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Elemental and Environmental Chemistry

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Electronic Configuration
- Each electron in an atom or a monoatomic ion has the potential energy which originates from the attraction between its negative and positive charges in the nucleus
  - Electrons have energy values unique to that element
    - Each allowed energy level for an electron is represented by a main shell using numbers 1,2,3,4,5,6 and lowercase letters s,p,d,f.
      - In the most stable, ground state, the electrons occupy subshells with the lowest available energy levels , they are allocated inn order of increasing energy levels
        1s<2s<2p>3s>3p<4s<3d<4p<5s
        Each subshell can accommodate a maximum number of electrons. S, 2, p, 6, d, 10, f, 14
        The number of electrons is equal to the atomic number of the element
        Two electron configurations that don't conform to the general rule within the first 38 elements of the periodic table are chromium and copper atoms.
        Positive ions of the main group (1-8) elements, have less electrons than the atom of the element, by the number equal to the numerical value of the charged ion
        Negative ions of the main group (1-8) of elements have more electrons than the atom of the elements by the number equal to he numerical value of the charged ion
        The number of the period in which an element is placed is equal to the highest numbered main shell that is occupied by electrons
        The number if electrons occupying the highest numbered shell (valence shell) determines the vertical column (group) to which a main grip element is assigned
        For any main group element, the group number equals the numbers of electron occupying its outer shell.
        For all elements in group 1, there is one electron in the highest energy shell: (Li)1s2>2s1, (Rb) 1s2>2s2>2p6>3s2>3p6>4s2>3d10>5s1
        For all elements in group 2, there are two highest energy electrons in an s subshell: (Mg) 1s2>2s2>2p6>3s2>3p6>4s2, (Ca) 1s2>2s2>2p6>3s2>3p6>4s2
        Groups 1 and 2 form the s block
        For all elements in group 3, there is one electron in the highest energy p subshell: (B) 1s2>2s2>2p1, (Ga), 1s2>2s2>2p6>3s2>3p6>4s2>3d10>4p1
        For all elemens in group 4, there are two electrons in the highest p subshell: (Si) 1s2>2s2>2p6>3s2>3p2, (Ge) 1s2>2s2>2p6>3s2>3p6>4s2>3d10>4p2
        For elements in groups 5-8 there are 3-6 electrons in the highest energy p subshell (He) being an exception at the top of group 8
        Groups 3-8 form the p block
        The transition elements, the highest energy electrons are in a d subshell (Ti) 1s2>2s2>2p6>3s2>3p6>4s2>3d2, (Co) 1s2>2s2>2p6>3s2>3p6>4s2>3d7
        The transition elements form the d block
        For the lanthanides and actinides, the highest energy electron are in the f subshell: (Sm) (lanthanide) 1s2>2s2>2p6>3s2>3p6>4s2>3d10>4p6>5s2>4d10>5p6>6s2>4f6
        The lanthanides and actinides form the f block
        Chemical Properties of Elements and the Periodic Table
        The elements in each group of the s or p blocks display similar chemical properties to each other.
        They react with each other elements and compounds forming compounds that conform to a common formula pattern
        E.g. All Group 1 elements react with Cl to form a chloride of a formula Mcl.
        They also react with water to form a hydroxide of formula MOH and Hydrogen H+
        Each element in Group 5 forms a compound wit hydrogen of formula type XH3.
        Electronic Configurations of the s and p block elements can be used as a basis for explaining and predicting their chemical properties
        The connection between the electronic configuration of an element and its position on the periodic table can be used to make predictions about the properties of an element, e.g. metal/metalloid/non-metal nature, the charge of its monatomic ion (s) and its likely oxidation state (s) in its compounds
        Atoms of metals lose their electrons in chemical reactions, non metals gain or share electrons in chemical reactions and metalloids lose or share their electrons in chemical reactions
        The similarity of chemical properties within each particular group is explained in terms of the similarities of their electronic configurations
        The electronic configurations of the resulting ions are more stable than the configurations of the atoms from which they have been formed
        Similarly, when atoms share electrons they acquire more stable electronic configurations
        The Octet Ru;e For period 1 and 2 elements an electronic configuration in which the valence shell is 'complete' with the maximum number of electrons is more stable than a configuration with an incomplete valence shell
        E.g. Mg 2+ 1s2>2s2>2p6 has a more stable configuration than the Mg atom 1s2>2s2>2p6>3s2.
