1. How To Prepare (c) 2006, Mark Rosengarten  DO NOT CRAM. Get your studying done with by the night before. Get a good night’s sleep and have breakfast the morning of the exam.  Use a review book with old exams, answers and explanations in it. Take the old tests and grade yourself. The questions you don’t understand why you got wrong make sure to see your teacher about.  Actively participate in any and all review classes and activities offered by your teacher.  Study vocabulary. Identify key words and use flash cards to help you remember what the meaning of those words are and the concepts behind them.

2. Outline for Review (c) 2006, Mark Rosengarten 1) The Atom (Nuclear, Electron Config)The Atom 2) Matter (Phases, Types, Changes)Matter 3) Bonding (Periodic Table, Ionic, Covalent)Bonding 4) Compounds (Formulas, Reactions, IMAF’s)Compounds 5) Math of Chemistry (Formula Mass, Gas Laws, Neutralization, etc.) Math of Chemistry 6) Kinetics and Thermodynamics (PE Diagrams, etc.)Kinetics and Thermodynamics 7) Acids and Bases (pH, formulas, indicators, etc.)Acids and Bases 8) Oxidation and Reduction (Half Reactions, Cells, etc.)Oxidation and Reduction 9) Organic Chemistry (Hydrocarbons, Families, Reactions)Organic Chemistry

3. The AtomAtom (c) 2006, Mark Rosengarten 1) Nucleons – click here for website on nucleonsNucleonsclick here 2) Isotopes – click here for website on isotopesIsotopesclick here 3) Natural RadioactivityNatural Radioactivity 4) Half-LifeHalf-Life 5) Nuclear PowerNuclear Power 6) Electron ConfigurationElectron Configuration 7) Development of the Atomic ModelDevelopment of the Atomic Model

4. Nucleons (c) 2006, Mark Rosengarten  Protons: +1 each, determines identity of element, mass of 1 amu, determined using atomic number, nuclear chargeatomic numbernuclear charge  Neutrons: no charge, determines identity of isotope of an element, 1 amu, determined using mass number - atomic number (amu = atomic mass unit)mass number  32 16 S and 33 16 S are both isotopes of S  S-32 has 16 protons and 16 neutrons  S-33 has 16 protons and 17 neutrons  All atoms of S have a nuclear charge of +16 due to the 16 protons. website

5. Isotopes (c) 2006, Mark Rosengarten  Atoms of the same element MUST contain the same number of protons.  Atoms of the same element can vary in their numbers of neutrons, therefore many different atomic masses can exist for any one element. These are called isotopes.  The atomic mass on the Periodic Table is the weight-average atomic mass, taking into account the different isotope masses and their relative abundance.weight-average atomic mass  Rounding off the atomic mass on the Periodic Table will tell you what the most common isotope of that element is.common isotope of that element

6. Weight-Average Atomic Mass (c) 2006, Mark Rosengarten  WAM = ((% A of A/100) X Mass of A) + ((% A of B/100) X Mass of B) + …  What is the WAM of an element if its isotope masses and abundances are:  X-200: Mass = 200.0 amu, % abundance = 20.0 %  X-204: Mass = 204.0 amu, % abundance = 80.0%  amu = atomic mass unit ( 1.66 × 10 -27 kilograms/amu) website

7. Most Common Isotope (c) 2006, Mark Rosengarten  The weight-average atomic mass of Zinc is 65.39 amu. What is the most common isotope of Zinc? Zn-65!  What are the most common isotopes of:  CoAg  SPb  FACT: one atomic mass unit (1.66 × 10 -27 kilograms) is defined as 1/12 of the mass of an atom of C-12.  This method doesn’t always work, but it usually does. Use it for the Regents exam.

8. Natural Radioactivity (c) 2006, Mark Rosengarten  Alpha Decay Alpha Decay  Beta Decay Beta Decay  Positron Decay Positron Decay  Gamma Decay Gamma Decay  Charges of Decay Particles Charges of Decay Particles  Natural decay starts with a parent nuclide that ejects a decay particle to form a daughter nuclide which is more stable than the parent nuclide was. website

9. AlphaAlpha Decay (c) 2006, Mark Rosengarten  The nucleus ejects two protons and two neutrons. The atomic mass decreases by 4, the atomic number decreases by 2.  238 92 U 

10. BetaBeta Decay (c) 2006, Mark Rosengarten  A neutron decays into a proton and an electron. The electron is ejected from the nucleus as a beta particle. The atomic mass remains the same, but the atomic number increases by 1.  14 6 C 

11. PositronPositron Decay (c) 2006, Mark Rosengarten  A proton is converted into a neutron and a positron. The positron is ejected by the nucleus. The mass remains the same, but the atomic number decreases by 1.  53 26 Fe 

12. GammaGamma Decay (c) 2006, Mark Rosengarten  The nucleus has energy levels just like electrons, but the involve a lot more energy. When the nucleus becomes more stable, a gamma ray may be released. This is a photon of high-energy light, and has no mass or charge. The atomic mass and number do not change with gamma. Gamma may occur by itself, or in conjunction with any other decay type.

13. Charges of Decay Particles (c) 2006, Mark Rosengarten

14. Half-Life (c) 2006, Mark Rosengarten  Half life is the time it takes for half of the nuclei in a radioactive sample to undergo decay.  Problem Types:  Going forwards in time Going forwards in time  Going backwards in time Going backwards in time  Radioactive Dating Radioactive Dating website

15. Going Forwards in Time (c) 2006, Mark Rosengarten  How many grams of a 10.0 gram sample of I-131 (half-life of 8 days) will remain in 24 days?  #HL = t/T = 24/8 = 3  Cut 10.0g in half 3 times: 5.00, 2.50, 1.25g

16. Going Backwards in Time (c) 2006, Mark Rosengarten  How many grams of a 10.0 gram sample of I-131 (half-life of 8 days) would there have been 24 days ago?  #HL = t/T = 24/8 = 3  Double 10.0g 3 times: 20.0, 40.0, 80.0 g

17. Radioactive Dating (c) 2006, Mark Rosengarten  A sample of an ancient scroll contains 50% of the original steady-state concentration of C-14. How old is the scroll?  50% = 1 HL  1 HL X 5730 y/HL = 5730y

18. Nuclear Power (c) 2006, Mark Rosengarten  Artificial Transmutation Artificial Transmutation  Particle Accelerators Particle Accelerators  Nuclear Fission Nuclear Fission  Nuclear Fusion Nuclear Fusion

19. Artificial Transmutation (c) 2006, Mark Rosengarten  40 20 Ca + _____ -----> 40 19 K + 1 1 H  96 42 Mo + 2 1 H -----> 1 0 n + _____  Nuclide + Bullet --> New Element + Fragment(s)  The masses and atomic numbers must add up to be the same on both sides of the arrow. Website

20. Particle Accelerators (c) 2006, Mark Rosengarten  Devices that use electromagnetic fields to accelerate particle “bullets” towards target nuclei to make artificial transmutation possible!  Most of the elements from 93 on up (the “transuranium” elements) were created using particle accelerators.  Particles with no charge cannot be accelerated by the charged fields. website

21. Nuclear FissionFission (c) 2006, Mark Rosengarten  235 92 U + 1 0 n  92 36 Kr + 141 56 Ba + 3 1 0 n + energy  The three neutrons given off can be reabsorbed by other U-235 nuclei to continue fission as a chain reaction  A tiny bit of mass is lost (mass defect) and converted into a huge amount of energy. website

22. Chain Reaction (c) 2006, Mark Rosengarten

23. Nuclear FusionFusion (c) 2006, Mark Rosengarten  2 1 H + 2 1 H  4 2 He + energy  Two small, positively-charged nuclei smash together at high temperatures and pressures to form one larger nucleus.  A small bit of mass is destroyed and converted into a huge amount of energy, more than even fission. website

24. Electron Configuration (c) 2006, Mark Rosengarten  Basic Configuration Basic Configuration  Valence Electrons Valence Electrons  Electron-Dot (Lewis Dot) Diagrams Electron-Dot (Lewis Dot) Diagrams  Excited vs. Ground State Excited vs. Ground State  What is Light? What is Light?

