Module 10: Understanding chemical reactions

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Formulae of elements and simple compounds

Word equations and simple balanced equations

 ionic equations

10.01 recall that all atoms of the same element have the same number of protons
Atoms contain protons, neutrons and electrons. 


The number of protons in an atom tells us the name of the element.  For example chlorine has two types of atom. One has 18 neutrons and the other has 20 neutrons but both have the same number, 17, of protons. 
Task 10.01

10.02 Relative charges and relative masses of protons, neutrons and electrons

Task 10.02

10.03 understand the terms atomic number and mass number
Atomic number = number of protons in the nucleus of an atom.  (found in periodic table)
Mass number is the number of protons and neutrons in the nucleus of an atom added together.

Mass number = atomic number + number of neutrons
Number of neutrons = Mass number - atomic number

Atoms are shown using symbols, mass numbers and atomic numbers as below:

   mass number                                      37
                      symbol      e.g. chlorine       Cl
atomic number                                      17

Task 10.03

10.04 understand that isotopes are atoms of the same element with the same number of protons and electrons, but different numbers of neutrons
Isotopes are atoms of the same element which have different numbers of neutrons in their nuclei. As these are the same element the atoms all have the same number of protons.  For example hydrogen has 3 isotopes.  Each atom has 1 proton but a different number of neutrons.

Task 10.04 Draw symbols for the isotopes of hydrogen, isotopes carbon-12, carbon-13, carbon-14, chlorine-35 and chlorine-37 given the atomic number of chlorine is 17 and the atomic number of carbon is 6.

10.05 calculate the relative atomic mass of an element from relative masses and abundances of its isotopes
Use relative mass of isotopes and their relative abundance.  E.g. Chlorine has two 
isotopes with mass numbers 35 and 37.
                                                                           _{35                       37}
                                                 75% is    Cl and 25% is    Cl
                                                              ^{17                            17}
Let there be 100 atoms
Total mass of 100 atoms = (75 * 35) + (25 * 37) = 3550
Average mass of an atom (relative atomic mass of chlorine) = Total mass /Number of atoms
=3550/100
so relative atomic mass of chlorine = 35.5
Task 10.05 Calculate the relative atomic mass of the following:
bromine; 50% Br 79 and 50% Br 81
gallium 60% Ga 69 and 40% Ga 71
neon 90% Ne 20 and 10% Ne 22
silver 50% Ag 107 and 50% Ag 109
boron 20% B 10 and 80% B 11

Chemical bonds
10.06 recall that some elements combine by means of chemical reactions to form compounds, for example, water, carbon dioxide, sodium chloride and iron sulfide

Task 10.06 Write word equations and balanced chemical equations for making the above compounds.

10.07 recall that an ion is an atom or group of atoms with a positive or negative charge
Atoms have no charge.  A charged particle is called an ion.  If an atom loses an electron, it becomes a positively charged (+) ion. An ion that is positively charged is known as a cation. If an atom gains an electron, it becomes a negatively charged (-) ion. An ion that is negatively is known as an anion.  The negative and positive ions attract each other to form an ionic bond.

Typical ions include: lithium Li⁺, sodium Na⁺, copper Cu²⁺, calcium Ca²⁺, iron Fe²⁺, iron Fe³⁺, chloride Cl^-, bromide Br^-, hydrogen H⁺, oxide O^2-, sulfide S^2-, hydroxide OH^-, ammonium NH₄⁺, carbonate CO₃^2-, sulfate SO₄^2-.

Task 10.07 Complete the gaps in the text below:
_____ have no charge.  A charged particle is called an ___.  If an atom loses an ________, it becomes a positively charged (+) ion. An ion that is positively charged is known as a ______. If an atom gains an electron, it becomes a negatively charged (-) ion. An ___ that is negatively is known as an anion.  The negative and positive ions attract each other to form an _____ bond.

10.08 recall that ionic bonds are formed between atoms of a metal and a non-metal, for example, sodium and chlorine forming sodium chloride
Task 10.08 write out the compounds which have ionic bonds by looking for metals and non-metals in the following: sodium chloride, hydrogen chloride, copper bromide, iron oxide, carbon monoxide, magnesium oxide, hydrogen oxide, NaCl, H₂O, FeS, CaBr₂.

10.09 recall that chemical bonding involves the transfer or sharing of electrons
There are two types of chemical bonding and both use electrons.  Ionic bonds form when electrons are transferred.  Covalent bonds form when electrons are shared.


