chemistry sac 2

carbon materials

it is the building block of life because of its ability to form many complex but stable molecules.

Some forms of pure carbon that exists include: diamond, graphite, graphene and amorphous carbon.

they are allotropes of carbon

• all made up pure carbon but have very different properties

allotropes

An allotrope is a different structural form of the same element in the same physical state.

diamond

a form of carbon where every carbon atom has four single covalent bonds to other carbon atoms arranged into a 3D covalent network lattice. This structure is very strong.

properties of diamond

• high melting point

• hard

• brittle

• no electrical conductivity

• high thermal conductivity

• insoluble

high melting point

Diamond is a giant covalent lattice, where each carbon atom is bonded to four others in a tetrahedral arrangement. These strong covalent bonds extend throughout the entire 3D structure. A large amount of energy is required to break these bonds and disrupt the structure of diamond. This results in a high melting point

hard

Each carbon atom in diamond forms four strong covalent bonds with other carbons. This creates a rigid, 3D network. The strong bonding throughout the structure makes diamond extremely hard

brittle ( hard but able to break easily)

the 3D rigid lattice structure of diamond does not allow it to be bent. When force is applied, the bonds break suddenly, causing diamond to shatter.

no electrical conductivity ( flow of electricity through the movement of charged particles)

In diamond, all four valence electrons of each carbon atom are used in bonding. This means that there are no free or delocalised electrons available to conduct electricity. Therefore, there is no movement of charged particles to conduct electricity.

high thermal conductivity (ability to transfer heat)

the strong covalent bonds between the carbon atoms allow atoms to vibrate and transfer heat. The rigid lattice allows these vibrations to travel without scattering resulting in a high thermal conductivity.

insoluble ( unable to be dissolved/ no aqueous state)

Diamond’s atoms are held together by strong covalent bonds in a 3D network. Solvents cannot break these strong bonds. As a result, diamond is insoluble in both water and organic solvents.

the strength of the covalent bonds in the structure of diamond cannot be overcome by the intermolecular bonds of a solvent like water

applications

• cutting tools

  the hardness of diamond resists wear and enhances durability

• thermal conductor in electrical components

  diamond’s strong covalent bonding enables thermal conductivity

• optical components

  diamond lasers can be produced due to diamond’s ability to transmit heat and light very effectively

• abrasive

  diamond is able to induce friction on other objects without experiencing wear itself due to its hardness

graphite

where carbon atoms are covalently bonded to 3 other carbon atoms, but only in 2D layers. There are also delocalised electrons moving around the different layers

this structure is called a covalent layer lattice

dispersion forces between graphite

properties of graphite

• high melting point

• high thermal conductivity

• soft, slippery feeling

• lower density than diamond

• insoluble

high melting point - important

Graphite has strong covalent bonds between carbon atoms within each layer. A large amount of energy is required to break these covalent bonds throughout the structure and disrupt the structure of graphite, resulting in a high melting point.

melting involves breaking covalent bonds not just separating layers

high thermal conductivity

Each carbon atom in graphite forms three covalent bonds, leaving one delocalised electron per atom. When heat is applied, these delocalised electrons gain kinetic energy and move freely and quickly within the layers, transferring energy by colliding with other electrons and atoms in the lattice structure. This allows them to pass on the thermal energy from an area of high temp to low temp

soft, slippery feeling

Graphite is a covalent layer lattice. Strong covalent bonds are present between carbons within layers. Between layers, only weak dispersion forces are present. As there are no strong covalent bonds holding the layers together, the layers can slide over each other easily, making graphite soft and slippery.

lower density than diamond

In graphite, carbon atoms are bonded in flat layers, with more space between layers due to weak intermolecular forces of dispersion. However diamond, has a tightly packed 3D structure, where atoms are closely bonded in all directions by covalent bonds. This more open structure in graphite results in fewer atoms per unit volume, so it has a lower density than diamond

insoluble

graphite has strong covalent bonds within its layers which are difficult to break. Most solvents cannot disrupt these strong bonds. Therefore, graphite is insoluble.

the strength of the covalent bonds in the structure of graphite cannot be overcome by the intermolecular bonds of a solvent like water

applications of graphite

• carbon brushes in electrical motors

  graphite is able to conduct electricity, and therefore can transfer current from a stationary wire to moving parts in electrical motors

• electrode in batteries

  graphite is an inert (unreactive) and electrically conductive material.

• industrial lubricant

  layers of graphite are able to slide over each other, which, when used as a lubricant, reduces friction in machinery

graphene

• single layer of graphite.

• known as carbon nanomaterial

• similar but slightly different properties to graphite because it is only a single layer.

