Studied by 8 people
Ionic bonding
Metals react with non-metals
Metal atoms lose 1 or more electrons, giving them a + charge
Non metal atoms gain 1 or more electrons, giving them a - charge
They attract each other as they are oppositely charged, and have strong electrostatic forces between ions called an ionic bond and both have full outer shells
Electrolysis
Contains two electrodes:
Anode: The positive electrode.
Cathode: The negative electrode.
These electrodes are connected to a power source.
The electrolyte (liquid or molten ionic compound) provides ions to carry the charge.
Positively charged ions (cations) in the electrolyte move toward the negatively charged cathode.
Negatively charged ions (anions) move toward the positively charged anode.
At the cathode, reduction occurs (gain of electrons) and unless the substance is gold, silver, platinum or copper, hydrogen (H+) gets produced.
At the anode, oxidation occurs (loss of electrons) and unless this substance is a halogen (group 7), oxygen gas (O2) is produced.
Covalent bonding
Occurs between non-metal atoms where electrons are shared to achieve full outer electron shells. Each shared pair of electrons constitutes a covalent bond. Covalent compounds typically have lower melting and boiling points compared to ionic compounds. They do not conduct electricity in solid or molten state, are often insoluble in water, and can exist in any state at room temperature. For some compounds, more electrons may need to be shared. They can have up to 3 bonds per atom.
Simple molecular structures
Structures formed by covalent bonds between non-metal atoms, typically consisting of small groups of atoms (molecules) held together by weak intermolecular forces. These structures can exist in various states at room temperature, ie. solids, liquids, and gases. Simple molecular substances generally have low melting and boiling points due to weak intermolecular forces that require little energy to overcome.
Giant covalent structures
Large networks of atoms connected by covalent bonds in a lattice. Unlike simple molecular structures, they consist of a vast number of atoms, resulting in strong intermolecular forces that contribute to high melting and boiling points. These structures do not conduct electricity due to the lack of free-moving charged particles and are insoluble in water, although graphite and graphene can conduct electricity, as they only use 3 electrons in bonding, leaving an electron which is free to move and carry a charge.
Metallic bonding
Metallic bonding is the chemical bonding that occurs between metal atoms. It involves a lattice of positively charged metal ions surrounded by a ‘sea’ of delocalized electrons. These free-moving electrons are responsible for many of the properties of metals, such as conductivity.
Properties of ionic compounds
High melting and boiling points due to strong electrostatic forces between oppositely charged ions in the lattice.
Ionic compounds are hard because of the strong ionic bonds but brittle due to the rigid structure; shifting layers cause repulsion between like charges.
Many ionic compounds dissolve in water as the ions interact with water molecules.
Solid State: Ionic compounds do not conduct electricity because ions are fixed in place.
Molten or Dissolved State: They conduct electricity as ions are free to move.
They form well-organized lattices.
Polymers
Polymers consist of long chains of repeating units, leading to large molecular structures.
Thermoplastics: Flexible and can be remolded.
Thermosets: Rigid and durable due to cross-linking .
Melting points are generally lower than ionic or metallic compounds but vary depending on structure and bonding.
Most polymers do not conduct electricity, except for certain conductive polymers.
Polymers are less dense than metals or ceramics.
Durable but Can Degrade Over Time: Polymers are resistant to many chemicals but may degrade due to UV light, heat, or environmental factors.
Found in plastics, textiles, adhesives, and more due to their versatility.
Properties of metals
High Electrical and Thermal Conductivity: Due to the presence of free-moving (delocalized) electrons, which kinetic energy and electrical charge can pass through easily.
Malleable and Ductile: Can be hammered into shapes or drawn into wires without breaking, due to metallic bonds allowing atoms to roll over each other without breaking.
High Melting and Boiling Points: Resulting from strong metallic bonds between metal ions and free electrons.
Lustrous Appearance: Metals are shiny due to their ability to reflect light.
Dense: Most metals have a high density due to closely packed atoms.
Properties of alloys
Stronger and Harder: Alloying increases strength by distorting the layers in the lattice structure as the different metal ions are different sizes
Improved Corrosion Resistance: Some alloys resist rusting or oxidation better than pure metals (e.g., stainless steel).
Tailored Properties: Alloys can be designed for specific purposes (e.g., brass for musical instruments, bronze for statues).