25. Basic Configuration (c) 2006, Mark Rosengarten  The number of electrons is determined from the atomic number.  Look up the basic configuration below the atomic number on the periodic table. (PEL: principal energy level = shell)  He: 2 (2 e - in the 1st PEL)  Na: 2-8-1 (2 e - in the 1st PEL, 8 in the 2nd and 1 in the 3rd)  Br: 2-8-18-7 (2 e - in the 1st PEL, 8 in the 2nd, 18 in the 3rd and 7 in the 4th)

26. Valence Electrons (c) 2006, Mark Rosengarten  The valence electrons are responsible for all chemical bonding.valence electrons  The valence electrons are the electrons in the outermost PEL (shell).  He: 2 (2 valence electrons)  Na: 2-8-1 (1 valence electron)  Br: 2-8-18-7 (7 valence electrons)  The maximum number of valence electrons an atom can have is EIGHT, called a STABLE OCTET.STABLE OCTET

27. Electron-Dot Diagrams (c) 2006, Mark Rosengarten  The number of dots equals the number of valence electrons.  The number of unpaired valence electrons in a nonmetal tells you how many covalent bonds that atom can form with other nonmetals or how many electrons it wants to gain from metals to form an ion.  The number of valence electrons in a metal tells you how many electrons the metal will lose to nonmetals to form an ion. Caution: May not work with transition metals.  EXAMPLE DOT DIAGRAMS EXAMPLE DOT DIAGRAMS Click here for website on valence electrons and electron dot diagramsvalence electrons and electron dot diagrams

28. Example Dot Diagrams (c) 2006, Mark Rosengarten Carbon can also have this dot diagram, which it has when it forms organic compounds.

29. Excited vs. Ground State (c) 2006, Mark Rosengarten  Configurations on the Periodic Table are ground state configurations.  If electrons are given energy, they rise to higher energy levels (excited state).  If the total number of electrons matches in the configuration, but the configuration doesn’t match, the atom is in the excited state.  Na (ground, on table): 2-8-1  Example of excited states: 2-7-2, 2-8-0-1, 2-6-3 website

30. What Is Light? (c) 2006, Mark Rosengarten  Light is formed when electrons drop from the excited state to the ground state.  The lines on a bright-line spectrum come from specific energy level drops and are unique to each element.bright-line spectrum  EXAMPLE SPECTRUM EXAMPLE SPECTRUM

31. (c) 2006, Mark Rosengarten This is the bright-line spectrum of hydrogen. The top numbers represent the PEL (shell) change that produces the light with that color and the bottom number is the wavelength of the light (in nanometers, or 10 -9 m). No other element has the same bright-line spectrum as hydrogen, so these spectra can be used to identify elements or mixtures of elements. website

32. Development of the Atomic Model (c) 2006, Mark Rosengarten  Thompson Model Thompson Model  Rutherford Gold Foil Experiment and Model Rutherford Gold Foil Experiment and Model  Bohr Model Bohr Model  Quantum-Mechanical Model Quantum-Mechanical Model

33. Thompson Model (c) 2006, Mark Rosengarten  The atom is a positively charged diffuse mass with negatively charged electrons stuck in it. website

34. Rutherford Model (c) 2006, Mark Rosengarten  The atom is made of a small, dense, positively charged nucleus with electrons at a distance, the vast majority of the volume of the atom is empty space. Alpha particles Alpha particles shot at a thin sheet of gold foil: most go through (empty space). Some deflect or bounce off (small + charged nucleus). website

35. Bohr Model (c) 2006, Mark Rosengarten  Electrons orbit around the nucleus in energy levels (shells). Atomic bright-line spectra was the clue. Animation

36. Quantum-Mechanical Model (c) 2006, Mark Rosengarten  Electron energy levels are wave functions.  Electrons are found in orbitals, regions of space where an electron is most likely to be found.orbitals  You can’t know both where the electron is and where it is going at the same time.  Electrons buzz around the nucleus like gnats buzzing around your head.

37. Matter (c) 2006, Mark Rosengarten 1) Properties of PhasesProperties of Phases 2) Types of MatterTypes of Matter 3) Phase ChangesPhase Changes

38. Properties of Phases (c) 2006, Mark Rosengarten  Solids: Crystal lattice (regular geometric pattern), vibration motion only Solids  Liquids: particles flow past each other but are still attracted to each other. Liquids  Gases: particles are small and far apart, they travel in a straight line until they hit something, they bounce off without losing any energy, they are so far apart from each other that they have effectively no attractive forces and their speed is directly proportional to the Kelvin temperature (Kinetic-Molecular Theory, Ideal Gas Theory) Gases

39. Solids (c) 2006, Mark Rosengarten The positive and negative ions alternate in the ionic crystal lattice of NaCl.

40. Liquids (c) 2006, Mark Rosengarten When heated, the ions move faster and eventually separate from each other to form a liquid. The ions are loosely held together by the oppositely charged ions, but the ions are moving too fast for the crystal lattice to stay together.

41. Gases (c) 2006, Mark Rosengarten Since all gas molecules spread out the same way, equal volumes of gas under equal conditions of temperature and pressure will contain equal numbers of molecules of gas. 22.4 L of any gas at STP (1.00 atm and 273K) will contain one mole (6.02 X 10 23 ) gas molecules. Since there is space between gas molecules, gases are affected by changes in pressure.

42. Types of Matter (c) 2006, Mark Rosengarten  Substances (Homogeneous)  Elements (cannot be decomposed by chemical change): Al, Ne, O, Br, H Elements  Compounds (can be decomposed by chemical change): NaCl, Cu(ClO 3 ) 2, KBr, H 2 O, C 2 H 6 Compounds  Mixtures Mixtures  Homogeneous: Solutions (solvent + solute) Homogeneous  Heterogeneous: soil, Italian dressing, etc. Heterogeneous

43. Elements (c) 2006, Mark Rosengarten  A sample of lead atoms (Pb). All atoms in the sample consist of lead, so the substance is homogeneous.  A sample of chlorine atoms (Cl). All atoms in the sample consist of chlorine, so the substance is homogeneous. website

44. Compounds (c) 2006, Mark Rosengarten  Lead has two charges listed, +2 and +4. This is a sample of lead (II) chloride (PbCl 2 ). Two or more elements bonded in a whole- number ratio is a COMPOUND.  This compound is formed from the +4 version of lead. This is lead (IV) chloride (PbCl 4 ). Notice how both samples of lead compounds have consistent composition throughout? Compounds are homogeneous! website

45. Mixtures (c) 2006, Mark Rosengarten  A mixture of lead atoms and chlorine atoms. They exist in no particular ratio and are not chemically combined with each other. They can be separated by physical means.  A mixture of PbCl 2 and PbCl 4 formula units. Again, they are in no particular ratio to each other and can be separated without chemical change. website

46. Phase Changes (c) 2006, Mark Rosengarten  Phase Change Types Phase Change Types  Phase Change Diagrams Phase Change Diagrams  Heat of Phase Change Heat of Phase Change  Evaporation Evaporation

47. Phase Change Types (c) 2006, Mark Rosengarten website

48. Phase Change Diagrams (c) 2006, Mark Rosengarten AB: Solid Phase BC: Melting (S + L)Melting CD: Liquid Phase DE: Boiling (L + G)Boiling EF: Gas Phase Notice how temperature remains constant during a phase change? That’s because the PE is changing, not the KE. website

49. Heat of Phase Change (c) 2006, Mark Rosengarten  How many joules would it take to melt 100. g of H 2 O (s) at 0 o C?  q=mH f = (100. g)(334 J/g) = 33400 J  How many joules would it take to boil 100. g of H 2 O (l) at 100 o C?  q=mH v = (100.g)(2260 J/g) = 226000 J website

50. Evaporation (c) 2006, Mark Rosengarten  When the surface molecules of a gas travel upwards at a great enough speed to escape.  The pressure a vapor exerts when sealed in a container at equilibrium is called vapor pressure, and can be found on Table H.vapor pressure, and can be found on Table H.  When the liquid is heated, its vapor pressure increases.  When the liquid’s vapor pressure equals the pressure exerted on it by the outside atmosphere, the liquid can boil.  If the pressure exerted on a liquid increases, the boiling point of the liquid increases (pressure cooker). If the pressure decreases, the boiling point of the liquid decreases (special cooking directions for high elevations).

51. Reference Table H: Vapor Pressure of Four Liquids (c) 2006, Mark Rosengarten website

52. Bonding (c) 2006, Mark Rosengarten 1) The Periodic TableThe Periodic Table 2) Ions Ions 3) Ionic BondingIonic Bonding 4) Covalent BondingCovalent Bonding 5) Metallic BondingMetallic Bonding

53. The Periodic Table (c) 2006, Mark Rosengarten  Metals Metals  Nonmetals Nonmetals  Metalloids Metalloids  Chemistry of Groups Chemistry of Groups  Electronegativity Electronegativity  Ionization Energy Ionization Energy Video

54. Metals (c) 2006, Mark Rosengarten  Have luster, are malleable and ductile, good conductors of heat and electricitylustermalleableductile  Lose electrons to nonmetal atoms to form positively charged ions in ionic bonds Lose electrons to nonmetal atoms to form positively charged ionsionic bonds  Large atomic radii compared to nonmetal atoms  Low electronegativity and ionization energyelectronegativityionization energy  Left side of the periodic table (except H)

55. Nonmetals (c) 2006, Mark Rosengarten  Are dull and brittle, poor conductors  Gain electrons from metal atoms to form negatively charged ions in ionic bonds Gain electrons from metal atoms to form negatively charged ions  Share unpaired valence electrons with other nonmetal atoms to form covalent bonds and moleculescovalent bonds  Small atomic radii compared to metal atoms  High electronegativity and ionization energyelectronegativityionization energy  Right side of the periodic table (except Group 18)

56. Metalloids (c) 2006, Mark Rosengarten  Found lying on the jagged line between metals and nonmetals flatly touching the line (except Al and Po).  Share properties of metals and nonmetals (Si is shiny like a metal, brittle like a nonmetal and is a semiconductor).