• delocalised electrons move incredibly fast when voltage is applied, making it a very good conductor of electricity.

properties of Graphene

• thin and light

• high thermal conductivity

• high melting point

• electrical conductivity

• flexible

thin and light

graphene consists of a single-layered network of carbon atoms. This one atom thick structure make it extremely thin, with minimal mass.

high thermal conductivity

Only three bonds are formed, similar to graphite. the movement of delocalised electrons increases as heat energy is applied. This results in electrons moving quicker, bumping into other electrons and thereby conducting thermal energy

high melting point

Graphene has a giant covalent strcutre with each carbon atom bonded to three others. Breaking these strong covalent bonds across the entire sheet requires a large amount of energy.

electrical conductivity

Each carbon is only bonded to three other carbons, leaving one delocalised electron. these electrons are free to move across the 2D sheet, allowing charge to flow easily.

flexible

since graphene is very thin, it is easy to bend. Additionally because of the strength of the covalent bonds within graphene, it is difficult to break graphene through bending. Therefore, graphene is very flexible.

applications of graphene

• solar cells

  graphene’s flexibility could produce a breakthrough in the solar industry, with solar cells being able to be applied to any surface, including curved ones

• filtration

  graphene membranes have been shown to be an effective filter

• drug delivery in medicine

  graphene- based carriers of drugs for cancer therapy have been shown to target cancer cells more effectively.

amorphous carbon

• can be formed from the burning wood and other plant matter when air ( oxygen gas) is not plentiful

• known as carbon black (soot)

• usually does not have a consistent structure and can be used as a fuel

Covalent bonding

forces of attraction formed when one or more pairs of electrons are shared between two nuclei.

Occurs between non metals

Forming bonds to achieve 8 electrons in its outer shell(octet rule)

octect rule

• states that elements lose or gain electrons to achieve noble gas configuration ie 8 electrons in the valence shell.

• Non-metallic atoms tend to have high electronegativities (they attract electrons easily).

• Have four or more outer shell electrons other than H.

• When two nonmetals react, both need to gain electrons to complete their outer shells.

• Pairs of electrons are shared to achieve a stable state (8 electrons in the outer shell)

ionic vs covalent

Ionic bonding: transfer of electrons between nonmetals and metals  

 

Covalent: sharing of electrons between non-metals and non-metals  

Molecule

A neutral, discrete group of atoms bonded together in a fixed ratio (of known formula) by covalent bonds. Molecules contain no free-moving charged particles, making them unable to conduct electricity in solid or liquid form.

compound

Feature	Molecule	Compound
Definition	A group of two or more atoms chemically bonded together.	A substance made of two or more different elements chemically bonded together in a fixed ratio.
Bond type	Usually covalent, but can include ionic or metallic in broader definitions.	Can be covalent or ionic.
Element types	Can be made of the same element (e.g. O₂, N₂) or different ones.	Always made of different elements (e.g. NaCl, CO₂).
Example	O₂ (oxygen molecule), H₂O (water)	NaCl (sodium chloride), CO₂ (carbon dioxide)

if a covalent bond - molecule and compound

everything else is just a compound

valency

• determines number of covalent bonds formed  

• Only the valence electrons are shared 

• The electrons that are shared are attracted to nucleus of both atoms  

 

• There may be other valence electrons that are not involved in any covalent bonds = non-bonding electrons = lone pairs  

no of bonds

Single bonds  

• Each atom shares one electron  

• One pair (2 electrons) are shared between atoms 

 

Double bonds  

• Each atom shares two electrons  

• Two pairs of electrons (4 in total) are shared between atoms  

 

 

Triple bond  

• Each atom shares three electrons  

• Three pairs of electrons (6 in total) are shared between atoms  

Lewis structure

• Can be used to show how electrons are shared 

• Nucleus and inner shell electrons are replaced by atom symbol  

 

types of molecules

• Diatomic molecules: molecules that contain two atoms  

• Two of the same: covalent molecular elements, they are made up of identical atoms held together by covalent bonds  

• Polyatomic molecules- molecules that contain more than two atoms  

• Covalent molecular compounds – they are made up of different elements held together by covalent bonds  

inter vs intra molecular bonding

ce, water, water vapour  

Are all still h20 molecules  

H20 has covalent bonding, this bonding is not disrupted as water changes states, intramolecular bonding  

 

Intermolecular bonding is disrupted.  

