Chapter overview
TABLE 2.1.3 FORMULAS OF A RANGE OF COMMON NEGATIVE IONS, INCLUDING POLYATOMIC IONS
+1 +2 +3 +4
Lithium Li+ Sodium Na+ Potassium K+ Hydrogen H+ Copper(I) Cu+ Silver Ag+ Ammonium NH4+ Hydronium H3O+
Magnesium Mg2+ Calcium Ca2+ Strontium Sr2+ Barium Ba2+ Iron(II) Fe2+ Copper(II) Cu2+ Zinc Zn2+ Lead(II) Pb2+
Aluminium Al3+ Chromium(III) Cr3+ Iron(III) Fe3+
Tin(IV) Sn4+ Lead(IV) Pb4+
TABLE 2.1.3 FORMULAS OF A RANGE OF COMMON NEGATIVE IONS, INCLUDING POLYATOMIC IONS
−1 −2 −3
Fluoride F− Chloride Cl− Bromide Br− Iodide I− Hydroxide OH− Nitrate NO3−
Permanganate MnO4− Cyanide CN−
Hydrogen sulfate HSO4− Hydrogen carbonate HCO3− Ethanoateion CH3COO−
Oxide O2− Sulfate SO42− Sulfide S2− Carbonate CO32− Chromate CrO42− Dichromate Cr2O72− Thiosulfate S2O32−
Nitride N3− Phosphate PO43−
All ionic formulas are made up of positive and negative ions. The total positive charge will always equal the total negative charge because, overall, a compound is neutral. Sometimes multiple ions will be needed to achieve this neutrality.
Ionic formulas are written as empirical formulas. An empirical formula is the lowest whole number ratio of the atoms in a compound.
All areas of knowledge have conventions or rules, sets of common understandings to make communication easier. For example, every language has grammar conventions. To name and write a balanced formula for a particular ionic compound, the following conventions are used.
• When naming an ionic compound, the positive ion is generally written fi rst, followed by the negative ion. For example, NaCl is called sodium chloride.
(The main exceptions here are salts of organic acids, such as sodium ethanoate, CH3COONa (see chapters 9 and 11).)
• As compounds do not carry an overall charge, it is necessary to balance the charges of the anion and cation components, for example NaCl, MgO, KOH, CaCO3, HNO3.
• If more than one of each ion is required to balance the overall charge, a subscript is used to indicate the number of each species required, for example H2SO4, AlCl3, K2S, PbCl2, K2Cr2O7.
Naming ionic compounds Writing ionic formulas
• In the case of some polyatomic ions, it may be necessary to use brackets to ensure no ambiguity is present. For example, CuOH2 is incorrect; it should be Cu(OH)2. Similarly Al2SO43 is incorrect; it should be Al2(SO4)3.
• For metals that are able to form ions of different charges, Roman numerals are used to indicate the relevant charge on the ion; for example, FeCl2 would be named as iron(II) chloride and FeCl3 would be named as iron(III) chloride.
• The name of the ion depends on its composition. For example, the ending -ate for a polyatomic ion indicates the presence of oxygen; the ending -ide indicates that the ion is made up of a single atom with a negative charge such as sulfi de, S2−, chloride, Cl−.
• The name of the ion often gives a clear indication of the atoms present.
The hydroxide ion is made up of hydrogen and oxygen, OH−, while
hydrogencarbonate ions are made up of hydrogen, carbon and oxygen, HCO3−.
Worked example 1 Solution Write balanced chemical
formulas for the following compounds.
a Potassium bromide b Sodium nitrate c Magnesium sulfi de d Sodium oxide e Calcium chloride f Iron(III) sulfate
Worked example 2 Solution
State the name of each of the following compounds, given their formulas.
a NaOH a Sodium hydroxide
b KMnO4 b Potassium permanganate
c Cu2SO4 c Copper(I) sulfate
d Ca(NO3)2 d Calcium nitrate
e FePO4 e Iron(III) phosphate
f (NH4)2S f Ammonium sulfi de
So far we have seen that the positive charge on the cation must be balanced by the negative charge on the anion to produce a neutral compound. Ionic compounds do not exist in nature as separate units of, for example, one sodium ion and one chloride ion (NaCl). In fact, the ions that make up ionic compounds arrange themselves into a regular pattern, a lattice structure (an ionic
lattice), containing many millions of ions that extend in all three dimensions.
No fi xed number of ions is involved, but the ratio of cations to anions is constant for a given compound and is shown in the empirical formula.