        These complete stable electronic configurations are the same of the noble gases of periods 1 and 2 and are often referred to as 'noble gas configurations'
        For period 2 and 3 noble gases, the electronic configurations show 8 electrons in the valence shell and hen atoms attain 8 electrons in their valence shell, they're said to have conformed to the 'octet rule'
        Expansion of the Octet: Atoms of period 3 elements from the p block often conform to the octet rule by accepting electrons to form negatively charged monoatomic ions or by sharing electrons with other atoms. E.g. S (atom) 1s2>2s2>3s2>3p4, forms the sulphide ion s2- 1s2>2s2>2p6>3s2>3p6
        Atom of the period 3 elements form groups 5-7 can share all of their valence shell electrons and as a consequence acquire more than 8 electrons in their valence shells.
        The extra electrons above the octet are accommodated in the previously unoccupied 3d subshell, this is referred to as the expansion of the octet
        Valence Electrons: For s block elements, the valence shell electrons are the highest energy s subshell electrons and for p block elements, the valence shell electrons are the 'outer' s and p subshel electrons.
        Chemical Reactions Involving the s-block Elements: Group 1 elements readily form comp[ounds by reacting with non metal compounds such as halogens, oxygen and sulfur and other oxidising agents such as water and the hydronium ion present in dilute acid solutions.
        In these reactions, the Group 1 elements lose their s1 valence shell electrons to for an M+ ion.The electron lost is gained by the other reactant. The product compounds are ionic with formulae such as MCl, M2S, M2O and MOH. The electronic configuration of the M+ ion is more stable than the M atom
        The K atom has the configuration 1s2>2s2>2p6>3s2>3p6>4s2. The outer shell for this ion is now the 3rd shell and it's complete. This ion conforms to the octet rule.
        The charge on the monatomic ions of the Group 1 elements is always, therefore theoxidation state of the Group 1 elements in their compounds is always +1, they are classified as metals, as they lose their electrons in chemicals reactions.
        Examples of this: 2Na + Cl2 >>>2NaCl 2K + S >>>K2S
        Group 2 elements: Also readily form compounds by reacting with non-metal elements, in these reactions, the atoms of the Group 2 elements lose their valence electrons to form an M2+ ion. The electrons lost are gained by the other reactant, to form product compounds that are ionic with formulae such as MCl2, MS, MO and M(OH)2. The M2+ ion is more stable than the M atom
        The charge of the monoatomic ions of the Group 2 elements is always 2+, therefore the oxidation state of the Group 2 elements in their compounds is always +2. Except for Be, the Group 2 elements are classified as metals because their atoms lose electrons in chemical reactions. Be is classified as a metalloid.
        For Example: 2Ca+ O2 >>> 2CaO, Mg + S >>> MgS
        Chemical Reaction Involving the p-block Elements: Group 3 elements, in their compounds- wither exhibit a covalence of 3 as in the case of B in its compounds, e.g. BCl3. When B forms compounds, its atoms share s2p1 valence shell electrons to form covalent bonds with other non-metal atoms.
        The 'Covalence' of an element is equal to the number of electrons that its atoms share when forming covalent bonds with other atoms. When B shares its 3 valence electrons with, e.g., electrons from three chlorine atoms, then B is exhibiting a covalence of 3.
        Or they exist as triple positive ions, Al3+, in compounds with non-metals. These ions are formed when atoms of the Group 3 elements lose their s2p1 valence shell electrons in electron transfer reactions.