10.10 explain the formation of simple ionic compounds (for example, sodium chloride) in terms of transfer of electrons

animated ion formation (needs powerpoint)
Task 10.11a  Draw atoms and ions for lithium, potassium, fluorine, magnesium, oxygen, sulfur and aluminium.
Task 10.11b Draw diagrams of ionic bonding in LiF, KF, LiCl, NaF, MgCl₂, AlF₃, MgO, MgS, Na₂O and Al₂O₃.

10.11 describe the structure of ionic compounds as a lattice structure, consisting of a regular arrangement of ions, held together by strong forces between them, forming crystals
Strong forces exist between positive and negative ions.  These are the ionic bonds.  There is a very ordered pattern for the ions.  This gives the solid crystals nice shapes with definite angles between sides.

Task 10.11 Sodium chloride crystals are perfect cubes.  Draw a large cube and put in 9 ions on each face to show how the ions are arranged to give this shape.

10.12 describe and explain the physical properties of giant ionic structures, including sodium chloride and magnesium oxide
Each ion is firmly held in place by strong ionic bonds so they have high melting and boiling points. If melted, charged ions become free to carry an electric current. The ions also become free if dissolved in water so solutions are also electrolytes. The solids are insulators because the ions are not free to move and cannot carry a current.  Sodium chloride NaCl, and magnesium oxide MgO are good examples.
Task 10.12 Explain why NaCl has a high melting point.  Explain why molten MgO conducts electricity.  Explain why aqueous NaCl solution is an electical conductor but solid NaCl is an insulator.

10.13 recall that covalent bonds are formed between atoms of some non metals to produce molecules (including hydrogen, nitrogen, oxygen, chlorine and hydrogen chloride)
Task 10.13 write out the compounds which have covalent bonds by looking for metals and non-metals in the following: sodium chloride, nitrogen, hydrogen chloride, copper bromide, iron oxide, oxygen, carbon monoxide, magnesium oxide, hydrogen oxide, H₂, NaCl, H₂O, FeS, CaBr₂,Cl₂.

10.14 explain the formation of simple covalent molecules (e.g. hydrogen, hydrogen chloride, water, methane, carbon dioxide) in terms of shared electrons between non-metal atoms, using dot and cross diagrams
Non-metal atoms join using covalent bonds. When a covalent bond is formed, atoms share their electrons. The atoms then have full shells. One covalent bond needs one shared electron from each atom. Each atom involved has to make enough covalent bonds to fill up its outer shell.  Sharing electrons is called covalent bonding.  Below is a diagram to show hydrogen gas (H₂).



Task 10.14: Draw atoms of F, H, O, N and C.
Draw a dot and cross diagram for fluorine F₂,  hydrogen fluoride HF, water H₂O, ammonia NH₃, methane CH₄, oxygen O₂, nitrogen N₂,CO₂ and ethene C₂H₄.  

10.15 describe the physical properties of simple molecular compounds
Some covalent compounds are simple molecules.  They have simple molecular structures. Compounds like this have low melting and boiling points and most are gases or liquids at room temperature. This is because of weak forces between the molecules. Molecular substances do not conduct electricity, because there are no ions.  E.g. water.

Task 10.15 Draw diagrams to show how the molecular structures for the following might look: fluorine F₂,  hydrogen fluoride HF, water H₂O, ammonia NH₃, methane CH₄, oxygen O₂, nitrogen N₂

10.16 understand that covalent bond formation can result in simple molecules (eg hydrogen, iodine) and giant structures (eg diamond and graphite)

Some covalent compounds have giant structures.  In giant covalent structures all the atoms are bonded to each other by strong covalent bonds so they have very high melting and boiling points. They do not usually conduct electricity even if in the liquid state. Diamond and graphite are two examples, which are made from carbon atoms.  These two different types of the same element are called allotropes.

Diamond: Each carbon atom forms four covalent bonds in a very rigid giant covalent structure.

Graphite: exists as layers of carbon atoms each held in place by three strong covalent bonds.  Each layer is held to the one above it by weak bonds.

Task 10.16 Silicon dioxide SiO2 has a giant atomic structure.  In it each silicon atom has 4 covalent bonds to oxygen atoms and each oxygen atom has 4 covalent bonds to silicon atoms.  Draw a possible structure for SiO2.  Describe the melting point of SiO2 and give a reason for your answer.  Explain why SiO2 is an electrical insulator.

10.17 describe and explain the differences between the physical properties of simple molecular substances and those with giant molecular structures

Giant molecular structures have very high melting points because all atoms are held firmly in place by strong covalent bonds. In graphite each carbon atom is held in place by three strong covalent bonds which gives graphite a high melting point. In diamond 4 strong covalent bonds holds each atom in place. This makes diamond very hard. Graphite has weak bonds between layers so the layers slip over each other making graphite soft.

