Elements Of Life

Where do elements come from
1. Nuclear Fusion
  1. Two light atomic nuclei fuse together to create a single heavier nucleus of a new element, releasing enormous amounts of energy
  2. Requires high temperature and pressure to overcome the repellent forces of the positive nuclei
  3. Can only occur in the centre of stars as the nuclei are moving with much more energy
2. Isotopes
  1. Atoms of the same element with different mass numbers. This causes a difference in mass
  2. The relative atomic mass (Ar) is the average of the relative isotopic masses, taking into account their abundances.
    1. This is measured by mass spectrometry
      1. Sample atoms are ionised and seperated according to their mass to charge ratios.
      2. The separated ions are detected, along with their abundance
Spectroscopy
1. Absorption spectra
  1. Glowing stars emit all light frequencies
  2. In the photosphere small molecules absorb some of the emitted radiation
    1. Lines appear black when light has been absorbed
2. Emission spectra
  1. When molecules absorb energy they are raised from their ground state to an exited state
    1. They lose energy by emitting electromagnetic radiation in the form of visible light
  2. Ultraviolet light emission spectrum is the Lyman
  3. Hydrogen emission spectrum in visible light is the Balmer series
  4. Lines become closer together at higher frequencies
    1. Speed of light = wavelength x frequency
  5. Bohr's theory
    1. Electron in the hydrogen atom only exists in certain definite energy levels or electron shells
    2. A photon of light is emitted or absorbed when the electron changes from one energy level to another
    3. The energy of the photon is equal to the difference between the two energy levels /\E
    4. Since E=hv it follows that the freqeuncy of the emitted or absorbed light is related to /\E by /\E=hv
  6. Unique to each element as there are different gaps between energy levels
3. Under certain conditions a substance can absorb or emit electromagnetic radiation
4. Flame tests
  1. Li+ Bight red
  2. Na+ yellow
  3. K+ lilac
  4. Cu2+ blue green
  5. Ba2+ apple green
  6. Ca2+ brick red
Electron Structure
1. Shells
  1. (n=1) 2 electrons
  2. (n=2) 8 electrons
  3. (n=3) 18 electrons
  4. (n=4) 32 electrons
2. Sub shells
  1. S
    1. 2 electrons
    2. 1 s orbital
  2. P
    1. 3 p orbitals
    2. 6 electrons
  3. D
    1. 5 D orbitals
    2. 10 electrons
  4. F
    1. 7 f orbitals
    2. 14 electrons
3. Electron configurations
  1. Orbitals are filled in order of increasing energy
  2. Orbitals are filled up singly, before pairing up
  3. The S orbital in the 4th shell fills up before the D orbital in the 3rd shell as the energy levels are lower
    1. The 3d sub shell is written alongside other n=3 sub shells even though it is filled after 4s
Periodicity
1. The occurance of periodic patterns
2. Metals to non metals across the group
3. First ionisation energy
4. Melting and boiling points
Covalent Bonding
1. Molecule Shape
  1. Electron repulsion - electron pairs try to be as far from each other as possible
  2. Planar Triangular
    1. 120
    2. E.g. Boron Flouride BF3
    3. 3 Groups of electrons No lone pairs
  3. Tetrahederal
    1. four groups of electrons round an atom
    2. 109.5
    3. E.g. Methane, CH4
  4. Bent
    1. Two lone pairs and Two single covalent bonds
    2. E.g, water
    3. 104.5
  5. Linear
    1. Two single or double covalent bonds around the central aton
    2. 180
    3. E.g. BeCl2
  6. Bipyrimidal
    1. Five groups of electrons round a central atom
    2. E.g. Phosphorous pentachloride PCl5
    3. Either 120 or 90, depending on the position within the molecule
  7. Octohederal
    1. Six groups of electrons round a central atom
    2. E.g. SF6
    3. 90
  8. Pyramidal
    1. 107
    2. Ammonia
    3. 3 groups of electrons One lone pair
2. Elements achieve a full outer shell by sharing electrons
3. Shown by dot and cross diagrams