        The charge of the monoatomic ions of the Group 3 elements is usually 3+, the oxidation state of he Group 3 elements is usually 3+, in compounds their oxidation state is 3-. The elements range from non-metal B at the top of the group, to metalloids Al and Ga, in the middle and metals at the bottom of the Groiup
        Group 4 elements in their elements- either exhibit a covalence of 4, e.g. C and Si in their compounds, e.g. CCl4 and SiH4. When C and Si form compounds, their atoms share the s2p2 valence shell electrons yto form covalence with other non-metal atoms.
        Or they exist as 2+ or 4+ ions, such as, Pb2+, Sn2+ or Pb4+, in compounds with non-metals. 4+ ions are formed when the atoms of Group 4 elements lose the s2p2 valence shell electrons in electron transfer reactions. 2+ ions are formed when the atoms of Group 4 elements lose only the p2 electrons of the s2p2 valence electrons in electron transfer reactions
        The charge on the monoatomic ions of the Group 4 elements is usually 2+ or 4+, the oxidation state of the Group 4 elements in their compounds is either +4, +2 or -4. The elements range from non-metal, C and Si, at the top of the group to metalloids for the rest of the Group.
        Group 5 elements, in their compounds- either exhibit a covalence of 3 , e.g. N in all of its compounds, NH3, and of P and As, in some of their compounds, e.g. AsCl3. In these compounds, N, P and As atoms share only only the p3 electrons from the s2p3 valence shell configuration to form covalent bonds with other non-metals. Sharing this way, means they are conforming to the 'Octet Rule'
        Or they exhibit a covalence of 5, e.g. P and As in some their compounds, e.g. AsCl5 and P4010. In these compounds, P and As atoms share all of the s2p3 electrons from the s2p3 valence shell configuration to form covalent bonds wit other non metal atoms. In sharing in this way they are expanding the octet.
        The charge on the monoatomic ions of the Group 5 elements is 3- The oxidation state of the Group 5 elements in their compounds is either +5, 3+ or 3-. The elements range from non-metal, N and P, at the top of the Group , to the metalloids As and Sb in the middle and the metal Bi at the bottom of the Group.
        Group 6 elements in their elements- either exhibit a covalence of 2, e.g. O in all of its compounds such as H2O, and of S and Se in some of their compounds e.g. SF2. In these compounds, these compounds only share their p4 electrons from the s2p4 valence shell configuration to form covalent bonds with other non-metal atom. In sharing this way they are conforming to the 'Octet Rule'.
        Or they exhibit a covalence of 4, as in the case of S and Se in some of their compounds, e.g. SO2 and SeF4. In these compounds the S and Se atoms share all of the p4 electrons from the s2p4 valence she'll configuration to form covalent bonds with other non metal atoms .In sharing this way they are 'Expanding the Octet'. Note: Oxygen does not have a covalence of 4.
        Or they exhibit a covalence of 6, e.g. S and Se in some of their compounds, SF6 and SeO3. In these compounds , the S and Se atoms share all of the s2p4 electrons from their valence shell configuration to form covalent bonds with other non-metal atoms. In sharing this way, they are 'Expanding the Octet'.Note: Oxygen does not exhibit a covalence of 6.
        Or they exist as 2- ions, e.g. o2- and S2- in compounds with metals. 20 ions are formed when the atoms of group 6 elements gain two electrons into the p subshell, changing the valence shell configuration from s2p4 to s2p6. The resulting ions conform to the 'Octet Rule'.
        The charge on the monoatomic ions of Group 6 elements is 2-, the oxidation state of the Group 6 elements in their compounds is +6,+4,+2 or -2. AN exception is O with an oxidation number of -1 in H2O2. The elements range from non-metals, O and S, at the top of the group, too the metalloids Se and Te in the middle and the metal polonium at the bottom of the Group.
