Ionic and Covalent Bonds

Chemical Bonds

• Three types of strong chemical bonds exist: covalent, ionic, and metallic.
• Covalent bonds form between non-metals.
• Ionic bonds form between a non-metal (negative ion) and a metal (positive ion).
• Metallic bonds form within metallic elements and alloys (mixtures of metallic elements).
• A covalent bond is a shared pair of electrons.
• An ionic bond involves the transfer of electrons from a metal to a non-metal.
• A metallic bond involves the sharing of delocalised electrons between atoms.

Ionic Bonding

• Metals in groups 1, 2, or 3 typically have 1, 2, or 3 electrons in their outermost shell.
• Non-metals are closer to having a full outermost shell, needing only 1, 2, 3, or 4 more electrons.
• The group number indicates the number of outermost electrons, and the period number indicates the number of electron shells.
• Atoms aim to achieve a full outermost shell, similar to noble gases (group 0).
• Metals transfer outermost electrons to non-metals, forming oppositely charged ions.
• Metals lose electrons, resulting in a positive charge.
• The charge corresponds to the group number; for example, Aluminum (Al) in group 3 becomes Al3+.
• Non-metals gain electrons, resulting in a negative charge.
• The charge of the ion is calculated as 8 minus the group number; for example, Oxygen (O) in group 2 becomes O2-.
• Oppositely charged ions are held together by electrostatic forces of attraction.
• Naming ionic compounds involves using the name of the metal and the name of the non-metal with the ending "-ide" if the compound contains one type of ion (e.g., MgCl2 = magnesium chloride).
• If the compound contains two or more types of ions, the ending "-ate" is used (e.g., CaCO3 = calcium carbonate).
• Carbonate is CO3 with a 2- charge.
• Sulfate is SO4 with a 2- charge.
• Nitrate is NO3 with a - charge.
• Hydroxide is OH with a - charge.
• Ionic compounds are held together by strong electrostatic forces of attraction between oppositely charged ions, acting in all directions within the lattice.

Limitations of Ball and Stick Diagrams

• Does not show the electrostatic forces of attraction acting in all directions.
• Shows a lot of space between the ions.
• Diagrams show any group 1 and group 7 elements as there is a 1:1 ratio of ions

Covalent Bonding

• Covalent bonds form when non-metals share pairs of electrons.
• Some covalently bonded substances consist of small molecules such as methane (CH4), water (H2O), hydrochloric acid (HCl), nitrogen (N2), ammonia (NH3), chlorine (Cl2), and hydrogen (H2).
• A covalent bond can be represented by a straight line (e.g., H—O—H).
• Each straight line represents a covalent bond.
• Nitrogen gas (N2) exhibits a triple bond.
• Covalent bonds are strong.
• Oxygen makes 2 bonds to have 8 electrons in the outer shell _ Hydrogen makes 1 bond to have 2 electrons in the outer shell
• Some covalently bonded substances are very large molecules, e.g. polymers.
• Some covalently bonded substances have giant covalent structures, e.g. diamond and silicon dioxide (silica).

Metallic Bonding

• Metals consist of giant structures of positive ions in a regular pattern.
• Metals have a few (1-3) electrons in their outermost shell.
• Metals want to lose those electrons to have the same electron configuration as a noble gas.
• Electrons in the outer shells are delocalised within the structure.
• Sharing delocalised electrons results in strong metallic bonds.
• Electrostatic forces of attraction between positive ions and delocalised electrons make the metallic bonds strong.
• Delocalised electrons are free to move through the structure.

States of Matter

• Three states of matter exist: solid, liquid, and gas.
• The amount of energy needed to change state from solid depends on the strength of the forces between the particles of the substance (often intermolecular forces).
• Stronger forces result in a higher melting and boiling point.
• Melting point is the temperature at which melting and freezing occur.
• Boiling point is the temperature at which boiling and condensing occur.
• (s) represents solid substances
• (l) represent liquid substances
• (g) represents gas substances
• (aq) represents aqueous solutions (substances dissolved in water)

Properties of Ionic Compounds

• Ionic compounds have regular structures, also known as giant ionic lattices.
• Strong electrostatic forces of attraction exist in all directions between oppositely charged ions.
• Ionic compounds have high melting and boiling points due to the significant energy required to overcome the many strong bonds.
• When melted or dissolved in water, ionic compounds can conduct electricity because ions are free to move, enabling charge flow.
• Ionic compounds cannot conduct electricity when solid because the ions are not mobile.

