DP Chemistry HL S2.3 The Metallic Model

DP Chemistry HL S2.3 The Metallic Model - Full notes

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DP Chemistry HL | S2.3 The Metallic Model

DP CHEMISTRY · STRUCTURE

S2.3

The Metallic Model

Full notes for HL · IB-style

Contents

1	Metallic bonding
2	Physical properties of metals
3	Factors affecting the strength of metallic bonding
4	Trends in melting point across a period and down a group
5	Transition elements and metallic bonding
6	Metals in use
7	Describing metallic bonding precisely
8	Factors affecting the strength of metallic bonding
9	Electrical and thermal conductivity
10	Malleability and ductility
11	Alloys: structure and properties
12	Density and appearance of metals
13	Trends in metallic character
14	Transition metals: bonding and properties
15	Metals and non-metals compared
16	Extracting and recycling metals
17	Modelling metals: limitations
18	Mixed review: worked problems
19	Additional exam-style practice
20	Practical: comparing metals
21	Metallic bonding and the bonding continuum
22	Sustainability of metals
23	Common misconceptions and links
24	Exam-style question with markscheme
25	Check your understanding
26	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
describe metallic bonding as attraction between cations and delocalized electrons
explain electrical and thermal conductivity, malleability and ductility of metals
explain how charge and ionic radius affect the strength of metallic bonding
explain trends in melting points of metals across period 3 and down group 1
explain why transition elements have high melting points and strong metallic bonding
relate the properties of metals to their uses

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

Guiding questions

What determines the metallic properties of an element? How does the metallic model explain the behaviour of metals?

How this topic is examined

S2.3 is a short subtopic examined through explanations of metallic properties and trends in melting points, often combined with ionic and covalent bonding questions.

Use the phrase 'electrostatic attraction between a lattice of cations and delocalized electrons' in every definition.

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 Metallic bonding

Cations in a sea of electrons.

Think about it

A copper wire bends without breaking and carries electricity across a whole city. What is it about copper's structure that allows both?

Metal atoms have few valence electrons, held relatively weakly (low ionization energies). In a metal, the valence electrons are delocalized: they are not attached to any one atom but are free to move throughout the structure. The atoms become positive ions (cations) arranged in a regular lattice.

Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a 'sea' of delocalized electrons. It is non-directional: the electrons hold the whole lattice together.

Figure 1. Metallic bonding: a lattice of cations (+) surrounded by delocalized electrons (red dots).

Quick check

1. Define metallic bonding.

2. What is meant by delocalized electrons?

Answers: 1. attraction between a lattice of cations and delocalized electrons 2. electrons free to move throughout the structure, not attached to one atom

2 Physical properties of metals

The model explains each property.

Electrical conductivity: delocalized electrons are free to move; when a potential difference is applied, they flow through the metal, carrying charge. Metals conduct in both solid and liquid states.

Thermal conductivity: delocalized electrons transfer kinetic energy rapidly through the metal, and closely packed ions pass on vibrations.

Malleability and ductility: metals can be hammered into shape (malleable) and drawn into wires (ductile). When a force is applied, layers of cations slide over each other; because the delocalized electrons move with them, the attraction is maintained in the new position and the metal does not break. (Contrast ionic solids, where sliding brings like charges together.)

High melting points (for most metals): strong metallic bonding requires much energy to overcome. Lustre: delocalized electrons reflect light, so freshly cut metals are shiny.

Property	Explanation
Conducts electricity	delocalized electrons move when voltage applied
Conducts heat	electrons transfer kinetic energy; ions vibrate
Malleable and ductile	layers of ions slide; electrons keep holding them together
High melting point	strong attraction between ions and electrons
Lustrous	delocalized electrons reflect light

Common mistake

Saying metals conduct because ions move. In solid metals, the ions are fixed in the lattice; only the delocalized electrons move.

Quick check

1. Why are metals malleable?

2. Why do metals conduct electricity when solid?

Answers: 1. layers slide without breaking metallic bonding 2. delocalized electrons are mobile

3 Factors affecting the strength of metallic bonding

Charge and size.

