DP Chemistry HL S2.1 The Ionic Model

DP Chemistry HL S2.1 The Ionic Model - Full notes

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DP Chemistry HL | S2.1 The Ionic Model

DP CHEMISTRY · STRUCTURE

S2.1

The Ionic Model

Full notes for HL · IB-style

Contents

1	Forming ions
2	Predicting ionic charges
3	Polyatomic ions and ionic formulas
4	Ionic bonding and lattice structure
5	Lattice enthalpy
6	Physical properties of ionic compounds
7	Predicting ion charges from the periodic table
8	Polyatomic ions
9	Writing formulas of ionic compounds
10	Electron transfer diagrams
11	The ionic lattice
12	Lattice enthalpy and the strength of ionic bonding
13	Physical properties of ionic compounds
14	Dissolving and hydration
15	Electronegativity and bond type
16	Naming ionic compounds
17	Ionic equations
18	Tests for ions
19	Ionic radius
20	Ionic compounds in everyday life
21	Mixed review: worked problems
22	Additional exam-style practice
23	Hydrated salts and water of crystallization
24	Common misconceptions and links
25	Exam-style question with markscheme
26	Check your understanding
27	Summary and topic map

These are full teaching notes. Read each section in order, try the Think about it question before you read on, and do every Quick check before moving to the next section.

Before you start

Rate yourself on each statement now, and again after you've worked through the notes. Shade R (not yet), A (getting there) or G (confident).

I can…	R	A	G
explain how and why atoms form ions
predict the charges of ions from the position of elements in the periodic table
write formulas and names for ionic compounds, including polyatomic ions
describe ionic bonding and the structure of ionic lattices
explain how ionic radius and charge affect lattice enthalpy
explain the volatility, electrical conductivity and solubility of ionic compounds

There is no additional HL content in this subtopic. SL and HL students study the same material.

Guiding questions

What determines the ionic nature and properties of a compound? How does the structure of an ionic lattice explain its properties?

How this topic is examined

S2.1 is examined through formula writing, explanations of ionic properties (melting point, conductivity, solubility) and comparisons of lattice enthalpies.

Explain properties in terms of strong electrostatic attractions between oppositely charged ions in a giant lattice.

Command term	What the examiner wants
State / Identify	a brief answer, no explanation
Outline / Describe	a brief or detailed account, without reasons
Explain	a detailed account with causes, reasons or mechanisms
Calculate / Determine	a numerical answer with working, correct significant figures and units
Deduce / Predict	reach a conclusion from given information
Draw / Sketch	a clear labelled diagram, structure or graph
Evaluate / Discuss	weigh strengths and limitations; reach a judgement

How to use these notes

Each section teaches one part of the syllabus in full. You'll see these boxes throughout; each has one job:

Box	What it's for
Think about it	a question to get you thinking before a section; the answer comes in the section
Key point	the one idea you must not forget
Common mistake	an error examiners see every year, and how to avoid it
In the real world	where the chemistry shows up outside the classroom
Try this	something quick you can do yourself to test the idea
Worked example	full solutions with IB-style marks (M1 method, A1 answer, R1 reasoning)
Quick check	short questions to test yourself at the end of each section

1 Forming ions

Atoms gain or lose electrons to achieve stable configurations.

Think about it

Sodium is a soft, reactive metal and chlorine is a toxic gas. Why does their compound, salt, behave so differently?

Atoms of the noble gases have full outer shells and are very unreactive. Other atoms tend to gain or lose electrons to achieve a stable electron configuration, often the same as a noble gas (the octet).

Metals, with few outer electrons and low ionization energies, lose electrons to form positive ions (cations): Na → Na+ + e−. Non-metals, with nearly full outer shells and high electron affinities, gain electrons to form negative ions (anions): Cl + e− → Cl−.

Ionic bonding typically forms when the difference in electronegativity between two elements is large (generally greater than about 1.8), as between a metal and a non-metal. Electrons are transferred from the metal to the non-metal.

Quick check

1. Why do metals form positive ions?

2. What determines whether bonding is ionic?

Answers: 1. they lose their few outer electrons easily (low ionization energy) 2. large electronegativity difference

2 Predicting ionic charges

The group number tells you the charge.