Compound Cation Anion Formula
a K+ Br− KBr
b Na+ NO3− NaNO3
c Mg2+ S2− MgS
d Na+ (× 2) O2− Na2O
e Ca2+ Cl− (× 2) CaCl2 f Fe3+ (× 2) SO42− (× 3) Fe2(SO4)3
The structure of ionic compounds
4.1.8
Describe the lattice structure of ionic compounds. © IBO 2007 WORKSHEET 2.2
Charges on ions
CHAPTER 2 BONDING
chloride ion a
sodium ion
chloride ion b
sodium ion
c
Figure 2.1.4 (a) A close-packed model of part of the sodium chloride crystal lattice. (b) A ball-and-stick model of the sodium chloride crystal lattice. (c) Scanning electron micrograph of sodium chloride crystals.
The most stable arrangement of ions for any particular ionic compound will be the one in which the positively charged ions are packed as closely as possible to the negatively charged ions, and the ions with the same charge are as far apart as possible. This arrangement serves to maximize the electrostatic attraction between the positive and negative ions and minimize the repulsion between like charged ions, thus lowering the overall chemical potential energy of the lattice. There are a number of different ion arrangements that can be generated to meet these criteria, depending on the relative sizes of the ions present and their ratio in the compound. Each arrangement will result in the particular lattice structure found for that compound.
Sodium chloride provides a good example of an ionic lattice. Its structure is shown in fi gure 2.1.4. Each positive sodium ion is surrounded by six chloride ions, and each chloride ion is surrounded by six sodium ions. A sodium chloride crystal is cubic in shape.
1 Identify the bonding type in each of the following cases as ionic, covalent or metallic.
a Silver b Hydrogen and oxygen c Magnesium and chlorine d Copper and sulfur e Carbon and oxygen f Tin and copper g Aluminium and sulfur
2 Draw electron shell diagrams for each of the following pairs of elements to represent how they interact to form ions. Clearly state the formula of the resulting compound.
a Sodium and fl uorine to produce sodium fl uoride b Calcium and sulfur to produce calcium sulfi de c Calcium and chlorine to produce calcium chloride d Aluminium and oxygen to produce aluminium oxide
PRAC 2.1
Modelling sodium chloride
DEMO 2.1
Conductivity of ionic compounds Section 2.1 Exercises
3 State balanced chemical formulas for each of the following compounds.
a Potassium nitrate b Calcium chloride c Sodium hydroxide d Copper(II) sulfate e Ammonium sulfi de f Aluminium nitrate
4 State the names of the following compounds.
a KCl b BaSO4 c HNO3
d Al2O3 e SnI2 f Cu3(PO4)2
5 Determine the charge on the positive ion (cation) for each of the following compounds.
a Mn(SO4)2 b Co(OH)3 c NiCO3 d Pt(NO3)2 e AuBr3 f Ga2(OH)2
6 Determine the charge on the negative ion (anion) for each of the following compounds.
a KBrO3 b Na2SiO3 c NH4AsO2 d Ag2SeO3 e LiTeO4 f AlIrCl6
7 Describe the structure of sodium chloride, with the help of a labelled diagram.
Observations of the physical properties of metals have led chemists to develop theories to explain these observations. For example, electrical conductivity requires the presence of charged particles that are free to move. Hardness and high melting points imply strong bonding. Malleability (the ability to be bent without breaking) and ductility (the ability to be drawn into a wire) suggest that there is regularity in the structure. From this information, chemists have devised a model to explain and represent the structure and bonding in metal elements in a simple way.
In this model, metal ions, formed when atoms lose their valence electrons, are arranged in a three-dimensional lattice. This array of ions is surrounded by freely moving electrons that form a ‘sea’ of mobile electrons. These electrons are said to be delocalized, as they are not confi ned to a particular location but can move throughout the structure. Electrons are attracted to positively
charged ions. This electrostatic attraction holds the lattice together, and prevents the ions pushing each other apart due to the electrostatic repulsion of like charges. This type of bonding is called metallic bonding.
Why do the metal atoms release their valence electrons to form the sea of electrons? Metal atoms achieve greater stability by releasing their valence electrons. Without their valence electrons, the metal atoms achieve a noble gas confi guration—an outer-shell octet of electrons. When non-metals are present, these valence electrons are transferred to the non-metal atoms, giving rise to the ionic
2.2 METALLIC BONDING
The nature of the metallic bond
4.4.1
Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons.
© IBO 2007
Figure 2.2.1 The structure of a metal may be described as a lattice of cations surrounded by a ‘sea’
of electrons.