57. Chemistry of Groups (c) 2006, Mark Rosengarten  Group 1: Alkali Metals Group 1: Alkali Metals  Group 2: Alkaline Earth Metals Group 2: Alkaline Earth Metals  Groups 3-11: Transition Elements Groups 3-11: Transition Elements  Group 17: Halogens Group 17: Halogens  Group 18: Noble Gases Group 18: Noble Gases  Diatomic Molecules Diatomic Molecules website

58. Group 1: Alkali MetalsAlkali Metals (c) 2006, Mark Rosengarten  Most active metals, only found in compounds in nature  React violently with water to form hydrogen gas and a strong base: 2 Na (s) + H 2 O (l)  2 NaOH (aq) + H 2 (g)  1 valence electron  Form +1 ion by losing that valence electron  Form oxides like Na 2 O, Li 2 O, K 2 O

59. Group 2: Alkaline Earth MetalsAlkaline Earth Metals (c) 2006, Mark Rosengarten  Very active metals, only found in compounds in nature  React strongly with water to form hydrogen gas and a base:  Ca (s) + 2 H 2 O (l)  Ca(OH) 2 (aq) + H 2 (g)  2 valence electrons  Form +2 ion by losing those valence electrons  Form oxides like CaO, MgO, BaO

60. Groups 3-11: Transition MetalsTransition Metals (c) 2006, Mark Rosengarten  Many can form different possible charges of ions  If there is more than one ion listed, give the charge as a Roman numeral after the name  Cu +1 = copper (I) Cu +2 = copper (II)  Compounds containing these metals can be colored.

61. Group 17: HalogensHalogens (c) 2006, Mark Rosengarten  Most reactive nonmetals  React violently with metal atoms to form halide compounds: 2 Na + Cl 2  2 NaCl  Only found in compounds in nature  Have 7 valence electrons  Gain 1 valence electron from a metal to form -1 ions  Share 1 valence electron with another nonmetal atom to form one covalent bond.

62. Group 18: Noble GasesNoble Gases (c) 2006, Mark Rosengarten  Are completely nonreactive since they have eight valence electrons, making a stable octet.  Kr and Xe can be forced, in the laboratory, to give up some valence electrons to react with fluorine.  Since noble gases do not naturally bond to any other elements, one atom of noble gas is considered to be a molecule of noble gas. This is called a monatomic molecule. Ne represents an atom of Ne and a molecule of Ne.

63. Diatomic Molecules(elements) (c) 2006, Mark Rosengarten  Br, I, N, Cl, H, O and F are so reactive that they exist in a more chemically stable state when they covalently bond with another atom of their own element to make two-atom, or diatomic molecules.  Br 2, I 2, N 2, Cl 2, H 2, O 2 and F 2  The decomposition of water : 2 H 2 O  2 H 2 + O 2

64. Electronegativity (c) 2006, Mark Rosengarten  An atom’s attraction to electrons in a chemical bond.  F has the highest, at 4.0  Fr has the lowest, at 0.7  If two atoms that are different in EN (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion.  If the two atoms have an END of less than 1.7, they will share their unpaired valence electrons…covalent bond! website

65. Ionization Energy (c) 2006, Mark Rosengarten  The energy required to remove the most loosely held valence electron from an atom in the gas phase.  High electronegativity means high ionization energy because if an atom is more attracted to electrons, it will take more energy to remove those electrons.  Metals have low ionization energy. They lose electrons easily to form (+) charged ions.  Nonmetals have high ionization energy but high electronegativity. They gain electrons easily to form (-) charged ions when reacted with metals, or share unpaired valence electrons with other nonmetal atoms. website

66. Ions (c) 2006, Mark Rosengarten  Ions are charged particles formed by the gain or loss of electrons.  Metals lose electrons (oxidation) to form (+) charged cations. Metals lose electrons  Nonmetals gain electrons (reduction) to form (-) charged anions. Nonmetals gain electrons  Atoms will gain or lose electrons in such a way that they end up with 8 valence electrons (stable octet).  The exceptions to this are H, Li, Be and B, which are not large enough to support 8 valence electrons. They must be satisfied with 2 (Li, Be, B) or 0 (H). website

67. Metal Ions (Cations-positive ion)Cations-positive ion (c) 2006, Mark Rosengarten  Na: 2-8-1  Na +1 : 2-8  Ca: 2-8-8-2  Ca +2 : 2-8-8  Al: 2-8-3  Al +3 : 2-8 Note that when the atom loses its valence electron, the next lower PEL becomes the valence PEL. Notice how the dot diagrams for metal ions lack dots! Place brackets around the element symbol and put the charge on the upper right outside!

68. Nonmetal Ions (Anions-negative ion)Anions-negative ion (c) 2006, Mark Rosengarten  F: 2-7  F -1 : 2-8  O: 2-6  O -2 : 2-8  N: 2-5  N -3 : 2-8 Note how the ions all have 8 valence electrons. Also note the gained electrons as red dots. Nonmetal ion dot diagrams show 8 dots, with brackets around the dot diagram and the charge of the ion written to the upper right side outside the brackets.

69. Ionic Bonding (c) 2006, Mark Rosengarten  If two atoms that are different in ELECTRONEGATIVITY (END) from each other by 1.7 or more collide and bond (like a metal atom and a nonmetal atom), the one with the higher electronegativity will pull the valence electrons away from the atom with the lower electronegativity to form a (-) ion. The atom that was stripped of its valence electrons forms a (+) ion.ELECTRONEGATIVITY  The oppositely charged ions attract to form the bond. It is a surface bond that can be broken by melting or dissolving in water.  Ionic bonding forms ionic crystal lattices, not molecules. website

70. Example of Ionic BondingIonic Bonding (c) 2006, Mark Rosengarten

71. Covalent Bonding (c) 2006, Mark Rosengarten  If two nonmetal atoms have an END of 1.7 or less, they will share their unpaired valence electrons to form a covalent bond.  A particle made of covalently bonded nonmetal atoms is called a molecule.  If the END is between 0 and 0.4, the sharing of electrons is equal, so there are no charged ends. This is NONPOLAR covalent bonding.  If the END is between 0.5 and 1.7, the sharing of electrons is unequal. The atom with the higher EN will be  - and the one with the lower EN will be  + charged. This is a POLAR covalent bonding. (  means “partial”) Video

72. Examples of Covalent Bonding (c) 2006, Mark Rosengarten website

73. Metallic Bonding (c) 2006, Mark Rosengarten  Metal atoms of the same element bond with each other by sharing valence electrons that they lose to each other.  This is a lot like an atomic game of “hot potato”, where metal kernals (the atom inside the valence electrons) sit in a crystal lattice, passing valence electrons back and forth between each other).  Since electrons can be forced to travel in a certain direction within the metal, metals are very good at conducting electricity in all phases. website

74. Compounds (c) 2006, Mark Rosengarten 1) Types of CompoundsTypes of Compounds 2) Formula WritingFormula Writing 3) Formula NamingFormula Naming 4) Empirical FormulasEmpirical Formulas 5) Molecular FormulasMolecular Formulas 6) Types of Chemical ReactionsTypes of Chemical Reactions 7) Balancing Chemical ReactionsBalancing Chemical Reactions 8) Attractive ForcesAttractive Forces

75. Types of Compounds (c) 2006, Mark Rosengarten  Ionic: made of metal and nonmetal ions. Form an ionic crystal lattice when in the solid phase. Ions separate when melted or dissolved in water, allowing electrical conduction (electrolytes-video). Examples: NaCl, K 2 O, CaBr 2 Ionicelectrolytes-video  Molecular: made of nonmetal atoms bonded to form a distinct particle called a molecule. Bonds do not break upon melting or dissolving, so molecular substances do not conduct electricity. EXCEPTION: Acids [H + A - (aq)] ionize in water to form H 3 O + and A -, so they do conduct. Molecular  Network: made up of nonmetal atoms bonded in a seemingly endless matrix of covalent bonds with no distinguishable molecules. Very high m.p., don’t conduct. Network website

76. Ionic Compounds (c) 2006, Mark Rosengarten website

77. Molecular Compounds (c) 2006, Mark Rosengarten website

78. Network Solids (c) 2006, Mark Rosengarten Network solids are made of nonmetal atoms covalently bonded together to form large crystal lattices. No individual molecules can be distinguished. Examples include C (diamond) and SiO 2 (quartz). Corundum (Al 2 O 3 ) also forms these, even though Al is considered a metal. Network solids are among the hardest materials known. They have extremely high melting points and do not conduct electricity.

