Level 2 Bonding and Structure etc.

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• Large networks of non metal atoms held together by covalent bonds.Can be in linear chains, two dimensional layers or three dimensional networks. 

1. Allotropes

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   • Two or more different physical forms of the same element of carbon and some carbon compounds are commonly used to demonstrate the structure bonding an properties of covalent networks.  
   1. Structure Bonding and properties of 2D networks

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      • Molecular solid  Amde up of Molecules   Covalent Bonds  Strong bonds Soft- not hard  Brittle  Not Ductile  Insulator- does not conduct  High melting point  Insoluble- due to strong bonds 
      1. High Melting Point

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         • 2 dimensional networks have high melting points. Melting a 2 dimensional network solid involves breaking the strong covalent bonds that hold atoms together. 
         1. Conductive

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            • 2Dnetworks are electrical conductors as each atom bonds to 3 other atoms.Therefore each carbon atom provides a deloicalised electron that is able to carry a charge through the solid 
            1. Insoluble

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               • Both one dimensional and two dimensional networks are large uncharged molecules and have weak attractions to water molecules so they are insoluble in polar solvents. 
               1. Soft and Slippery

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                  • Eg Graphite has a soft and slippery feel because sheets of graphite can move over each other due to the weak intermolecular forces holding the sheets together
   2. Structure and bonding of 3D Covalent Networks

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      • Diamond (an allotrope of carbon) and silicon dioxide are two common examples of 3D dimensional covalent networks. They have similar properties due to their bonding and structure. 
      1. High melting point

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         • High Melting points:- Strong covalent bonds holding the atoms together so high heat is required. 
         1. Not Conductive

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            • There are no free charged particles within the network to conduct therefore it is an electrical insulator 
            1. Insoluble

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               • -Insoluble because the forces of attarction between the atoms of the molecule are greater than that between the atoms and solvent molecules. 
               1. Hard

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                  • They are very hard due to the 3 dimensional tetrahedral arangement of the atoms covalently bonded to each other. 
      2. Fullerenes

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         • Carbon atoms are bonded to 3 others. Delocalised electron. Relatively recent discovery. Forms a spherical ball shape. 
         1. Melting Point

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            • Are extremely stable and are able to withstand high temperatures and pressures 
            1. Electrically Conductive

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               • Carbon Atoms are covalently bonded to 3 others leaving a delocalised electron that is able to carry an electrical current. Their hollow structure allows other atoms to be house dwithin them (called doping). This has many potential uses including using them as low temepratre super conductors. 
               1. Soft and Slippery

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                  • In a solid form they are discrete molecules that are able to move past neighbouring molecules. 
                  1. Applications

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                     • Potential for being used as lubricants, catalysts,In fuel cells and super computers. 
   3. 1D Covalent Netwtorks (Linear Chains)

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      • Carbon polymers form long linear chains.  Eg polyethene molecules 
      1. Low Melting Point

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         • Neighbouring chains are held together with weak intermolecular forces. Only a small amount of energy is required to seperate the chains and therefore change their state.  However the longer the carbon chain the greater the forces of attraction therefore melting point increases as chain length increases 
         1. Not condictive

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            • 1D carbon chains are electricl insulators as there are no free charged particles in the molecule
            1. Insoluble

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               • Both one dimensional and two dimensional networks are large uncharged molecules and have weak attractions to water molecules so they are insoluble in polar solvents. 
               1. Flexible

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                  • eg plastic. The linear chains are able to slide past each other

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• -Metals can be found on the left hand side of the periodic table. -some examples include iron aluminium, and copper -All metals have partially fixed shells. -Metals have a lot of use in industry and society due to their useful properties. 

1. Structure

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   • All metals share a smiliar structure Metals are poitively chatged as they contain  Positively charged metal nuclei are arranged into a lattice-The nuclei form a regular pattern in 3 dimensions. -Each metal atom provides one or more valence electrons that are delocalised: free to move throughout the lattice -The delocalised electrons are attracted to the positively charged nuclei. 
   1. Metallic Bonding

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      • Metallic bonding is strong electrostatic attraction between positively charged nuclei and the 'sea' of delocalised valence eletrons that surround them The bonding is desribed as non directional as the nuclei and delocalised valence electrons experience attractive forces in all directions.
      1. Conductivity

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         •  Because the valence electrons are free to move throughout the metal lattice they are able to carry an electrical current.  Therefore all metals have high electrical conductivity. Metals are able to condict in solid and liquid form. Copper is used for electrical wiring in housesAluminium is used to costruct high voltage cables. gold is used for electrical connectors in electronic devices such as cellphones and computers. 
         1. Malleability

