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CHEMISTRY

CHEMISTRY. Patterns and Compounds. 5.1 Looking for Patterns in Chemical Reactivity. Valence Shell – The outermost electron shell that an atom has. Valence Electrons – the electrons found in the atom ’ s valence shell. Steam Boat Demo. Preparation

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CHEMISTRY

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  1. CHEMISTRY Patterns and Compounds

  2. 5.1 Looking for Patterns in Chemical Reactivity • Valence Shell – The outermost electron shell that an atom has. • Valence Electrons – the electrons found in the atom’s valence shell.

  3. Steam Boat Demo • Preparation • Using a dropper fill the tubes and diaphragm with water • Place the boat in the water • Place and lit a candle under the diaphragm • Explain diaphragm expansion • When heated the water particles move faster • This increases the space between particles

  4. This causes an increase in pressure, pushing the tin up • A change of state occurs (vaporization), the water turns to steam (volume increases 1600 times) • What causes flattening of the diaphragm? • As the steam escapes, a vacuum is created. • This causes water to be “sucked” in and the cycle repeats itself causing a “putt, putt” sound

  5. What propels the boat? • The escaping steam leaves the exhaust pipe with great force. • For every action, there is an equal and opposite reaction. • This pushes the boat forward. • External Combustion Engine? • The heat source is “outside” of the engine • Contrast with internal combustion engine • Controlled explosion in a cylinder

  6. Chemical Reactions • Atoms • NEVER gain or lose protons. (if they did we could make gold) • Almost NEVER gain or lose neutrons. (OK radioactive elements do) • Often gain or lose valence electrons. The number of valence electrons give the element it’s physical and chemical properties.

  7. Atomic Diagrams (draw Na diagrams below) Bohr-Rutherford Diagram /Lewis Dot Diagram (Electron Dot Diagram) • Since only the valence e- matter for properties, it becomes a waste of time to draw other electrons. • Lewis Dot Diagrams are faster and easier to draw, and they show all the important information.

  8. Drawing Lewis Diagrams: 1)Write the element’s symbol 2)Draw the valence electrons clockwise around the symbol.

  9. Read and complete p. 145 in the space below

  10. IONS • Keep an eye-on those electrons!! • Key Eye-dea: Atoms don’t want valence shells partly full.

  11. Cations • positively charged atoms that have lost electron(s). • Alkali Metals (1 valence e-) • easily LOSE the e- • Very reactive • BBC/OU Open2.net - The World Around Us - Alkali metals • Alkaline Earth Metals (2 valence e-) • Not as easy to lose 2 e-. • Reactive

  12. Cations • Reactivity increases down the family (group) because of the higher distance from e- to proton(+). • It’s easier to lose the e- if it is farther from the nucleus (higher energy level). • Metals form cations.

  13. Anions • Negatively charged atoms that have gained one or more electron(s) • Halogens (7 valence e-) TAKE electrons forcefully!! • Non-metals form anions. • Chalcogens (6 valence e-) • more difficult to take two electrons

  14. Anions • Noble Gases – don’t react with other atoms because their valence shells are already full.

  15. REMEMBER: Ca+ions are posi+ive. Anions are negative.

  16. Assignment Do pg 146 #1 – 6 and BLM’s 5-1, 5-2, 5-6, 5-7, 5-9 and 5-10

  17. 5.2 - Atoms vs. Ions Atoms • # of protons = # of electrons • neutral charge Ions • atoms which have gained or lost electrons • have charges on them

  18. 5.2 - Atoms vs. Ions Ex. If Mg loses 2 electrons it becomes (Remember metaLs Lose electrons) Ex. If Cl gains 1 electron it becomes (non-metals gain electrons)

  19. Chemical Bonding Note: Atoms always bond with each other through the valence electrons.

  20. Chemical Bonding There are 3 ways in which atoms bond: • Ionic bonds – between metals and non-metals. • Metals happily transfer their electrons to non-metals which happily accept the electrons. • Electrons transfer making two oppositely charged IONS which stick together.

  21. Chemical Bonding • Ionic bonds (cont’d) • Ionic compounds are formed. • Ionic compounds: • have high melting points. (strong ionic bonds) • dissolve easily in water. • are electrolytes – materials that conduct electricity when molten or when dissolved (aqueous).

  22. Lewis Diagrams for Bonding: Ionic Bonding – atoms become ions. • Metal atom donates e- to non-metal atom. X denotes donated e-. . X. LiH Fill in the dots below [ Li ]+[ H ]-

  23. Lewis Diagrams for Bonding: If the metal needs to donate 2 e-, it may need to find 2 atoms to accept one each. Fill in the missing electrons below. Cl Mg Cl [ Cl ]-[ Mg ]2+[ Cl ]-

  24. Lewis Diagrams for Bonding: OR find a non-metal that will take 2 e- Fill in the missing electrons. [Mg]2+[O]2- The charge that the ion has is often referred to as the combining capacity of the ion. Try: HCl

  25. Chemical Bonding • Covalent bonds – between two non-metals. • Electrons are shared by the non-metals. • Sharing electrons allows each atom to have a full shell for short periods of time.

  26. Chemical Bonding • Covalent bonds (cont’d) • Molecular compounds are formed. • molecular compounds: • have low melting points (weaker bonds) • don’t conduct electricity • don’t dissolve as easily in water.

