NCEA Level 1 Science/Properties and changes of matter

Vapors of hydrogen chloride in a beaker and ammonia in a test tube meet to form a cloud of a new substance, ammonium chloride

Introduction

Metals are commonly found all around us in our everyday lives. Usually, we see them as compounds such as stainless steel (made of iron, nickel and chromium). In chemistry, metals are elements found on the left and middle of the periodic table. Metals in the middle are called transition metals.

Physical properties

Most metals have the following physical properties:

• electrical conductivity (has free electrons)
• thermal conductivity (heat conductor)
• density (tightly packed atom structure)
• ductility (able to be drawn into wires ie. electrical wires)
• lustre (shiny)
• malleability (able to be beaten into shapes)

Metal are usually solid at room temperature (20˚C) with an exception to mercury which has a melting point of -39˚C. Metals are also usually grey or silver with the exception of copper.

• 
  Electrical conductivity
• 
  Thermal conductivity
• 
  Density
• 
  Ductility
• 
  Lustre
• 
  Malleability

Reaction of metals

Magnesium fire. The first stream comes for a CO₂ fire extinguisher, the second from a powder extinguisher, and the third from a foam one. Since the camera sets the sensibility according to the most lightning source, keep an eye to the setting to have an idea of the light emited by the fire.

Activity (reactivity) series

Metal	Na	Ca	Mg	Al	Zn	Fe	Pb	Cu
Symbol	Na+	Ca2+	Mg2+	Al3+	Zn2+	Fe2+	Pb2+	Cu2+
Reactivity

The reactivity series shows how reactive a metal is to oxygen, water or acids. Sodium and calcium are the most reactive. Sodium is the most reactive because it is 1 electron from having a full valence energy level.

Metals and oxygen

Helpful Hint!
Metals + Oxygen → Metal oxide

When metals react with oxygen, they form a metal oxide coating.

With word equations, the names of the reactants and products are required. For chemical equations, the drop and swap rule should be followed. In addition, the equation needs to be balanced so the number of reactants is equal to the number of products.

Metal	Reaction	Comments
Sodium	Sodium + Oxygen → Sodium oxide 4Na + O2 → 2Na2O	Upon exposure to air, sodium reacts immediately with oxygen so it is necessary to place it under oil to prevent contact with air.
Calcium	Calcium + Oxygen → Calcium oxide 2 Ca + O2 → 2CaO	Calcium is highly reactive to oxygen since it is two electrons away from having a full outer shell.
Magnesium	Magnesium + Oxygen → Magnesium oxide 2Mg + O2 → 2MgO	In the magnesium oxide reaction, a bright white light is observed when the magnesium is burnt in the bunsen flame. The product, magnesium oxide is the white powder that results.
Aluminium	Aluminium + oxygen → Aluminium oxide 4Al +3 O2 → 2Al2O3	The aluminium oxide layer is complete leaving no gaps thus prevents the aluminium from contact with oxygen, water or acids. This insoluble coating makes aluminium appear to be unreactive despite aluminium’s high ranking on the activity series.
Zinc	Zinc + oxygen → Zinc oxide 2Zn + O2 → 2ZnO	Zinc is quite reactive it will easily oxidise.
Iron	Iron + Oxygen → Iron oxide 4Fe +3 O2 →2Fe2O3	The iron oxidation reaction is commonly known as rusting. Rusting slowly occurs in the presence of oxygen and water forming iron oxide.
Lead	Lead + Oxygen → Lead oxide 2Pb + O2 → 2PbO	Lead (II) oxide occurs when lead is heated to 600°C and is red or yellow. Like lead, lead oxide is toxic and is a component in lead paint.
Copper	Copper + Oxygen → Copper Oxide 2Cu + O2 → 2CuO	Copper being at the end of the activity series is quite unreactive. It will take a long time to react with oxygen. Being a transition metal, it will change colour when it becomes a compound. Copper turns from brownish-red to black when it becomes copper oxide.

Rusting

A heart-shaped patch of rusted metal, showing through cracked and flaking paint; a metal wall, near railway tracks.

The iron oxide reaction is commonly known as rusting.

Rusting occurs in the presence of oxygen (not air which is mostly nitrogen) and water. Both elements are required for a metal to rust or to oxidise. With those conditions present, iron will combine with air and water to form hydrated iron oxide.

Rusting is a very expensive problem because it corrodes the metal eventually completely dissolving it. Rusting can be prevented by:

• alloying

By alloying iron with chromium or nickel or both, stainless steel can be formed which is chemically resistant to corrosion. This is the most expensive method by it makes the metals absolutely rustproof.