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            • Describes how metals are able to be shaped using force. Unlike ionic substances that shatter with sufficient force, particles within metals are able to move past one another and remain strongly bonded together. When the metal nuclei are forced into a new position they still expereince strong electrostatic bonding to the 'sea' of valence electrons due to the non-directional nature of the bonds. 
            1. Ductility

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               • Describes how metals are able to be drawn into thin wires 
               1. Hardness

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                  • Hardness is the property of a metal that gives it the ability to resist being permanently deformed (bent, having its shape change or being broken) when a force is applied.
                  1. Solubility

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                     • Metals are insoluble in any type of solvent. The metallic bond is too strong to be overcome by the attractive forces between the solvent and the particles that make up a metal. 
                     1. Melting points

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                        • Metals have relatively high melting points due to the nature of the metallic bond. There are strong forces of attraction between the delocalised electrons and the positive nuclei of the atoms. These forces of attraction are difficult to overcome.   Eg: Aluminium 660       Copper 1083       Silver 961       Gold 1063        Nickel 1455

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• *Molecular substances consist of 2 or more atoms joined by covalent bonds. *Almost always made up of non metal atoms, found on the right hand side of the periodic table, eg Nitrogen, Carbon, Oxygen etc.  

1. Covalent Bonds

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   • *A  chemical bond formed when 2 non-metal atoms share a pair of electrons. *Convalent intermolecular bonds are very strong bonds that hold the atoms together within a molecule. *Arise from the sharing of a pair or pairs of valence electrons in order to achieve a full valence shell.
2. Shapes of Molecules

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   • Lone pairs and covalent bonds are made up of electrons. These negatively charged areas repel each other and will move around until they find an orientation with equally repulsive forces from adjacent electrons. Therefore the shape of a molecule is governed by the arrangement of the electron pairs around the central atom. You must consider both the bonded and unbonded pairs of electrons. There are 6 basic shapes. 
   1. Tetrahedral

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      • 4 regions of negative charge around the central atom. All regions bonded. Bond angle of 109°
      1. Trigonal Planar

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         • 3 regions of negative charge around the central atom. All regions bonded. Bond angle of 120° 
         1. Linear

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            • 2 regions of negative charge arund the cnetral atom. All regions are bonded. Bond angle of 180°
            1. Bent (3 regions)

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               • 3 regions of negative charge around the central atom. 2 bonded, 1 unbonded. Bond angle of 120°
               1. Bent (4 regions)

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                  • 4 regions of negative charge around the central atom.2 bonded, 2 undonded. Bond angle of 109° 
                  1. TrigonalPyramidal

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                     • 4 regions of  negative charge around the central atom. 3 bonded, 1 unbonded. Bond angle of 109°
   2. Determining the shape
      1. Lews Diagrams

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         • Show only the valence electrons that occurr in that atom. Remebering that the number of valence electrons is deterined from the elements atomic number. 
3. Electronegativity

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   • A measure of an atoms tnedency to attrcat electrons to itself. Metal atoms have low electronegativity Electon in group 17 have high electronegativity and strongly attrat their valence and bonding electrons. 
   1. Polar and Non polar bonds

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      • Polar covalent bonds are a covalent bond between 2 stoms at different electronagativity, The electrons in the bond ar enot shared equally between the atoms.  Non polar covalent bonds are a covalent bond between two atoms of the same electronegativity. he electrons in the bond are shared equally between the atoms. 
      1. Polarity of Molecules

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         • The polarity of molecules depends on wether the molecule contains polar bonds and the arrangement of these bonds  Carbon dioxide is a molecule with a polar C-O bond due to the difference in electronegativity between different C and O atoms. However because this molecule is linear and symmetrical the effect of these polar bonds cancel out and the molecule is therefore non polar 
         1. Solubility

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            • The rule 'Like dissolves like' describes the fact that Polar molecules will dissolve in polar solvents an non polar molecules will dissolve in Non polar solvents. Insoluble molecules are immiscible meaning that they are unable to form one continous mixture when mixed  
            1. Melting and Boiling points

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               • Covalent Bonds occuring inside molecules are intramolecular. Th eattractive forces between neighbouring molecules are called intermolecular forces.These are relatively weak and so they only need a small amount of heat energy input to break the weak bonds holding molecules together. For this reason Molecules have relatively low melting and boiling points.  
               1. As Insulators

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                  • -Moleculare substances are electrical insulators They consist of atoms bonded tightly by covalent bonds.  There are no free ions or electrons available to carry/ conduct an electrical current. 

