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Ionic Bonding Metallic Bonding Covalent and Da*Ve Covalent Metallic bonding AS Bonding & Periodicity Test: Answers Q1 Ti is a giant metallic lace. There is strong electrostac arac2on between posi2ve metal ions and ‘sea’ of delocalised electrons. Layers of atoms slide over one another when hammered, but, in the new metallic shape, the same strong metallic bonding (and bond strength) is re- formed between metal ions and delocalised electrons. Q2 Na conducts and NaCl does not conduct. Delocalised electrons flow through the metal and transfer charge. Ions are relavely fixed in solid NaCl, so it does not conduct. However, aqueous and molten NaCl can conduct as Na+ Cl- are mobile and able to transfer electrical charge. Ionic Bonding Q1 MgO is a giant ionic lace. There is strong arac2on between oppositely charged Mg2+ and O2- ions. Iodide ion has larger atomic radius. Thus, there is increased separaon between Na+ and I- nuclei, Q2 which decreases the strength of the ionic arac2ve force between ions, thus decreasing the overall strength of the ionically bonded giant lace. Covalent and dave covalent bonding PH + H+ à PH + 3 4 Q1 + PH4 formed by lone pair from P + PH3 is pyramidal atom being donated to H . This is a steric 4, 1 LP coordinate (dave covalent) bond. bond angle 107° Steric 4, No LP, bond angle 109.5° How bond shapes are formed: Q2 + Equal repulsion between H3O is pyramidal, bond angle 107°; steric 4, 1LP bonding electron pair ligands Lone pair from O donated to H+ − Q3 + F− Lone pair from F− donated to B atom to form dave covalent bond − BF3 trigonal planar BF4 tetrahedral steric 3, No LP, bond angle 120° steric 4, No LP, bond angle 109.5° Q4 M 267 = Al Cl + Cl− r 2 6 1 − − Lone pair from Cl donated to Al atom to form dave covalent bond in AlCl4 Bond Polarity Electronegavity is the ability of an atom to aract electron density from a pair of electrons Q1 contained in a covalent bond Q2 CH4 intermolecular force= van der Waals ; NH3 intermolecular force= hydrogen bonding Large electronegavity difference between N and H (3.0-2.1=0.9) Q3 This forms a dipole charge separaon Nδ− and Hδ+ within the N-H covalent bond + Lone pair of electrons on N atom is aracted to (hydrogen bonds with) Hδ of adjacent NH3 molecule Q4 NO has covalent bonding as there is only a small electronegavity difference between N and O (difference is 0.5 units according to table of electronegavity values) Forces between molecules The hydrogen bonding intermolecular forces that exist between methanol molecules are Q1 much stronger than the van der Waals forces that exist between oxygen molecules, hence methanol has a higher b.p. oxygen. Q2 Intermolecular forces: F2 is van der Waals, CH3F is dipole-dipole, HF is hydrogen bonding. Large electronegavity difference between F and H This forms a dipole charge separaon Fδ− and Hδ+ within the H-F covalent bond Lone pair of electrons on F atom is aracted to (hydrogen bonds with) Hδ+ of adjacent HF molecule Q3 The molecular mass (size) of hydrogen halides increases as the group is descended, hence there is increasing strength of van der Waals forces of arac2on. Conversely, the difference in electronegavity between hydrogen and halogen decreases as the group is descended, hence there is decreased dipole-dipole intermolecular forces of arac2on. Overall, these trends culminate in an increased boiling point as the hydrogen halide group is descended. Q4 Hydrogen bonding exists between NH3 and H2O or Q5 Hydrogen bonding exists between HF molecules The hydrogen bonding intermolecular forces (IMF) that exist between HF molecules are much stronger than the van der Waals forces that exist between F molecules. 2 2 Hence, lower energy is required to break IMF of liquid F2 so F2 has lower boiling point. Bonding and physical proper2es and states of maer Iodine is molecular crystal and graphite is giant covalent macromolecule Q1 Graphite: composed of layers of hexagonal covalently bound carbon atoms that are separated by weak van der Waals forces between the layers. Strong covalent bonds needs to be broken to melt graphite, hence it has a high mel2ng point. The hydrogen bonding intermolecular forces (IMF) that exist between H2O molecules are Q2 much stronger than the van der Waals and dipole-dipole forces that exist between H2S molecules. The increased molecular size of the compounds formed between hydrogen and group VI elements as the Group is descended means there are increased van der Waals arac2ve forces between molecules, hence a trend towards increased boiling point. Shapes of simple molecules Cl2O Q1 Steric 4, two LP Angular, 104.5° bond angles BeCl2 Steric 2, no LP Linear, 180° bond angle .. Q2 NHF2 Steric 4, one LP pyramidal, 107° bond angle BF3 Steric 3, no LP Trigonal planar, 120° bond angles - Q3 − N NH2 Steric 4, two LP