79. Formula Writing (c) 2006, Mark Rosengarten  The charge of the (+) ion and the charge of the (-) ion must cancel out to make the formula. Use subscripts to indicate how many atoms of each element there are in the compound, no subscript if there is only one atom of that element.  Na +1 and Cl -1 = NaCl  Ca +2 and Br -1 = CaBr 2  Al +3 and O -2 = Al 2 O 3  Zn +2 and PO 4 -3 = Zn 3 (PO 4 ) 2  Try these problems! Try these problems! website

80. Formulas to Write (c) 2006, Mark Rosengarten  Ba +2 and N -3  NH 4 +1 and SO 4 -2  Li +1 and S -2  Cu +2 and NO 3 -1  Al +3 and CO 3 -2  Fe +3 and Cl -1  Pb +4 and O -2  Pb +2 and O -2

81. Formula Naming (c) 2006, Mark Rosengarten  Compounds are named from the elements or polyatomic ions that form them.  KCl = potassium chloride  Na 2 SO 4 = sodium sulfate  (NH 4 ) 2 S = ammonium sulfide  AgNO 3 = silver nitrate  Notice all the metals listed here only have one charge listed? So what do you do if a metal has more than one charge listed? Take a peek!Take a peek!

82. The Stock System (c) 2006, Mark Rosengarten  CrCl 2 = chromium (II) chloride Try  CrCl 3 = chromium (III) chloride Co(NO 3 ) 2 and  CrCl 6 = chromium (VI) chloride Co(NO 3 ) 3  FeO = iron (II) oxideMnS = manganese (II) sulfide  Fe 2 O 3 = iron (III) oxideMnS 2 = manganese (IV) sulfide  The Roman numeral is the charge of the metal ion! website

83. Empirical Formulas (c) 2006, Mark Rosengarten  Ionic formulas: represent the simplest whole number mole ratio of elements in a compound.  Ca 3 N 2 means a 3:2 ratio of Ca ions to N ions in the compound.  Many molecular formulas can be simplified to empirical formulas  Ethane (C 2 H 6 ) can be simplified to CH 3. This is the empirical formula…the ratio of C to H in the molecule.  All ionic compounds have empirical formulas. website

84. Molecular Formulas (c) 2006, Mark Rosengarten  The count of the actual number of atoms of each element in a molecule.  H 2 O: a molecule made of two H atoms and one O atom covalently bonded together.  C 2 H 6 O: A molecule made of two C atoms, six H atoms and one O atom covalently bonded together.  Molecular formulas are whole-number multiples of empirical formulas:  H 2 O = 1 X (H 2 O)  C 8 H 16 = 8 X (CH 2 )  Calculating Molecular Formulas Calculating Molecular Formulas

85. Types of Chemical Reactions (c) 2006, Mark Rosengarten  Redox Reactions: driven by the loss (oxidation) and gain (reduction) of electrons. Any species that does not change charge is called the spectator ion. Redox Reactions  Synthesis Synthesis  Decomposition Decomposition  Single Replacement Single Replacement  Ion Exchange Reaction: driven by the formation of an insoluble precipitate. The ions that remain dissolved throughout are the spectator ions.precipitate  Double Replacement Double Replacement website

86. Synthesis (c) 2006, Mark Rosengarten  Two elements combine to form a compound  2 Na + O 2  Na 2 O  Same reaction, with charges added in:  2 Na 0 + O 2 0  Na 2 +1 O -2  Na 0 is oxidized (loses electrons), is the reducing agentoxidized  O 2 0 is reduced (gains electrons), is the oxidizing agentreduced  Electrons are transferred from the Na 0 to the O 2 0.  No spectator ions, there are only two elements here.

87. Decomposition (c) 2006, Mark Rosengarten  A compound breaks down into its original elements.  Na 2 O  2 Na + O 2  Same reaction, with charges added in:  Na 2 +1 O -2  2 Na 0 + O 2 0  O -2 is oxidized (loses electrons), is the reducing agent  Na +1 is reduced (gains electrons), is the oxidizing agent  Electrons are transferred from the O -2 to the Na +1.  No spectator ions, there are only two elements here.

88. Single Replacement (c) 2006, Mark Rosengarten  An element replaces the same type of element in a compound.  Ca + 2 KCl  CaCl 2 + 2 K  Same reaction, with charges added in:  Ca 0 + 2 K +1 Cl -1  Ca +2 Cl 2 -1 + 2 K 0  Ca 0 is oxidized (loses electrons), is the reducing agent  K +1 is reduced (gains electrons), is the oxidizing agent  Electrons are transferred from the Ca 0 to the K +1.  Cl -1 is the spectator ion, since it’s charge doesn’t change.

89. Double Replacement (c) 2006, Mark Rosengarten  The (+) ion of one compound bonds to the (-) ion of another compound to make an insoluble precipitate. The compounds must both be dissolved in water to break the ionic bonds first.  NaCl (aq) + AgNO 3 (aq)  NaNO3 (aq) + AgCl (s)  The Cl -1 and Ag +1 come together to make the insoluble precipitate, which looks like snow in the test tube. precipitate  No species change charge, so this is not a redox reaction.  Since the Na +1 and NO 3 -1 ions remain dissolved throughout the reaction, they are the spectator ions.  How do identify the precipitate? How do identify the precipitate?

90. Identifying the Precipitate (c) 2006, Mark Rosengarten  The precipitate is the compound that is insoluble. AgCl is a precipitate because Cl - is a halide. Halides are soluble, except when combined with Ag + and others.precipitate website

91. Balancing Chemical Reactions (c) 2006, Mark Rosengarten  Balance one element or ion at a time  Use a pencil  Use coefficients only, never change subscripts(formulas)  Revise if necessary  The coefficient multiplies everything in the formula by that amount  2 Ca(NO 3 ) 2 means that you have 2 Ca, 4 N and 12 O.  Examples for you to try! Examples for you to try! website

92. Reactions to Balance (c) 2006, Mark Rosengarten  ___NaCl  ___Na + ___Cl 2  ___Al + ___O 2  ___Al 2 O 3  ___SO 3  ___SO 2 + ___O 2  ___Ca + ___HNO 3  ___Ca(NO 3 ) 2 + ___H 2  __FeCl 3 + __Pb(NO 3 ) 2  __Fe(NO 3 ) 3 + __PbCl 2

93. Attractive Forces (c) 2006, Mark Rosengarten  Molecules have partially charged ends. The  + end of one molecule attracts to the  - end of another molecule.  Ions are charged (+) or (-). Positively charged ions attract other to form ionic bonds, a type of attractive force.  Since partially charged ends result in weaker attractions than fully charged ends, ionic compounds generally have much higher melting points than molecular compounds.  Determining Polarity of Molecules Determining Polarity of Molecules  Hydrogen Bond Attractions Hydrogen Bond Attractions

94. Determining Polarity of Molecules (c) 2006, Mark Rosengarten ----------------------------------------------------------------------------- website

95. Hydrogen Bond Attractions (c) 2006, Mark Rosengarten A hydrogen bond attraction is a very strong attractive force between the H end of one polar molecule and the N, O or F end of another polar molecule. This attraction is so strong that water is a liquid at a temperature where most compounds that are much heavier than water (like propane, C 3 H 8 ) are gases. This also gives water its surface tension and its ability to form a meniscus in a narrow glass tube.hydrogen bond website

96. Math of Chemistry (c) 2006, Mark Rosengarten 1) Formula MassFormula Mass 2) Percent CompositionPercent Composition 3) Mole ProblemsMole Problems 4) Gas LawsGas Laws 5) NeutralizationNeutralization 6) Concentration Concentration 7) Significant Figures and RoundingSignificant Figures and Rounding 8) Metric ConversionsMetric Conversions 9) CalorimetryCalorimetry

97. Formula Mass (c) 2006, Mark Rosengarten  Gram Formula Mass = sum of atomic masses of all elements in the compound Gram Formula Mass  Round given atomic masses to the nearest tenth  H 2 O: (2 X 1.0) + (1 X 16.0) = 18.0 grams/mole  Na 2 SO 4 : (2 X 23.0)+(1 X 32.1)+(4 X 16.0) = 142.1 g/mole  Now you try:  BaBr 2  CaSO 4  Al 2 (CO 3 ) 3 websiteVideo

98. Percent Composition (c) 2006, Mark Rosengarten The mass of part is the number of atoms of that element in the compound. The mass of whole is the formula mass of the compound. Don’t forget to take atomic mass to the nearest tenth! This is a problem for you to try.This is a problem for you to try website

99. Practice Percent Composition Problem (c) 2006, Mark Rosengarten  What is the percent by mass of each element in Li 2 SO 4 ?