For main-group elements, the charge on the ion can be predicted from the group: group 1 forms 1+ ions; group 2, 2+; group 13, 3+ (Al3+); group 15, 3− (N3−); group 16, 2− (O2−, S2−); group 17, 1− (halides).

Transition elements can form ions with different charges (variable oxidation states), shown by Roman numerals: iron(II), Fe2+, and iron(III), Fe3+; copper(I), Cu+, and copper(II), Cu2+.

Group	1	2	13	15	16	17
Ion charge	1+	2+	3+	3−	2−	1−
Example	Na+	Mg2+	Al3+	N3−	O2−	Cl−

3 Polyatomic ions and ionic formulas

The overall charge must be zero.

Polyatomic ions are groups of covalently bonded atoms with an overall charge. You must know these:

Ion	Formula	Ion	Formula
ammonium	NH4+	hydroxide	OH−
nitrate	NO3−	hydrogencarbonate	HCO3−
carbonate	CO32−	sulfate	SO42−
phosphate	PO43−

To write the formula of an ionic compound, balance the charges so the total is zero. For aluminium oxide, Al3+ and O2−: the lowest common multiple of 3 and 2 is 6, so 2 Al3+ and 3 O2−: Al2O3. Brackets are used when more than one polyatomic ion is needed: calcium hydroxide is Ca(OH)2, ammonium sulfate is (NH4)2SO4.

Names: the cation is named first; the anion of a single element ends in -ide (chloride, oxide, nitride).

WORKED EXAMPLE 1 Writing formulas

Write formulas for magnesium nitride, iron(III) sulfate and ammonium phosphate.

Solution

Mg2+ and N3−: Mg3N2A1

Fe3+ and SO42−: Fe2(SO4)3A1

NH4+ and PO43−: (NH4)3PO4A1

Quick check

1. Write the formula of calcium nitrate.

2. Name Na2CO3.

Answers: 1. Ca(NO3)2 2. sodium carbonate

4 Ionic bonding and lattice structure

Oppositely charged ions in a giant three-dimensional array.

Ionic bonding is the electrostatic attraction between oppositely charged ions. It is non-directional: each ion attracts all oppositely charged ions around it.

Ionic compounds form a giant ionic lattice: a regular, repeating three-dimensional arrangement of cations and anions extending throughout the crystal. In sodium chloride, each Na+ is surrounded by six Cl− ions, and each Cl− by six Na+ (coordination number 6). The formula NaCl represents the ratio of ions, not a molecule; the term formula unit is used.

Figure 1. A layer of the sodium chloride lattice: each ion is surrounded by oppositely charged ions.

Quick check

1. Define ionic bonding.

2. What is the coordination number of Na+ in NaCl?

Answers: 1. electrostatic attraction between oppositely charged ions 2. 6

5 Lattice enthalpy

A measure of the strength of an ionic lattice.

The lattice enthalpy (ΔHlat) is the enthalpy change when one mole of a solid ionic compound is separated into its gaseous ions: NaCl(s) → Na+(g) + Cl−(g). It is always endothermic (positive), since energy is needed to overcome the attractions. A larger value means stronger ionic bonding.

Lattice enthalpy depends on two factors. Ionic charge: higher charges give stronger attraction and larger lattice enthalpy (MgO is much larger than NaCl). Ionic radius: smaller ions can approach more closely, giving stronger attraction (NaCl is larger than KCl). Larger lattice enthalpy generally means a higher melting point.

Compound	Ions	Lattice enthalpy / kJ mol−1	Melting point / °C
NaF	Na+, F−	+930	993
NaCl	Na+, Cl−	+790	801
KCl	K+, Cl−	+720	770
MgO	Mg2+, O2−	+3791	2852

WORKED EXAMPLE 2 Comparing lattice enthalpies

Explain why MgO has a much larger lattice enthalpy than NaCl.

Solution

Mg2+ and O2− have double the charges of Na+ and Cl−R1

and smaller ionic radiiR1

so the electrostatic attraction is much strongerR1

Quick check

1. Is lattice enthalpy exothermic or endothermic as defined?

2. Why is the lattice enthalpy of NaCl greater than KCl?

Answers: 1. endothermic 2. Na+ smaller than K+; ions closer; stronger attraction

6 Physical properties of ionic compounds

Structure explains properties.