  27. Chemical Bonding • Covalent bonds (cont’d) • there are 2 types of covalent bonds: • polar covalent – the atoms are different elements. Electrons are not shared evenly by the atoms. One atom gets the electrons for a longer time. Ex) H2O • H is the positive pole and O is the negative pole.

  28. Chemical Bonding • Covalent bonds (cont’d) • Non-polar covalent – the electrons are shared evenly because both atoms are the same element. Ex) N2, O2, F2, Cl2, Br2, I2 are the only examples. They are called diatomic molecules. See the “7” pattern in the periodic table. • HOFBrINCl the clown

  29. Lewis Diagrams for Bonding: Covalent Bonding – non-metals share electrons. Line up atoms so that they can fill each others outer shells. Circle shared pairs. H2O H O H O has 6 valence e-, H has 1

  30. Lewis Diagrams for Bonding: Try: CH4, Br2, O2 We will draw covalent bonds as follows: H H – C – H Br – Br O = O H

  31. Chemical Bonding • Metallic bonding – between atoms that are metals. • Metals allow their electrons to be juggled back and forth from one atom to the next. • Electricity: e- easily move from atom to atom.

  32. Comparing Types of Bonding

  33. http://rlv.zcache.com/war_of_the_atoms_poster-p228361748241900516tdcp_400.jpghttp://rlv.zcache.com/war_of_the_atoms_poster-p228361748241900516tdcp_400.jpg

  34. http://www.lab-initio.com/screen_res/nz063.jpg

  35. Assignment Do P 154 2-4 and BLM 5.8 – 5.11 & 5.13

  36. 5.3 Naming Binary Ionic Compounds • Binary – having only 2 types of atoms • Ionic Compounds – electrons are transferred from metals to non-metals. The ions stick together. (pg 148-149 diagrams)

  37. Reading the Name from the formula: • Binary compounds usually end it “ide”. • Write cation first and anion second adding “ide” to the anion. Examples: • CaCl2 is called Calcium Chloride • MgBr2 is called Magnesium Bromide

  38. Writing the formula from the name: • Determine that the compound is ionic. • (metal and a non-metal.) Eg. Calcium & Chlorine • Temporarily mark the combining capacity (charge) on each atom. Eg. Ca2+ Cl - • Note: in case of a charge of 1 the number 1 is assumed

  39. Writing the formula from the name: • Do a “crossover” with the combining capacities. • The 1 from Cl moves to a subscript on the Ca, and the 2+ on one Ca ion moves to the subscript on the Cl ion. • Write the number of atoms needed as a subscript. Note: Cations get written first. • CaCl2 Ionic “Dating” Activity

  40. Transition Metals Each of these metals can form more than one cation. Eg. Fe can form Fe2+ and Fe3+. They are located in the middle region of the periodic table.

  41. Classical Naming 1) Use the latin name for the atom. Fe is ferrum in Latin. a. Use suffix “ous” for the lesser charge ion. Fe2+ is ferrous. b. Use suffix “ic” for the greater charged ion. Fe3+ is ferric. Eg. Ferrous Sulfate would contain the Fe2+ ion.

  42. Stock Naming System: • Use roman numerals to show the charge on the cation. • Eg. FeCl3 is iron(III) chloride Assignment: BLM 5.14 – use page 156 #1 – 13 as a guide

  43. Naming Compounds

  44. Naming Polyatomic Ionic Compounds • Poly – many • Polyatomic/Complex Ions – Ions which have more than one type of element in them (see top of yellow sheet) • the ionic charge is on the entire group of atoms. Ex) (NO3)- is a group of atoms which has an extra electron. • Note: the nitrogen and oxygen atoms have covalent bonds.

  45. Naming Polyatomic Ionic Compounds • Common polyatomic ions: • ammonium, (NH4)+ • nitrate, (NO)3- • sulfate, (SO)42- • carbonate, (CO3)2- • phosphate, (PO4)3- • hydroxide, (OH)- • perchlorate, (ClO)4-

  46. Reading the name from the formula: • In compounds where the anion is polyatomic, the names end in “ate” or “ite”. • A few exceptions Hydroxide OH- • Write cation first and anion second. Examples: • Ca(SO4) is called calcium sulfate • (NH4)2S is called ammonium sulfide • (Note: “ide” because anion isn’t polyatomic)

  47. Reading the name from the formula: • If the cation is a transition metal: • Determine the charge on the cation by reverse crossover. E.g. Fe2O3 Fe2O3 2- 3+ Iron (III) Oxide

  48. Reading the name from the formula: • Give the charge of the cation in the name. • Use stock naming (Roman Numerals) on cation. • Roman numeral gives the charge of the cation E.g. Iron(II) Oxide is the Fe2+ ion FeO Iron (III) Oxide is the Fe 3+ ion Fe2O3 • Use Classical naming • “ous” ending for lesser charge. • “ic” ending for greater charge.

  49. Reading the name from the formula: Example: CuCO3 Cu(CO3)2- what charge does Cu have for the crossover to make CuCO3? Cu2+ • So CuCO3 is Copper (II) Carbonate (Note: Cu+ exists, so we must have the greater ion) or Cupric Carbonate.

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