CHAPTER 2 BONDING bonding discussed in section 2.1. When only metal atoms are present, the ‘lost’
valence electrons simply become delocalized within the metallic lattice.
The uses of metals by humans in the past as well as the present centre on two important properties of metals: their ability to conduct electricity and their malleability. You would be reading this book by candlelight if metals had not proved able to conduct electricity, and our bridges and buildings would be far less impressive than they are today.
The electrical conductivity of metals can be explained by the presence of the sea of delocalized electrons that surrounds the lattice of positive metal ions. In the solid state these electrons can move freely and will respond to the application of a potential difference. When a metal is connected to a power supply, electrons enter one end of the metal (the end connected to the negative terminal). The same number of electrons then exits from the other end of the metal (moving towards the positive terminal). Delocalized electrons move freely through the structure, but metal ions vibrate, causing a barrier to the smooth fl ow of electrons. Some energy is therefore lost, causing the metal to heat as current passes through it.
(e.g. from a battery) metal rod delocalized electrons
metal cations
Figure 2.2.2 Delocalized electrons can move freely through the metallic structure, conducting electricity.
The malleability of a metal is its ability to be beaten or bent into shape without breaking. Once again the sea of delocalized electrons is responsible for this property of metals. When a metal is bent, its lattice of positive ions is displaced and there is a possibility of positive ions coming into contact with other positive ions and repelling each other. The constant movement of the delocalized electrons prevents this from occurring, so the metal bends without breaking.
M+
electron sea electron sea flows between cations in their new positions
Figure 2.2.3 When a force is applied to the metal lattice, layers of metal ions are displaced, but the sea of electrons prevents repulsion and the subsequent breakage of bonds.
DEMO 2.2
Modelling the structures of metals Important properties of metals
4.4.2
Explain the electrical conductivity and malleability of metals. © IBO 2007
PRAC 2.2
Growing metal crystals
CHEM COMPLEMENT
Metals with a memory Metals are malleable and we can readily change their shapes.
Some metals can actually
‘remember’ their original shape and return to it under certain conditions! These ‘memory metals’ are alloys that have found use in a wide range of applications. One example, Nitinol, an alloy of nickel and titanium, is produced at high temperatures as a fine metal screen. It is then cooled and reshaped into a wire. This wire can be inserted into the bloodstream. As it warms in the body it ‘remembers’ its original shape and resumes the form of a fine metal screen. This screen can prevent movement of a blood clot through the circulatory system. Another type of memory metal is the brass used in safety devices and automatic switches.
The use of memory metals in orthodontic work make
attachment of wires to the teeth easy. The alloy used is flexible at room temperature. As it warms in the mouth, the wire becomes less flexible and resumes its original shape, pulling the teeth into position.
1 Describe what each of the following properties suggests about the nature of the bonding in metals.
a Metals conduct electricity.
b Most metals have high melting points.
c Metals are malleable.
2 Describe how the structure of a metal changes when a force is applied to the metal.
3 a When referring to the ‘electron sea’ model for metallic bonding, describe what is meant by the following terms:
i delocalized electrons ii lattice of cations.
b Explain why the positive ions in a metal do not repel each other.
4 In order for a substance to conduct electricity it must have charged particles that are free to move. State the name of the particles that are responsible for the conduction of electricity in a metal.
5 Use the ‘electron sea’ model for metallic bonding to explain each of the following observations.
a When you touch a piece of metal on a cold day it feels cold.
b An empty aluminium drink can be crushed without breaking the metal.
c Tungsten used in light globe fi laments has a melting point of over 3000°C.
6 In terms of the ‘electron sea’ model, explain why the electrical conductivity of metals decreases as the metal is warmed.
A chemical bond forms when outer-shell electrons come close enough to each other to interact and rearrange themselves into a more stable arrangement—
one with a lower chemical energy. This chemical energy is the sum of the chemical potential energy of the particles and their kinetic energy. The electrostatic attraction between the positively charged nuclei and the negatively charged electrons is the most signifi cant source of the chemical potential energy. As the two (or more) atoms approach one another, the positively charged nuclei repel one another, as do the negatively charged
electrons (see fi gure 2.3.1). These repulsion forces increase the potential energy of the system. At the same time, however, the oppositely charged particles are attracting one another, causing the potential energy to decrease. At some point, the distance separating the atoms or ions will be such that the repulsive forces of the particles with the same charge exactly balance the attractive forces between oppositely charged particles. It is in this stable arrangement that a chemical bond is formed (see fi gure 2.3.2).
Section 2.2 Exercises
THEORY OF KNOWLEDGE