Angular, 104.5° bond angles Q4 + AsCl4 ion has a bond angle of 109.5°, because the shape formed has equal repulsion between 4 bonding (electron pair) ligands. Q5 As + AsF5 ClF2 Steric 5, no LP Steric 4, two LP 3 Trigonal Bipyramidal, 90° and 120° bond angles Angular, 104.5° bond angles Period 3: decreasing atomic radius, increasing 1st IE & trends in m.p. Decreasing atomic radius: Q1 increasing nuclear charge, similar electron shielding ∴ increased effec2ve nuclear charge density ∴ greater arac2on of valence electron(s) by nucleus ∴ decreased atomic radius Q2 1st IE: the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous mono-posi2ve ions. X(g) à X+(g) + e- Q3 Mg+ (g) à Mg2+(g) + e- More energy is required to remove an electron from a posi2vely charged ion than a neutral ion. Q4 Also, Mg+ has a smaller atomic radius than Mg, so greater energy is needed to overcome the increased nuclear arac2ve force that is holding the outer shell electron closer to the nucleus. Increasing 1st IE: Q5 increasing nuclear charge, similar electron shielding ∴ increased effec2ve nuclear charge density ∴ greater arac2on of valence electron(s) by nucleus and decreased atomic radius ∴ more energy needed to remove outershell (valence) electron Sulfur: less than expected increase in 1st IE: Q6 Electron pair repulsion in the 3p suborbital means less energy is required to remove the valence e- Aluminium: less than expected increase in 1st IE: Q7 The valence electron is located in a higher energy 3p orbital, which is further away from the nucleus and its nuclear arac2ve force, hence less energy is required to remove the valence e- Q8 X is Aluminium (considerable increase in 1st IE between 3rd and 4th electron removal) S and P exist as simple molecular structures with weak van der Waals intermolecular forces Q9 8 4 between molecules. However, as S8 is a larger sized molecule than P4, it has stronger vdW. Thus, more energy is needed to break the IMF of S8 so it has a higher mel2ng point. Silicon is a giant covalent macromolecule. Each silicon atom is covalently bonded to 4 other Q10 silicon atoms in a diamond cubic crystal structure. Considerable energy is needed to break the strong covalent bonds within this giant structure arrangement, hence silicon has a high m.p. Al and Na exist as giant laces of metallic bonded atoms. However, Al, compared to Na, has a: Q11 • higher metal ion charge (+3 vs. +1) • more delocalised electrons per metal ion • smaller atomic radius than Na (closer distance between metal ion and delocalised e- sea). In combinaon, these factors ensure Al achieves stronger electrostac arac2on between the more posi2vely charged metal ions and higher e- density delocalised e- sea, compared to Na. Hence, metallic bonding is stronger in Al and more energy is needed to overcome this arac2ve force and melt Al compared to Na. 4 Extended wri2ng ques2on: Bonding • NaCl is an ionic compound. Oppositely charged ions are held together by electrostac forces in a giant ionic cubic lace. • Water is composed of covalent molecules that are held together by intermolecular hydrogen bonds. In each water molecule, a single oxygen atom is covalently bonded to two hydrogen atoms forming an angular shaped molecule. As ice, water molecules form a hexagonal crystal lace where two hydrogen bonds form per molecule of water. • Overall, ionic bonding within the lace of NaCl is stronger than the intermolecular forces holding water molecules together in ice. Hence, as more energy is needed to separate the separate the Na+ Cl- ions compared to breaking down hydrogen bonds in ice, NaCl has a higher m.p. than ice. • SF4 is steric 5, 4 bonding ligands and one lone pair • Structure is based on equal repulsion between bonding ligands and lone pair • Hence, molecule structure is see-saw shape. The electron pair distribu2on is trigonal bipyramidal with the lone pair located on the axial or equatorial locaon. • There is slight lone pair-bond pair repulsion that pushes the S-F bonds closer together, hence bond angles are slightly <120° and <180°, respec2vely. Extended wri2ng: Periodicity Elements in the p block have their outer electron(s) in p orbitals or levels or sub-shells e.g. Carbon 1s2 2s2 2p2 Periodicity: paern in the change in the proper2es of a row of elements that is repeated in the next row (underneath or above the ‘index’ row). Atomic radius decreases and electronegavity increases, across Period 3. Both trends share common reasons: increasing nuclear charge, similar electron shielding ∴ increased effec2ve nuclear charge density ∴ greater arac2on of valence electron(s) by nucleus and decreased atomic radius. Conduc2vity: Metals Na, Mg, Al conduct electricity due to free delocalised electrons being able to transfer electrical charge through the metallic bonded structure.
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