        Group 7 elements in their compounds- either exhibit a covalence of 1 in their compounds, e.g. HBr and Ccl4. In these compounds, the Group 7 atoms share one of the p5 electrons from the s2p5 valence shell configuration to form covalent bonds with other non-metal atoms. In sharing this way they are conforming to the 'Octet Rule'
        Or they exhibit a covalence of 3,5, or 7 and in sharing in this way they are 'Exapanding the Octet'. Note: F does not exhibit a covalence of 3,5, or 7
        In compounds, the atoms of the Group 7 elements share electrons in the following ways: Covalence of 3: there of he p5 electrons are shared, Covalence of 5: all five of the p5 electrons are shared, Covalence of 7: all seven of the s2p5 electrons shared.
        Or they exist as 1- ions in compounds with teals. 1- ions are formed wen the atoms of Group 7 elements gain on electron into the p subshell, changing the valence shell configuration demo s2p5 to s2p5. The resulting ions conform to the 'Octet Rule'
        The charge on monoatomic ions of the Group 7 elements is 1-, the oxidations state in their compounds is +7, +5, +3, +1 or -1. The elements are all non-metals.
        Limitations of the Covalent Bonding Model. There are some examples of the s and p block elements exhibiting oxidation states (and covalences) that are different from those given in the summaries above. Nitrogen, oxidation state +4, e.g. NO2, Nitrogen, +2, NO and Chlorine +4, ClO2
        Electronegativities of the Elements: The ability of an atom to attract electrons to itself is called its electronegativity. The higher the electronegativity, the stronger the attraction for electrons. Metals have lower electronegativity values than metalloids, which have lower electronegativities than non-metals. Two trends of the s and p block elements are made apparent. Group 1- Group 7 electronegativities increase across each period and Period - Period 4 electronegativities decrease down each group.
The Acidic/Basic Nature of Oxides
- An oxide is a two-element compounds with one of the elements being O, e.g. CO (carbon monoxide), Fe2O3 (iron (3) oxide, and NO2 (nitrogen dioxide).
  - Oxides can be classified as acidic, amphoteric or basic depending on their reactivity, or lack of reactivity, with acids and bases.
    - Acidic Oxides: Acidic oxides react with OH- to produce oxyanions (negatively charged ions of the element combined with H and O, e.g. of oxyanions CO3 (2- (carbonate), SO4 (2-) (sulcate), AlO2 (-) (aluminate)
      - If soluble in H2O, acidic oxides react with H2O to form oxyacids (acids consisting of the element combined with H and O), e.g. of oxyacids H2CO3 (carbonic acids), H2SO4 (sulphuric acid) and H3PO4 (-) (orthophosphoric acid)
        These oxyacids consist of molecules with covalent OH- as part of their structure
        Acids undergo partial or complete ionisation with H2O to Produce H3O (+)
        Acidic oxides are the oxides of non-metals, oxides of metals with high electronegativities and are covalent molecular e.g. CO2 and SO3 or continuous covalent compounds e.g. SiO3
        Basic Oxides: They react with acids or H(+) ions to produce positively charged metal ions and H2O molecules, the solid appears to dissolve in the acid.
        If soluble in H2O, basic oxides react with H2O to form metal ions and OH- ions. Basic oxides are the oxides of metals, they are the oxides of metals with low electronegativity and are ionic compounds consisting of metal ions and O (2-) (oxide ions)
        Amphoteric Oxides: They dislay a basic character by reacting with acids (H+) ions to produce positively charged monoatomic ions and H2O. When reacting in this way, the solid oxides appear to dissolve in the acid.
        They also display an acidic character by reacting with Oh- to produce oxyanions and H2O. When reacting in this way, the solid oxides appear to dissolve in the OH- solution
        Molecular Substances: Compounds and elements consisting of molecules with 10 or less atoms per molecule (small molecule) e.g. CO2, SO2, H2S, Cl2, NH3,O3, CFCL3 and H2SO4
        Small molecules are formed when atoms of non-metal elements covalently bond to each other.
        H atoms can also form small molecules with atoms of the elements mention above.