 

 

Intra 

• Bonds Within the molecule 

• e.g covalent bonds  

Inter  

• Bonds between molecules 

• e.g h bonding  

• Responsible for molecules states 

• e.g bonds between two water molecules  

 

shapes of molecules

• Shapes of molecuels affect how it interacts with other molecules 

• Shape affects properties such as melting point, boiling point, hardness and solubility  

bonding and non bonding electrons

Atoms in covalent molecules are most stable when they have eight electrons in their valence shell (octet rule)  

 

These eight electrons are arranged into four pairs of electrons and can be  

Bonding: if shared between two atoms, forming a covalent bond  

Lone pair: non-bonding electrons that are not shared, belongs to a single atom  

vesper theory

The shape of a molecule will depend on both the position of the covalent bonds and the lone pairs

Vsepr: the valence shell electron pair repulsion: states that Negatively charged electron pairs in the outer shell repel each other, and are arranged as far away from each other as possible

The 3d shape of a molecule : distance between pair of electorns is maximised

Both bonding and lone pair electrons are treated the same way to determine shape Lone pair is ignored when naming the shape

The lone pair occupies slightly more space than the bonding electorns The covalent bonds are pushed closer together

electronegativity

• Polarity is dependent on electronegativity  

• Electronegativity is the tendency of an atom to attract electrons to itself  

• Determines the way electrons are distributed in a molecule  

• Increases from left to right and bottom to top in the periodic table: due to more protons in nucleus, a greater electrostatic attraction

polar and non polar molecules

• Polarity describes the electrical charge around a molecule  

• Polar molecules have an uneven distribution of electron density (charge) 

• Distinct positive and negative poles with molecule (partial charges)  

 

• Non poplar molecules  

• Even distribution of electron density (charge) 

• No significant charge separation  

• Molecule is symmetrical in shape  

polar

If the two atoms have different electronegativity 

The electrons will stay closer to the most electronegative atom 

 

The imbalance in the electron distributes results in a region of the molecule being partial negative and the other partial positive 

Molecules with an imbalanced electron distribution are said to be polar  

Non polar

When two atoms from covalent bond, they compete for the electrons being shared 

If the two atoms have the same electronegativity  

Then we consider the electrons to be shared equally, and the bond is said to be nonpolar  

dipole

When atoms in a molecule share electrons unequally this is known as dipole (two oppositely charged poles at each end).  

Hf has permanent dipole (electrons always spend more time on F) due to the different electronegativities of hydrogen and fluorine  

polarity

Polarity   The difference between the two molecules electronegativities  is more than0.5 then its polar   0.5 - 2

The polarity depends on the difference between the electronegativities of the two atoms     The larger the difference in electronegativity, the larger the dipole and the greater the polarity of the bond

intermolecular forces

The forces of attraction between molecules  

They hold molecules together  

100x weaker than intramolecular forces e.g covalent bond s 

Responsible for physical state of the substance ( solid, liquid or gas)  

The stronger the intermolecular forces the higher the melting points  

 

 

dipole dipole

Present when molecules are polar  

Occur between partially positive dipole on one molecule and partially negative dipole on another molecule  

 

 

• Relatively weak. The more polar a molecule is the stronger the forces are  

• The stronger the dipole dipole force the higher the melting and boiling points of the substance  

hydrogen bonding

Only occurs between polar molecules in which a h atom is covalently bound to an N, O or F  

The N,O,F are highly electronegative and attract the electrons  

Creating a significant positive charge over the H atom  

 

The partial positively charged h is then attracted to the lone pairs of electrons of the N,O,or F of neighbouring molecules  

 

 

• Approximately ten times stronger than a dipole dipole bond  

• But about 1/10 the strength of an ionic or covalent bond  

• Their presence results in high melting and boiling points  

• Two key requirements for hydrogen bonding  

• A hydrogent atom covalently bonded to an O,N or F atom  

• A lone pair of electrons  on o n or f atom of a neighbouring molecules  

dispersion forces

Occurs between all atoms/ molecules, including single atom (such as noble gas), both polar and non-polar molecules 

 

Caused by temporary dipoles or instantaneous dipoles which are result of random movement of the electrons surrounding the molecule 

 

Electrons are constantly moving in a molecule. Nonpolar molecules, electrons spend an equal amount of time around each atom. Occasionally electrons gather more closely together at one end of the molecule causing one end to become negative and the other end to become positive. This is known as a temporary dipole 

 

These temorary dipoles can then induce (create) dipoles in the neighbouring molecules  

Two factors influence the strength of dispersion forces  

1. The number of electrons  

• The more electrons the stronger the dispersion force ( larger atoms / larger molecules)  

2. The shape of the molecule 

Shape affects how closely the molecules can get to each other. Molecules that form long chains will have stronger dispersion forces than compact molecules