The strength of metallic bonding depends on: the charge on the cations (number of delocalized electrons per atom): more electrons and higher charge give stronger attraction; and the radius of the cations: smaller ions have a higher charge density and the delocalized electrons are closer to the nucleus, giving stronger attraction.

Stronger metallic bonding generally means a higher melting point, greater hardness and greater density.

4 Trends in melting point across a period and down a group

More electrons and smaller ions mean stronger bonding.

Across period 3 (Na, Mg, Al): the melting point increases, because the charge on the ion increases (Na+, Mg2+, Al3+), so there are more delocalized electrons per atom, and the ionic radius decreases. The attraction between ions and delocalized electrons is stronger.

Down group 1 (Li to Cs): the melting point decreases, because the ionic radius increases while the charge stays at 1+. The delocalized electrons are further from the nucleus, so the attraction is weaker. Caesium melts at only 28 °C.

Figure 2. Melting points of group 1 metals decrease down the group.

Figure 3. Melting points increase from Na to Al.

WORKED EXAMPLE 1 Explaining a trend

Explain why magnesium has a higher melting point than sodium.

Solution

Mg forms Mg2+ ions, Na forms Na+ ions: Mg has two delocalized electrons per atom, Na oneR1

Mg2+ is smaller than Na+R1

so the electrostatic attraction between cations and delocalized electrons is stronger in MgR1

Quick check

1. Why does melting point decrease down group 1?

2. Why does aluminium have a higher melting point than sodium?

Answers: 1. larger ions; delocalized electrons further away; weaker attraction 2. 3+ charge, more delocalized electrons, smaller ion

5 Transition elements and metallic bonding

Many delocalized electrons make strong metals.

Transition elements (such as iron, chromium, tungsten) generally have high melting points, high densities and great strength. This is because both the 4s and 3d electrons can be delocalized, giving a large number of delocalized electrons per atom and strong metallic bonding. Tungsten has the highest melting point of any metal (3422 °C), which is why it was used for light bulb filaments.

Transition metals also show variable oxidation states, form coloured compounds and complex ions, and act as catalysts (S3.1).

6 Metals in use

Properties determine uses.

Metal	Property	Use
Copper	excellent electrical conductor, ductile	electrical wiring
Aluminium	low density, good conductor, resists corrosion	aircraft, overhead power cables, cans
Iron (as steel)	strong, cheap	construction, vehicles
Gold	unreactive, lustrous, good conductor	jewellery, electrical contacts
Tungsten	very high melting point	filaments, cutting tools

Pure metals are often too soft for many uses; mixing them with other elements to form alloys makes them harder (S2.4).

7 Describing metallic bonding precisely

Cations in a sea of electrons.

Metallic bonding is the electrostatic attraction between a lattice of positive metal ions (cations) and delocalized electrons. The valence electrons are not attached to any one atom and move freely throughout the structure. The bonding is non-directional.

WORKED EXAMPLE 2 Definition (basic)

Define metallic bonding.

Solution

Electrostatic attraction between a lattice of cationsR1

and a sea of delocalized electronsR1

Common mistake

Describing metals as made of 'atoms' surrounded by electrons, or of 'positive and negative ions'. Metals contain positive ions and delocalized electrons; there are no negative ions.

Quick check

1. Why is metallic bonding described as non-directional?

Answers: 1. the delocalized electrons attract all surrounding cations equally in all directions

8 Factors affecting the strength of metallic bonding

Charge, size, number of electrons.

Metallic bonding is stronger when the cations have a higher charge (more delocalized electrons per ion) and a smaller radius (delocalized electrons closer to the nucleus). This shows in melting points.

Metal	Ion	Ionic radius / pm	Melting point / °C
Na	Na+	102	98
Mg	Mg2+	72	650
Al	Al3+	54	660
K	K+	138	63

WORKED EXAMPLE 3 Comparing (exam level)

Explain why magnesium has a higher melting point than sodium.

