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5. Bonding Flashcards

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1
Q

Structure of Metallic Substances

A

Metals are made of a lattice of positive metal ions (cations) surrounded by a sea of delocalised electrons.

Valence electrons are released from metal atoms and are free to move throughout the lattice.

Metallic bonding = electrostatic attraction between positive metal ions and the mobile electrons.

Bonding is:

Non-directional: Acts in all directions around each ion.

Delocalised: Electrons are not tied to a specific atom.

🧠 Think: Metal cations held together by mobile electrons like glue.

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2
Q

Explain why metals conduct electricity

A

Metals have free-moving valence electrons (also called a “sea of electrons”).

These electrons can move easily between atoms, allowing electricity to flow.

When voltage is applied:

Electrons are repelled from the negative terminal of a power source.

They flow through the metal and are attracted to the positive terminal.

This electron movement is what forms an electric current.

Metals conduct electricity well because their electrons are delocalised and mobile..)

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3
Q

Explain why metals conduct heat

A
  • Electrons gain kinetic energy in hotter areas of the metal.
  • They quickly transfer heat energy to other areas by moving freely.
  • Metal ions also vibrate, helping spread thermal energy through the lattice.
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4
Q

Explain why metals are malleable and ductile

A
  • Metals are malleable (can be hammered) and ductile (can be stretched).
  • Reason: Non-directional bonding:

When force is applied, layers of ions can slide past each other.

The sea of electrons readjusts, keeping the structure stable.

🧠 Like ball bearings in oil: movement without breaking.

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5
Q

Explain why metals have a high melting point

A
  • Strong electrostatic attraction between cations and electrons requires a lot of energy to break.
  • More valence electrons → stronger bonding → higher melting point.

🧠 Mercury (Hg) is an exception — it’s a liquid at room temperature due to weak bonding.

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6
Q

Explain why metals have high density

A
  • Most metals have closely packed lattices → particles are tightly arranged.
  • Result: High density, heavier than non-metals of similar volume.
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7
Q

Explain why metals are solid at room temperature

A
  • Strong metallic bonding holds the lattice together → most metals are solids at room temperature.

❗Exception: Mercury (MP = -39°C) is a liquid at room temperature due to weak metallic bonding.

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8
Q

What are Delocalize electrons

A
  • Free-moving valence electrons not tied to a single atom.
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9
Q

What are Metallic Lattices

A
  • A 3D arrangement of metal cations with electrons moving around them
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10
Q

What is Non-directional bonding

A
  • A force of attraction that acts equally in all directions.
  • Metals experience a non-directional force of attraction, which acts equally in all directions between positive metal ions and the sea of delocalised electrons.
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11
Q

What is a Sea of Electrons

A
  • Mobile, negative electrons holding the lattice together.
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12
Q

Electrical Conductivity Comparison of Metallic Substances in Different States

A

Solid: ✅ Conduct
→ Sea of delocalised electrons moves freely

Molten: ✅ Conduct
→ Delocalised electrons still mobile

Aqueous: ❌ Do not dissolve
→ Metals don’t form ions in water (except reactions)

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13
Q

What is an Ionic Compound

A

Formed by a metal reacting with a non-metal (or ammonium, NH₄⁺).

The metal loses electrons → forms a positive ion (cation).

The non-metal gains electrons → forms a negative ion (anion).

Both ions achieve a noble gas configuration (full outer shell) — this follows the octet rule.

The oppositely charged ions are held together by strong electrostatic forces, forming an ionic bond.

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14
Q

Structure of Ionic Compounds

A

Ions arrange in a 3D ionic lattice (e.g. NaCl shown in Figure 8).

Positive and negative ions are arranged in a regular, repeating pattern.

Each positive ion is surrounded by negative ions and vice versa.

This arrangement maximises attractive forces and minimises repulsion.

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15
Q

Ionic Bonding

A

An ionic bond is the electrostatic attraction between oppositely charged ions.

These forces are strong and non-directional — they act equally in all directions.

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16
Q

Electron Transfer Diagrams & Lewis Structures

A

Electron transfer diagrams show how electrons move from the metal to the non-metal.

Lewis structures (Figure 7) show ions in square brackets:

Dots and crosses represent original electrons.

Only valence electrons are shown.

Square brackets indicate electrons are no longer shared.

📌 Used to show charge and structure of ions in compounds like MgCl₂, Al₂O₃.

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17
Q

📊 Properties of Ionic Compounds

A

❌ Poor Conductors of Electricity (Solid State)
In a solid, ions are held in fixed positions → no free-moving charged particles.

Electrons are also localized → can’t carry charge.

Therefore, ionic solids do not conduct electricity.

