Atoms and elements 

Elements are substances which cannot be broken down any further. In fact, elements are made entirely out of one type of atom. The word ‘atom’ comes from a Greek word for indivisible.

Two or more atoms joined together are called a molecule. You can have molecules containing two of the same atom: oxygen gas is made of two oxygen atoms joined together. We write this O2. The number at the bottom right is the number of atoms of that type. H2O has two atoms of H (hydrogen) and one of O (oxygen).

Molecular substances are molecules which don’t form ions e.g. CO2. Ionic substances form a lattice of positive and negative ions. Some of these come apart to form separate ions when they dissolve in water e.g. NaCl. Some are insoluble e.g. CaCO3.

The first twenty elements of the periodic table are shown below. You are expected to know the names give the symbol, and be able to give the electron configuration (video tutorial)

Electron orbit review sheet

Some prior knowledge is expected in this unit:

Structure of the atom: This is covered in the Year 9 unit on Atoms

Ions and ionic formula: This is covered in the Year 10 unit on Periodic Trends and bonding

You are expected to be familiar with this material and it will not be revisited here.

Ions

The table of ions you will be provided with in assessments is shown below:

This table will not have names on it. You are expected to memorise these; they are shown below:

Note that lithium, iron III and fluoride ions do not occur in any of the reactions covered in this course.

The hydrogen carbonate ion has a historic name, the bicarbonate ion, that is still widely used. Some of the other ions in this table also have historic names but these are used less often.

You are expected to be able to apply the no-name version of this table to CORRECTLY write the formula for any compound of these ions e.g. ammonium sulfate (NH4)2SO4   or barium nitrate     Ba(NO3)2 . If ANY part of a formula is wrong, the formula is wrong. This includes:

• incorrect numbers in the subscripts

• missing or un-needed brackets

• numbers in the wrong place e.g. superscript when it should be subscript

• charges leftover in the formula (there should be no "+" or "-" signs)

• incorrect capitilisation e.g. ba instead of Ba or So4 instead of SO4

If you need reminders about how to write ionic formulae go back to the Year 10 page and revisit it; in particular you could watch this video.  Correct formulae are essential for balance equations and balanced equations are required for Excellence.

Another video tutorial on chemical formula    and a more advanced one

Ionic Formula Practice - PDF Ionic Formula Practice - Google Doc

Ionic Formula Mistakes - PDF Ionic Formula Mistakes - Google Doc

Naming and Formula - PDF Naming and Formula - Google Doc

More Formula Practice

Chemical equations

A chemical equation is a way of describing a chemical reaction. During a chemical reaction, one or more substances  (the reactatnts) will undergo a chemical change to form one or more new substances (the products),

A chemical equation has the reactants on the left, an arrow to show the chemical change and the products on the right. A common mistake is to use an equals sign instead of an arrow.

If you are writing digitally and there is no arrow character available, use two dashes and shift-period (>) to produce a pseudo-arrow: --> 

Word equations have the names of the reactants and products. Formula equations have the chemical formula for every chemical, both reactant and product. Formula equations MUST be balanced. A balanced equation has the same number of atoms on each side.

Words and formula should NEVER be mixed up in the same equation.

Example: calcium metal burns in air to produce calcium oxide

Equations: 

calcium + oxygen --> calcium oxide 

2Ca + O2 --> 2CaO 

The top equation is the word equation and the bottom one the formula equation. A common error would be for students to forget that oxygen is O2  not O. All element gases we deal with in the laboratory are diatomic (made of molecules with two atoms). A good tip is to write the formula for every separate chemical (what we call a 'species') FIRST then do the balancing. This is because students would otherwise be tempted to balance by changing CaO to CaO2 . This chemical (calcium peroxide) does exist, but it is not the chemical formed in this reaction because it does not form under these conditions (it forms when calcium is burned in pure oxygen at high temperature). Therefore balancing the equation by changing the formula from calcium oxide to calcium peroxide would be wrong. The formula of every chemical in the formula equation must match its name in the word equation.

