Module 2 Chemistry OCR

Module 2 Chemistry OCR

1. SHELLS

   • 1s shell -2 elec
     • 2nd shell- 8 elec
       • 3rd shell-18 elec
         • 4th shell- 32 elec
   • sub-shells
     • S- 2 elec
       • P- 6 elec
         • D- 10 elec
           • F- 14 elec
     • ORBITAL - can hold up to 2 electrons, each sub-shell has different orbitals s,p,d,f
       • s orbital has a spherical shape
       • P orbital has a 3D dumbbell shape at right angles to one another
       • orbits are regions around the nucleus that have electron density
       • electron repel each other and have a property called a spin
         • in orbitals they have opposite spins to counteract the repulsion
     • Exceptions
       • Cr = 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 5d5
         • Cu = 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 2d10
   • carbon- 12= international standard
     • the mass of carbon-12 is exactly 12
     • the mass of 1/12 of an atom of carbon 12 is exactly 1
     • RELATIVE ISOTOPIC MASS- the mass of an atom of an isotope compared to 1/12 of the mass of an atom of carbon-12
       • RELATIVE ATOMIC MASS- is the weighted average mass of an atom of an element compared with 1/12 of the mass of am atom of carbon-12
       • RELATIVE MOLECULAR MASS- the weighted average mass of an molecule of a compound compared with 1/12 of the mass of an atom of carbon-12
         • ISOTOPES-atoms of the same element with different masses with the same number of protons and neutrons
         • RELATIVE FORMULA MASS- the average mass of the formula unit of a compound compared with 1/12 the mass of an atom of carbon-12
   • Mass spectrometer = Device to find out what samples are made up of by measuring mass of the components
     • RAM = (1st RIM x % abundance) + (2nd RIM x 2nd % abundance Total % abundance
     • The x-axis on a mass spectrometer graph is the mass/charge (m/z)
       • The y-axis is the relative abundance
     • If there are are only two bi-atomic isotopes and there are three lines, it is where the bi-atomic atom is made up of the two different isotopes
     • 1) Vaporise and ionise ions to form positive ions
       • 2) Accelerate ions.
         • 3) Deflect ions. Heavier ions will be more difficult to deflect so isotopes will be seprated
           • 4) Ions are detected on a mass spectrum

2. SHAPES OF MOLECULELS

   • the shape of the molecule depends on the number of electron pairs surrounding the central atom
     • electron pairs repel each other and move as far apart as possible
       • Lone pairs of electrons repel more than bonded pairs
   • molecules with bonded pairs
     • 2- linear
       • 3- trigonal planar
         • 4- tetrahedral
           • CH4
           • 109.5
           • 5- Trigonal bipyramid
             • 6- ocrahedral
               • SF6
               • 90
             • PCl5
             • 120 & 90
         • BF3
         • 120
       • BeCl2
       • 180
   • molecules with lone pairs
     • 0- tetrahedral
       • 1- pyramidal
         • 2- non-linear
           • H20
           • 104.5
         • NH3
         • 107
       • CH4
       • 109.5

