S2.3 Metallic Bonding - Retake

1. What is metallic bonding?

• Definition: Metallic bonding is the electrostatic attraction between a lattice of positive metal ions (cations) and a sea of delocalised electrons.

• Key points:

  • Outer (valence) electrons are removed from individual metal atoms and become delocalised.

  • These electrons are free to move throughout the entire structure, not tied to any one nucleus.

  • Positive metal ions stay fixed in position in a giant, regular, close-packed lattice.

  • Attraction is non-directional — the electrons are equally attracted to all neighbouring cations, so the bonding doesn’t depend on a fixed atom-to-atom alignment.

  • The result is a very strong, cohesive structure that explains many metallic properties.

2. How does metallic bonding differ from ionic bonding?

• Both involve electrostatic forces, but:

  • Ionic: between different species — cations and anions.

  • Metallic: between same species — metal cations and delocalised electrons.

• In metallic bonding, all atoms in the pure element are the same type, and all share their outer electrons into one communal electron sea.

• Metallic bonding’s non-directional nature allows atoms to move past one another without the bonding being destroyed, unlike ionic lattices, which shatter if layers shift.

3. Why do metals form metallic bonds?

• Metals have low ionisation energies — their outer electrons are relatively easy to remove.

• Removing electrons costs energy, but this is compensated by:

  • The large release of energy when strong cation–electron attractions form.

  • The stability of the delocalised electron arrangement.

• Close-packed arrangements maximise the overlap between electron cloud and cation cores, strengthening the bonding.

4. What is the structure of a metallic lattice?

• Cations arranged in close-packed or layered arrangements:

  • Hexagonal close-packed (hcp) or face-centred cubic (fcc) structures are common.

• Delocalised electrons move freely between cations, filling the spaces in the lattice.

• Positive cations repel each other, so the lattice is spaced evenly — the electron sea counteracts this repulsion, “gluing” the structure together.

5. Why are metals malleable and ductile?

• Malleable: hammered into thin sheets.

• Ductile: drawn into wires.

• When a force is applied, layers of cations slide over one another:

  • Delocalised electrons instantly redistribute to maintain bonding.

  • No large-scale build-up of repulsion occurs because the electron sea flows into the new positions.

• The lattice is rearranged but not broken — this is unique to non-directional bonding.

6. Why are metals strong and hard?

• Cations are closely packed and surrounded by a high density of electrons.

• Strong electrostatic attractions act in all directions, resisting attempts to pull the structure apart.

• Significant forces are needed to deform or fracture the lattice.

7. Why can metals conduct electricity in both solid and liquid states?

• Solid: electrons are free to move through the fixed cation lattice.

• Liquid: cations move too, but the electron sea still allows current to flow.

• Applying a voltage:

  • Electrons are repelled from the negative terminal and attracted to the positive.

• Across a period: more valence electrons → higher delocalised electron density → greater conductivity (Na → Mg → Al).

8. Why are metals good thermal conductors?

• Two pathways for heat transfer:

  1. Vibrating cations transfer kinetic energy to neighbours through collisions.

  2. Delocalised electrons absorb kinetic energy and travel through the lattice, passing energy on rapidly.

• This dual mechanism makes metals excellent conductors of heat compared to other solids.

9. Why do metals have high melting and boiling points?

• Breaking metallic bonds means overcoming strong cation–electron attractions throughout the giant lattice.

• Across a period: cation charge increases, radius decreases, more delocalised electrons — all strengthen bonding, raising melting points.

• Down a group: larger radii weaken attractions, lowering melting points.

10. How do metallic properties influence uses?

• Malleability/ductility → shaping into wires (Cu) or sheets (Al foil).

• High conductivity → electrical wiring (Cu, Al).

• Corrosion resistance → packaging (Al), cutlery (stainless steel).

• Strength → building frameworks (steel).

• Thermal conductivity → cookware (Cu, Al).

• Choice of metal depends on the property–use match, cost, and reactivity.

11. What factors determine metallic bond strength?

1. Charge on cations — higher charge means stronger electrostatic pull on electrons.

2. Number of delocalised electrons — more electrons increase electron density.

3. Ionic radius — smaller cations have stronger attractions because electrons are closer to the nucleus.

12. What are the melting point trends in s- and p-block metals?

• Across a period: Na (Group 1, +1) → Mg (+2) → Al (+3): charge ↑, size ↓, delocalised electrons ↑ → stronger bonds → higher melting points.