✅ Good Conductors of Electricity (Molten or Aqueous)
In molten form (melted) or aqueous solution (dissolved in water):

Ions are mobile and free to move.

Positive ions move toward the negative electrode.

Negative ions move toward the positive electrode.

This allows electric current to be conducted.

📌 Figure 8 shows how ions flow in both molten and aqueous states.

🧱 Hard and Brittle
Ionic solids are hard due to strong bonding.

But they are also brittle:

If a force shifts the lattice and like charges align, repulsion occurs.

This causes the lattice to fracture.

📌 Repulsive forces between like charges are strong and destabilise the structure.

🔥 High Melting and Boiling Points
Due to strong electrostatic attraction between oppositely charged ions.

Large amounts of energy are needed to:

Break the lattice,

Separate the ions.

Therefore, ionic compounds have high melting and boiling points.

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18
Q

What is an Ionic Bond

A

Electrostatic attraction between positive and negative ions.

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19
Q

What is an Ionic Lattice

A

3D repeating structure of alternating charged ions. - Consist of a lattice of cations and anions

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20
Q

What is an Electrostatic Force

A

The strong attraction between opposite charges.

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21
Q

Electrical Conductivity Comparisons of Ionic Compounds in Different States

A

✅ Ionic Compounds:

Solid: ❌ Do not conduct
→ Ions are fixed in lattice, cannot move

Molten (liquid): ✅ Conduct
→ Lattice breaks, ions free to move

Aqueous (dissolved in water): ✅ Conduct
→ Ions separate and move freely

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22
Q

Properties of Non-Metallic Substances

A
  • Made of molecules: groups of two or more non-metal atoms bonded covalently in fixed proportions.
  • Molecules are held together internally by strong covalent bonds.
  • Different molecules are held together weakly by intermolecular forces (e.g., dispersion, dipole-dipole, H-bonds).
  • Boiling water breaks intermolecular forces, not covalent bonds.
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23
Q

What is a Covalent Bond

A
  • Strong bond formed between non-metal atoms.
  • Involves sharing electrons so each atom achieves a noble gas configuration (Octet Rule).

Can form:

Single bonds: 1 pair shared (e.g., F₂)

Double bonds: 2 pairs shared (e.g., O₂)

Triple bonds: 3 pairs shared (e.g., N₂)

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24
Q

What are Diatomic Molecules

A

Diatomic molecules = made of two atoms.

Can be:

  • Elements (e.g., F₂, Cl₂, O₂)
  • Compounds (e.g., HF, HCl)

Example 1: Fluorine atoms each share 1 electron → single bond (F₂)

Example 2: H shares 1e⁻ with F → forms HF (single bond)