100. Mole Problems (c) 2006, Mark Rosengarten  Grams Moles Grams Moles  Molecular Formula Molecular Formula  Stoichiometry Stoichiometry

101. Grams Moles (c) 2006, Mark Rosengarten  How many grams will 3.00 moles of NaOH (40.0 g/mol) weigh?  3.00 moles X 40.0 g/mol = 120. g  How many moles of NaOH (40.0 g/mol) are represented by 10.0 grams?  (10.0 g) / (40.0 g/mol) = 0.250 mol Video

102. Molecular Formula (c) 2006, Mark Rosengarten  Molecular Formula = (Molecular Mass/Empirical Mass) X Empirical Formula  What is the molecular formula of a compound with an empirical formula of CH 2 and a molecular mass of 70.0 grams/mole?  1) Find the Empirical Formula Mass: CH 2 = 14.0  2) Divide the MM/EM: 70.0/14.0 = 5  3) Multiply the molecular formula by the result: 5 (CH 2 ) = C 5 H 10

103. Stoichiometry (c) 2006, Mark Rosengarten  Moles of Target = Moles of Given X (Coefficent of Target/Coefficient of given)  Given the balanced equation N 2 + 3 H 2  2 NH 3, How many moles of H 2 need to be completely reacted with N 2 to yield 20.0 moles of NH 3 ?  20.0 moles NH 3 X (3 H 2 / 2 NH 3 ) = 30.0 moles H 2 website

104. Gas Laws (c) 2006, Mark Rosengarten  Make a data table to put the numbers so you can eliminate the words.  Make sure that any Celsius temperatures are converted to Kelvin (add 273).  Rearrange the equation before substituting in numbers. If you are trying to solve for T 2, get it out of the denominator first by cross-multiplying.  If one of the variables is constant, then eliminate it.  Try these problems! Try these problems! website

105. Gas Law Problem 1 (c) 2006, Mark Rosengarten  A 2.00 L sample of N 2 gas at STP is compressed to 4.00 atm at constant temp-erature. What is the new volume of the gas?  V 2 = P 1 V 1 / P 2  = (1.00 atm)(2.00 L) / (4.00 atm)  = 0.500 L

106. Gas Law Problem 2 (c) 2006, Mark Rosengarten  To what temperature must a 3.000 L sample of O 2 gas at 300.0 K be heated to raise the volume to 10.00 L?  T 2 = V 2 T 1 /V 1  = (10.00 L)(300.0 K) / (3.000 L) = 1000. K

107. Gas Law Problem 3 (c) 2006, Mark Rosengarten  A 3.00 L sample of NH 3 gas at 100.0 kPa is cooled from 500.0 K to 300.0 K and its pressure is reduced to 80.0 kPa. What is the new volume of the gas?  V2 = P 1 V 1 T 2 / P 2 T 1  = (100.0 kPa)(3.00 L)(300. K) / (80.0 kPa)(500. K)  = 2.25 L

108. Neutralization (c) 2006, Mark Rosengarten  10.0 mL of 0.20 M HCl is neutralized by 40.0 mL of NaOH. What is the concentration of the NaOH?  #H MaVa = #OH MbVb, so Mb = #H MaVa / #OH Vb  = (1)(0.20 M)(10.0 mL) / (1) (40.0 mL) = 0.050 M  How many mL of 2.00 M H 2 SO 4 are needed to completely neutralize 30.0 mL of 0.500 M KOH? website

109. Concentration (c) 2006, Mark Rosengarten  Molarity Molarity  Parts per Million Parts per Million  Percent by Mass Percent by Mass  Percent by Volume Percent by Volume

110. Molarity (c) 2006, Mark Rosengarten  What is the molarity of a 500.0 mL solution of NaOH (FM = 40.0) with 60.0 g of NaOH (aq)?  Convert g to moles and mL to L first!  M = moles / L = 1.50 moles / 0.5000 L = 3.00 M  How many grams of NaOH does it take to make 2.0 L of a 0.100 M solution of NaOH (aq)?  Moles = M X L = 0.100 M X 2.0 L = 0.200 moles  Convert moles to grams: 0.200 moles X 40.0 g/mol = 8.00 g website

111. Parts Per Million (c) 2006, Mark Rosengarten  100.0 grams of water is evaporated and analyzed for lead. 0.00010 grams of lead ions are found. What is the concentration of the lead, in parts per million?  ppm = (0.00010 g) / (100.0 g) X 1 000 000 = 1.0 ppm  If the legal limit for lead in the water is 3.0 ppm, then the water sample is within the legal limits (it’s OK!)

112. Percent by Mass (c) 2006, Mark Rosengarten  A 50.0 gram sample of a solution is evaporated and found to contain 0.100 grams of sodium chloride. What is the percent by mass of sodium chloride in the solution?  % Comp = (0.100 g) / (50.0 g) X 100 = 0.200%

113. Percent By Volume (c) 2006, Mark Rosengarten  Substitute “volume” for “mass” in the above equation.  What is the percent by volume of hexane if 20.0 mL of hexane are dissolved in benzene to a total volume of 80.0 mL?  % Comp = (20.0 mL) / (80.0 mL) X100 = 25.0%

114. Sig Figs and Rounding (c) 2006, Mark Rosengarten  How many Significant Figures does a number have? How many Significant Figures does a number have?  What is the precision of my measurement? What is the precision of my measurement?  How do I round off answers to addition and subtraction problems? How do I round off answers to addition and subtraction problems?  How do I round off answers to multiplication and division problems? How do I round off answers to multiplication and division problems?

115. How many Sig Figs?  Start counting sig figs at the first non-zero.  All digits except place-holding zeroes are sig figs. Measurement# of Sig Figs 234 cm3 67000 cm2 _ 45000 cm 4 560. cm3 560.00 cm5 Measurement# of Sig Figs 0.115 cm3 0.00034 cm2 0.00304 cm3 0.0560 cm3 0.00070700 cm5 (c) 2006, Mark Rosengarten website

116. What Precision? (c) 2006, Mark Rosengarten  A number’s precision is determined by the furthest (smallest) place the number is recorded to.  6000 mL : thousands place  6000. mL : ones place  6000.0 mL : tenths place  5.30 mL : hundredths place  8.7 mL : tenths place  23.740 mL : thousandths place

117. Rounding with addition and subtraction (c) 2006, Mark Rosengarten  Answers are rounded to the least precise place.

118. Rounding with multiplication and division (c) 2006, Mark Rosengarten  Answers are rounded to the fewest number of significant figures. website

119. Metric Conversions (c) 2006, Mark Rosengarten  Determine how many powers of ten difference there are between the two units (no prefix = 10 0 ) and create a conversion factor. Multiply or divide the given by the conversion factor. How many kg are in 38.2 cg? (38.2 cg) /(100000 cg/kg) = 0.000382 km How many mL in 0.988 dL? (0.988 dg) X (100 mL/dL) = 98.8 mL

120. Calorimetry (c) 2006, Mark Rosengarten  This equation can be used to determine any of the variables here. You will not have to solve for C, since we will always assume that the energy transfer is being absorbed by or released by a measured quantity of water, whose specific heat is given above.  Solving for q Solving for q  Solving for m Solving for m  Solving for  T Solving for  T website

121. Solving for q (c) 2006, Mark Rosengarten  How many joules are absorbed by 100.0 grams of water in a calorimeter if the temperature of the water increases from 20.0 o C to 50.0 o C?  q = mC  T = (100.0 g)(4.18 J/g o C)(30.0 o C) = 12500 J

122. Solving for m (c) 2006, Mark Rosengarten  A sample of water in a calorimeter cup increases from 25 o C to 50. o C by the addition of 500.0 joules of energy. What is the mass of water in the calorimeter cup?  q = mC  T, so m = q / C  T = (500.0 J) / (4.18 J/g o C)(25 o C) = 4.8 g

123. Solving for  T (c) 2006, Mark Rosengarten  If a 50.0 gram sample of water in a calorimeter cup absorbs 1000.0 joules of energy, how much will the temperature rise by?  q = mC  T, so  T = q / mC = (1000.0 J)/(50.0 g)(4.18 J/g o C) = 4.8 o C  If the water started at 20.0 o C, what will the final temperature be?  Since the water ABSORBS the energy, its temperature will INCREASE by the  T: 20.0 o C + 4.8 o C = 24.8 o C