• galvanising

Metals that react more readily with corroding substances than iron will sacrifice themselves to protect iron (these metals can be found in the activity series above). This sacrificial corrosion is commonly used to protect iron when used on rooftops in the form of galvanised iron (iron galvanised with zinc). In the presence of corrosive substances, zinc forms an electric potential with iron causing it to sacrifice itself and protect the iron as long as the zinc remains. However, if another lesser reactive metal touches this zinc coating, the zinc will dissolve to prevent the corrosion of this lesser reactive metal. When the zinc has corroded, the iron will begin to corrode.

Fence which has rusted

• coating

The coating must completely cover over the iron as well as being impermeable (preventing contact with oxygen and water). This works as long as the coating completely covers it. It is the least expensive therefore the most common. Paint is a commonly used coating to coat metals like cars. The protective layer may also be an unreactive metal such as chromium and tin (used in cans called tin cans). However, this means that the electric potential is set up to protect the layer so if the coating is broken, rusting will occur at an accelerated rate on the iron.

Special oxide layers

Though aluminium is quite reactive chemically, aluminium forms a thin, transparent and complete oxide layer that protects it from further corrosion. Thus under normal atmospheric conditions, it appears to be unreactive.

Zinc and lead which are less reactive than aluminium also form similar protective layers.

Metals and Water

Helpful Hint!
Metal + water → metal hydroxide + hydrogen

Metals react with water to form metal hydroxide (liquid) and hydrogen gas. The metal hydroxide is a base. Metals such as sodium react violently with water while copper is quite unreactive.

Examples:

Sodium + water → Sodium hydroxide + hydrogen

Mg + H₂O → Mg(OH)₂ → H₂

Zn + H₂O → Zn(OH)₂ + H₂

Calcium + water → calcium hydroxide + hydrogen

Pb + 2H₂O → Pb(OH)₂ + H₂

Iron + water → Iron hydroxide + hydrogen

Al + H₂O → Al(OH)₃ + H₂O

H is the element hydrogen and H₂ is the gas.

The test for hydrogen gas is the pop test. The pop test involves collecting igniting a sample of gas. Hydrogen burns with a pop sound.

Metals and Acids

Helpful Hint!
Metal + acid → metal ion (salt) + hydrogen

When metals react with acids, a metal ion (salt) is formed and hydrogen gas is produced.

Examples:

Zinc + hydrochloric acid → zinc chloride + hydrogen

Ca + H₂SO₄ + CaSO₄ + H₂

Magnesium + hydrochloric acid → magnesium chloride + hydrogen

Mg + HCl → MgCl₂ + H₂

Sodium + sulphuric acid → sodium sulphate + hydrogen

Al + H₂SO₄ → Al₂(SO₄)₃ + H₂

Lead + hydrochloric acid → lead chloride + hydrogen

Cu + H₂SO₄ → CuSO₄ + H₂

Balancing Equations

Helpful Hint!
reactants → products

A balanced equation is where the number of atoms for each element is the same on each side of the arrow. To balance an equation, we put large numbers in front of elements of compounds.

Example

2Mg + O2 → 2MgO

Effects on indicators

Acid

An acid is a compound that contains hydrogen and can release hydrogen ions (H⁺) in water, producing a concentration of hydrogen ions which is more than what is found in pure water. Acids taste sour and can be corrosive.

Examples of organic acids (acids with carbon):

• 
  Coke (carbonic acid)
• 
  Citrus fruits (citric acid)
• 
  Vinegar (acetic acid)
• 
  Bee Sting
• 
  Leaves (Oxalic acid)
• 
  Aspirin (Salicylic Acid)
• 
  Stomach acid (hydrochloric acid)

Hydrochloric acid is common laboratory reagent.

The inorganic acids used in laboratories include

• Hydrochloric acid (HCl)
• Sulphuric acid (H₂SO₄)

Hydrogen Chloride (HCl) is an acid because it dissolves in water to give a solution of hydrogen ions and chloride ions known as hydrochloric acid. A clue to whether a substance is an acid is whether its formula begins with H, as in HCl and H₂SO₄. However, this is not true in all cases.

Acids are corrosive because when dissolved in water (i.e. hydrogen chloride dissolved to become hydrochloric acid), they release hydrogen ions. These hydrogen are highly reactive as they seek to bond with another compound to form become stable. When acid spills into skin, the acid breaks down the skin by joining onto compounds that make up the skin. However, in a school laboratory environment, the acid is usually highly diluted but will still have the potential to cause harm if not quickly treated.

Bases

Alkalis are a class of bases that taste bitter and feel slippery. When dissolved in water, a base produces an excess of hydroxide ions (OH^-.)

Example of common household bases:

• Soap
• Oven Cleaners
• Cleaning Products
• Indigestion tablets
• Laundry detergents
• Household cleaners
• Dishwashing liquid

Common bases include:

• Sodium hydroxide (NaOH)
• Calcium hydroxide (Ca(OH)₂)
• Ammonia (NH₃)

Neutralisation

When acids and alkalis react with each other, a neutralisation reaction occurs to give a salt and water. The point at which the mixture becomes neutral can be found with an indicator such as litmus.

