Unit 1 As chemistry

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  • Acids & Salts, Redox, Shells and Structure, Bonding (Physical properties) and Periodicity
  • Acids and Salts Molecular formula Name CH3COOH Ethanoic acid HCOOH Methanoic acid HNO3 Nitric acid HCL Hydrochloric acid C6H8O7 Citric acid H2SO4 Sulphuric acid H3PO4 Phosphoric acid
  • Reactions of acids and Redox Acid + carbonate salt +CO 2 H2 O Acid + Base salt + H2 O Acid + Alkali salt + H2 O Acid + Metal salt + H2 OIL RIG: AN ACID IS A PROTON DONOR A BASE IS A PROTON ACCEPTOR Oxidation Is Loss, Reduction Is Gain (of electrons.)
  • Reactions of acids and Redox Acids release H+ ions when they are in aqueous solutions. The common bases are metal oxides, metal hydroxides and ammonia. An alkali is a soluble base that releases OH ions in aqueous solutions. Metals generally form ions by losing electrons with an increase in oxidation number to form positive ions. Non-metals generally react by gaining electrons with a decrease in oxidation number to form negative ions.
  • Reactions of acids and Redox Name Base Metal Oxides MgO and CuO Metal Hydroxides NaOH and Mg (OH)2 Ammonia NH 3 Amine CH 3 NH 2 Magnesium Hydroxide Mg (OH)2 Calcium Hydroxide Ca(OH)2
  • Key Definitions Definitions: Hydrated Refers to a crystalline compound containing water molecules Anhydrous Refers to a substance that contains no water molecules Water of crystallisation Refers to water molecules that form and essential part of the crystalline structure of a compound
  • Shells, atoms, bonds and structure First ionisation energy: The amount of energy required to remove one electron from each atom in one mole of gaseous atoms, to form one mole of gaseous 1+ ions. Electrons are arranged in sub shells (s-, p- and d-sub-shells.) One orbital is a region that can hold up to two electrons of opposite spins. For example: Aluminium. The electrons always fill a subshell by them selves first, rather like finding a seat on a bus; youre not going to sit next to someone you dont know. Successive ionisation energy for Sodium: Ionisation energy is all about removing electrons. The equation is always the same, just with a change in element obviously. The definition is key, and is worth 3 marks in exam papers.
  • Shells, atoms, bonds and structure Factors effecting ionisation energys: Atomic radius: Greater atomic radius = smaller nuclear attraction on outer electrons. Nuclear charge: Greater nuclear charge = greater attractive force on the outer electrons. Electron shielding: This is when the inner shells repel the outer shells of electrons. More inner shells = smaller nuclear attraction on the outer electrons.
  • Bonding: Ionic, Covalent and Metallic Ionic bonding: This is the electrostatic attraction between oppositely charged ions. Covalent bonding: This can be described as a shared pair of electrons. A dative covalent bond is where a pair of electrons is provided by one of the atoms. Metallic bonding: This is the attraction of positive ions to a sea of delocalised electrons. Sea of delocalised electrons
  • Bonding: The forces and attractions Electronegativity and polarity: Non-polar bonds: In a molecule of H2 the bond is identical, as both atoms have a fair share of electrons (100% perfect covalent bond). This results in a non-polar molecule. This is also true for Cl2 and other diatomic atoms. Polar bonds: In a molecule of HCL, one atom is more electronegative than the other; in this case, it is the Cl. This means it has a greater attraction for the bonding pair, which draws the electrons closer to it, creating a permanent dipole. This results in a polar bond. Polar molecules: HCL is a polar molecule with polar bonds because of the charge difference across the whole molecule. This is also a non- symmetrical molecule. Electronegativity: This increases across the periods and decreases down the groups. If there is a small difference in the electronegativity, it will result in a polar covalent bond. If there is a big difference in electronegativity, one atom will effectively capture the electrons and result in forming an ionic bond.
  • Bonding: The forces and attractions Intermolecular forces: An intermolecular force: This is the attractive force between neighbouring molecules. Permanent dipoles: A permanent dipole-dipole force is a weak attractive force between permanent dipoles in neighbouring polar molecules. Van der Walls: These are an attractive force between induced dipoles in neighbouring molecules. These attractions are caused by electron wobbling within the molecule. This induces a charge in neighbouring molecules as they attract to each other. These forces increase with increasing number of electrons, which increase the attractive forces between the molecules. Hydrogen bonding: This temporary attraction between the slightly negative oxygen and the slightly positive hydrogen in a water molecule. The bond runs from one of the two lone pairs on the oxygen to the hydrogen. This gives water its special properties of high surface tension and ice being less dense than water. This is because the rigid bonds in the ice collapse when it melts, allowing the water molecules to move closer together.
  • Bonding and physical properties Structure Type of bonding Properties Malleable or ductile? Soluble or insoluble? Simple molecular Within in the structure there are strong covalent bonds, but the molecules are held together by weak van der Waals forces. They have low melting and boiling points. Weak forces dont need a lot of energy to be broken. They dont conduct electricity as there are no free ions to move in the structure. They are soluble in non-polar solvents as weak van der Waals forces form between the solvent and simple molecular structure. This formation of bonds weakens the lattice. Giant covalent Within the structure there are strong covalent bonds. High melting and boiling point. High temperatures needed to break the strong covalent bonds. Non-conductors of electricity due to no free ions. Giant covalent structures are soluble in both polar and non- polar solvents because the covalent bonds are too strong to be broken. Giant ionic Every ion in a giant ionic lattice is surrounded by an ion of the opposite charge. Strong electrostatic forces which need energy to break. When a solid, it is a non-conductor of electricity. When molten or dissolved in in water, the ions can move about and conduct electricity. Giant ionic lattices can dissolve in polar solvents such as water. Giant metallic In a giant metallic lattice, positive ions hold a fixed position while the delocalised electrons are free to move throughout the structure. High melting points and boiling points because the attraction between the positive ions and negative electrons is strong and needs high temperatures to break the bonds. They also have good electric conductivity due to the free moving electrons. Giant metallic lattices are both malleable and ductile, meaning it can be stretched out into a wire (ductile), or hammered into shape (malleable).
  • Reasons for high or low melting and boiling points in periods 2 and 3 Period 2 Li Be B C N2 O2 F2 Ne Period 3 Na Mg Al Si P4 S8 Cl2 Ar Structure Giant metallic Giant covalent Simple molecular structures Forces Strong forces between ions and delocalised electrons Strong forces between atoms Weak forces between molecules Bonding Metallic bonding Covalent bonding Van der Waals forces
  • Periodicity and first ionisation energy