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Chemical Bonds Standard 2: chapter 7, 8, 9

Siddall. Chemistry. Vocabulary leave enough space for definition AND an example Metallic bond Alloy Ionic bond Cation Anion Crystal Covalent bond Polar covalent bond Diatomic molecule Electron dot structure. Chemical Bonds Standard 2: chapter 7, 8, 9.

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Chemical Bonds Standard 2: chapter 7, 8, 9

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  1. Siddall. Chemistry. Vocabulary leave enough space for definition AND an example Metallic bond Alloy Ionic bond Cation Anion Crystal Covalent bond Polar covalent bond Diatomic molecule Electron dot structure Chemical Bonds Standard 2:chapter 7, 8, 9

  2. Standard 2d: Intermolecular forces Gases: • Particles have no attraction to each other • Extremely low melting point • low boiling point • Particles move rapidly and randomly • no fixed shape • no fixed volume Liquids: • Particles are weakly attracted to each other • Low melting point • Particles move around each other freely • no fixed shape • fixed volume Solids: • Particles are strongly attracted to each other • High melting point • Particles vibrate in place • fixed shape • fixed volume

  3. study question 1 • If a substance has no fixed shape it could be: _________ or _________ • If a substance has no fixed shape and no fixed volume it would be: ___________

  4. Physical state: The state of a material depends on the balance between: • the kinetic energy of the particles. • the attractions between particles • Kinetic energy > attractions = gas • Kinetic energy < attractions = liquid • Kinetic energy << attractions = solid

  5. study question 2 • If the kinetic energy of particles in a substance is much greater than the forces between particles, the substance is a __________

  6. Phase changes. • Melting a solid or evaporating a liquid requires energy to overcome the forces holding the particles together. • Freezing a liquid or condensing a gas is caused by removing energy so attractive forces between particles dominate

  7. Physical state gas evaporating Condensing liquid Energy added/absorbed intermolecular attractions are overcome Energy released/removed intermolecular attractions take over melting freezing solid

  8. study question 3 Which processes occur when energy is removed from a substance?

  9. Honors Only: Volatility = degree of change. • A substance with high volatility will change easily from solid to liquid or liquid to gas. For example: • carbon dioxide, oxygen are very volatile compounds (gas at room temperature) • Water is less volatile (liquid at room temperature) • Iron, salt are considered non-volatile compounds (solids at room temperature)

  10. Study question 4 • List the following compounds as ‘extremely volatile’, ‘somewhat volatile’ or ‘non-volatile’. Explain each choice. • Methane (natural gas) • Alcohol • Calcium carbonate (rocks)

  11. Standard 2a: Types of Bonds

  12. study question 5 • What type of bonds are formed when non-metal atoms share electrons?

  13. Bonding in Metals Metal atoms share valence electrons. • Atoms are very close together ஃMetals are solid compounds • Electrons move around (sea of electrons) ஃ Metals conduct electricity ஃ Metals are malleable. • ex: lead Alloy: mixture of different metals with specific properties superior to individual metals. e.x. steel frame construction.

  14. study question 6 • Explain why metals are solid and why they conduct electricity.

  15. Standard 2c: Ionic bonds. An Ionic bond is formed between metal and non-metal atoms • Each atom gains or loses electrons in order to form an octet • An ion is a charged particle • A cation = positive ion = a Metal. • e.x. Na+, Ca2+, Al3+ • An anion = negative ion = a Non-metal • e.x. Cl-, S2-, P3-

  16. study question 7 For the compound KF: • Which atom is the cation? • Which atom is an anion?

  17. Crystal Lattice Structure • All ionic compounds form a crystal lattice structure • formed by a very large network of electrostatic attractions (positive and negative ions attracted to each other). • Lattice energy: The energy needed to break the electrostatic attractions holding the lattice structure. • Ionic compounds are always solid because of the strong electrostatic attractions between ions

  18. Crystal Lattice Structure. + + + + +

  19. study question 8 Why do ionic compounds form crystal lattice structures?

  20. Properties of ionic compounds Electrostatic attractions are very strong therefore ionic compounds: • are solids at room temperature. • have very high melting points.