• Down a group: radius ↑, attraction ↓ → melting point decreases.

13. Why do transition metals have higher melting points than s-block metals?

• Transition metals can delocalise d-electrons as well as s-electrons, increasing electron density.

• This creates stronger cation–electron attraction, so more energy is needed to melt them.

• All Period 4 transition metals melt at higher temperatures than Na or Mg.

14. Why are transition metals good conductors?

• Large number of delocalised electrons from both s and d orbitals.

• Electrons are mobile and less prone to scattering due to the regular, close-packed lattice.

• Silver, copper, and gold are the most conductive; copper is favoured in industry for its conductivity–cost balance.

15. Why is beryllium an exception among s-block metals?

• Be²⁺ has a small radius and high +2 charge, creating a very high charge density.

• Strong metallic bonding gives it a higher melting point than expected for Group 2.

• Its bonding strength rivals some transition metals.

16. How do experiments demonstrate metallic properties?

• Conductivity tests: metals conduct when solid — unlike ionic solids.

• Hammer/malleability test: metals bend rather than shatter.

• Thermal conduction test: heating one end of a metal rod quickly warms the other end.

• Density: generally high due to close-packed lattices.

17. How does metallic bonding allow alloy formation?

• Non-directional bonding means adding other atoms doesn’t disrupt cation–electron attractions.

• Substitutional alloy: atoms of similar size replace each other in the lattice (e.g., brass = Cu + Zn).

• Interstitial alloy: smaller atoms fit into gaps between larger cations (e.g., steel = Fe + C).

• Alloys can be stronger or more corrosion-resistant because lattice distortion reduces layer sliding.

18. What factors affect transition metal bonding strength?

• More delocalised electrons (from d and s orbitals) → higher electron density.

• Higher cation charges → stronger attractions.

• Smaller radii → shorter electron–nucleus distances, stronger bonds.

19. How does metallic bonding vary across a period in the d-block?

• Trends are less regular than in s/p blocks because:

  • Different numbers of d-electrons participate in bonding.

  • Stability of half-filled and filled d-subshells influences bonding.

• Result: melting points fluctuate rather than follow a simple increase.

20. How does metallic bonding sometimes gain covalent character?

• Very small, highly charged cations (e.g., Be²⁺, Al³⁺) can strongly polarise the electron sea.

• Electron density becomes distorted towards certain cations, giving partial directional (covalent-like) bonding.

• This strengthens the bond but may alter metallic properties.

21. Why are some metals better conductors than others?

• More delocalised electrons per atom → better conductivity.

• Fewer lattice imperfections → less scattering of electrons.

• Silver has the highest conductivity because its single s-electron per atom moves with minimal scattering.

22. How do periodic trends in metallic bonding explain melting point anomalies?

• In s/p-block metals, melting points track charge density (charge ÷ size).

• In transition metals, anomalies occur due to:

  • Electron configuration effects (e.g., half-filled d⁵ in Mn weakens bonding).

  • Varying d-electron involvement.

23. How do experimental data support metallic bonding theory?

• Electrical resistivity vs temperature: resistivity increases with temperature because ion vibrations scatter electrons more.

• Melting point patterns: data align with predictions from cation charge and radius.

• X-ray diffraction: confirms close-packed, ordered metallic lattice arrangements.

1) Guiding question — What determines the metallic nature and properties of an element?

• Definition / core model (IB wording): A metallic bond is the electrostatic attraction between a lattice of cations and delocalized electrons. This “sea of electrons” model is the basis for all metallic behaviour.

• What determines metallic character: (a) the availability of delocalized electrons (number of electrons that can become mobile), (b) the charge and size of the metal cations (charge density), and (c) the packing/coordination and crystal structure of the metal lattice. Together these control bond strength (cohesive energy), and hence melting point, hardness, conductivity and malleability. fileciteturn7file0

• How to connect to properties quickly (memorize): more delocalized electrons and smaller cations → stronger metallic bonding → higher melting point, greater hardness, and higher electrical/thermal conductivity

2) Explain electrical conductivity, thermal conductivity and malleability of metals

• Electrical conductivity: In metals many electrons are delocalized and free to move. Under an applied potential difference these electrons drift and carry charge — therefore metals conduct electricity in the solid and liquid states. The conductivity magnitude depends on the density of mobile electrons and scattering processes (impurities, temperature). fileciteturn7file4

• Thermal conductivity: Heat is transferred both by vibrating metal ions (phonons) and — importantly in metals — by delocalized electrons that rapidly transport kinetic energy through the lattice. This is why many metals are good thermal conductors (e.g., copper, silver).