High melting and boiling points (low volatility): a large amount of energy is needed to overcome the many strong electrostatic attractions throughout the lattice. Ionic compounds are solids at room temperature.

Electrical conductivity: ionic compounds do not conduct when solid, because the ions are fixed in the lattice and cannot move. They conduct when molten or dissolved in water, because the ions are free to move and carry charge.

Solubility: many ionic compounds dissolve in polar solvents such as water: polar water molecules surround ions (hydration) and the energy released helps overcome the lattice. They are generally insoluble in non-polar solvents such as hexane.

Brittleness: if a force shifts a layer of ions, ions of the same charge come next to each other and repel, so the crystal shatters.

Property	Explanation
High melting point	many strong electrostatic attractions to overcome
No conduction as solid	ions fixed in lattice
Conducts molten or aqueous	ions free to move
Soluble in water	ions hydrated by polar water molecules
Brittle	shifting layers brings like charges together; repulsion

Common mistake

Saying ionic compounds conduct because of free electrons. Conduction is by mobile ions, only when molten or dissolved.

Quick check

1. Why does solid NaCl not conduct electricity?

2. Why are ionic compounds brittle?

Answers: 1. ions cannot move 2. like charges repel when layers shift

7 Predicting ion charges from the periodic table

Group number tells the charge.

Group 1 metals form +1 ions, group 2 form +2, group 13 (Al) form +3. Group 15 non-metals form −3 ions, group 16 form −2 and group 17 form −1. Transition elements form more than one ion, shown with Roman numerals: iron(II), Fe2+, and iron(III), Fe3+.

WORKED EXAMPLE 3 Charges (basic)

Predict the ions formed by Li, Ba, N, Se and I.

Solution

Li+, Ba2+A1

N3−, Se2−, I−A1

Quick check

1. Predict the ion formed by strontium.

2. Name the ions Cu+ and Cu2+.

Answers: 1. Sr2+ 2. copper(I), copper(II)

8 Polyatomic ions

Groups of atoms with a charge.

Polyatomic ions stay together in reactions and formulas. Use brackets when more than one is needed: Ca(OH)2, (NH4)2SO4.

Ion	Formula	Ion	Formula
ammonium	NH4+	hydroxide	OH−
nitrate	NO3−	sulfate	SO42−
carbonate	CO32−	hydrogencarbonate	HCO3−
phosphate	PO43−	ethanoate	CH3COO−

WORKED EXAMPLE 4 Formulas (standard)

Write the formulas of aluminium sulfate and ammonium phosphate.

Solution

Al2(SO4)3A1

(NH4)3PO4A1

Quick check

1. Write the formula of calcium nitrate.

2. Write the formula of iron(III) hydroxide.

3. Name Mg(HCO3)2.

Answers: 1. Ca(NO3)2 2. Fe(OH)3 3. magnesium hydrogencarbonate

9 Writing formulas of ionic compounds

Balancing charges.

The total positive charge must equal the total negative charge. A quick method is to 'swap and drop' the charges, then simplify: Al3+ and O2− give Al2O3.

WORKED EXAMPLE 5 Formulas (basic)

Write formulas for magnesium nitride, potassium oxide and calcium phosphate.

Solution

Mg3N2; K2OA1 A1

Ca3(PO4)2A1

Common mistake

Writing CaCl for calcium chloride. Ca2+ needs two Cl− ions: CaCl2.

Quick check

1. Write the formula of sodium sulfide.

2. Write the formula of chromium(III) oxide.

3. Write the formula of barium hydroxide.

Answers: 1. Na2S 2. Cr2O3 3. Ba(OH)2

10 Electron transfer diagrams

Dot-and-cross for ionic compounds.

In ionic bonding, metal atoms transfer electrons to non-metal atoms. Draw the outer electrons of each atom (dots and crosses), show the transfer with arrows, and draw the ions in square brackets with their charges and complete outer shells.

WORKED EXAMPLE 6 Magnesium chloride (standard)

Describe the formation of MgCl2 in terms of electrons.