124. Kinetics and Thermodynamics (c) 2006, Mark Rosengarten 1) Reaction RateReaction Rate 2) Heat of ReactionHeat of Reaction 3) Potential Energy DiagramsPotential Energy Diagrams 4) EquilibriumEquilibrium 5) Le Châtelier’s PrincipleLe Châtelier’s Principle 6) Solubility CurvesSolubility Curves

125. Reaction Rate (c) 2006, Mark Rosengarten  Reactions happen when reacting particles collide with sufficient energy (activation energy) and at the proper angle.activation energy  Anything that makes more collisions in a given time will make the reaction rate increase.  Increasing temperature  Increasing concentration (pressure for gases)  Increasing surface area (solids)  Adding a catalyst makes a reaction go faster by removing steps from the mechanism and lowering the activation energy without getting used up in the process.catalyst website

126. Heat of Reaction (c) 2006, Mark Rosengarten  Reactions either absorb PE (endothermic, +  H) or release PE (exothermic, -  H) Exothermic, PE  KE, Temp  Endothermic, KE  PE, Temp  Rewriting the equation with heat included: 4 Al(s) + 3 O 2 (g)  2 Al 2 O 3 (s) + 3351 kJ N 2 (g) + O 2 (g) +182.6 kJ  2 NO(g)

127. Potential Energy Diagrams (c) 2006, Mark Rosengarten  Steps of a reactions:  Reactants have a certain amount of PE stored in their bonds (Heat of Reactants)  The reactants are given enough energy to collide and react (Activation Energy)  The resulting intermediate has the highest energy that the reaction can make (Heat of Activated Complex)  The activated complex breaks down and forms the products, which have a certain amount of PE stored in their bonds (Heat of Products)  H products - H reactants =  H EXAMPLESEXAMPLES website Video

128. Making a PE Diagram (c) 2006, Mark Rosengarten  X axis: Reaction Coordinate (time, no units)  Y axis: PE (kJ)  Three lines representing energy (H reactants, H activated complex, H products )  Two arrows representing energy changes:  From H reactants to H activated complex : Activation Energy  From H reactants to H products :  H  ENDOTHERMIC PE DIAGRAM ENDOTHERMIC PE DIAGRAM  EXOTHERMIC PE DIAGRAM EXOTHERMIC PE DIAGRAM

129. Endothermic PE Diagram (c) 2006, Mark Rosengarten If a catalyst is added?

130. Endothermic with Catalyst (c) 2006, Mark Rosengarten The red line represents the catalyzed reaction.

131. Exothermic PE Diagram (c) 2006, Mark Rosengarten What does it look like with a catalyst?

132. Exothermic with a CatalystCatalyst (c) 2006, Mark Rosengarten The red line represents the catalyzed reaction. Lower A.E. and faster reaction time!

133. Equilibrium (c) 2006, Mark Rosengarten When the rate of the forward reaction equals the rate of the reverse reaction.

134. Examples of Equilibrium (c) 2006, Mark Rosengarten  Solution Equilibrium: when a solution is saturated, the rate of dissolving equals the rate of precipitating.  NaCl (s)  Na +1 (aq) + Cl -1 (aq)  Vapor-Liquid Equilibrium: when a liquid is trapped with air in a container, the liquid evaporates until the rate of evaporation equals the rate of condensation.  H 2 O (l)  H 2 O (g)  Phase equilibrium: At the melting point, the rate of solid turning to liquid equals the rate of liquid turning back to solid.  H 2 O (s)  H 2 O (l)

135. Le Châtelier’s Principle (c) 2006, Mark Rosengarten  If a system at equilibrium is stressed, the equilibrium will shift in a direction that relieves that stress.  A stress is a factor that affects reaction rate. Since catalysts affect both reaction rates equally, catalysts have no effect on a system already at equilibrium.  Equilibrium will shift AWAY from what is added  Equilibrium will shift TOWARDS what is removed.  This is because the shift will even out the change in reaction rate and bring the system back to equilibrium  NEXT NEXT Video website

136. Steps to Relieving Stress (c) 2006, Mark Rosengarten  1) Equilibrium is subjected to a STRESS.  2) System SHIFTS towards what is removed from the system or away from what is added.  The shift results in a CHANGE OF CONCENTRATION for both the products and the reactants.  If the shift is towards the products, the concentration of the products will increase and the concentration of the reactants will decrease.  If the shift is towards the reactants, the concentration of the reactants will increase and the concentration of the products will decrease.  NEXT NEXT

137. Examples (c) 2006, Mark Rosengarten  For the reaction N 2 (g) + 3H 2 (g)  2 NH 3 (g) + heat  Adding N 2 will cause the equilibrium to shift RIGHT, resulting in an increase in the concentration of NH 3 and a decrease in the concentration of N 2 and H 2.  Removing H 2 will cause a shift to the LEFT, resulting in a decrease in the concentration of NH 3 and an increase in the concentration of N 2 and H 2.  Increasing the temperature will cause a shift to the LEFT, same results as the one above.  Decreasing the pressure will cause a shift to the LEFT, because there is more gas on the left side, and making more gas will bring the pressure back up to its equilibrium amount.  Adding a catalyst will have no effect, so no shift will happen.

138. Solubility Curves (c) 2006, Mark Rosengarten  Solubility: the maximum quantity of solute that can be dissolved in a given quantity of solvent at a given temperature to make a saturated solution. Solubility  Saturated: a solution containing the maximum quantity of solute that the solvent can hold. The limit of solubility.  Supersaturated: the solution is holding more than it can theoretically hold OR there is excess solute which precipitates out. True supersaturation is rare. Supersaturated  Unsaturated: There are still solvent molecules available to dissolve more solute, so more can dissolve. Unsaturated  How ionic solutes dissolve in water: polar water molecules attach to the ions and tear them off the crystal. How ionic solutes dissolve in water:  LIKE DISSOLVES LIKE LIKE DISSOLVES LIKE website

139. Solubility (c) 2006, Mark Rosengarten SolubilitySolubility: go to the temperature and up to the desired line, then across to the Y-axis. This is how many g of solute are needed to make a saturated solution of that solute in 100g of H 2 O at that particular temperature. At 40 o C, the solubility of KNO 3 in 100g of water is 64 g. In 200g of water, double that amount. In 50g of water, cut it in half. Video

140. Supersaturated (c) 2006, Mark Rosengarten If 120 g of NaNO 3 are added to 100g of water at 30 o C: 1) The solution would be SUPERSATURATED, because there is more solute dissolved than the solubility allows 2) The extra 25g would precipitate out 3) If you heated the solution up by 24 o C (to 54 o C), the excess solute would dissolve.

141. Unsaturated (c) 2006, Mark Rosengarten If 80 g of KNO 3 are added to 100g of water at 60 o C: 1) The solution would be UNSATURATED, because there is less solute dissolved than the solubility allows 2) 26g more can be added to make a saturated solution 3) If you cooled the solution down by 12 o C (to 48 o C), the solution would become saturated

142. How Ionic Solutes Dissolve in Water (c) 2006, Mark Rosengarten Water solvent molecules attach to the ions (H end to the Cl -, O end to the Na + ) Water solvent holds the ions apart and keeps the ions from coming back together

143. Acids and Bases (c) 2006, Mark Rosengarten 1) Formulas, Naming and Properties of AcidsFormulas, Naming and Properties of Acids 2) Formulas, Naming and Properties of BasesFormulas, Naming and Properties of Bases 3) NeutralizationNeutralization 4) pHpH 5) IndicatorsIndicators 6) Alternate TheoriesAlternate Theories

144. Formulas, Naming and Properties of Acids (c) 2006, Mark Rosengarten  Arrhenius Definition of Acids: molecules that dissolve in water to produce H 3 O + (hydronium) as the only positively charged ion in solution.  HCl (g) + H 2 O (l)  H 3 O + (aq) + Cl -  Properties of Acids Properties of Acids  Naming of Acids Naming of Acids  Formula Writing of Acids Formula Writing of Acids

145. Properties of Acids (c) 2006, Mark Rosengarten  Acids react with metals above H 2 on Table J to form H 2 (g) and a salt.  Acids have a pH of less than 7.  Dilute solutions of acids taste sour.  Acids turn phenolphthalein CLEAR, litmus RED and bromthymol blue YELLOW.  Acids neutralize bases.  Acids are formed when acid anhydrides (NO 2, SO 2, CO 2 ) react with water for form acids. This is how acid rain forms from auto and industrial emissions.

