Ionic bonding

Ions are electrically charged particles formed when atoms lose or gainelectrons. They have the same electronic structures as noble gases.

Metal atoms form positive ions, while non-metal atoms form negative ions. The strong electrostatic forces of attraction between oppositely charged ions are called ionic bonds.

Ions

How ions form

Ions are electrically charged particles formed when atoms lose or gain electrons. This loss or gain leaves a complete highest energy level, so the electronic structure of an ion is the same as that of a noble gas - such as a helium, neon or argon.

Metal atoms and non-metal atoms go in opposite directions when they ionise:

• Metal atoms lose the electron, or electrons, in their highest energy level andbecome positively charged ions
• Non-metal atoms gain an electron, or electrons, from another atom tobecome negatively charged ions

Positively charged sodium and aluminium ions

Negatively charged oxide and chloride ions

How many charges?

There is a quick way to work out what the charge on an ion should be:

• The number of charges on an ion formed by a metal is equal to the group number of the metal
• The number of charges on an ion formed by a non-metal is equal to the group number minus eight
• Hydrogen forms H⁺ ions.

 

	Group 1	Group 2	Group 3	Group 4	Group 5	Group 6	Group 7	Group 0
Example element	Na	Mg	Al	C	N	O	Cl	He
Charge	1+	2+	3+	Note 1	3-	2-	1-	Note 2
Symbol of ion	Na+	Mg2+	Al3+	Note 1	N3-	O2-	Cl-	Note 2

Note 1: carbon and silicon in Group 4 usually form covalent bonds by sharing electrons.

Note 2: the elements in Group 0 do not react with other elements to form ions.

Representing positive ions

You need to be able to show the electronic structure of some common metal ions, using diagrams like these:

Lithium, Li

Lithium is in Group 1. It has one electron in its highest energy level. When this electron is lost, a lithium ion Li⁺ is formed.

Sodium, Na

Sodium is also in Group 1. It has one electron in its highest energy level. When this electron is lost, a sodium ion Na⁺ is formed.

Neon atom

Note that a sodium ion has the same electronic structure as a neon atom (Ne).

But be careful - a sodium ion is not a neon atom. This is because the nucleus of a sodium ion is the nucleus of a sodium atom and has 11 protons - but the nucleus of a neon atom has only 10.

Magnesium, Mg

Magnesium is in Group 2. It has two electrons in its highest energy level. When these electrons are lost, a magnesium ion Mg²⁺ is formed.

A magnesium ion has the same electronic structure as a neon atom (Ne).

Calcium, Ca

Calcium is also in Group 2. It has two electrons in its highest energy level. When these electrons are lost, a calcium ion Ca²⁺ is formed.

A calcium ion has the same electronic structure as an argon atom (Ar).

Representing negative ions

You need to be able to show the electronic structure of some common non-metal ions, using diagrams like these:

Fluorine, F

Fluorine is in Group 7. It has seven electrons in its highest energy level. It gains an electron from another atom in reactions, forming a fluoride ion, F^-.

Note that the atom is called fluorine, but the ion is called fluoride.

Neon atom

Note that a fluoride ion has the same electronic structure as a neon atom (Ne).

Once again, a fluoride ion is not a neon atom, because thenucleus of a fluoride ion is the nucleus of a fluorine atom, with 9 protons, and not of a neon atom, with 10.

Chlorine, Cl

Chlorine is in Group 7. It has seven electrons in its highest energy level. It gains an electron from another atom in reactions, forming a chloride ion, Cl^-.

Oxygen, O

Oxygen is in Group 6. It has six electrons in its highest energy level. It gains two electrons from one or two other atoms in reactions, forming an oxide ion, O^2-.

When metals react with non-metals, electrons are transferred from the metal atoms to the non-metal atoms, forming ions. The resulting compound is called anionic compound.

Consider reactions between metals and non-metals, for example:

• sodium + chlorine → sodium chloride
• magnesium + oxygen → magnesium oxide
• calcium + chlorine → calcium chloride

In each of these reactions, the metal atoms give electrons to the non-metal atoms. The metal atoms become positive ions and the non-metal atoms become negative ions.

There is a strong electrostatic force of attraction between these oppositely charged ions, called an ionic bond. The animation shows ionic bonds being formed in sodium chloride, magnesium oxide and calcium chloride.

Group 1 and Group 7

The elements in Group 1 of the Periodic Table are called the alkali metals. They form ionic compounds when they react with non-metals. Their ions have a single positive charge. For example, sodium forms sodium ions, Na+.

The elements in Group 7 of the Periodic Table are called the halogens. They form ionic compounds when they react with metals. Their ions have a single negative charge. For example, chlorine forms chloride ions, Cl–.

Sodium chloride

Sodium chloride, NaCl, forms when sodium and chlorine react together. It contains oppositely charged ions held together by strongelectrostatic forces of attraction – the ionic bonds. The ions form a regular lattice in which the ionic bonds act in all directions.

Dot-and-cross diagrams

You need to be able to draw dot-and-cross diagrams to show the ions in some common ionic compounds.