• Malleability / ductility: Metallic bonds are non‑directional. When layers of positive ions slide past one another, the electron sea simply re‑arranges and continues to bind the ions in their new positions. Because like‑charge repulsion does not suddenly appear (unlike ionic lattices), metals deform rather than shatter.

3) Relate characteristic properties of metals to their uses

Give concise property → use pairs with 1‑line justification:

• High electrical conductivity (delocalized electrons) → electrical wiring (Cu), busbars (Al/Cu) — electrons carry charge efficiently.

• High thermal conductivity → cookware, heat sinks (Cu, Al) — fast heat distribution via electrons and phonons.

• Malleability & ductility (non‑directional bonding) → sheet metal, wire drawing — can be shaped without fracture.

• High melting point & strength (transition metals) → structural components, turbine blades, tools — strong cohesive energy from many delocalized electrons (including d‑electrons).

• Corrosion resistance / inertness (noble metals or alloys) → electronic contacts, jewellery — metals that resist oxidation (Au, Pt) or that form stable passivating layers (Cr in stainless steel).

4) Tool 1 / Inquiry 2 / Structure 3.1 — What experimental data demonstrate metal properties and trends?

• Electrical conductivity measurements: resistivity (Ω·m) or conductivity (S·m⁻¹) of metals (four‑probe or circuit methods) shows metals conduct in solid state and ranks conductors (Ag > Cu > Au typically). Use tabulated conductivity values as evidence.

• Thermal conductivity measurements: steady‑state heat flow or transient methods showing metals conduct heat well; data correlate with electrical conductivity (Wiedemann–Franz concept at a high level).

• Mechanical testing: tensile tests, hardness, and impact tests demonstrate malleability/ductility vs brittleness; stress–strain curves show yield and ductile failure behaviour for metals.

• Melting‑point tables across periods/groups: numeric mp data (e.g., Na 371 K, Mg 923 K, Al 933 K) illustrate the effect of increasing delocalized electrons/charge density across a period and of increasing atomic size down a group.

• Simple classroom demos: conductivity circuit (solid vs molten), hammering metal, and heating to show thermal conduction — all explicitly recommended in the syllabus as evidence linking bonding → properties.

5) Reactivity 3.2 — Trends in reactivity of metals predicted from the periodic table

• General rule (groups): metals become more reactive down a group because outer electrons are further from the nucleus and ionization energy decreases → electrons are lost more easily (alkali metals + water more vigorous down group).

• Across a period: metallic character generally decreases left→right; ionization energies rise, so elements become less readily oxidized.

• Transition metals: reactivity is less regular — d‑electrons, variable oxidation states and complex ion formation complicate simple periodic predictions; experimentally derived reactivity series and standard electrode potentials are used.

6) Explain trends in melting points of s‑ and p‑block metals

• Across a period (s→p metals): as you move right, number of delocalized electrons per atom increases (Na 1e⁻ → Mg 2e⁻ → Al 3e⁻), and nuclear charge increases without a new shell → electron density rises and metallic bonding strengthens → melting point tends to increase (e.g., Na → Mg → Al).

• Down a group: atomic/ionic radius increases (more shells) → electron density per unit volume decreases → weaker metallic bonding → melting point decreases (e.g., Na → K → Rb).

• p‑block complexities / anomalies: some p‑block “metals” (Ga, Sn, Pb) show structure‑dependent behaviour and partial covalency, causing departures from a simple trend — mention that IB expects a “simple treatment” in terms of charge/electron density, but be ready to note exceptions (Be unusually high mp for its group due to small size/high charge density).

7) Simple treatment: charge of cations & electron density (memorisable statement)

• One‑line concept to memorise: The strength of metallic bonding increases with increasing positive charge on the cations and with increasing electron density (more delocalized electrons per unit volume), and decreases as ionic radius increases. Use this to justify mp and conductivity trends.

8) Structure 2.4 — What features of metallic bonding make alloy formation possible?