Solution

Mg (2,8,2) loses two electrons, one to each Cl (2,8,7)R1

Ions Mg2+ (2,8) and 2Cl− (2,8,8), held by electrostatic attractionR1

Quick check

1. Describe the formation of Na2O.

2. Describe the formation of aluminium fluoride.

Answers: 1. two Na atoms each lose one electron to one O atom: 2Na+ O2− 2. Al loses three electrons, one to each of three F atoms: Al3+ 3F−

11 The ionic lattice

A giant structure.

Ionic compounds form giant three-dimensional lattices of alternating positive and negative ions held together by strong electrostatic attraction in all directions. The formula shows the simplest ratio of ions, not a molecule. In NaCl, each Na+ is surrounded by six Cl− ions and vice versa (coordination number 6).

WORKED EXAMPLE 7 Lattice (standard)

Explain why sodium chloride has no molecular formula.

Solution

It is a giant lattice, not discrete moleculesR1

NaCl is the simplest ratio of ions (1 : 1)R1

Quick check

1. State the coordination number of each ion in the NaCl lattice.

Answers: 1. 6

12 Lattice enthalpy and the strength of ionic bonding

Charge and size.

Lattice enthalpy is the energy needed to separate one mole of an ionic solid into gaseous ions. It is larger for ions with higher charges and smaller radii, because the electrostatic attraction is stronger. This explains differences in melting point.

Compound	Lattice enthalpy / kJ mol−1	Melting point / °C
NaCl	790	801
KCl	720	770
MgO	3791	2852
CaO	3401	2613

WORKED EXAMPLE 8 Comparing (exam level)

Explain why MgO has a much higher melting point than NaCl.

Solution

Mg2+ and O2− have double the charges of Na+ and Cl−R1

Mg2+ and O2− are also smaller, so stronger electrostatic attractionR1

More energy needed to break the latticeR1

Quick check

1. Predict which has the higher lattice enthalpy: NaF or NaI. Explain.

Answers: 1. NaF: F− is smaller than I−

13 Physical properties of ionic compounds

Linking structure to properties.

Property	Explanation
high melting and boiling points	strong electrostatic forces throughout the lattice
conduct when molten or dissolved	ions are free to move
do not conduct when solid	ions are fixed in the lattice
brittle	a shift brings like charges together, which repel
often soluble in water	ions are attracted to polar water molecules (hydration)

WORKED EXAMPLE 9 Conductivity (standard)

Explain why solid sodium chloride does not conduct electricity but molten sodium chloride does.

Solution

Solid: ions fixed in position, cannot moveR1

Molten: ions free to move and carry chargeR1

Common mistake

Saying that ionic compounds conduct because 'electrons move'. In ionic compounds, the charge carriers are mobile ions.

Quick check

1. Explain why ionic crystals shatter when hit.

2. Explain why ionic compounds are generally insoluble in hexane.

Answers: 1. layers shift; like charges align and repel 2. non-polar solvent cannot stabilize ions

14 Dissolving and hydration

Ions surrounded by water.

When an ionic solid dissolves in water, polar water molecules surround the ions: the δ− oxygen atoms point towards cations and the δ+ hydrogen atoms towards anions. This hydration releases energy that helps overcome the lattice enthalpy.

WORKED EXAMPLE 10 Hydration (standard)

Describe how water molecules arrange themselves around Na+ and Cl− ions.

Solution

O atoms (δ−) face Na+A1

H atoms (δ+) face Cl−A1

Quick check

1. Why does silver chloride not dissolve well in water although it is ionic?

Answers: 1. its lattice energy is larger than the energy released by hydration

15 Electronegativity and bond type

Ionic or covalent?

The greater the difference in electronegativity between two elements, the more ionic the bond. A difference above about 1.8 is generally considered predominantly ionic. The bonding triangle (S2.4) places compounds on a continuum between ionic, covalent and metallic.

WORKED EXAMPLE 11 Predicting (standard)

Use electronegativities (Na 0.9, Cl 3.2, C 2.6, H 2.2) to classify bonds in NaCl and CH4.

Solution

NaCl: Δ = 2.3: ionicA1

C–H: Δ = 0.4: covalent (essentially non-polar)A1

Quick check

1. Classify the bond in AlCl3 (Al 1.6, Cl 3.2).

Answers: 1. Δ = 1.6: polar covalent, with partial ionic character

16 Naming ionic compounds

Systematic names.

