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Elements Of Life

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Elements Of Life

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