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25
What are Lewis Structures
Only valence electrons shown (outermost shell). Dots (:) represent electrons. Lines can replace shared pairs.
26
What are Polyatomic Molecules
A polyatomic molecule is a molecule made up of more than two atoms covalently bonded together. All atoms are usually non-metals, and the molecule has no overall charge. Example: H₂O, CO₂, NH₃
27
Electrical Conductivity Comparisons of Covalent Substances in Different States
**Solid State:** - Do NOT conduct electricity - No free electrons or ions present **Molten (Liquid) State:** - Still do NOT conduct electricity - Most covalent substances remain non-conductors (no ions formed) - Exception: Some acids (e.g., HCl) will ionise if melted **Aqueous Solution:** - Most do NOT conduct electricity - Exception: Some covalent molecules (like HCl, NH₃, SO₂) ionise in water to form ions → can conduct - Majority (e.g., sugar, ethanol, oxygen gas) do NOT conduct (no ions formed in water)
28
What are Polyatomic Ions
- A polyatomic ion is a group of covalently bonded atoms that acts as a single ion and carries a net positive or negative charge due to the loss or gain of electrons. Example: OH⁻, SO₄²⁻, NH₄⁺
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Covalent Network Substances - Definition
- Covalent network substances are non-metallic elements or compounds made of atoms connected by strong covalent bonds in a continuous 3D lattice. Examples include: - Diamond and graphite (carbon) - Silicon dioxide (SiO₂) - Silicon carbide (SiC) - Boron
30
Covalent Network Substances - Structure
- Atoms are covalently bonded in a giant lattice. - No discrete molecules. - E.g., Quartz (SiO₂): Each silicon atom is covalently bonded to oxygen atoms in a 3D arrangement.
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Properties of Covalent Network Substances
**Non Conductors of Heat/Electricity:** - No mobile charge carriers (ions or free electrons. **Very hard and brittle:** - Strong covalent bonds need a large amount of energy to break. - Covalent network substances are brittle because when a strong force is applied, the rigid bonds break along a plane, causing the structure to shatter rather than bend. **Very high melting and boiling points:** - Massive energy required to break covalent lattice
32
Diamond Covalent Network Substance - Structure, Bonding, Properties
**Structure:** - Each C atom covalently bonded to 4 other C atoms (tetrahedral 3D lattice). **Bonding:** - Strong covalent C–C bonds throughout. **Properties:** - Extremely hard (hardest natural substance) - Very high melting point - Non-conductor of electricity (no free electrons) Use: - Cutting tools, jewelry
33
Graphite Covalent Network Substances - Structure, Bonding, Properties
**Structure:** - Each C atom bonded to 3 others, forming 2D layers (hexagonal sheets). **Bonding:** - Strong covalent bonds within layers, weak forces between layers; delocalised electrons between layers. **Properties:** - Soft, slippery (layers slide over each other) - Good conductor of electricity (delocalised electrons) - High melting point Use: - Lubricants, pencils, electrodes
34
Silicon Dioxide (SiO2) Quartz Covalent Network Substances - Structure, Bonding, Properties
**Structure:** - Each Si atom bonded to 4 O atoms, each O bonded to 2 Si atoms (3D network). **Bonding:** - Strong Si–O covalent bonds throughout. **Properties:** - Very hard and brittle - Very high melting and boiling points - Non-conductor of electricity Use: - Glass, sand, quartz
35
Silicon (Si) Covalent Network Substances - Structure, Bonding, Properties
**Structure:** - Each Si atom covalently bonded to 4 other Si atoms (diamond-like lattice). **Bonding:** - Strong Si–Si covalent bonds. **Properties:** - Hard and brittle - Very high melting point - Semiconductor (conducts when doped) Use: - Microchips, solar panels
36
Silicon Carbide (SiC) Covalent Network Substances - Structure, Bonding, Properties
**Structure:** - Each Si atom bonded to 4 C atoms (and each C to 4 Si), 3D network. **Bonding:** - Strong Si–C covalent bonds. **Properties:** - Extremely hard - Very high melting point - Chemically inert, non-conductor (unless doped) Use: - Abrasives, cutting tools
37
Boron Covalent Network Substances - Structure, Bonding, Properties
**Structure:** - Complex 3D covalent network (main form: α-rhombohedral). **Bonding:** - Strong B–B covalent bonds. **Properties:** - Very hard and brittle - Very high melting point - Non-conductor (unless doped) Use: -Specialty materials, semiconductors
38
Properties of Covalent Bonds
**Melting Point:** _Covalent Molecular Substances:_ - Low melting and boiling points. - Held together by weak intermolecular forces (e.g., dispersion, dipole-dipole, hydrogen bonding). - Only small energy is needed to separate molecules. _Covalent Network Substances:_ - Very high melting and boiling points. - Atoms are bonded by strong covalent bonds in a 3D lattice. - A large amount of energy is needed to break the network. **Conductor of Electricity:** _Covalent Molecular Substances:_ - Do not conduct electricity. - No mobile ions or free electrons to carry charge. _Covalent Network Substances:_ - Diamond, SiO₂: Do not conduct – all electrons are used in bonding. - Graphite: Does conduct – delocalised electrons between layers can move freely. **Hardness:** _Covalent Molecular Substances:_ - Usually soft or brittle. - Easy to break weak intermolecular forces between molecules. _Covalent Network Substances:_ - Very hard (e.g., diamond – hardest natural material). - Strong covalent bonds throughout the structure make them rigid and strong. - Graphite is soft despite being a network solid – layers slide due to weak intermolecular forces.
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Network Substances Definition
Network substances are solids made of a continuous lattice (network) of strongly bonded atoms or ions extending in all directions. They can be covalent, ionic, or metallic depending on the bonding type.
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Network Substances Types