They do not usually conduct electricity even when molten because there are no charged particles to carry the current. There are free electrons between layers in graphite so it conducts electricity.
Simple molecular substances like water have weak bonds between molecules so melt at low temperatures because little energy is needed to separate the molecules.  Giant covalent structures like diamond have strong covalent bonds holding each atom in place.  They melt at high temperatures because a lot of energy is needed to break these strong bonds.

Task 10.17 Silicon carbide SiC, has a giant atomic structure but carbon dioxide CO2, has a molecular structure.  Draw diagrams to show how their structures might look.  Explain why SiC has a high melting point whereas CO2 has a low one.  Explain why both od these substances are electrical insulators.

Energy transfers
10.18 recall that changes of temperature often accompany reactions
Many reactions give off heat such as the burning of wood which causes a temperature increase. Other reactions take in energy and cause a temperature fall.

10.19 recall that an exothermic reaction is one in which thermal energy is given out
An Exothermic reaction is one which gives out energy to the surroundings, usually in the form of heat and usually shown by a rise in temperature.  An example of an exothermic process is the burning of fuels.

10.20 recall that an endothermic reaction in one in which thermal energy is taken in
An Endothermic reaction is one which takes in energy from the surroundings, usually in the form of heat and usually shown by a fall in temperature. An example of an endothermic is photosynthesis. This is because it takes in energy from the sun in the form of light.

10.21 understand that the breaking of bonds is endothermic and that the making of bonds is exothermic
During a chemical reaction, old bonds are broken and new ones formed. Energy must be supplied to break existing bonds and so this is endothermic. Energy is released when new bonds are formed and so the formation of bonds is exothermic.
Hydrogen reacts with oxygen to form water.
O=O + H-H + H-H --->O +O + H +H + H +H --->H-O-H + H-O-H
Old bonds break            single atoms with            new bonds form
in oxygen and                 no bonds                       in water molecules
hydrogen molecules                                              exothermic
endothermic

Using chemical equations
10.22 calculate the relative formula masses of simple compounds, given relative atomic masses
Add up the relative atomic mass (found in periodic table) of each atom in the compound.
e.g.  Al₂0₃    relative atomic masses of Al = 27, O = 16 (found in periodic table).  The formula shows 2 atoms of aluminium and 3 atoms of oxygen so:
formula mass of       = (2*27) + (3*16) =54 + 48 = 102
Task 10.22 Work out the relative formula masses of the following: MgO, FeS, O₂, H₂O, CaBr₂, Na₂S, CaCO₃, NaOH, HCl, (NH₄)₂SO₄.  Relative atomic masses Mg=24, O=16, Fe=56, S=32, Ca=40, Br=80, C=12, Na=23, H=1, Cl=35.5.

10.23 use chemical equations quantitatively to determine the masses of substances used and produced
You can work out ratio of the masses of products and reactants by simply multiplying the number of moles shown in the equation by the formula mass of each substance.
mgburnpic.JPG      mgburnvid.3gp

Example: What mass of magnesium oxide can be made from 12g of magnesium?
equation   2Mg_(s) + O_2(g) -->  2MgO_(s)
amounts   2 moles   1 mole     2 moles
masses     2*24      1{16*2}      2{24+16}
                =48g         =32g        =80g so
                48g Mg  forms          80g MgO
                1g Mg    forms          80/48 g MgO
               12g Mg  forms          12*80/48 g MgO = 20g
Also note that the ratio of amounts of reactants and products in the equation above can be written as:

Amount of Mg/amount of O₂ =2/1
Or
Amount of O₂/amount of MgO = 1/2
Task 10.23

10.24 determine the empirical formulae of simple compounds from reacting masses
elements reacting                          magnesium           chlorine
symbols of elements                            Mg                     Cl
masses reacting (from experiment)       2.4g                   7.1g
molar mass (look up relative atomic     24g/mol              35.5g/mol
mass in periodic table)
amounts (amount = mass/molar mass)  2.4g/24g/mol      7.1g/35.5g/mol  
                                                       =  0.1mol               0.2mol
ratio of atoms (divide by smallest)         1            :           2
formula                                                          MgCl₂

Task 10.24
Work out formulae of compounds formed when the following react:
56g of iron and 32g of sulphur (Fe =56, S =32)
2g of hydrogen and 16g of oxygen (H=1, O=16)
14g of lithium and 16g of oxygen (Li=7)
32g of copper and 8g of oxygen (Cu=64)
6.4g of copper and 0.8g of oxygen.