3. BONDING STRUCTURE AND PERIOCIC TABLE

   • ionic
     • electrostatic attraction between oppositely charged ions (metals and non-metals)
       • ionic lattices
         • structure compromising of hundreds of thousands of ions
           • an ion is surrounded by oppositely charged ions, forming a giant ionic lattice
     • properties of ionic compounds
       • high melting and boiling point due to lot of energy needed to break the strong electrostatic forces holding the ions rigidly in the solid lattice
       • electrical conductivity- (solid lattice)- ions are in a fixed position and there are no MOBILE charge carriers so ionic compounds are a non conductor of electricity in the solid state. When melted or dissolved in water the solid lattice breaks down and the ions are free to move as mobile charge carriers therefore ionic compounds are electricity conductors in a liquid or aqueous state
       • An ionic lattice often dissolves in polar solvents (e.g) water
   • Dative covalent bond
     • e.g- and ammonium ion NH4+ (containing3 covalent bonds and one dative bond
     • When the shared pair of electrons are donated by only one bonded atom
   • covalent
     • a shared pair of electrons between atoms of non metals
     • a molecule is the smallest part of a covalent compound that can take part in a chemical equation
     • an electron that takes part in forming a chemical bonds is called a valence electron
     • simple covalent
       • small molecules with weak intermolecular forces
       • low melting and boiling points due to less energy required to break the weak intermolecular forces
       • no mobile charged particles so the structures are non-conductors of electricity
       • van der waals forces form between a simple molecular structure and a non-polar solvent which weakens the structure are often soluble in non polar solvents
     • giant covalent
       • e.g- diamond, graphite and SiO2
       • high melting and boiling point due to lots of energy required to break the strong covalent bonds in the lattice
         • except fro graphite there are no mobile charged particles so they are non conductors
           • the strong covalent bonds in a lattice are too strong to be broken by polar or non-polar solvents
     • Electrostatic attraction between the bonded pair of electrons and the nucleus of the atom.
   • metallic
     • the electrostatic attraction between the positive metal ions and delocalised electrons
     • the delocalised electrons in metals are able to move throughout the structure can cannot assign to s positive metal ion

4. ELECTRONEGATIVITY AND POLARITY

   • ELECTRONEGATIVITY- the measure of the attraction of an atom in a molecule for the pair of electrons in a covalent bond
     • generally the smallest atoms are the most electronegative and the most electronegative atoms are those of highly reactive non-metallic elements- e.g.. O, Fe and Cl
   • Polar and non-polar molecules
     • polar bonds
       • the bonded electrons are shared equally between both atoms, and the bonded atoms have similar electonegativities
       • Sodium chloride can be dissolved in polar solvents (water molecules). Water molecules attract sodium and chloride ions which breaks down the ionic lattice. Sodium ions are attracted to the oxygen which is slightly negative. Chloride ions are attracted to the hydrogen of water which is slightly negative.
     • non-polar bonds
       • the bonded electrons are shared equally between both atoms and the bonded atoms have similar electronegativities
       • Simple molecular compounds dissolve in non-polar solvents. Intermolecular forces form between the molecules and the solvent. The interactions break the existing intermolecular forces and the compound dissolves. When a simple molecular substance is added to a polar solvent there is little interaction because the intermolecular bonding within the polar solvent is too strong to be broken.
     • symmetrical and unsymmetrical molecules
       • Dipole = Separation of opposite charges. The more electronegative atom will have a slightly negative charge as electrons are loser to it.
         • each bond is polar but the dipoles act in different directions, the overall effect of the dipoles to cancel out

5. INTERMOLECULAR FORCES

   • Van der waals forces (induced dipole-dipole)
     • weak intermolecular forces that act between all molecules
     • caused by the uneven distribution of electrons in molecules which inures a dipole in an neighbouring molecule which is induces more dipoles
     • van der waals forces increase in strength with increasing number of electrons
       • the greater the number of electrons in each molecule the larger the oscillating and induced dipoles. The greater the attractive forces between molecules and the greater the van der waals forces
   • Permanent dipole-dipole
     • the delta positive/negative charges on a polar molecule attract oppositely charged dipoles on another polar molecule
       • this gives a weak intermolecular force called a permanent dipole-dipole interaction- eg- HCL
         • stronger than van der waals forces although they are both weak
   • Hydrogen bonds
     • strong intermolecular bond between polar molecules
     • special properties of water due to hydrogen bonds are -that ice is less dense than liquid water this is because particles in solids are usually packed closely together than in liquids however hydrogen bonds hold water molecules apart in an open lattice structure -ice has a high melting point and water has a high billing point because of the strong hydrogen bonds between water molecules
     • Form between hydrogen atom and lone pair of electrons in electronegative atom. NOF & H