Reduced Malleability and Ductility: Compared to pure metals, alloys are often less malleable as the layers can’t slide over each other as easily.
Lower Electrical Conductivity: Alloys generally conduct electricity less effectively than pure metals due to the distorted layers meaning the charge cannot pass through as easily as in pure metals, where the layers are regular and straight.
Properties and structure of diamond (giant covalent)
Hardness: Diamond is the hardest natural substance as each carbon atom is bonded to four others and there are many strong covalent bonds that require lots of energy to break.
High Melting Point: Extremely high due to strong covalent bonds throughout the lattice.
Electrical Insulator: No free electrons or ions to conduct electricity as each carbon atom is bonded to 4 others.
Thermal Conductor: Excellent thermal conductivity due to the strong lattice vibrations.
Transparent and Lustrous: Light passes through diamond so it looks very nice.
Brittle: Despite its hardness, diamond can shatter under a sharp impact because it is not malleable.
Structure and properties of graphite
Structure:
Graphite is a form (allotrope) of carbon where each carbon atom is covalently bonded to three other carbon atoms in a hexagonal arrangement, forming flat layers/giant lattice.
The layers are held together by weak intermolecular forces, allowing them to slide over each other.
Each carbon atom has one free (delocalized) electron, which moves freely between the layers.
Properties:
Soft and Slippery: Layers can slide over each other, making graphite an excellent lubricant and used in pencils.
High Melting Point: Strong covalent bonds within layers require a lot of energy to break.
Electrical Conductor: Each carbon atom is bonded to only 3 others, delocalized electrons can move freely through the structure, allowing graphite to conduct electricity.
Opaque and Black: Graphite has a dark, metallic appearance.
Insoluble in Water: Covalent bonds and the structure make it insoluble.
Lightweight: Graphite is less dense compared to diamond due to the spacing between layers.
Structure and properties of graphite
Structure:
A single layer of carbon atoms arranged in a two-dimensional honeycomb lattice.
Each carbon atom is covalently bonded to three others, leaving one delocalized electron per atom.
Properties:
High Electrical Conductivity: Delocalized electrons allow efficient charge flow.
Extremely Strong: Stronger than steel due to its covalent bonding and lightweight structure.
Thin and Transparent: Only one atom thick but still visible.
Flexible: Can bend without breaking.
High Thermal Conductivity: Excellent at transferring heat.
Structure and properties of fullerenes
Structure:
Molecules of carbon atoms arranged in hollow shapes, such as spheres, tubes, or ellipsoids.
The most well-known fullerene is C₆₀ (Buckminsterfullerene), which has a spherical shape.
Properties:
Conductivity: Can conduct electricity but less so than graphene.
Lightweight and Strong: Suitable for nanotechnology and material science.
Catalytic Properties: Used in reactions as catalysts.
Solubility: Some fullerenes are soluble in organic solvents.
Versatile Uses: Found in drug delivery systems, electronics, and lubricants.
Uses of nanoparticles
Medicine:
Drug Delivery: Nanoparticles can deliver drugs directly to specific cells, improving treatment efficiency (e.g., cancer therapy).
Diagnostics: Used in imaging techniques like MRI for improved precision.
Electronics:
Enhance the performance of electronic devices (e.g., smaller transistors, flexible displays).
Used in manufacturing conductive inks for printed circuits.
Cosmetics:
Sunscreens: Titanium dioxide and zinc oxide nanoparticles block harmful UV rays without leaving a white residue.
Skincare: Nanoparticles improve the absorption of active ingredients.
Environment:
Water Purification: Nanoparticles can remove contaminants like heavy metals and pathogens.
Pollution Control: Used in air filters and catalytic converters to reduce emissions.
Energy:
Solar Cells: Increase efficiency by improving light absorption.
Batteries: Nanoparticles enhance energy storage and lifespan.
Materials Science:
Stronger and Lighter Materials: Carbon nanotubes and graphene improve the strength of materials while reducing weight.
Self-Cleaning Surfaces: Nanoparticles create coatings that repel dirt and water.
Food Industry:
Packaging: Nanoparticles are used in packaging to extend shelf life by preventing microbial growth.
Food Additives: Enhance texture or nutritional value.
Medicine and Antimicrobials:
Silver Nanoparticles: Used in wound dressings, clothing, and medical devices for their antibacterial properties.
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