  21. study question 9 Which of the following are solid at room temperature? • CaO • CO • NO2 • Na2O

  22. naming ionic compounds e.x. Na2O • Use cation name • modify anion name (ide) • Do NOT use prefixes • Name = sodium oxide More examples: CaF2 = calcium fluoride K2O = potassium oxide

  23. study question 10 Name the following • MgCl2 • Al2O3 • NaBr

  24. Weird things • Polyatomic ions: act as one charged particle in an ionic bond • Anion names are not modified for polyatomic ions • Example: NH4OH = Ammonium hydroxide • Example: Al(NO3)3 = aluminum nitrate

  25. study question 11 Name the following • NaOH • K2SO4 • Mg(NO3)2

  26. Writing ionic formulas from names • example: • Sodium hydroxide (Na+ and OH-) • Charges must cancel out = NaOH • example: • Magnesium hydroxide (Mg2+ and OH-) For each Mg2+ there must be 2 x OH- = Mg(OH)2 note: use parenthesis only when showing more than one polyatomic ion

  27. study question 12 Write formulas for the following compounds: • Aluminum hydroxide • Potassium oxide • Magnesium nitrate

  28. Standard 2b: covalent bonds O C Covalent (molecular) compounds: • Formed when non-metal atoms bond. • Bonds between atoms are strong • But many covalent compounds are liquids or gases because molecules are not strongly attracted to each other • ex: H2O, CO2 • Properties: • many covalent molecules have very low melting points and high volatility • Many covalent molecules are gases or liquids O O H H

  29. study question 13 Identify the covalent compounds: • CO2 • CaO • MgCl2 • CCl4

  30. Naming covalent compounds. • e.x. CO2 = 1 carbon + 2 oxygen • Name = carbon dioxide • Modify name of second atom (ide). • Add pre-fix to indicate number of atoms.

  31. Prefixes • mono • di • tri • tetra • penta • hexa • hepta • octa • nona • deca

  32. Examples: • CCl4 = • Carbon tetrachloride. • N2O3 = • Dinitrogen trioxide. • Exception: The ‘mono’ prefix is usually omitted from the first atom • NO = nitrogen monoxide

  33. Diatomic molecules. hydrogen nitrogen oxygen fluorine chlorine bromine iodine = molecules formed from 2 atoms • H2 • N2 • O2 • F2 • Cl2 • Br2 • I2 NOTE: Nitrogen = N2 N2 is a molecule N = a nitrogen atom News flash: you must know these

  34. study question 14 Name the following: • CO • CO2 • Cl2 • NO • N2O

  35. diagrams show: Chemical symbol Valence electrons Standard 2e: Lewis Dot Diagrams • Each atom has 4 valence electron orbitals (one s orbital and 3 p orbitals) • Each orbital can hold 2 electrons. • Electrons like to be alone. • electrons pair up if necessary.

  36. Examples of Lewis Dot diagrams e.x. Nitrogen atom e.x. Sulfur atom • • • electron S • N • • • • • • • Electron orbitals

  37. study question 15 • Draw a Lewis Dot Diagram for an oxygen atom • Draw the Lewis Dot Diagram for a chlorine atom

  38. Creating an Octet Non-metal atoms form covalent bonds in order to share electrons and create an octet. Example: H2 • Each hydrogen has one electron. • Each hydrogen needs two electrons (like He). • H H H H • Covalent bond = 2 shared electrons (show in between atoms)

  39. study question 16 • Draw the Lewis dot diagram for a chlorine molecule (Cl2)

  40. e.x. Oxygen molecule (O2). • Oxygen atom: • needs 2 more electrons (to have an octet) • forms 2 bonds (using 2 unpaired electrons) •• • • • • • •• O O • • •

  41. •• •• Double bond O •• O •• •• •• One bond • The oxygen molecule still has the same total number of electrons • But each atom ‘thinks’ it has an ‘octet’.

  42. study question 17 • Draw the Lewis dot diagram for N2

  43. Rules for Dot Diagrams: • Count total number of valence electrons for all atoms. • Determine number of bonds needed for each atom. • Allocate unpaired electrons to bonds. • Allocate unshared (paired) electrons to orbitals so each atom has an octet. • Re-count total number of electrons in diagram.

  44. e.x.: CH4 (methane molecule) • Hydrogen • has 1 electron • needs 1 electron • forms 1 bond. • Carbon • has 4 electrons • needs 4 electrons • forms 4 bonds. • •H •H • • C •H • •H Total number of electrons = 8

  45. H octet • The total number of electrons did not change. • Each atom ‘thinks’ it has an octet. • • Single bond. • • H C H • • • • Helium electron configuration H

  46. study question 18 • Draw the correct Lewis Dot Diagram for H2O

  47. Danger! HONORS STUDENTS ONLY BEYOND THIS POINT.

  48. VESPR TheoryValence Shell Electron Pair Repulsion • An atom with no unshared electrons forms four bonds: 4 single bonds = tetrahedral 2 single bonds & 1 double bond = trigonal planar 2 double bonds = linear 1 single bond & 1 triple bond = tetrahedral

  49. Study question 19 • Draw the lewis dot diagram for CF4 and determine the shape of the molecule

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