Solution

Mg2+ has a higher charge than Na+ and contributes two delocalized electronsR1

Mg2+ is smallerR1

Stronger attraction between cations and delocalized electronsR1

Quick check

1. Explain the trend in melting points down group 1.

2. Predict which has the higher melting point: Ca or K.

Answers: 1. decrease: larger ions, weaker attraction 2. Ca (2+ ion, smaller)

9 Electrical and thermal conductivity

Mobile electrons.

Metals conduct electricity because delocalized electrons can move through the lattice when a potential difference is applied. They conduct heat well because the mobile electrons transfer kinetic energy rapidly, and because the closely packed ions pass vibrations along.

WORKED EXAMPLE 4 Conductivity (standard)

Explain why aluminium is used for overhead power cables.

Solution

Good conductor: delocalized electronsR1

Low density (light) and resistant to corrosionR1

Quick check

1. Why does copper conduct heat well?

Answers: 1. delocalized electrons carry kinetic energy; closely packed ions transmit vibrations

10 Malleability and ductility

Layers slide without breaking bonds.

When a force is applied, layers of cations slide over each other. The delocalized electrons move with them, so the attraction is maintained in the new position: the metal changes shape without breaking. Ionic solids, in contrast, shatter because like charges come together.

WORKED EXAMPLE 5 Comparing (exam level)

Explain why copper can be drawn into wires but sodium chloride cannot.

Solution

In copper, layers of cations slide; delocalized electrons maintain bondingR1

In NaCl, a shift brings ions of like charge together, which repel and the crystal shattersR1

Quick check

1. Define malleable and ductile.

Answers: 1. can be hammered into shape; can be drawn into wires

11 Alloys: structure and properties

Mixing metal atoms.

An alloy is a mixture of a metal with other metals or non-metals. Atoms of different sizes disrupt the regular layers, making it harder for layers to slide, so alloys are usually harder and stronger than the pure metal, and often more resistant to corrosion.

Alloy	Composition	Property / use
steel	iron + carbon	strong; construction
stainless steel	iron + chromium + nickel	corrosion resistant; cutlery
brass	copper + zinc	hard, attractive; fittings, instruments
bronze	copper + tin	hard, resists corrosion; statues
solder	tin + others	low melting point; joining

WORKED EXAMPLE 6 Alloy hardness (standard)

Explain why brass is harder than copper.

Solution

Zinc atoms are a different size from copper atomsR1

They distort the layers, so layers cannot slide easilyR1

Quick check

1. Why is pure gold rarely used for jewellery?

2. Why is stainless steel resistant to rusting?

Answers: 1. too soft; alloying makes it harder 2. chromium forms a protective oxide layer

12 Density and appearance of metals

Closely packed ions.

Metals are generally dense because their ions are closely packed. Their shiny lustre arises because delocalized electrons absorb and re-emit light across the visible spectrum.

WORKED EXAMPLE 7 Lustre (standard)

Explain why freshly cut sodium is shiny but quickly becomes dull.

Solution

Delocalized electrons reflect lightR1

Sodium reacts with oxygen to form a dull oxide layerR1

In the real world

Titanium alloys combine low density with high strength and corrosion resistance, making them ideal for aircraft parts, artificial joints and high-performance bicycles.

13 Trends in metallic character

Across periods and down groups.

Metallic character (the tendency to lose electrons) decreases across a period and increases down a group. Metalloids such as silicon and germanium lie along the boundary and have intermediate properties, such as semiconductivity.

WORKED EXAMPLE 8 Trend (standard)

Explain why metallic character increases down group 14 (C to Pb).

Solution

Outer electrons further from the nucleus and more shieldedR1

Lost more easilyR1

Quick check

1. Name a metalloid and give a use.

Answers: 1. silicon: semiconductors in computer chips

14 Transition metals: bonding and properties

d-electrons join the sea.