        There is a small number of non-metal elements and non-metal-non-metal compounds that are non molecular, e.g. SiO2, C (diamond) and SiC (silicon carbide). These are continuous lattices of atoms bonded together by covalent bonds, 'Continuous Covalent Lattices'.
        Properties of Molecular Elements and Compounds: Elements and compounds usually contain small molecules which have low melting and boiling points and are usually bases or liquids at room temperature. Solids generally have melting points below 200 degrees Celsius. Molecular element and compounds are usually poor conductors of electricity in the solid, liquid or gaseous state.
        Chemical Bonding in Molecular Elements and Compounds: There are two types of bonding, Covalent bonding (intramolecular/primary bonds), opera rate within molecules, they hold atoms to each other in molecules
        They are formed when atoms share valence electrons, this bond arises from the electrostatic attraction between the shared electrons and the positively charged of the nuclei of the atoms. They're classified as strong bonds, and require a large amount of energy to break them
        Secondary interactions ((intermolecular interactions) are the forces of attraction that operate between neighbouring molecules, hold molecules to each other in liquid and solid states
        When molecular substances are vaporised, sufficient energy must be provided to the molecules in order to overcome the secondary interactions holding them together
        Boiling points of molecular points provide an indication of the strengths of secondary interactions, the stronger the secondary interactions the higher the temperature to which the molecules must be raised in order to overcome the secondary interactions and become independent molecules in the vapour phase
        Different types of Secondary Interactions: There are three types that can operate between molecules, dispersion forces, dipole-dipole interactions and hydrogen bonding. The type or combination of secondary interactions which operate for a particular molecular substance depends upon the polarity of its molecules.
        Two factors must be considered together when determining how polar a molecule is: the polarity of the primary bonds in the molecule and the spatial arrangement of the atoms in the molecule (the shape of the molecule).
        The Polarity of Covalent Bonds: The higher the electronegativity of an atom, the stronger its attraction for the shared electrons. A covalent one formed between two atoms of different elements, the atoms will have different electronegativities and so there ail be an unequal attraction for the shared electrons.
        The atoms of the more electronegative element attains a partial negative charge, as it aquires, by means of its stronger attraction, a greater share of the negatively charged electron pair than the other atom does
        On the other hand, the other atom attains an equal and opposite charge. The resulting bond is said to be polar, as here has been a charge separation as a result of unequal electron sharing.
        The greater the difference in electronegativities between two atoms sharing electrons, the greater the magnitude of the positive and negative charges attained by the atoms and the greater the polarity of the covalent bond.
        The polarity of the O-H bond is greater than the polarity of the H-Cl none because of the difference in electronegativities of the H and O is greater than the difference in the electronegativities of H and Cl, this in turn is a greater consequence of O having a greater electronegativity than Cl.
        Electronegativity Trends Within the Main Groups of The Periodic Table: Electronegativity increases across a period from Group 1- Group 7 and increases from the bottom to the top of a Group
        The Shapes of Molecules: The shape of a molecule can be described and predicted using the Valence Shell Electron Pair Repulsion theory
        Repulsion occurs between pairs of valence electrons in molecules, also bonding pairs of electrons, bonding and non-bonding (lone0 pairs of electrons, non-bonding pairs of electrons and lastly, single, double and triple pairs of electrons (single, double and triple)
        To predict or explain the shape of a molecule of a two-element compound it is necessary to establish, which is the central atom of the molecule, the number of bonding and non-bonding pairs of valence electrons around the central atom and the spatial positioning of these valence electrons as a result of repulsion
        To establish the central atom of a molecule of a binary compound, the covalence of the atom of the element has to be determined, which is indicated by the number f pairs of electrons that an atom of the element can share with neighbouring atoms in a molecule and is numerically equal to the magnitude of the oxidation number of the element
        The Polarity of Molecules, Diatomic Molecules: Consists of two atoms covalently bonded to each other e.g. N2, HF and NO. If the bonding in a diatomic molecule is polar, then the molecule itself is polar. Conversely, id the bonding in the diatomic molecule is non-pole, then the molecule itself is non-polar.