Name the cation first, then the anion. Monatomic anions end in -ide (chloride, oxide, nitride). Oxyanions usually end in -ate. Use Roman numerals for metals with variable charge: iron(III) chloride, FeCl3.

WORKED EXAMPLE 12 Naming (basic)

Name CuO, Fe2(SO4)3 and NH4NO3.

Solution

copper(II) oxideA1

iron(III) sulfate; ammonium nitrateA1

Quick check

1. Name MnO2.

2. Name Pb(NO3)2.

Answers: 1. manganese(IV) oxide 2. lead(II) nitrate

17 Ionic equations

Showing only the ions that react.

In precipitation and acid–base reactions, spectator ions do not change and are omitted from the ionic equation. For example, adding silver nitrate to sodium chloride: Ag+(aq) + Cl−(aq) → AgCl(s).

WORKED EXAMPLE 13 Precipitation (standard)

Write the ionic equation for the reaction of barium chloride solution with sodium sulfate solution.

Solution

Ba2+(aq) + SO42−(aq) → BaSO4(s)A2

Quick check

1. Write the ionic equation for the neutralization of HCl by NaOH.

2. Identify the spectator ions when copper(II) sulfate reacts with sodium hydroxide.

Answers: 1. H+(aq) + OH−(aq) → H2O(l) 2. Na+ and SO42−

18 Tests for ions

Identifying ions in solution.

Ion	Test	Result
Cl−	HNO3 then AgNO3(aq)	white precipitate
Br−	HNO3 then AgNO3(aq)	cream precipitate
I−	HNO3 then AgNO3(aq)	yellow precipitate
SO42−	HCl then BaCl2(aq)	white precipitate
CO32−	add dilute acid	fizzing; gas turns limewater milky
NH4+	warm with NaOH(aq)	gas turns damp red litmus blue

WORKED EXAMPLE 14 Identifying (standard)

A white solid gives a lilac flame and a cream precipitate with acidified silver nitrate. Identify it.

Solution

Lilac flame: K+A1

Cream precipitate: Br−: KBrA1

19 Ionic radius

Ions are different sizes from atoms.

Positive ions are smaller than their atoms: they lose the outer shell, and the remaining electrons are pulled in more strongly. Negative ions are larger than their atoms: extra electrons increase repulsion. In an isoelectronic series (same electron number), radius decreases as nuclear charge increases: O2− > F− > Na+ > Mg2+.

WORKED EXAMPLE 15 Isoelectronic series (exam level)

Explain why Na+ is smaller than F− although both have 10 electrons.

Solution

Na+ has 11 protons, F− has 9R1

Greater nuclear charge pulls the same electrons closerR1

Quick check

1. Arrange in order of increasing radius: Cl−, K+, S2−, Ca2+.

Answers: 1. Ca2+ < K+ < Cl− < S2−

20 Ionic compounds in everyday life

From salt to ceramics.

Ionic compounds include table salt, limestone (CaCO3), gypsum (CaSO4·2H2O), many minerals and ceramics. Their high melting points make them useful as refractory materials (MgO lines furnaces), and their ions are essential nutrients (Na+, K+, Ca2+).

In the real world

Lithium-ion batteries rely on the movement of Li+ ions between electrodes; the ionic conductivity of the electrolyte limits how fast the battery can charge.

Quick check

1. Why is magnesium oxide used to line furnaces?

Answers: 1. very high melting point due to strong ionic bonding

21 Mixed review: worked problems

Bringing the ideas together.

WORKED EXAMPLE 16 Mixed problem (exam level)

Element X is in group 2 and Y in group 15. (a) Write the formula of the compound. (b) Give the electron configurations of the ions if X is calcium and Y nitrogen. (c) Predict two physical properties.

Solution

(a) X3Y2A1

(b) Ca2+ [Ar]; N3− [Ne]A1

Quick check

1. How many ions are in one formula unit of Al2(SO4)3?

Answers: 1. 5

22 Additional exam-style practice

Paper 1 and Paper 2.

WORKED EXAMPLE 17 Paper 1 question (standard)

Which compound has the highest lattice enthalpy? A. NaCl B. KBr C. MgO D. CaS

Solution

C: small, doubly charged ionsA1

WORKED EXAMPLE 18 Paper 2 question (exam level)

(a) Explain why the melting point of CaO (2613 °C) is higher than that of KCl (770 °C). (b) State the conditions under which CaO conducts electricity. (c) Write the formula of calcium nitride.