Balancing equations tutorial video 1               Balancing equations video tutorial 2

Balancing equations worksheet (Google doc)                                     Balancing equation game

Types of reaction

There are quite a few different types of reaction, but we are dealing with four particular reaction types in this unit:

• combination reactions: this is when two different reactants combine to make a single product

• decomposition reactions: this is the reverse of a combination reaction, when a single reactant breaks down to form two or more different products.

• precipitation reactions: also called exchange reactions; this is when two solutions, containing two sets of ions each, are mixed and the ions 'exchange places' to produce a product that is insoluble . The product is in the form of a suspended solid called a precipitate

• displacement reactions: this occurs when a reactive metal is put into a solution containing ions of a less chemically reactive metal. Again they 'swap places' so that the reactive metal dissolves and forms ions and the less reactive one precipitates, plates or crystallizes out as pure metal.

In the next section we are going to look at each of these reaction types one by one.

Some common combination reactions

Although there are many possible combination reactions, you won’t be asked to do most of them because they are too dangerous e.g. sodium metal burns in chlorine gas to form sodium chloride (salt)

Sodium + chlorine --> sodium chloride

2Na + Cl2 --> 2NaCl

This is a combination reaction because we are combining two elements to make a single compound. For safety reasons, we can't do combination reactions that explode or produce poisonous gases. Below I am going to illustrate some of the safer and more common combination reactions you could encounter.

Reactions with oxygen

Iron will burn in oxygen as steel wool (very fine iron wire) or iron filings. Both of these forms have a very high surface area for a given mass of iron. This helps the reaction proceed, because the heat produced by the reaction helps keep it going.

The reaction requires heat to get it going - this is called activation energy. It also requires pure oxygen rather than air. This is because air is only 20% oxygen - if you use pure oxygen the particle collisions happen 5 times as often. This generates more heat which speeds the reaction up.

The word equation is

iron + oxygen --> iron II oxide

2Fe + O2 -->  2FeO

Iron is a greyish metal. Iron II oxide is black or blue black. The reaction will take place in air if you heat the iron wool continuously e.g. over a bunsen burner to provide enough energy for it (in pure oxygen the heat from the reaction is enough to keep it going).

Quite a few metals will burn in oxygen or air as powders. However, most of these are too dangerous (in the Ministry of Education's view) to be permitted in the school laboratory so you won't be asked to do them in a practical A selection are shown below.

Magnesium powder used to be used as a flash for cameras. A somewhat safer way of demonstrating the reaction of magnesium and oxygen is to burn magnesium ribbon over a bunsen burner:

Decomposition Reactions

These occur when one compound breaks down into two or more elements or compoinds

i.e.                    AB              ---->       A      +     B

There are lots of different types of decomposition, but only a few you need to know for this course:

• heating a metal carbonate (except sodium or potassium) causes it to decompose into the metal oxide plus carbon dioxide

• heating the hydrogen carbonates of either sodium or potassium causes them to decompose into the corresponding plain carbonate plus carbon dioxide and water vapour

• heating a metal hydroxide (except for sodium and potassium again) causes it to decompose into the metal oxide plus water vapour.

• hydrogen peroxide decomposes into water and oxygen gas in the presence of a suitable catalyst or other initiator such as ultraviolet light.

The first three are referred to as thermal decomposition and the last one is catalytic decomposition.

In most cases, heating a substance does not produce much by way of visible changes for white powders, although you may be able to see a disturbance as the gas evolves. However, some coloured powders will change colour.

Copper carbonate is teal green in its anhydrous form but a lighter green in its more common 'basic' form. You are more likely to get the light green form because it is easier to obtain (and cheaper). Copper hydroxide is a bluish colour. All three of these undergo thermal decomposition to black copper oxide when heated,

Pure copper carbonate gives of carbon dioxide only. Basic copper carbonate gives off both carbon dioxide and water vapour. Copper hydroxide gives off water vapour only.