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• *Positive ions are called Cations  *Negative ions are called anions. *When anions bond with cations this creates an ionic compound. *The ions within these are arranged in a 3D lattice with each cation surrounded by anions and vice versa. *The lattice is held together by forces of attraction between the + and - called an electrostatic force. 

1. Melting Points

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   • Ionic substances have high melting points. The strong electrostatic forces holding anions and cations together are not easily overcome. They require a lot of heat energy to break the lattice. The same goes for ionic boiling points. 
   1. Solubility

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      • Ions are charged particles that are attracted elecyrostaticly to polar solvents, most commonly water. The negative charged region of the solvent will be attracted to the cations of the Ionic molecule and vice versa with Anions. Solvents that are non polar will not dissolve Ionic substances. Substances with a very high lattice energy 
      1. Conductivity

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         • *Ionic substances are able to conduct electricity only when the ions from which they are made are free to move. *As solids the ions are fixed in place. *Dissolving an ionic substance in a solvent like water frees the ions *Meaning Solutions of ionic substances will condict electricity.
         1. Hardness

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            • Ionic compounds are very hard due to their closely-packed 3D lattice structure and relatively strong ionic bonds between the charged ions. Diamond and Aluminium Oxide are examples of very strong Ionic Compounds.  
            1. Brittleness

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               • Ionic solids are brittle meaning they do not deform when a force is applied to them .They maintain their shape until the force exceeds the strength of the ionic bonds holding the ions in place. Then they shatter. 
      2. Lattice energy

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         • *The amount of energy needed to break the lattice structure is called Lattice energy.  *As the lattice energy increases the solubility ionic compound decreases  *Because amount of energy required to seperate the ions will be greater. This means some ionic substances are insoluble.  
2. Polyatomic Ions

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   • Ions made of more than one atom are called plyatomic ions.  Know the following-  NH4+ = ammonium OH- =Hydroxide NO3- = NitrateHCO3- = Hydrogen Carbonate SO42- = Sulfate CO32- = Carbonate 
3. Ionic Bonds

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   • A chemical bond found in Ionic substances that is formed by the electrostatic attraction between oppositely charged ions (Cations and Anions)
   • Cation, usually a metal atom that has lost one or more electrons has a positive charge.  Anion, usually a non metal atom or group of atoms that Cl2 + 2Na= 2NaCl

Annotations:

• Chemical reactions involve the breaking and formation fo chemical bonds. The general rule is that energy is Absorbed in the breaking of chemical bonds and energy is released when bonds are formed. 

1. Exothermic Reactions

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   • Reactions that release heat. the energy required tobreak the bonds in the reactants is exceeded by thr energy released when chemical bonds are formed in the product(s) The reaction between Hydrogen Gas and oxygen is exothermic . Energy is absorbed to break the 2 hydrogen binds linking hydrogen atoms and the 2 bond linking the oxygen atoms. Energy is released when the water molecules form. 4 covalent bonds forming between Hydrogen and oxygen atoms release more energy than is required to break the bonds in the reactants. This excess energy is released into the environment. 
   1. Endothermic Reactions

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      • Can be considered the opposite to exothermic . They absorb energy from the environemtn in which they happen to break bonds. More energy is required to break that bonds in the water molecules than is released when the reactants form. The excess energy is 'stored' in the chemical bonds of the reactants and can be released when hydrogen is burned as a fuel.  
      1. Enthalpy

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         • A term used to describe the energy within a chemical system.  It cannot be measured directly so we instead discuss 'energy changes' within chemical reactions. Energy diagrams show the change in enthalpy for a specific chemical reaction. 
         1. Breaking Bonds

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            • When a substance changes its state its because chemical bonds between the particles of the substance have been either weakened or strengthened. Melting and evaporating are endothermic as energy is absorbed to weaken the bonds.  Condensing and solidifying are exothermic processes as energy is released to strengthen the bonds.  
            1. Bond Enhalpy

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               •  The energy needed to break the covalent bonds in gaseous molecules in known as bond enthalpy. 
               1. Mass , Moles and Molar mass

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                  • n=m/M       M=m/n       m= nxM n= no. of moles (mol) m= mass (g) M=Molar mass (g mol-1)These formula can be used to calculat ethe number of moles in a known mass of a substance Eg: ------------>
                  • Calculate the number of moles in 128g of Methane (CH4)  M(CH4)= 1C (12.0) + 4H(4) = 16.0 gmol-1  n=m/M n= 128/16 n=8 moles  

Annotations:

• Atoms of a chemical element can exist in different forms. These are called isotopes. They have the same atomic number but a different mass number. Same number of protons and electrons but different number of neutrons.