The neutralization reaction then becomes:

H⁺ + OH^- → H₂O

However, due the presence of other compounds in the acid and base reaction, the actual reaction forms a salt.

Acid + base → salt + water

Examples

Sodium hydroxide (also known as caustic soda) is mixed with hydrochloric acid to form water and sodium chloride.

Sodium hydroxide + hydrochloric acid → water + sodium chloride NaOH + HCl → H₂O + NaCl

The driving force behind this reaction is the formation of a solvent which in this case is water. Though water molecules have been formed, the sodium (Na⁺) and the chloride (Cl^-) ions are still separated. By evaporating the water, sodium chloride (NaCl) will form as the ions bond together to form a compound.

Indicators

An indicator is any chemical or natural product that changes colour in an acidic, basic or neutral solution. Common chemical indicators are:

• universal indicator
• &litmus (dye and paper)

The most useful of these is the universal indicator. Universal indicator is a mixture of several different chemical indicators. The range of colours is possible because colour of the universal indicator solution at a particular pH is derived from the colours of the individual indicators. The overall effect is a gradual colour change over a large range of pH.

Litmus and many other indicators are derived from plant juices. Litmus is made from lichens. Many vegetable juices such as beetroot juice can also change colour depending on a solution’s acidity, alkalinity or neutrality.

Colour change at different levels of acidity and alkalinity

Indicator	Acids	Colour	Neutral	Colour	Base	Colour
Universal Indicator	Red, orange, yellow		Green		Blue, purple
Litmus	Red				Blue

Examples of Colour Changes

Solution	Colour change	Colour	Solution Type
Jif	Blue		Base
Milk	Yellow		Weak acid
Oven cleaner	Blue		Base
Coffee	Yellow		Weak acid
Lime juice	Red		Acid
Sprite	Red		Acid
Vinegar	Red		Acid
Dish washing liquid	Blue		Base
Tea	Yellow		Weak acid

pH values

The pH (potential for hydrogen) scale is the measure of how acid or alkaline a solution is.

Universal Indicator

The universal indicator colours determine the strength of an acid or base. Moving away from the pH of 7 will increase the strength of the acid or base.

pH range	Description	Colour	Colour sample
0	Strong acid	Red
1	Strong acid	Red
2	Strong acid	Red
3	Mild Acid	Orange
4	Mild Acid	Orange
5	Weak Acid	Yellow
6	Weak Acid	Yellow
7	Neutral	Green
8	Weak base	Blue
9	Weak base	Blue
10	Mild base	Blue
11	Mild base	Blue
12	Strong base	Purple
13	Strong base	Purple
14	Strong base	Purple

Acids and carbonates

Acids and Metal Oxides

When metal oxides react with acids, a salt is formed and water is produced.

Helpful Hint!
metal oxide + acid → salt + water

Examples

MgO + 2HCl → MgCl₂ + H₂O

Magnesium oxide + hydrochloric acid → Magnesium chloride + water

FeO + 2HCl → FeCl₂ + H₂O

Sodium oxide + sulphuric acid → sodium sulphate + water

Al₂O₃ + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂O

Lead oxide + hydrochloric acid → lead chloride + water

CuO + H₂SO₄ → CuSO₄ + H₂O

Metals oxides are basic and can be neutralised with acids.

Acids and Hydroxides

The acids with metal hydroxide reaction also produces a salt and water.

Helpful Hint!
metal hydroxide + acid → salt + water

Examples

Cu(OH)₂ + H₂SO₄ → CuSO₄ + H₂O

Copper hydroxide + Sulfuric acid → Copper Sulfate + water

Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O

Magnesium hydroxide + Hydrochloric acid → Magnesium chloride + water

2Al(OH)₂ +3H₂SO₄ → 2Al₂(SO₄)₃ + 6H₂O

Acids and Carbonates

Helpful Hint!
metal carbonate + acid → salt + carbon dioxide + water

Metal carbonates react with acids to produce a salt, carbon dioxide gas and water.

Examples

ZnCO₃ + HCl → ZnCl₂ + CO₂ + H₂O

The test for CO₂ is limewater turns cloudy.

Acids and Hydrogen Carbonates

Helpful Hint!
metal hydrogen carbonate + acid → salt carbon dioxide + water

Metal hydrogen carbonates (bicarbonates) also react with acids to produce salt, carbon dioxide and water.

Examples

Calcium bicarbonate + Sulphuric acid → Calcium sulphate + carbon dioxide + water

NaHCO₃ + HCl → NaCl + CO₂ + H₂O

Baking Soda

NaHCO₃ is commonly known as baking soda. Baking powder is NaHCO₃ with tartaric acid added to improve the taste.

Metals

Metal	Symbol	Example
Sodium	Na
Calcium	Ca
Magnesium	Mg
Aluminium	Al
Zinc	Zn
Iron	Fe
Lead	Pb
Copper	Cu