CuCO3 (s)  -->   CuO (s)  + CO2 (g)

Cu(OH)2 (s)  -- >   CuO (s)  + H2O (g)

Although the 'basic' copper carbonate has a more complex equation, you don't need to do it (just the top equation)

Testing for carbon dioxide:

To test the gas given off to ensure that it really is carbon dioxide it needs to be bubbled through limewater. The limewater goes cloudy:

Decomposition of hydroxides - testing for water

When hydroxides decompose under heat they give off water - as a gas, because of the heat. An example is calcium hydroxide:

Ca(OH)2 (s) → CaO (s) + H2O(g)  

Decomposition of hydrogen peroxide

Hydrogen peroxide has the chemical formula H2O2.  The oxygen in it has a negative 1 charge (the peroxide ion) rather than the normal negative 2 charge. This ion is highly unstable, and can undergo a chemical reaction called disproportionation where half the peroxide ions gain one electron, becoming negative two oxide ions and forming water, and the other half of the peroxide ions lose their one electron to reach 'zero charge' to form oxygen gas. The chemical equation for this is:

hydrogen peroxide → water + oxygen

2 H2O2 (l) → 2 H2O (l) + O2 (g) 

Because hydrogen peroxide is so unstable it can undergo this reaction 'spontaneously', and normally very slowly. The reaction can be greatly sped up by a catalyst such as manganese dioxide. Note that the catalyst IS NOT part of the reaction so is not written into the chemical equation. In practice, many things can catalyse this reaction including enzymes (peroxidases), silver wire and quite a number of metals (silver, nickel, platinum). The video below shows oxygen being generated from hydrogen peroxide catalytically disproportionating under the catalysis of manganese dioxide, and relighting a glowing splint:

Displacement Reactions

Displacement reactions are a type of reaction that occurs when a piece of metal is placed into a solution containing positive ions (cations) that are less reactive than the metal.

Reactivity for metals is dependent on how easily they give away electrons.  The table below gives various elements in order of how easily they form cations and is called an 'Activity Series'. Calcium, on the left, is the most reactive and silver, on the right, the least.

For example, if you place a piece of magnesium into copper sulfate solution, the magnesium metal forms ions more readily than copper. This will cause the magnesium metal to give its electrons to the copper ions, forming copper metal. This happens on the surface of the magnesium because this is where the electrons are swapped, so a layer of copper forms on the magnesium.

In the video above we can see the red, metallic copper initially forming a black coating. This is made of tiny silver particles. As the silver crystals grow, they turn into feathery silvery looking crystals. Because these are good conductors of electrons, silver ions can continue to crystallize out on these as they receive their electrons, which are conducted through from the copper. Copper atoms pass their electrons along before they jump into solution. As the number of copper ions in solution increases, the solution turns blue due to this being the colour of aqueous Cu2+

 ions. The equation is

copper (s) + silver nitrate (aq)  →    copper nitrate (aq) +  silver (s)

Cu (s) + 2AgNO3 (aq)  →    Cu(NO3)2 (aq) +  2Ag (s)

or, in net ionic form

Cu (s) + 2Ag+ (aq)  →    Cu2+ (aq) +  2Ag (s)

Note that copper is the only metal that will change a solution' s colour significantly, but that this is fairly slow and requires observation for a period of time (10 min or so).

Explanations of why a displacement reaction occurs must refer to the transfer of electrons.

Metals and acids

Metals reacting with acids is also a type of displacement reaction. Metals that are to the left of hydrogen on the reactivity series will fizz in acid to produce hydrogen gas. Hydrogen is a metal - but only in extreme environments like the interior of Jupiter. In lab conditions it forms a diatomic molecular gas. Like other metals, it can give away electrons and is between lead and copper in its willingness to do so.

Calcium and magnesium are so reactive that they can react even with the tiny proportion of hydrogen ions present in neutral water - calcium does so quite vigorously. This is why we saw the bubbles forming as a by-reaction with the magnesium ribbon in copper sulfate.

Because Acids and Bases are in a different Achievement Standard, and the reaction of acids with metals is a displacement reaction (not an acid-base one), the strange logic of the NZQA means that metal-acid reactions aren't covered in either of the two inorganic Level 1 Chemistry standards. Therefore you won't be asked about them in either this internal or the matching external. Go figure.

Precipitation reactions

Precipitation reactions occur when two solutions (clear liquids) are mixed and it goes cloudy. The cloudiness is caused by the formation of an insoluble chemical compound in the form of tiny particles in suspension. These particles are often denser than the liquid they are in, so will fall down and settle on the bottom like rain. This process is called precipitation and the solid that has formed (whether or not it settles out) is called a precipitate.