Properties of Small Molecules

• Small molecules typically have covalent bonding and exist as liquids or gases at room temperature.
• Small molecules commonly have low melting and boiling points due to weak intermolecular forces.
• Intramolecular covalent bonds are not broken during a change of state; only weak intermolecular forces are overcome.
• As the size of the molecule increases, so does the substance's melting and boiling point.
• Small covalent molecules do not conduct electricity because they have no overall electrical charge, and there are no delocalised electrons or mobile ions available to carry a charge.

Polymers

• Polymers are very large/long covalent molecules.
• Atoms in polymers are linked to other atoms by strong covalent bonds.
• Intermolecular forces between polymers are relatively strong.
• Polymers are solid at room temperature due to relatively strong intermolecular forces.

Giant Covalent Structures

• Substances with giant covalent structures are solid at room temperature and have high melting points.
• All atoms are bonded to other atoms by strong covalent bonds.
• Strong covalent bonds must be overcome when the substance is melted or boiled.
• Diamond and silica (silicon dioxide) are examples of giant covalent structures.

Properties of Metals and Alloys

• Metals have giant regular structures with strong metallic bonding
• Metals have high melting and boiling points due to strong metallic bonding
• Pure metals are arranged in layers and are malleable, can be bent to shape, as the layers can slide
• Alloys are stronger as the different sized atoms distort the layers, making it harder for them to slide
• Metals are good conductors of electricity because the delocalised electrons in the metal carry electrical charge
• Metals are good conductors of thermal energy because energy is transferred by the delocalised electrons

Diamond

• Made up entirely of carbon, a giant covalent molecule
• Each carbon atom is covalently bonded to 4 other carbon atoms
• Covalent bonds are strong
• Very hard and has a very high melting point
• Cannot conduct electricity as every carbon's (four) electrons are in covalent bonds with other carbon atoms

Graphite

• Allotrope of diamond
• Each carbon atom is covalently bonded to 3 other carbon atoms
• Each carbon atom has one unbonded electron which forms layers of hexagonal rings
• This electron is delocalised and able to carry an electrical charge
• Can conduct electricity
• There are no covalent bonds between the layers and so they can slide and this is what makes graphite soft and slippery
• Weak intermolecular forces are present between the layers and are overcome when graphite is melted or boiled, this makes graphite brittle
• Has a lower melting and boiling point than diamond
• Allotrope: different structural forms of the same element in the same physical state

Graphene and Fullerenes

• Graphene is a single layer of graphite
• Graphene is useful in electronics and composites as it can conduct electricity. Each carbon atom has a delocalised electron.
• Fullerenes are molecules of carbon with hollow shapes like hexagonal rings.
• The rings may be composed of 5, 6 or 7 carbon atoms.
• The first fullerene to be discovered was Buckminsterfullerene (C60), containing 60 carbon atoms.
• Buckminsterfullerene has a spherical shape.
• Carbon nanotubes are cylindrical fullerenes with very high length-to-diameter ratios.
• Fullerenes are useful for nanotechnology, electronics, and materials.

Sizes of Particles and Their Properties

• Nanoscience refers to structures that are 1–100 nm in size
• Nanoparticles are smaller than fine particles (PM2.5) and coarse particles (PM10)
• PM2.5 (tiny particles that reduce visibility
• PM10 is commonly referred to as dust.
• Nanoparticles have large surface area to volume ratios
• As the side of a cube decreases by a factor of 1 the surface area to volume ratio increases by a factor of 10
• Nanoparticles may have properties different from those for the same materials in bulk due to their high surface area to volume ratio
• Smaller quantities of nanoparticles are need to be effective than for materials with normal particle sizes

Uses of Nanoparticles

• Medicine - drug transporters
• Electronics – chips
• Cosmetics – for better coverage
• Sun cream- for better coverage
• Deodorants
• Catalysts - can be cheaper and more efficient
• Advantages – better coverage, less waste, less materials to save resources
• Disavantages – potential cell damage to body, harmful effects on the Environment, more safety testing required is expensive and time consuming
• New applications for nanoparticulate materials are an important area of research