Transition metals have high melting points and high densities because both their 4s and 3d electrons can become delocalized, giving strong metallic bonding. They are often hard and strong, and many are good catalysts.

WORKED EXAMPLE 9 Melting point (exam level)

Explain why iron (1538 °C) has a much higher melting point than calcium (842 °C).

Solution

Iron delocalizes both 4s and some 3d electronsR1

More delocalized electrons and smaller ions: stronger metallic bondingR1

Quick check

1. Why are transition metals generally denser than group 1 metals?

Answers: 1. smaller atoms with greater mass, closely packed

15 Metals and non-metals compared

Contrasting bonding.

Property	Metals	Non-metals
bonding	metallic	covalent
conductivity	good	poor (except graphite)
oxides	basic	acidic
ions	positive	negative
solid state	malleable, ductile	brittle (if solid)

WORKED EXAMPLE 10 Classifying (standard)

Element Z has a high melting point, conducts in the solid state and forms a basic oxide. Classify Z.

Solution

MetalA1

16 Extracting and recycling metals

Linking reactivity and bonding.

Metals are extracted from ores by reduction. Less reactive metals (iron, zinc) are reduced with carbon; more reactive metals (aluminium) require electrolysis. Recycling metals saves energy and resources: recycling aluminium uses about 5% of the energy needed to extract it from bauxite.

In the real world

About 75% of all aluminium ever produced is still in use today because it can be recycled repeatedly without losing quality.

Quick check

1. Why is aluminium extracted by electrolysis rather than reduction with carbon?

Answers: 1. aluminium is more reactive than carbon

17 Modelling metals: limitations

What the sea model cannot explain.

The electron-sea model explains conductivity, malleability and lustre, but not why some metals are magnetic, why conductivity decreases with temperature in detail, or the different crystal structures of metals. More advanced band theory is needed.

WORKED EXAMPLE 11 Evaluating (standard)

State one strength and one limitation of the electron-sea model.

Solution

Explains electrical conductivityR1

Does not explain magnetism (or melting-point differences between transition metals)R1

18 Mixed review: worked problems

Bringing the ideas together.

WORKED EXAMPLE 12 Mixed problem (exam level)

Aluminium melts at 660 °C and sodium at 98 °C. (a) Describe the bonding in aluminium. (b) Explain the difference in melting points. (c) Explain why both conduct electricity.

Solution

(a) lattice of Al3+ ions in a sea of delocalized electronsA1

(b) Al3+ higher charge and smaller radius; three delocalized electrons per atomR1 R1

Quick check

1. Explain why mercury is a liquid at room temperature while tungsten melts at 3422 °C (qualitatively).

Answers: 1. metallic bonding much weaker in Hg; W has many delocalized d-electrons

19 Additional exam-style practice

Paper 1 and Paper 2.

WORKED EXAMPLE 13 Paper 1 question (standard)

Which explains why metals are malleable? A. strong ionic bonds B. layers of cations slide while delocalized electrons maintain bonding C. metals have low melting points D. electrons are fixed in place

Solution

BA1

WORKED EXAMPLE 14 Paper 2 question (exam level)

Steel is an alloy of iron and carbon. (a) Explain why steel is harder than pure iron. (b) Suggest why steel still conducts electricity.

Solution

(a) carbon atoms disrupt the regular layers, preventing slidingR1

(b) delocalized electrons are still presentR1

Quick check

1. Explain why potassium is softer than sodium.

Answers: 1. larger K+ ions: weaker metallic bonding

20 Practical: comparing metals

Hardness and conductivity tests.

Simple experiments compare metals: scratch tests for hardness, measuring resistance of wires of the same length and thickness, or timing heat transfer along rods. Control variables carefully (same dimensions, same temperature).

WORKED EXAMPLE 15 Fair test (standard)

List two variables to control when comparing the electrical resistance of metal wires.

Solution

Length and cross-sectional area (diameter)A1

TemperatureA1

Try this

Hold one end of copper, iron and aluminium rods of equal size in hot water and feel which end warms first. Relate your observation to delocalized electrons.

