AQA A-Level Chemistry (new spec) 1.3 Bonding Already Passed

AQA A-Level Chemistry (new spec) 1.3 Bonding Already Passed Ionic Bonding The electrostatic force of attraction between two oppositely charged ions formed by electron transfer Covalent Bonding A shared pair of electrons Dative Covalent Bonding (AKA Co-ordinate bonding) Formed when the shared pair of electrons in the covalent bond come from only one of the bonding atoms. Metallic Bonding The electrostatic force of attraction between the positive metal cations and the sea of delocalised electrons Factors affecting the strength of metallic bonding: The number of protons The more protons in the cations, the stronger the electrostatic force of attraction between the cations and the sea of delocalised electrons Factors affecting the strength of metallic bonding: Number of delocalised electrons per atom The more delocalised electrons, the stronger the electrostatic force of attraction Factors affecting the strength of metallic bonding: Size of ion The smaller the ion, the stronger the electrostatic force of attraction Electronegativity The relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond towards itself Why does electronegativity increase as you go across a period? -The number of protons increased -The atomic radius decreases because the electrons in the same shell are pulled in more Why does electronegativity decrease as you go down a group? -Distance between the nucleus and the outer electrons increases -Shielding increases Why aren't the noble gases electronegative? Because they don't form bonds Using electronegativity to predict bonding: Covalent If both atoms have a similar electronegativity, the pull on the electrons from them will be of a similar strength, making a non-polar covalent bond. If one atom has a stronger electronegativity than the other, the electrons will be pulled more towards one atom, making the bond polar-covalent Using electronegativity to predict bonding: Ionic If the electronegativity difference is really large, the sharing of electrons is so uneven that the more electronegative atom has full possession of the 2 electrons, creating an ionic bond Using electronegativity to predict bonding: Metallic If both atoms have a low electronegativity, neither can attract electrons, so the electrons don't remain localised to the bond at all, causing a sea of delocalised electrons and a metallic bond Orbitals & Covalent Bonds When a covalent bond is formed, the 2 outer orbitals overlap, forming a normal covalent bond. Some atoms promote electrons to give more unpaired electrons and to allow more covalent bonding. For example, carbon promotes one of the electrons in the 2s orbital to the 2p orbital, meaning there are 4 unpaired electrons, so it can form 4 covalent bonds Orbitals and Dative Covalent Bonds Any atom with filled valence shell (outer shell) orbitals can donate their electrons for the covalent bond. This includes group 5,6,7 and 0 Any atom which has an empty orbital in their valence shell can accept a pair of electrons. Sigma Bonds Where the atomic orbitals overlap directly along the internuclear axis. All single bonds are sigma bonds. Pi Bonds Where the atomic orbitals overlap above and below the internuclear axis. All double bonds contain a sigma and a pi bond. All triple bonds contain a sigma bond and 2 pi bonds Strength of covalent bonds is affected when.. The atoms are smaller because the closer the electrons are to the nuclei, the stronger the bond Molecular Shapes: 2 electron pairs Linear, 180 degrees Molecular Shapes: 3 electron pairs Trigonal Planar, 120 degrees Molecular Shapes: 2 bonding pairs, 1 lone pair Bent, 118 degrees Molecular Shapes: 4 electron pairs Tetrahedral, 109.5 degrees Molecular Shapes: 3 bonding pairs, 1 lone pair Trigonal Pyramidal, 107 degrees Molecular Shapes: 2 bonding pairs, 2 lone pairs Bent, 104 degrees Molecular Shapes: 5 electron pairs Trigonal Bipyramidal, 120 &90 degrees Molecular Shapes: 3 bonding pairs, 2 lone pairs Trigonal Planar, 120° OR T-Shape 89 degrees Molecular Shapes: 6 electron pairs Octahedral, 90 degrees Molecular Shapes: 5 bonding pairs, 1 lone pair Distorted square pyramid, 89 degrees Molecular Shapes: 4 bonding pairs, 2 lone pairs Square planar, 90 degrees How to work out shapes -Write group number of central atom -Add number of atoms around the central atom -Add/Subtract charge (If charge is negative, add it & if charge is positive, subtract it) -To find bonding pairs, divide step 3 by 2 -To find lone pair, subtract step 4 by the number of atoms around the central atoms Intermolecular Forces: Van Der Waals Forces Occur between all simple covalent molecules & the separate atoms in noble gases As the electrons move, parts of the molecule can become more or less electronegative. These dipoles can induce dipoles in adjacent molecules. Factors which affect the strength of VDW forces: N