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Chemistry Vocabulary and Intermolecular Forces Study Guide

Learn about chemical bonds, physical states, phase changes, types of bonds, ionic bonds, crystal lattice structures, and naming ionic compounds. Study question examples included.

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Chemistry Vocabulary and Intermolecular Forces Study Guide

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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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