• Non‑directional bonding: the electron sea binds cations without fixed pairwise directionality, allowing different types of atoms to replace host atoms or occupy interstices without destroying the bonding network.

• Substitutional alloys: atoms of similar size substitute for host atoms (e.g., zinc in copper → brass).

• Interstitial alloys: small atoms occupy interstitial sites (e.g., carbon in iron → steel).

• Why alloys are stronger: lattice distortions from size/charge differences impede dislocation motion, increasing hardness and yield strength; alloys also allow property tuning (corrosion resistance, melting point, conductivity). Explain using non‑directional bonding + microstructural effects.

9) Explain high melting point & electrical conductivity of transition elements

• High mp: transition metals can delocalize d‑electrons in addition to s‑electrons; this increases the total number of electrons contributing to the metallic bond and raises electron density and cohesive energy → higher melting points than many s‑block metals. Cite period‑4 transition metals as a classically higher set than Group 1/2.

• High electrical conductivity: delocalized d‑ and s‑electrons provide a large density of mobile charge carriers, so many transition metals (Ag, Cu, Au, etc.) are excellent conductors (note: Ag is the most conductive). Add that band structure and scattering still produce variation across the d‑block.

10) Structure 3.1 — Why is the trend in melting points across a period less evident across the d‑block?

d‑block complications: across the d‑block the filling of d orbitals, variations in the number of delocalizable electrons, changes in crystal structure (BCC/FCC/HCP) and electron pairing/exchange effects create non‑monotonic changes in cohesive energy.

Therefore melting points show irregular peaks and troughs (not the smooth increase/decrease seen in simple s/p blocks). IB expects you to state that the d‑block shows less evident trends because multiple competing electronic factors control bonding strength.

Explain: why copper is used for wiring

• Core physical reason: Copper has a very high electrical conductivity because it has a high density of delocalized (mobile) electrons that can move easily under an applied potential difference; those electrons are the charge carriers in the metal. As the S2.3 notes explain, when a potential difference is applied the delocalized electrons drift toward the positive terminal, producing current.

• Mechanical reason (ductility/ductile nature): Copper’s non‑directional metallic bonding and close‑packed crystal structure allow layers of cations to slide past each other without breaking the overall bonding (the electron sea re‑arranges), so copper is ductile and malleable — ideal for drawing into long wires and making reliable connections.

• Practical balance: Copper combines excellent conductivity, good ductility, corrosion resistance (reasonable), and moderate cost, which is why it is preferred over slightly better conductors like silver and more expensive alternatives like gold. IB/notes list copper as the most used cable metal for these combined reasons.

Three properties explained (write these short exam‑style bullets on the back):

1. Electrical conductivity: delocalized electrons are free to move; under a potential difference they drift and carry charge → metals conduct as solids.

2. Malleability / ductility: metallic bonding is non‑directional; when layers of cations slide the electron sea maintains attraction so the metal deforms rather than shatters.

3. Thermal conductivity: vibrating cations transfer kinetic energy to delocalized electrons, which rapidly transport energy through the lattice → good thermal conduction.

Describe substitutional vs interstitial alloys + one example each

• Substitutional alloy: atoms of similar size replace host metal atoms at lattice sites. Example: Brass = Cu with some Zn replacing Cu sites; Zn atoms substitute in and change mechanical and corrosion properties. (Use this if the question asks for an example of improved hardness or altered melting point.)

• Interstitial alloy: small atoms occupy interstitial holes between metal atoms, distorting the lattice. Example: Steel (Fe with small C atoms in interstices) — carbon atoms in octahedral/tetrahedral holes obstruct dislocation motion.

Exam‑style reasoning why an alloy is harder than its components (one crisp sentence):

“Alloying introduces atoms of different size or type that distort the metal lattice, impeding dislocation motion and making plastic deformation harder — therefore the alloy is harder and stronger than the pure metal.”

Deduce (two‑sentence explanation using mp data Na → Mg → Al)

Two‑sentence model answer:

“Melting points increase from Na → Mg → Al because the number of delocalized electrons per atom (1 → 2 → 3) increases across the period, raising electron density and strengthening metallic bonding. Additionally, the increasing nuclear charge draws electrons closer (smaller effective ionic radius), further increasing the electrostatic attraction between cations and the electron sea and so requiring more energy to melt.”