Solution

(a) higher ionic charges (2+, 2−) and smaller ions: stronger attractionR1 R1

(b) molten (ions mobile)A1

Quick check

1. State why ionic compounds have no overall charge.

Answers: 1. the total positive and negative charges are equal

23 Hydrated salts and water of crystallization

Water inside the lattice.

Many ionic crystals contain water molecules within the lattice, such as CuSO4·5H2O (blue). Heating removes the water, leaving the anhydrous salt (white CuSO4); adding water reverses the change and is used as a test for water.

WORKED EXAMPLE 19 Test for water (standard)

Describe the test for water using anhydrous copper(II) sulfate.

Solution

Add the liquid to white anhydrous CuSO4R1

Turns blue if water is presentA1

24 Common misconceptions and links

Students often think…	Actually…
NaCl exists as molecules	it is a giant lattice; NaCl is a formula unit
ionic solids conduct	only molten or aqueous
ionic bonds are between pairs of ions	each ion attracts all neighbouring opposite ions

How S2.1 connects to the rest of the course

Topic	Connection
S1.3 Electron configurations	ion formation
S2.2 Covalent model	electronegativity and bond type
S3.1 Periodic table	group and ionic charge
R1.2 Energy cycles (HL)	Born-Haber cycles

25 Exam-style question with markscheme

Try this under exam conditions (about 15 minutes) before reading the markscheme.

Figure 2. Lattice enthalpies of four ionic compounds.

(a) Write the formula of aluminium sulfate. [1]

(b) Explain the difference in lattice enthalpy between NaF and NaCl. [2]

(d) Explain the electrical conductivity of NaCl in different states. [3]

Markscheme

Part	Answer	Marks
(a)	Al2(SO4)3	A1
(b)	F− smaller than Cl−	R1
	ions closer, stronger attraction	R1
(c)	giant lattice with 2+ and 2− ions	R1
	very strong electrostatic attractions need much energy	R1
(d)	solid: ions fixed, no conduction	R1
	molten: ions mobile, conducts	R1
	aqueous: ions mobile, conducts	R1

Examiner's comments. In (d), answers mentioning electrons scored zero for the explanation.

26 Check your understanding

Section A: multiple choice

1. A group 16 element forms ions with charge

A. 2− B. 2+ C. 6+ D. 1−

2. The formula of calcium phosphate is

A. CaPO4 B. Ca2(PO4)3 C. Ca3(PO4)2 D. Ca(PO4)2

3. Ionic compounds conduct when

A. solid B. molten or dissolved C. never D. only solid

4. Lattice enthalpy increases with

A. larger ions B. lower charges C. higher temperature D. higher ionic charges and smaller ions

5. The ammonium ion is

A. NH3 B. NH4+ C. NO3− D. NH2−

6. Ionic compounds are brittle because

A. like charges repel when layers shift B. they contain electrons C. bonds are weak D. they are molecules

Section B: structured

7. Describe the structure of sodium chloride. [3]

8. Write formulas for potassium nitride and copper(II) hydroxide. [2]

9. Explain why ionic compounds are soluble in water but not hexane. [3]

Answers

1. A 2. C 3. B 4. D 5. B 6. A

7. Giant lattice; alternating Na+ and Cl−; each surrounded by six of the opposite.

8. K3N; Cu(OH)2.

9. Polar water hydrates ions, energy released; hexane non-polar, cannot interact with ions.

27 Summary and topic map

Idea	Remember
Ion formation	metals lose, non-metals gain electrons
Charges	group 1: 1+, 2: 2+, 13: 3+, 15: 3−, 16: 2−, 17: 1−
Formulas	balance charges; polyatomic ions in brackets
Ionic bonding	electrostatic attraction; giant lattice
Lattice enthalpy	higher charge, smaller ions: larger
Properties	high mp, conduct molten/aqueous, soluble in water, brittle

Before the exam, make sure you…

•balance charges for formulas
•explain properties using lattice and mobile ions
•compare lattice enthalpies with charge and radius
•learn polyatomic ions

The Ionic Model on one page

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Figure 3. Topic map for The Ionic Model.

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