In the video above, a clear liquid (sodium iodide solution) is being added to another clear liquid (lead nitrate solution). A bright yellow precipitate of lead iodide is forming because lead iodide is highly insoluble. Left for a while, this would settle to the bottom of the test tube because lead compounds are very heavy.  This is like rain or snow (precipitation; from Latin precipitum - to fall). Not all precipitates will do this. 

Ionic solids

Ionic compounds in solid form consist of positive and negative ions held together by the attraction of their opposite charges. Usually they are arranged in a regular pattern (thus forming crystals). Below is a diagram showing sodium ions (purple) and chloride ions (green) forming a crystal of sodium chloride (table salt).

The strong attractive forces between the ions hold them together and they can't easily be separated.

Water has an unusual property. The oxygen end of a water molecule attracts the shared electrons more than the hydrogens do, making it slightly negative. The hydrogen end becomes slightly positive. Water is thus a polar molecule - something explored in more depth in the Electricity topic.

This polarity gives water molecules the ability to arrange themselves around ions to 'neutralise' their charge:

This allows the ions to separate from each other, something they will do if they can. This is because of the natural tendency for things to 'spread out' if they are able to do so e.g. your washing dries on the clothesline because the water in it spreads out into the air (causing evaporation).

Not all ionic solids will dissolve in water to the same degree. This is because some ions attract each other much more strongly than they attract water molecules. The amount of a substance that will dissolve in an amount of solvent is called its solubility. 

For example, one gram of salt (sodium chloride) can be dissolved al little as 4 mL of water - salt is highly soluble. However, to dissolve 1 gram of barium sulfate would take over 400,000 litres of water. Therefore 4 mL of water would contain nearly undetectable  amounts of barium sulfate. Barium sulfate is highly insoluble.

Substances which dissolve in amounts we would find in our lab reagents are said to be soluble, and ones which don't are said to be insoluble for the purposes of our present unit of study.

Suppose we pour some sulfuric acid (a liquid containing sulfate ions) into barium chloride solution. Barium chloride is highly soluble, and the solution contains plenty of barium ions and chloride ions.

When we pour in the sulfuric acid, the sulfate ions meet the barium ions and 'grab on to them' far more strongly than the water molecules that surround them. As a result, a molecule of barium sulfate is formed. These molecules are highly charged and attract each other, so they rapidly grow into small particles of ionic solid made of tightly bound barium ions and sulfate ions. What we see is that a strongly white milky suspension forms - this is the precipitate of barium sulfate. After a few minutes it will settle to the bottom.

Suspensions are cloudy because the particles in them are large enough to scatter light. Dissolved ions are far smaller than light waves, so they don't scatter or refract them. As a result, solutions (containing dissolved ions) are clear and suspensions (containing suspended solids) are cloudy. 

Solutions can be clear and coloured if the ions absorb certain wavelengths of light e.g. copper ions in solution give the liquid a blue tinge. A solution of copper sulfate is described as a clear blue solution. We describe the ions in this solution as being aqueous, or the solution as being aqueous copper sulfate. In a chemical equation, we use the letters "aq" in brackets as a subscript to the right of a species to indicate that it is aqueous.

This means that the word equation for the reaction above would be:

sulfuric acid (aq)  +  barium chloride (aq)  →  barium sulfate (s)  + hydrochloric acid (aq)

H2SO4 (aq) +  BaCl2 (aq)  →   BaSO4 (s)  + 2HCl(aq)

or in net ionic form

Ba2+(aq)  +  SO42-(aq)  →   BaSO4 (s)

Note that the states MUST be present in the equation for it to be marked correct. Either net ionic or full formula equations can count for excellence provided they are correctly balanced.

Solubility rules

You are given the following data table to help you work out what is the precipitate in a given reaction:

This is provided in a booklet with the examination paper.

In addition, the following information helps you identify substances formed:

For some reason they have missed out copper carbonate (which is a blue or blue-green solid). The title does say 'selected' - what the basis is for selection I have no idea.