nitrate ion (NO₃^–).
• When these ions form a compound with a metal atom, the name of the metal comes first.
• E.g. KOH is potassium hydroxide, CaCO₃ is calcium carbonate.

Writing and balancing chemical equations

• These use the chemical symbols of each reactant and product.
• When balancing equations, there needs to be the same number of atoms of each element on either side of the equation.
• The following nonmetals must be written as molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂.
• Work across the equation from left to right, checking one element after another.
• If there is a group of atoms, for example a nitrate group (NO₃^–) that has not changed from one side to the other, then count the whole group as one entity rather than counting the individual atoms. For example:
  • NaOH + HCl → NaCl + H₂O
  • There are equal numbers of each atom on either side of the reaction arrow so the equation is balanced.

• Deduce the balanced equation for a chemical reaction, given relevant information

Using state symbols:
State symbols are written after formulae in chemical equations to show which physical state each substance is in:

Example 1: 
Aluminium (s)  +   Copper (II) Oxide (s)  →   Aluminium Oxide (s)  +   Copper (s)
Unbalanced symbol equation:    Al     +     CuO     →     Al₂O₃     +     Cu
ALUMINIUM: There is 1 aluminium atom on the left and 2 on the right so if you end up with 2, you must start with 2. To achieve this, it must be 2Al
2Al     +     CuO     →     Al₂O₃     +     Cu
OXYGEN: There is 1 oxygen atom on the left and 3 on the right so if you end up with 3, you must start with 3. To achieve this, it must be 3CuO.
2Al     +     3CuO     →     Al₂O₃     +     Cu
COPPER: There is 3 copper atoms on the left and 1 on the right. The only way of achieving 3 on the right is to have 3Cu
2Al     +     3CuO     →     Al₂O₃     +     3Cu
Example 2:
Magnesium Oxide (s)  +  Nitric Acid (aq)  →   Magnesium Nitrate (aq)  +   Water (l)
Unbalanced symbol equation:  MgO   +   HNO₃   →   Mg(NO₃)₂   +   H₂O
MAGNESIUM: There is 1 magnesium atom on the left and 1 on the right so there are equal numbers of magnesium atoms on both sides so these are kept the same
MgO   +   HNO₃   →   Mg(NO₃)₂   +   H₂O
OXYGEN: There is 1 oxygen atom on the left and 1 on the right so there is an equal number of oxygen atoms on both sides. It is therefore kept the same (remember that you are counting the nitrate group as separate group, so do not count the oxygen atoms in this group)
MgO   +   HNO₃   →   Mg(NO₃)₂   +   H₂O
HYDROGEN: There is 1 hydrogen atom on the left and 2 on the right. Therefore you must change HNO₃ to 2HNO₃
MgO   +   2HNO₃   →   Mg(NO₃)₂   +   H₂O

Balancing ionic equations

• In aqueous solutions ionic compounds dissociate into their ions, meaning they separate into the component atoms or ions that formed them.
• E.g. hydrochloric acid and potassium hydroxide dissociate as follows:
  • HCl →H⁺ + Cl^–
  • KOH → K⁺ + OH^–
• It is important that you can recognise common ionic compounds and their constituent ions.
• These include:
  • Acids such as HCl and H₂SO₄.
  • Group I and Group II hydroxides e.g. sodium hydroxide.
  • Soluble salts e.g. potassium sulfate, sodium chloride.
• Follow the example below to write ionic equations.

Example:        Write the ionic equation for the reaction of aqueous chlorine and aqueous potassium iodide.
Step 1:   Write out the full balanced the equation:      
2KI(aq)     +     Cl₂(aq)     →     2KCl(aq)    +     I₂(aq)
Step 2:   Identify the ionic substances and write down the ions separately:
2K⁺ (aq) + 2I^–(aq)     +     Cl₂(aq)     →     2K⁺(aq) + 2Cl^–(aq)    +     I₂(aq)
Step 3:   Rewrite the equation eliminating the ions which appear on both sides of the equation (spectator ions) which in this case are the K⁺ ions:
2I^–(aq)     +     Cl₂(aq)     →     2Cl^–(aq)    +     I₂(aq)

• When balancing equations you cannot change any of the formulae, only the amounts of each atom or molecule.
• This is done by changing the numbers that go in front of each chemical species.

• You need to be able to identify the products which are not ions in ionic equations.
• These are usually molecules such as water or bromine but they may also be precipitated solids.