Sodium chloride, NaCl

Sodium ions have the formula Na⁺, while chloride ions have the formula Cl^-. You need to show one sodium ion and one chloride ion. In the exam, make sure the dots and crosses are clear, but do not worry about colouring them.

Magnesium oxide, MgO

Magnesium ions have the formula Mg²⁺, while oxide ions have the formula O^2-. You need to show one magnesium ion and one oxide ion.

Calcium chloride, CaCl₂

Calcium ions have the formula Ca²⁺. Chloride ions have the formula Cl^-.

You need to show two chloride ions, because two chloride ions are needed to balance the charge on a calcium ion.

Formulae of ionic compounds

Ionic compounds are represented by formulae. Symbols and numbers show the atoms in the compound.

For example: ZnCO₃ is the formula for zinc carbonate.

One zinc atom (Zn) and one carbon atom (C) are chemically bonded with three oxygen atoms (O₃). Notice that we don't need to write a 1 next to the Zn or C.

Brackets in formulae

Sometimes brackets are used.

For example: Fe(OH)₃ is the formula for iron(III) hydroxide.

Iron(III) hydroxide consists of one iron atom joined with three oxygen and three hydrogen atoms. The formula is written like this because the oxygen and hydrogen atom often act together.

Constructing formulae

The formula of a compound can be worked out if the ions in it are known. For example, the compound formed from Na⁺ and SO₄^2- will consist of two Na⁺ ions to every one SO₄^2- ion so that the compound is neutral overall. The formula is therefore Na₂SO₄.

Here are the formulae of some common ions.

Positive ions (cations)

Name	Formula	Name	Formula	Name	Formula
ammonium	NH4+	magnesium	Mg2+	zinc	Zn2+
hydrogen	H+	calcium	Ca2+	lead	Pb2+
lithium	Li+	barium	Ba2+	iron(II)	Fe2+
sodium	Na+	silver	Ag+	iron(III)	Fe3+
potassium	K+	copper(II)	Cu2+	aluminium	Al3+

Negative ions (anions)

Name	Formula	Name	Formula	Name	Formula
fluoride	F-	hydrogen carbonate	HCO3-	sulfide	S2-
chloride	Cl-	hydroxide	OH-	sulfate	SO42-
bromide	Br-	nitrate	NO3-	carbonate	CO3-
iodide	I-	oxide	O2-

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The future of copper

We are running out of copper-rich ores. Research is being carried out to find new ways to extract copper from the remaining low-grade ores, without harming the environment too much. This research is very important, as traditional mining involves huge open-cast mines that produce a lot of waste rock.

Watch

You may wish to view thisBBC News item (2005) about a huge copper mine in Chile, South America.

Watch

Phytomining, bioleaching and scrap iron

Some plants absorb copper compounds through their roots. They concentrate these compounds as a result of this. The plants can be burned to produce an ash that contains the copper compounds. This method of extraction is calledphytomining.

Some bacteria absorb copper compounds. They then produce solutions called leachates, which contain copper compounds. This method of extraction is calledbioleaching.

Copper can also be extracted from solutions of copper salts using scrap iron. Iron is more reactive than copper, so it can displace copper from copper salts. For example:

iron + copper sulfate → iron sulfate + copper

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Copper

Copper is soft and easily bent and so is a good conductor of electricity, which makes it useful for wiring. Copper is also a good conductor of heat and it does not react with water. This makes it useful for plumbing, and making pipes and tanks.

Copper ores

Some copper ores are copper-rich – they have a high concentration of copper compounds. Copper can be extracted from these ores by heating them in a furnace, a process called smelting. The copper is then purified using a process called electrolysis.

Electricity is passed through solutions containing copper compounds, such as copper sulfate. During electrolysis, positively charged copper ions move towards the negative electrode and are deposited as copper metal.

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Extracting metals and making alloys

Metals are very useful. Ores are naturally occurring rocks that contain metal or metal compounds in sufficient amounts to make it worthwhile extracting them: most everyday metals are mixtures called alloys.

Methods of extracting metals

The Earth's crust contains metals and metal compounds such as gold, iron oxide and aluminium oxide, but when found in the Earth these are often mixed with other substances. To become useful, the metals have to be extracted from whatever they are mixed with. A metal ore is a rock containing a metal, or a metal compound, in high enough concentration to make it economic to extract the metal.

Ores are mined. They may need to be concentrated before the metal is extracted and purified. The economics of using a particular ore may change over time. For example, as a metal becomes rarer, an ore may be used when it was previously considered too expensive to mine.

Reactivity and extraction method

Metals are produced when metal oxides are reduced (have their oxygen removed). The reduction method depends on the reactivity of the metal. For example, aluminium and other reactive metals are extracted by electrolysis, while iron and other less reactive metals may be extracted by reaction with carbon or carbon monoxide.

Reactivity and extraction method

Metals (in decreasing order of reactivity)	Method of extraction
potassium sodium calcium magnesium aluminium	extract by electrolysis
carbon
zinc iron tin lead	extract by reaction with carbon orcarbon monoxide
hydrogen
copper silver gold platinum	extracted in various ways

The method of extraction of a metal from its ore depends on the metal's position in the reactivity series.