146. Naming of Acids (c) 2006, Mark Rosengarten  Binary Acids (H + and a nonmetal)  hydro (nonmetal) -ide + ic acid  HCl (aq) = hydrochloric acid  Ternary Acids (H + and a polyatomic ion)  (polyatomic ion) -ate +ic acid  HNO 3 (aq) = nitric acid  (polyatomic ion) -ide +ic acid  HCN (aq) = cyanic acid  (polyatomic ion) -ite +ous acid  HNO 2 (aq) = nitrous acid

147. Formula Writing of Acids (c) 2006, Mark Rosengarten  Acids formulas get written like any other. Write the H +1 first, then figure out what the negative ion is based on the name. Cancel out the charges to write the formula. Don’t forget the (aq) after it…it’s only an acid if it’s in water!  Hydrosulfuric acid: H +1 and S -2 = H 2 S (aq)  Carbonic acid: H +1 and CO 3 -2 = H 2 CO 3 (aq)  Chlorous acid: H +1 and ClO 2 -1 = HClO 2 (aq)  Hydrobromic acid: H +1 and Br -1 = HBr (aq)  Hydronitric acid:  Hypochlorous acid:  Perchloric acid:

148. Formulas, Naming and Properties of Bases (c) 2006, Mark Rosengarten  Arrhenius Definition of Bases: ionic compounds that dissolve in water to produce OH - (hydroxide) as the only negatively charged ion in solution.  NaOH (s)  Na +1 (aq) + OH -1 (aq)  Properties of Bases Properties of Bases  Naming of Bases Naming of Bases  Formula Writing of Bases Formula Writing of Bases

149. Properties of Bases (c) 2006, Mark Rosengarten  Bases react with fats to form soap and glycerol. This process is called saponification.  Bases have a pH of more than 7.  Dilute solutions of bases taste bitter.  Bases turn phenolphthalein PINK, litmus BLUE and bromthymol blue BLUE.  Bases neutralize acids.  Bases are formed when alkali metals or alkaline earth metals react with water. The words “alkali” and “alkaline” mean “basic”, as opposed to “acidic”.

150. Naming of Bases (c) 2006, Mark Rosengarten  Bases are named like any ionic compound, the name of the metal ion first (with a Roman numeral if necessary) followed by “hydroxide”. Fe(OH) 2 (aq) = iron (II) hydroxide Fe(OH) 3 (aq) = iron (III) hydroxide Al(OH) 3 (aq) = aluminum hydroxide NH 3 (aq) is the same thing as NH 4 OH: NH 3 + H 2 O  NH 4 OH Also called ammonium hydroxide.

151. Formula Writing of Bases (c) 2006, Mark Rosengarten  Formula writing of bases is the same as for any ionic formula writing. The charges of the ions have to cancel out.  Calcium hydroxide = Ca +2 and OH -1 = Ca(OH) 2 (aq)  Potassium hydroxide = K +1 and OH -1 = KOH (aq)  Lead (II) hydroxide = Pb +2 and OH -1 = Pb(OH) 2 (aq)  Lead (IV) hydroxide = Pb +4 and OH -1 = Pb(OH) 4 (aq)  Lithium hydroxide =  Copper (II) hydroxide =  Magnesium hydroxide =

152. Neutralization (c) 2006, Mark Rosengarten  H +1 + OH -1  HOH  Acid + Base  Water + Salt (double replacement)  HCl (aq) + NaOH (aq)  HOH (l) + NaCl (aq)  H 2 SO 4 (aq) + KOH (aq)  2 HOH (l) + K 2 SO 4 (aq)  HBr (aq) + LiOH (aq)   H 2 CrO 4 (aq) + NaOH (aq)   HNO 3 (aq) + Ca(OH) 2 (aq)   H 3 PO 4 (aq) + Mg(OH) 2 (aq)  website

153. pH (c) 2006, Mark Rosengarten  A change of 1 in pH is a tenfold increase in acid or base strength.  A pH of 4 is 10 times more acidic than a pH of 5.  A pH of 12 is 100 times more basic than a pH of 10. website

154. Indicators (c) 2006, Mark Rosengarten At a pH of 2: Methyl Orange = red Bromthymol Blue = yellow Phenolphthalein = colorless Litmus = red Bromcresol Green = yellow Thymol Blue = yellow Methyl orange is red at a pH of 3.2 and below and yellow at a pH of 4.4 and higher. In between the two numbers, it is an intermediate color that is not listed on this table. website

155. Alternate Theories (c) 2006, Mark Rosengarten  Arrhenius Theory: acids and bases must be in aqueous solution.  Alternate Theory: Not necessarily so!  Acid: proton (H +1 ) donor…gives up H +1 in a reaction.  Base: proton (H +1 ) acceptor…gains H +1 in a reaction.  HNO 3 + H 2 O  H 3 O +1 + NO 3 -1  Since HNO 3 lost an H +1 during the reaction, it is an acid.  Since H 2 O gained the H +1 that HNO 3 lost, it is a base. Website-video

156. Oxidation and Reduction (c) 2006, Mark Rosengarten 1) Oxidation NumbersOxidation Numbers 2) Identifying OX, RD and SI SpeciesIdentifying OX, RD and SI Species 3) AgentsAgents 4) Writing Half-ReactionsWriting Half-Reactions 5) Balancing Half-ReactionsBalancing Half-Reactions 6) Activity SeriesActivity Series 7) Voltaic CellsVoltaic Cells 8) Electrolytic CellsElectrolytic Cells 9) ElectroplatingElectroplating

157. Oxidation Numbers (c) 2006, Mark Rosengarten  Elements have no charge until they bond to other elements.  Na 0, Li 0, H 2 0. S 0, N 2 0, C 60 0  The formula of a compound is such that the charges of the elements making up the compound all add up to zero.  The symbol and charge of an element or polyatomic ion is called a SPECIES.  Determine the charge of each species in the following compounds:  NaClKNO 3 CuSO 4 Fe 2 (CO 3 ) 3 Video website

158. Identifying OX, RD, SI Species (c) 2006, Mark Rosengarten  Ca 0 + 2 H +1 Cl -1  Ca +2 Cl -1 2 + H 2 0  Oxidation = loss of electrons. The species becomes more positive in charge. For example, Ca 0  Ca +2, so Ca 0 is the species that is oxidized. Oxidation  Reduction = gain of electrons. The species becomes more negative in charge. For example, H +1  H 0, so the H +1 is the species that is reduced. Reduction  Spectator Ion = no change in charge. The species does not gain or lose any electrons. For example, Cl -1  Cl -1, so the Cl -1 is the spectator ion. website

159. Agents (c) 2006, Mark Rosengarten  Ca 0 + 2 H +1 Cl -1  Ca +2 Cl -1 2 + H 2 0  Since Ca 0 is being oxidized and H +1 is being reduced, the electrons must be going from the Ca 0 to the H +1.  Since Ca 0 would not lose electrons (be oxidized) if H +1 weren’t there to gain them, H +1 is the cause, or agent, of Ca 0 ’s oxidation. H +1 is the oxidizing agent.  Since H +1 would not gain electrons (be reduced) if Ca 0 weren’t there to lose them, Ca 0 is the cause, or agent, of H +1 ’s reduction. Ca 0 is the reducing agent.

160. Writing Half-Reactions (c) 2006, Mark Rosengarten  Ca 0 + 2 H +1 Cl -1  Ca +2 Cl -1 2 + H 2 0  Oxidation: Ca 0  Ca +2 + 2e -  Reduction: 2H +1 + 2e -  H 2 0 The two electrons lost by Ca 0 are gained by the two H +1 (each H +1 picks up an electron). PRACTICE SOME! Video

161. Practice Half-Reactions (c) 2006, Mark Rosengarten  Don’t forget to determine the charge of each species first!  4 Li + O 2  2 Li 2 O  Oxidation Half-Reaction:  Reduction Half-Reaction:  Zn + Na 2 SO 4  ZnSO 4 + 2 Na  Oxidation Half-Reaction:  Reduction Half-Reaction:

162. Balancing Half-Reactions (c) 2006, Mark Rosengarten  Ca 0 + Fe +3  Ca +2 + Fe 0  Ca’s charge changes by 2, so double the Fe.  Fe’s charge changes by 3, so triple the Ca.  3 Ca 0 + 2 Fe +3  3 Ca +2 + 2 Fe 0  Try these:  __Na 0 + __H +1  __Na +1 + __H 2 0  (hint: balance the H and H 2 first!)  __Al 0 + __Cu +2  __Al +3 + __Cu 0 website

163. Activity Series (c) 2006, Mark Rosengarten  For metals, the higher up the chart the element is, the more likely it is to be oxidized. This is because metals like to lose electrons, and the more active a metallic element is, the more easily it can lose them.metals  For nonmetals, the higher up the chart the element is, the more likely it is to be reduced. This is because nonmetals like to gain electrons, and the more active a nonmetallic element is, the more easily it can gain them.

