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ISOTOPES

ISOTOPES. Isotopes are atoms have same number of protons but different number of neutrons. They have identical chemical properties. same number of electrons. They have different physical properties.  different number of neutrons. 6.8 Isotopes. 6.9 RELATIVE MASSES OF ATOMS.

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ISOTOPES

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  1. ISOTOPES Isotopes are atoms have same number of protons but different number of neutrons. • They have identical chemical properties. • same number of electrons. • They have different physical properties.  different number of neutrons 6.8 Isotopes

  2. 6.9 RELATIVE MASSES OF ATOMS Relative isotopic mass  mass number Relative atomic mass = a% × MA + b% × MB + c% × MC where a%, b%, c% = percentage abundance of isotopes A, B and C respectively MA, MB, MC = isotopic masses of isotopes A, B and C respectively 6.9 Relative masses of atoms

  3. Chapter 6 The Periodic Table Periods A horizontal row of elements is called a period. Period number = number of occupied electron shells Groups A vertical column of elements is called a group. Group number = number of electrons in outermost shell 7.2 The periodic table

  4. Some of the groups have special names: Group I : Alkali metals Group II : Alkaline earth metals Group VII : Halogens Group 0 : Noble gases Elements within the same group of the Periodic Table have similarchemical properties because they have the same number of outermost shell electrons

  5. Reactivity of metals increases down the group. Why? Metal react by releasing electrons. As the size of atom increases, distance between nucleus and outermost electrons increases. Thus, attraction between outermost electrons and nucleus decreases. As a result, it is easier for a large atom to release their outermost electrons. Reactivity of non-metals decreases down the group.

  6. PREDICTING CHEMICAL PROPERTIES OF AN UNFAMILIAR ELEMENT Elements of the same group have similar chemical properties! Write an equation for the reaction between Rb(s) and water. 2Rb(s) + 2H2O(l)  2RbOH(aq) + H2(g) 7.5 Predicting chemical properties of an unfamiliar element [Extension]

  7. Chapter 7 Chemical bonding: Ionic bonding IONIC BOND is formed by the transfer of one or more electrons from one atom (or group of atoms) to another. An ionic bond is formed between metal and non-metal. Chapter 8 Chemical bonding: Ionic bonding

  8. Evidence for the presence of ions COLOUR OF IONS Q Predict the colour of each of the following solutions: (a) Magnesium nitrate solution (b) Sodium permanganate solution (c) Ammonium chromate solution (d) Iron(II) sulphate solution colourless purple yellow pale green Chapter 8 Chemical bonding: Ionic bonding

  9. Table 7.1 The colours of some ions in aqueous solution.

  10. filter paper moistened with sodium sulphate solution small potassium permanganate crystal anode cathode microscope slide purple spot Why do the ions move in an electric field? mobile ions

  11. Cations Anions Charge Formula Name Charge Formula Name 1+ Na+ K+ Cu+ Ag+ Hg+ sodium ion potassium ion copper(I) ion silver ion mercury(I) ion 1- H- Cl- Br- I- OH- NO3- NO2- HCO3- HSO4- CN - MnO4- ClO3- ClO- hydride ion chloride ion bromide ion iodide ion hydroxide ion nitrate ion nitrite ion hydrogencarbonate ion hydrogensulphate ion cyanide ion permanganate ion chlorate ion hypochlorite ion H+ NH4+ hydrogen ion ammonium ion 2+ Mg2+ Ca2+ Ba2+ Pb2+ Fe2+ Co2+ Ni2+ Mn2+ Cu2+ Zn2+ Hg2+ magnesium ion calcium ion barium ion lead(II) ion iron(II) ion cobalt(II) ion nickel(II) ion manganese(II) ion copper(II) ion zinc ion mercury(II) ion 2- O2- S2- SO42- SO32- SiO32- CO­­3­2- CrO42- Cr2O72- oxide ion sulphide ion sulphate ion sulphite ion silicate ion carbonate ion chromate ion dichromate ion 3+ Al3+ Fe3+ Cr3+ aluminium ion iron(III) ion chromium(III) ion 3- N3- P3- PO43- nitride ion phosphide ion phosphate ion NAMES AND FORMULAE OF COMMON IONS

  12. Q Draw electron diagrams (showing electrons in the outermost shell only) to show the bond formation in (a) potassium sulphide (b) calcium bromide. A 8.7 Ionic bonding and ionic substances

  13. Chapter 8 Chemical bonding: Covalent bonding COVALENT BOND is the strong directional electrostatic attraction between the shared electrons (negatively charged) and the two nuclei (positively charged) of the bonded atoms. A covalent bond is formed by the sharing of outermost shell electrons between two atoms. A covalent bond is usually formed between non-metals.  molecules Chapter 9 Chemical bonding: Covalent bonding