**Covalent Network Substances:** - Bonding: Atoms held by strong covalent bonds. - Examples: Diamond, silicon dioxide (SiO₂), graphite. Properties: - Very high melting/boiling points - Hard (except graphite – soft due to weak forces between layers) - Poor conductors (except graphite – delocalised electrons) **Ionic Network Substances:** - Bonding: Positive and negative ions held by strong electrostatic forces. - Examples: Sodium chloride (NaCl), magnesium oxide (MgO). - Properties: - High melting/boiling points - Hard and brittle - Conduct when molten or in solution (ions free to move) **Metallic Network Substances:** - Bonding: Positive metal ions in a "sea of delocalised electrons". - Examples: Copper (Cu), iron (Fe), aluminium (Al). - Properties: - High melting/boiling points - Malleable and ductile - Conduct electricity and heat well
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Network Substances Matter State
They are all solids
42
What are Nanoparticles
Nanoparticles are materials with sizes ranging from 1 to 100 nanometres (nm).
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Key properties of Nanoparticles
- High surface area to volume ratio - Size-dependent optical properties – Can show different colours due to light absorption in the visible spectrum - Enhanced reactivity and toughness - Properties depend on their unique size, shape, and structure
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Nanoparticles - Example (Gold)
- Exhibit size-dependent optical effects - Used in ancient artifacts such as the Lycurgus Cup (Roman, ~AD 400) - Cup changes colour: - Green in reflected light - Red in transmitted light - Caused by gold-silver alloy nanoparticles embedded in glass
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Nanoparticles Applications
- Catalysis (speeding up chemical reactions) - Imaging (e.g., medical imaging) - Medical applications (e.g., drug delivery) - Energy research (e.g., solar cells, fuel cells) - Environmental uses (e.g., water purification, pollutant detection)
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Environmental Concerns in regards to Nanoparticles
- Heavy metal nanoparticles (e.g., lead, mercury, tin) - Extremely stable and resistant to degradation - Potential to cause environmental toxicity - Because they do not break down easily
47
What are Allotropes
Allotropes are different structural forms of the same element in the same physical state - The atoms are arranged differently, giving each allotrope unique physical and chemical properties
48
What are Carbon Allotropes
Allotropes of Carbon: Carbon can exist in multiple allotropes due to its ability to form four covalent bonds and create various bonding networks (e.g., rings, layers, 3D frameworks).
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Why Allotropes Happen in Carbon
**Carbon's Valency (4) allows for:** - Single, double, and triple bonds - Chains, rings, layers, and 3D networks **This bonding flexibility enables it to form:** - Covalent network solids - Discrete molecules - 2D Sheets
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Summary of Major Carbon Allotropes
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What is an Alloy
An alloy is a mixture of a metal with one or more other elements, usually metals or carbon. - Example: - Steel = Iron + Carbon - Importance: Stronger and more resistant to corrosion and pure iron --> used in buildings, tools, and cars
52
How does atom size affect allow properties
- Different-sized atoms distort the regular metal lattice. - This disrupts layers from sliding easily → makes alloy harder and stronger than pure metal. - Called "lattice distortion".
53
Pure Metal vs. Alloy Properties
**Strength:** - Pure metals are softer, as layers slide easily - Alloys are harder, due to distorted layers **Corrosion:** - Pure metals often rusts/corrodes easily - Alloys are more corrosion-resistant **Malleability:** - Pure metals are more malleable (bendable) - Alloys are less malleable (more rigid) **Conductivity:** - Metals have High Conductivities - Alloys have slightly lower conductivity
54
Diamond Allotrope - Structure, Bonding, and Properties
**Structure:** 3D tetrahedral network **Bonding:** Each carbon forms 4 strong covalent bonds **Properties Explained:** - Hard + High MP: Rigid 3D structure, bonds require lots of energy to break - Non-conductive: No free electrons or ions to carry charge
55
Graphite Allotrope - Structure, Bonding, and Properties
**Structure:** Layers of hexagonal rings (graphene sheets) **Bonding:** - Within layers: Each carbon atom forms 3 covalent bonds - 1 valence electron is delocalised, moving between layers → conducts electricity - Between layers: Weak dispersion forces (no covalent bonding) **Properties Explained:** - Slippery (good lubricant): Weak interlayer attraction → layers slide easily - Conductive: Free delocalised electrons carry charge
56
Graphene Allotrope - Structure, Bonding, and Properties
**Structure:** Single layer of graphite (2D hexagonal sheet) **Bonding:** Each carbon bonded to 3 others, 1 delocalized electron **Properties Explained:** - Very strong: Strong covalent bonds across the sheet - Excellent conductor: Free electrons move easily across the sheet
57
Fullerenes (C60) Allotrope - Structure, Bonding, and Properties
**Structure: Spherical shape (buckyball)** **Bonding:** Covalent bonds in a closed cage (each C bonded to 3 others) **Properties Explained:** - Conductive: Some delocalised electrons can move - Used in nanotech: Unique shape, small size, and reactivity
58
Carbon Nanotubes Allotrope - Structure, Bonding, and Properties
**Structure:** Cylindrical tubes of graphene **Bonding:** Covalent bonds with delocalised electrons **Properties Explained:** - Very strong: Covalent bonds along the tube = high tensile strength - Conductive: Delocalised electrons travel along the tube like a wire
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Order of Strongest IMFs
Strongest → Weakest IMFs: - Ion-Ion - Ion-Dipole - Hydrogen Bonding - Dipole-Dipole - (Ion-Induced Dipole / Dipole-Induced Dipole) - Dispersion (London) Forces
63
Order of strongest Bonds
Covalent > Ionic > Metallic

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