        Two-element Polyatomic Molecules: Consists of three or more atoms, some common two-element e.g CO2, NH2, PCl3, CH4, SO3 and H2O
        Non-Polar Polyatomic Molecules: Fall into two distinct groups. Firstly, molecules in which the bonds are non polar e.g. P4, S8 AND O3
        The other group, much larger, with molecules in which bonds are polar, all have symmetrical spatial distribution of atoms about a central atom e.g. CO2, SO3, CH4 and BCl3
        Polar Polyatomic Molecules: Molecules with a spatial distribution of atoms about a central atom that results in the molecules having positive and negative ends, each molecule has only one angle of symmetry e.g. SO2, CHCl3, NH3 AND H2S
        Polarity of Three-Element Molecules: Involves the comparison of the electronegativities of the elements that make up the molecule, and the reactive polarities of the bonds.
        Polyatomic Molecules With Polar and Non-Polar Groups: Usually the non-polar part is a hydrocarbon chain or ring with the polar part being a functional group. Whether the molecule behaves overall as polar or non-olar depends on the size of the hydrocarbon part relative to the number and polarity of the functional groups. E.g. ethanol, polar, hep tan-1.amine, non-polar, benzoic acid, non-polar and ethanal, polar.
        Secondary Interactions- Dispersion Forces
        Electrons in molecules are constantly moving , as a result of the continuous changing of positions of the electrons, a momentary non-symentrical distribution of charge can arise, forming a 'temporary dipole
        This temporary dipole can induce a dipole in a neighbouring molecule. Neighboring dipoles align in such as way that electrostatic forces of attraction, 'dispersion forces, can operate between them
        Dispersion Forces and Molar Mass: The strength of these dispersion forces increases as the number of electrons in each molecule increases. With a greater number of electrons, there are larger positive and negative charges formed on the dipoles as more electrons can move from one end of a molecule. Larger positive and negative charges result in stronger forces of attraction between the molecules.
        For non-polar substances, the increase in boiling points as molar mass can be explained in terms of the increased strength of dispersion forces resulting from the larger positive and negative charges on the temporary charges on the temporary and induced dipoles. Molecular shape also affects the strength of dispersion forces.
        The greater the molar mass the greater the number of electrons per molecule.
        For non-polar molecules belonging to the same structural class, boiling points (and hence strengths of secondary interactions) increase as molar molar mass.
        Another factor affecting the strength of dispersion forces between non-polar hydrocarbon chains is the presence of double bonds in the chain. The greater the number of double bonds, the weaker the dispersion forces, e.g. triglycerides that are oils at 25 degrees Celsius, whereas saturated fats are solids.
        Secondary Interactions Between Polar Molecules: Permanent electric dipoles with clearly defined positive and negative wends, with secondary interaction that comprise of both dispersion forces and electrostatic attraction between the positive and negative charges of the permanent dipoles.
        For molecules of similar molar mass and shape, the dispersion forces are approximately the same strength. Generally, the additional presence of dipole-dipole increases the overall strength of the secondary forces with a consequent increase in the boiling point of the substance.
        Hydrogen Bonding: Some molecular substances have much higher boiling points than would be expected for their relatively small molar masses.
        Compounds with high boiling points have either H-F,O-H or N-H bonds as part of their structure, in each of their molecules, there is a hydrogen atom covalently bonded to an atom of one of the three most electronegative elements, F, O and N. The very strong secondary interactions operating in these compounds is called 'Hydrogen Bonding'.
        O-H Bonds are found in H2O, OH- (Alcohols), COOH (Carboxylic Acids), Carbohydrates and Amino Acids, N-H Bonds in NH3, Amines, Amides, Amino Acids, Proteins and Synthetic Polyamides and C=O in COOH (carboxylic acids), Amides, Amino Acids, Aldehydes, Ketones, Proteins and Synthetic Polyamides
        Hydrogen Bonds influence the melting and boiling points of compounds, their miscibility with each other, their ability to absorb H2O, the tertiary structure of polymers and rigidity of polymer solids.