21 Metallic bonding and the bonding continuum

Linking to S2.4.

Metallic bonding sits at one corner of the bonding triangle (S2.4). Compounds of two metals are metallic (alloys); a metal and a non-metal with a large electronegativity difference give ionic bonding; two non-metals give covalent bonding. Many substances have intermediate character.

WORKED EXAMPLE 16 Placing on the triangle (standard)

Where would an alloy of copper and zinc appear on the bonding triangle?

Solution

Near the metallic corner: low average electronegativity, small differenceR1

22 Sustainability of metals

Resources and reuse.

Many metals are finite resources. Mining has environmental impacts (habitat loss, energy use, waste). Strategies include recycling, using alternative materials, designing products for disassembly, and reducing use of scarce metals such as cobalt and indium.

Quick check

1. Give two benefits of recycling copper.

Answers: 1. saves energy; conserves ore and reduces mining waste

23 Common misconceptions and links

Students often think…	Actually…
metals conduct by moving ions	delocalized electrons move
metallic bonds are between pairs of atoms	attraction is between all cations and the electron sea
all metals have high melting points	group 1 metals have low melting points

How S2.3 connects to the rest of the course

Topic	Connection
S2.1 Ionic model	compare brittleness and conductivity
S2.4 Materials	alloys
S3.1 Periodic table	metallic character and transition elements
S1.3 Electron configurations	4s and 3d electrons

24 Exam-style question with markscheme

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

Figure 4. Melting points of group 1 metals.

(a) Describe the trend shown. [1]

(b) Explain this trend. [3]

(d) Explain why copper is used for electrical wiring. [2]

Markscheme

Part	Answer	Marks
(a)	melting point decreases down the group	A1
(b)	ionic radius increases	R1
	charge constant (1+)	R1
	weaker attraction between cations and delocalized electrons	R1
(c)	layers of cations slide over each other	R1
	delocalized electrons maintain bonding	R1
(d)	excellent conductor: delocalized electrons	R1
	ductile: can be drawn into wires	R1

Examiner's comments. In (b), the best answers mentioned both radius and constant charge.

25 Check your understanding

Section A: multiple choice

1. Metallic bonding is attraction between

A. cations and delocalized electrons B. anions and cations C. shared pairs and nuclei D. molecules

2. Metals conduct electricity because

A. ions move B. they are hard C. delocalized electrons move D. they are dense

3. Melting point increases from Na to Al because

A. atoms get bigger B. charge increases and ionic radius decreases C. fewer electrons D. bonding becomes covalent

4. Down group 1 melting point

A. increases B. stays the same C. rises then falls D. decreases

5. Transition metals have high melting points because

A. they are ionic B. 4s and 3d electrons are delocalized C. they have small atoms only D. they are gases

6. Metals are malleable because

A. layers of ions slide B. electrons are fixed C. bonds are directional D. ions repel

Section B: structured

7. Describe metallic bonding. [2]

8. Compare the electrical conductivity of metals and ionic compounds. [3]

9. Explain why aluminium is used for overhead power cables. [2]

Answers

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

7. Lattice of cations; delocalized electrons; electrostatic attraction.

8. Metals conduct solid and liquid via electrons; ionic only molten/aqueous via ions.

9. Good conductor (delocalized electrons); low density; resists corrosion.

26 Summary and topic map

Idea	Remember
Metallic bonding	cations + delocalized electrons
Properties	conductivity, malleability, ductility, lustre
Bond strength	higher charge, smaller ions: stronger
Period 3	Na < Mg < Al melting point
Group 1	melting point decreases down group
Transition metals	many delocalized electrons; high mp

Before the exam, make sure you…

•define metallic bonding precisely
•explain properties with mobile electrons
•use charge and radius for trends
•link properties to uses

The Metallic Model on one page

Cover this page with a sheet of paper and try to rebuild it from memory. Whatever you can't recall is what to revise next.

Figure 5. Topic map for The Metallic Model.

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