4. Electron pairs that form bonds are bonding pairs
5. Electron pairs not involved in bonding are lone pairs
6. when two pairs of electrons form a covalent bond, it is a double bond
  1. E.g. Oxygen, O2
7. Dative Covalent Bonding
  1. Both bonding electrons come from the same atom
  2. Shown by an arrow coming away from the donating atom
  3. E.g. Carbon Monoxide, CO
8. Simple Molecular
  1. E.g. CO2, Cl2
  2. Weak intermolecular bonds between molecules
  3. Low melting point
9. Giant Covalent Structure
  1. E.g. graphite, diamond
  2. Very high melting point
  3. Insoluble
  4. Group 4 elements
Moles
1. The Avagadro constant is 6.02 x 10^23
2. Empirical Formula
  1. Use mass given in the question to work out molecular formula from the empirical formula
  2. Simplest ratio for moles of atoms
3. Water of crystalisation
  1. Water in ionic substances, formed when the substance crystalises
Group 1
1. Reactivity increases down a group
2. First ionisation energy
  1. energy required to remove an electron from every atom from the outer shell in one mole of isolated gaseous atoms of an element
    1. One mole of gaseous elements with positive charge is formed
  2. energy is always needed as electrons are attracted to the positive nucleus
  3. Decreases down the group as electrons are more shielded by full shells
  4. Increases across the peroid as the nucleus has more of a positive charge
3. Lithium Sodium Potassium Rubidium Caesium Francium
4. Soft, weak, with a low melting point
5. Reactive with water and oxygen
Group 2
1. Carbonates
  1. Formulae: MgCO3 CaCO3 SrCO3 BaCO3
  2. Less soluble down the group
  3. Decompose when heated
    1. BaCO3 > BaO + CO2
    2. More difficult to decompose down the group
      1. More thermally stable
    3. Charge density: same charge over smaller size = stronger force on electrons
2. Hydroxides
  1. Increasingly soluble
    1. Mg(OH)2 to Ba(OH)2
  2. Formulae: Mg(OH)2 Ca(OH)2 Sr(OH)2 Ba(OH)2
3. Oxides
  1. Formulae: MgO, CaO, SrO, BaO
  2. Reacts with acids: MgO + H2OSO4 > MgSO4 + H2O
4. More reactive going down the group
  1. Electron sheilding
5. Magnesium Calcium Strontium Barium
Salts
1. Acid
  1. A proton donator
  2. Has a pH of less than 7
  3. Turns litmus paper red
  4. Reacts with carbonates to give carbon doixide
  5. Neutralised by bases
  6. H2SO4, HCl
2. Base
  1. Proton (H+) acceptor
  2. Reacts with an acid to produce water and a salt
  3. Alkali
    1. Dissolves in water to produce OH- ions
    2. All alkalis are bases but not all bases are alkalis
  4. Oxides, hydroxides, ammonia
3. Water is amphoteric (can act as a base or acid)
4. Bronsted-Lowry Theory: An acid is a H+ donator and a base is a H+ accpetor
Ionic Bonding
1. Ion: An atom which has lost or gained an electron and therefore has a charge
  1. E.g. Na+
2. Cations = positive
  1. H+ Hydrogen NH4+ ammonium Li+ Lithium
  2. Mg2+ magnesium Fe2+ Iron [II] Ca2+ Calcium
  3. Al3+ Aluminium Fe3+ Iron [III}
  4. Metals
3. Anions = negative
  1. Non metals
  2. F- Flourine Cl- Chlorine OH- hydroxide NO3- nitrate HCO3- hydrogen carbonate
  3. O2- Oxide CO3 2- Carbonate SO4 2- Sulphate
4. Held together in ELECTROSTATIC bonds by their opposite charges
5. Ionic salts
  1. Acid + alkali > salt + water
  2. Acid + carbonate > salt + water + carbon doixide
  3. Acid + metal > salt + hydrogen
6. Form regularly shaped crystals
  1. High melting points
7. Precipitates
  1. Formed when soluble reactants form an insoluble salt
  2. Non soluble ionic compounds
    1. Barium, calcium, lead, and silver sulphates
    2. Silver and lead halides
    3. Metal carbonates
    4. Metal hydroxides (except group 1 and ammonium hydroxide
  3. Spectator ions
    1. ions not involved in the reaction so are not included in the ionic equation
8. Electrons are transferred from metal ions to nonmetal ions

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Elements Of Life

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