        The Strength of Hydrogen Bonds: They are the strongest dipole-dipole interaction, but are considerably weaker than covalent bonds.
        Reasons for strength in Hydrogen Bonding: the positive charge on the H atoms, when bonded to a F, O or N atom is higher than the positive charge on other atoms in polar molecules. A covalently bonded H atoms is the only atom that jas its positive nucleus unshielded from interactions with electrons of other atoms in neighbouring molecules. The small size of the H atoms allows it to come into close contact with the negatively charged atoms of other molecules.
        Cycles In Nature
        Small Molecules, Ch4, NH3, CO2 and inorganic ions such as NO3 (2-), NH4 (+), CO3 (2-), and SO4 (2-) are produced an consumed in the atmosphere and biosphere nu chemical reactions that occur in interconnected cycles, named afternoon-metal elements, C, N, S and P.
        Organic compounds are derived from living organisms
        C, N, S and P are main components of living things, their waste products and many products derived from living things e.g. carbohydrates, proteins, fats and nucleic acids.
        When living things die, their remains either undergo natural decay, plant mater being mulched, or they are burt, tree being burnt after land clearance.
        Natural decay processes are chemical reactions that occur within the cells of bacteria and other microbes. Compounds comprised pf complex organic molecules such as polymers are broken down in compounds comprised f simpler, smaller molecules and ions. Microbes decompose the complex compounds, from these decomposition reactions, a energy release is produced, which bacteria use for growth and their own metabolism.
        Microbes are classified as aerobic, requiring O2 to function or anaerobic, functioning only in the complete absence of O2
        Aerobic Decay: Complex organic compounds are broken down under aerobic conditions, C, S, P and N constiuents of the molecules are converted to CO2, So4 (2-), PO4 (3-). In each of these product molecules/ ions, C, S, P and N exist in their maximum oxidation state
        Anaerobic Decay: When anaerobic conditions prevail e.g. at the bottom of stagnant swamps and deep down in landfills, the microbial decomposition of the complex organic compounds produces small molecules such as CH4, H2S, PH3 (phosphine) and NH3. In each of rhwaw product molecules, the C, S, P and N exists in their minimum oxidation state.
        Two processes of the C and O cycles are photosynthesis and respiration. Photosynthesis is a photochemical process whereby green plants convert CO2 and H2O into glucose and oxygen in the leaves using solar energy; it's a multi-step process that makes use of chlorophyll as a catalyst.
        Aeronic Respiration is the process by which many plants and animals derive energy from the burning of glucose, this is also a multi-step process
        The Nitrogen Cycle: In the atmosphere N naturally occurs as diatomic molecules, N2, in which atoms are held together by a very strong, covalent bond. This strong bond accounts for N's lack of reactivity, as a large amount of energy is required to break the bond to enable a reaction to occur.
        Nitrogen is removed from the atmosphere by fixation processes that involve the conversion of the N2 to compound of N such as NH3 and NO. These nitrogen compounds can then undergo further chemical transformations into other compounds such as proteins, N2O, NO3 (2-) and NO2 (-).
        Natural Nitrogen FIxation: Certain nitrogen-fixing bacteria, 'rhizobium', are found in nodules on the roots of legume plants, which contain the enzyme nitrogenase that catalyses the conversion of N2 from the atmosphere onto NH3 or NH4 (+) compounds, this enzyme is also found in blue-green algae.
        NNF is also initiated by lightning discharges and forest forces, the high-energy outputs associated with the events are sufficient to break the triple bond in the N2 molecule, leading to the reaction between N2 and O2 in the atmosphere, forming NO, which then reacts with more O2 forming NO2. NO2 dissolves in H2O in the atmosphere forming NO3 (-) which are deposited in the soil when it rains.