  14. 9.1 Covalent bonding and covalent substances

  15. Classification of substances according to structure • Molecular structures • Giant ionic structures • Giant covalent structures • Giant metallic structures

  16. Van der Waals’ forces covalent bond What forces will be broken when melting the following molecular crystals?

  17. The atoms within a molecule are strongly bonded together (by covalent bonds). However, each molecule is attracted to other neighbouring molecules by weak intermolecular forces only. In general, the larger the molecular size, the greater will be the van der Waals’ forces between molecules. 9.4 Weak intermolecular forces: van der Waals’ forces

  18. Simple molecular solids are soft Water : insoluble Non-aqueous solvents like chloroform: soluble Have low melting points and boiling points Non-conductors of electricity, whether as solids, liquids or in aqueous solution Properties of simple molecular substances

  19. NaCl Giant Ionic Structure Use ionic bonds to link up millions of ions to form a large solid. P. 19 / 5

  20. Caesium ion is larger in size than sodium ion. Therefore, the structure of caesium chloride CsCl is different from that of sodium chloride. NaCl + Cs − Cl caesium chloride Structure of caesium chloride Two interpenetrating simple cubic structure P. 20 / 5

  21. All ionic compounds are solids. Have high melting points and boiling points. Water: soluble; Non-aqueous solvents: insoluble. Conduct electricity when molten or in aqueous solution by mobile ions. Properties of ionic compounds P. 21 / 5

  22. Giant Covalent Structure Use covalent bonds to link up millions of atoms to form a large solid --- no molecule

  23. Diamond Silicon(IV) oxide

  24. 2. All (except graphite) are hard. Giant covalent structures are all solids withvery high melting points and boiling points. 1. 3. They are insoluble in any solvent. 4. All (except graphite) are non- conductors of electricity. Properties of giant covalent structures

  25. Strong covalent bonds (within layers) weak van der Waals’ forces (between layers) Graphite

  26. Properties of Graphite Very high melting points and boiling points: Strong covalent bonds between atoms. 2. Soft: weak van der Waals’ forces between layers 3. Conductor of electricity: delocalized electrons.

  27. positive ions delocalized electrons moving in all directions  ‘sea’ of electrons Giant metallic structures

  28. 4. 3. Metals are malleable and ductile. Most metals are solids with high melting points and boiling points. Metals are good conductors of heat. 2. Metals are good conductors of electricity. 1. Most metals have high densities. 5. Properties of metals explained by structure and bonding

  29. Higher level question: The melting point of metal is relatively low compared with its boiling point, why? e.g. m.p. of tin is 232 oC and b.p. is 2602 oC When melting a metal, the metallic bonds are only weakened (not broken). However, when boiling a metal, all the metallic bonds are broken.

  30. Non Octet Structure

  31. Boron trifluoride (BF3) • The boron atom contains only six outermost shell electrons. Shape?

  32. Boron trifluoride (BF3) Trigonal Planar

  33. Phosphorus pentachloride (PCl5) Shape?

  34. Five electron pairs around the central atom Phosphorus pentachloride (PCl5) trigonal bipyramid

  35. Everyday chemistry Sulphur hexafluoride (SF6) Shape?

  36. Sulphur hexafluoride (SF6) octahedral

  37. Shape of other molecules:CH4 NH3 H2O

  38. Methane (CH4) Tetrahedral

  39. Ammonia (NH3) Trigonal pyramidal

  40. Water (H2O) V-shaped

  41. Electronegativity of an atom represents the power of that atom to attract bonding electronsin a covalent bond. Electronegativity of an atom

  42. Electronegativity increases Electronegativity decreases Learning tip Why?

  43. R r Bond polarity Unequally shared! • When atoms of different electronegativities form a covalent bond, the bonding electrons are notshared equally.

  44. Bond polarity • This gives a polar bond and (a polar molecule).

  45. In general, Greater electronegativity difference More polar the bond

  46. Factors affecting the strength of van der Waals’ forces between molecules Molecular size Molecular shape Polarity of molecules

  47. The van der Waals’ forces between the molecules are stronger if: • the molecular size is larger, • the contact surface area between the molecules is greater, • the molecules are more polar.

  48. H2O HF H2Te AsH3 SbH3 H2Se NH3 HI Boiling point (°C) H2S SnH4 HCl HBr PH3 GeH4 SiH4 CH4 Period Evidence of hydrogen bonding

  49. Formation of hydrogen bonding • When a hydrogen atom is directly bonded to a highly electronegative atoms: fluorine, oxygen and nitrogen, a highly polar bond is formed.

  50. Electrostatic attractions exist between this partial positive charged H and the lone pair electrons on a highly electronegative atom of another molecule – hydrogen bond

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