Gold, because it is so unreactive, is found as the native metal and not as acompound. It does not need to be chemically extracted from its ore, but chemical reactions may be needed to remove other elements that might contaminate the metal.

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Quarrying

You need to be able to evaluate some of the effects of the limestone industry.

The main advantages and disadvantages of the limestone industry

Advantages	Disadvantages
Limestone is a valuable natural resource, used to make things such as glass and concrete.	Limestone quarries are visible from long distances and may permanently disfigure the local environment.
Limestone quarrying provides employment opportunities that support the local economy in towns around the quarry.	Quarrying is a heavy industry that creates noise and heavy traffic, which damages people's quality of life.

Links

You may wish to read thisBBC News item (2006) about a public inquiry into a limestone quarry in the Peak District National Park. Consider the arguments for and against quarrying limestone in a national park.

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Uses of limestone

Limestone is a type of rock, mainly composed of calcium carbonate. Limestone is quarried (dug out of the ground) and used as a building material. It is also used in the manufacture of cement, mortar and concrete.

Reactions with acids

Carbonates react with acids to produce carbon dioxide, a salt and water. For example:

calcium carbonate + hydrochloric acid → carbon dioxide + calcium chloride + water

CaCO₃ + 2HCl → CO₂ + CaCl₂ + H₂O

Since limestone is mostly calcium carbonate, it is damaged by acid rain. Sodium carbonate, magnesium carbonate, zinc carbonate and copper carbonate also react with acids: they fizz when in contact with acids, and the carbon dioxide released can be detected using limewater.

Calcium hydroxide

When limestone is heated strongly, the calcium carbonate it contains decomposes to form calcium oxide. This reacts with water to form calcium hydroxide, which is an alkali. Calcium hydroxide is used to neutralise excess acidity, for example, in lakes and soils affected by acid rain.

Cement, mortar and concrete

Cement is made by heating powdered limestone with clay. Cement is an ingredient in mortar and concrete:

• mortar, used to join bricks together, is made by mixing cement with sand and water
• concrete is made by mixing cement with sand, water and aggregate (crushed rock)

Advantages and disadvantages of various building materials

Limestone, cement and mortar slowly react with carbon dioxide dissolved in rainwater and wear away. This damages walls made from limestone, and leaves gaps between bricks in buildings. These gaps must be filled in or ‘pointed’. Pollution from burning fossil fuels makes the rain more acidic than it should be, and this acid rain makes these problems worse.

Concrete is easily formed into different shapes before it sets hard. It is strong when squashed, but weak when bent or stretched. However, concrete can be made much stronger by reinforcing it with steel. Some people think that concrete buildings and bridges are unattractive

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Products from calcium carbonate

For your exam, you need to know how calcium hydroxide is obtained from calcium carbonate.

Making calcium oxide

If calcium carbonate is heated strongly, it breaks down to form calcium oxide and carbon dioxide. Calcium oxide is yellow when hot, but white when cold.

Here are the equations for this reaction:

calcium carbonate  calcium oxide + carbon dioxide

CaCO₃ CaO + CO₂

This is a thermal decomposition reaction.

Making calcium hydroxide

Calcium oxide reacts with water to form calcium hydroxide, which is an alkali. Here are the equations for this reaction:

calcium oxide + water → calcium hydroxide

CaO + H₂O → Ca(OH)₂

A lot of heat is produced in the reaction, which may even cause the water to boil.

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Transition metals

The transition metals are placed in the periodic table in a large block between groups 2 and 3. Most metals (including iron, titanium and copper) are transition metals.

The transition metals

Common properties

The transition metals have these properties in common:

• they are metals
• they are good conductors of heat and electricity
• they can be hammered or bent into shape easily

The transition metals are useful as construction materials. They are also useful for making objects that need to let electricity or heat travel through them easily.

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Calcium carbonate

Limestone is mainly calcium carbonate, CaCO₃, which when heated breaks down to form calcium oxide and carbon dioxide. Calcium oxide reacts with water to produce calcium hydroxide. Limestone and its products have many uses, including being used to make cement, mortar and concrete.

Thermal decomposition

Calcium carbonate breaks down when heated strongly. This reaction is calledthermal decomposition. Here are the equations for the thermal decomposition of calcium carbonate:

calcium carbonatecalcium oxide + carbon dioxide

CaCO₃CaO + CO₂

Other metal carbonates decompose in the same way, including:

• sodium carbonate
• magnesium carbonate
• copper carbonate

For example, here are the equations for the thermal decomposition of copper carbonate:

copper carbonate copper oxide + carbon dioxide

CuCO₃CuO + CO₂

Metals high up in the reactivity series (such as sodium, calcium and magnesium) have carbonates that need a lot of energy to decompose them. Indeed, not all the carbonates of group 1 metals decompose at the temperatures reached by a Bunsen burner.

Metals low down in the reactivity series, such as copper, have carbonates that are easily decomposed. This is why copper carbonate is often used at school to show thermal decomposition. It is easily decomposed and its colour change, from green copper carbonate to black copper oxide, is easy to see.

The thermal decomposition of copper(II) carbonate is easily demonstrated

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