        Artificial Nitrogen Fixation: Large scale industrial production of NH3 by the Haber process, the combustion of fuels (including the burring of forest in land clearing) and increased cultivation of legume plants all remove large quantities of N2 from the atmosphere. The Haber process uses high temperatures to bring about the reaction between N2 and H2 to produce ammonia.
        The combustion of fuels in air in internal combustion engines and industrial furnaces produces the high temperatures needed for the atmospheric N2 and O2 to react, similarly to lightning discharges and forest fires.
        Chemical Conversion of Fixed Nitrogen: The oxidation of NH4(+) to NO3(-)- important process that occurs in H2O and soils, microbes bring about the two-step conversion of NH4 (+) to NO3 (-), it is primarily in the NO3 (-) form that N is absorbed by plants.
        The conversion of NO3(-) to protein in plants: Plants absurd N as aqueous NO3 (-), the converse process process is a multi-step process. The reduction of NO3 (-) to N2 or N2O:This is the mechanism by which N2 is returned into the atmosphere. The reductions are catalysed by enzymes present in microorganisms. Ni2O is released into the atmosphere.
        The conversion of organic N compounds into inorganic N compounds: The conversion of organic N compounds e.g. amino acids, NH4 (+) NO3 (-) happens when micro-organisms decompose the remains of dead plants or animals in both ski and aquatic environments. The actual products of the decomposition depend on the pH of the soil or H2O and the level of O2 available to the micro-organisms.
        Industrial Chemical Conversions of Fixed Nitrogen: The chemical industry uses the Haber process to fix N as NH3, which can be applied directly to soils as a plant fertiliser or it can be converted into other compounds, a number of which are used as fertilisers. The made fertiliser compounds made from ammonia are NH4NO3, (NH4)2SO4, (NH4)3PO4 and NH2CONH2; once they enter the soil, they enter into the nitrogen cycle.
        The Greenhouse Effect
        Hot objects emit electromagnetic radiation. The higher the temperature of an object, the high and the shorter the wavelength of the radiation emitted. The sun emits radiation that includes the ultraviolet, visible and infra-red radiation.
        Of the solar radiation that reaches the lower atmosphere, slightly less than half of it actually warms the earth's surface, the remainder is either reflected by the clouds and oceans or absorbed by the atmosphere .
        Certain gases in th slower atmosphere can absorb certain wavelengths of this infra-red radiation. THis process warms the earth's lower atmosphere and is referref to sd the greenhouse effect e.g. of these gases include CO2. CO2 and water vapour absorb infra-red radiation strongly in the two main bands, outside of these bands, most of the infra-red radiation emitted from the earth's surface escapes out through the atmosphere.
        Other naturally occurring greenhouse gases in the lower atmosphere are Ch4, N2O and O3. The synthetic halocarbons are also powerful greenhouse gases. The overall effect that the greenhouse effect provides an atmospheric 'blanket' for warming the earth.
        Only polar bonds can absorb infra-red radiation, therefore, the two most abundant gases in the atmosphere are O2 and N2, which don't absorb infra-red and are not greenhouse gases.
        Since the 1950s the concentrations of all the greenhouse gases have been increasing. Higher concentrations of greenhouse gases have resulted in increased absorption of infra-red radiation emitted from the earth's surface, which scientists have hypothesised has lead to the overall increased temperature of the earth, with potential long term effects on climate
        This is also come from the increased levels of human activities being undertaken to meet the living require,met of an ever-expanding world population.
        CO2, from the combustion of fossil fuels, human activity, transport, power production and burning of rain forests, CH4, anaerobic breakdown of plant material, human activity, rice farming, cattle and sheep farming, N2O, reduction of NO3 (-) by microbes, human activity, use of nitrogenous fertilisers in agriculture.

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Description: Mind Map on Elemental and Environmental Chemistry, created by kandsdecorso on 06/10/2016.
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Elemental and Environmental Chemistry

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