164. Metal Activity (c) 2006, Mark Rosengarten  Metallic elements start out with a charge of ZERO, so they can only be oxidized to form (+) ions.  The higher of two metals MUST undergo oxidation in the reaction, or no reaction will happen.  The reaction 3 K + FeCl 3  3 KCl + Fe WILL happen, because K is being oxidized, and that is what Table J says should happen.  The reaction Fe + 3 KCl  FeCl 3 + 3 K will NOT happen. 3 K 0 + Fe +3 Cl -1 3 REACTION Fe 0 + 3 K +1 Cl -1 NO REACTION

165. Voltaic Cells (c) 2006, Mark Rosengarten  Produce electrical current using a spontaneous redox reaction  Used to make batteries!batteries  Materials needed: two beakers, piece of the oxidized metal (anode, - electrode), solution of the oxidized metal, piece of the reduced metal (cathode, + electrode), solution of the reduced metal, porous material (salt bridge), solution of a salt that does not contain either metal in the reaction, wire and a load to make use of the generated current!salt bridge  Use Reference Table J to determine the metals to use  Higher = (-) anodeLower = (+) cathode Animation

166. Making Voltaic Cells (c) 2006, Mark Rosengarten Create Your Own Cell!!!! More Info!!!

167. How It Works (c) 2006, Mark Rosengarten  The Zn 0 anode loses 2 e -, which go up the wire and through the load. The Zn 0 electrode gets smaller as the Zn 0 becomes Zn +2 and dissolves into solution. The e - go into the Cu 0, where they sit on the outside surface of the Cu 0 cathode and wait for Cu +2 from the solution to come over so that the e - can jump on to the Cu +2 and reduce it to Cu 0. The size of the Cu 0 electrode increases. The negative ions in solution go over the salt bridge to the anode side to complete the circuit. Since Zn is listed above Cu, Zn 0 will be oxidized when it reacts with Cu +2. The reaction: Zn + CuSO 4  ZnSO 4 + Cu

168. You Start At The Anode (c) 2006, Mark Rosengarten

169. Make Your Own Cell!!! (c) 2006, Mark Rosengarten

170. Electrolytic Cells (c) 2006, Mark Rosengarten  Use electricity to force a nonspontaneous redox reaction to take place.  Uses for Electrolytic Cells:  Decomposition of Alkali Metal Compounds Decomposition of Alkali Metal Compounds  Decomposition of Water into Hydrogen and Oxygen Decomposition of Water into Hydrogen and Oxygen  Electroplating Electroplating  Differences between Voltaic and Electrolytic Cells:  ANODE: Voltaic (-) Electrolytic (+) ANODE  CATHODE: Voltaic (+) Electrolytic (-) CATHODE  Voltaic: 2 half-cells, a salt bridge and a load  Electrolytic: 1 cell, no salt bridge, IS the load Website & Video Video

171. Decomposing Alkali Metal Compounds (c) 2006, Mark Rosengarten 2 NaCl  2 Na + Cl 2 The Na +1 is reduced at the (-) cathode, picking up an e - from the battery The Cl -1 is oxidized at the (+) anode, the e - being pulled off by the battery (DC)

172. Decomposing Water (c) 2006, Mark Rosengarten 2 H 2 O  2 H 2 + O 2 The H + is reduced at the (-) cathode, yielding H 2 (g), which is trapped in the tube. The O -2 is oxidized at the (+) anode, yielding O 2 (g), which is trapped in the tube.

173. Electroplating (c) 2006, Mark Rosengarten The Ag 0 is oxidized to Ag +1 when the (+) end of the battery strips its electrons off. The Ag +1 migrates through the solution towards the (-) charged cathode (ring), where it picks up an electron from the battery and forms Ag 0, which coats on to the ring.

174. Organic Chemistry (c) 2006, Mark Rosengarten 1) Hydrocarbons Hydrocarbons 2) Substituted HydrocarbonsSubstituted Hydrocarbons 3) Organic FamiliesOrganic Families 4) Organic ReactionsOrganic Reactions

175. Hydrocarbons (c) 2006, Mark Rosengarten  Molecules made of Hydrogen and Carbon  Carbon forms four bonds, hydrogen forms one bond  Hydrocarbons come in three different homologous series:  Alkanes (single bond between C’s, saturated) Alkanes  Alkenes (1 double bond between 2 C’s, unsaturated) Alkenes  Alkynes (1 triple bond between 2 C’s, unsaturated) Alkynes  These are called aliphatic, or open-chain, hydrocarbons.  Count the number of carbons and add the appropriate suffix! website

176. Vocabulary (c) 2006, Mark Rosengarten

177. Nucleon – particle found in the nucleus of an atom – includes the proton and neutron only – equal to the mass number of an atom Isotope – atoms of the same element which have the same atomic number but different mass number Atomic Number - equal to the number of protons in the nucleus of an atom Mass Number - equal to the sum of the protons and neutrons in the nucleus of an atom. Nuclear charge - equal to the number of protons in the nucleus of an atom. Alpha Particle – A radioactive particle equivalent to a helium nucleus (2 protons, 2 neutrons) - Mass of 4 and a +2 charge Beta Particle – A radioactive particle equivalent to an electron. Has no mass and -1 charge Positron – A radioactive particle equivalent to an positively charged electron. Has no mass and +1 charge Gamma Rays – High energy light given off during a nuclear process – Have no mass or charge Fission – A nuclear reaction where a large nucleus breaks up into smaller ones. This is what happens in nuclear power plants Fusion - process where two or more small nuclei combine to form a larger nucleus. Fusion is the reverse process of nuclear fission

178. (c) 2006, Mark Rosengarten Valence Electron – Electrons in the outermost energy level (furthest away from the nucleus) – Generally the only electrons involved in chemical reactions. Electron Dot Diagram (EDD) – Symbol of an element surrounded by dots which represent valence elctrons Stable Octet - The octet rule is a rule that states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. Bright Line Spectrum - When electrons jump from the excited state to the ground state, the electrons emit energy in the form of light, producing a bright-line spectrum. Each element has its own unique bright-line spectrum. Orbital- Regions of the most probable electron location in the wave-mechanical model of the atom Solid – phase of matter with a definite shape and volume and low entropy – particles arranged in a regular geometric pattern

179. (c) 2006, Mark Rosengarten Liquid – phase of matter that has a definite volume but takes shape of its container Gas – phase of matter that takes the shape of & fills its entire container – has high entropy. Element - substances that are composed of atoms that have the same atomic number. Elements cannot be broken down by chemical change. Compound - substance composed of two or more different elements that are chemically combined in a fixed proportion. A chemical compound can be broken down by chemical means. Mixture - composed of two or more different substances that can be separated by physical means.When different substances are mixed together, a homogeneous or heterogeneous mixture is formed. Homogeneous Mixture – Components are evenly distributed – Also called solutions. Heterogeneous Mixture – Components are unevenly distributed

180. (c) 2006, Mark Rosengarten Solution - a homogeneous mixture of a solute dissolved in a solvent Melting – a phase change in which a solid changes to a liquid Boiling – the process of rapidly converting a liquid to its gaseous (vapor) state, typically by heating the liquid to a temperature called its boiling point. Boiling Point - temperature at which the vapor pressure is equal to the pressure of the gas above it. Freezing - a phase change in which a liquid cools and changes to a solid Condensation - a phase change in which a gas cools and changes to a liquid Sublimation - When a solid can change directly into a gas skipping the liquid phase Evaporation – a phase change in which a liquid changes to a gas Exothermic – process which releases energy causing the temperature of its surroundings to increase

181. (c) 2006, Mark Rosengarten Endothermic – process which absorbs energy causing the temperature of its surroundings to decrease Heat of Fusion – Amount of heat in Joules or KF required to melt 1 gram of ice to water Heat of Vaporization - Amount of heat in Joules or KF required to vaporize 1 gram of water to vapor Metals – Found on the left side of the Periodic table – Are malleable, ductile, lustrous, good conductors and form positive ions Malleable – can be pounded into thin sheets Ductile – can be stretched into wire Luster - shiny Nonmetals – Found on the right side of the Periodic table – Are brittle, dull, poor conductors and form negative ions Metalloids – have the properties of both metals and nonmetals – found along the stair-step line Ionization energy – amount of energy required to remove the most loosely held electron in an atom – Values found on table S

182. (c) 2006, Mark Rosengarten Electronegativity – The attraction a nucleus has for electrons in a bond – Values found on Table S – Fluorine has highest Alkali metals – Group 1 metals - Most active metals, only found in compounds in nature – Form +1 ions Alkaline Earth Metals – Group 2 metals - Very active metals, only found in compounds in nature – Form +2 ions Transition Metals – Groups 3-11 - Many can form different possible charges of ions - Compounds containing these metals can be colored. Halogens – Group 17 nonmetals – Most reactive nonmetals – Fluorine most active Noble Gases - Are completely nonreactive since they have eight valence electrons, making a stable octet. Diatomic Elements - Br 2, I 2, N 2, Cl 2, H 2, O 2 and F 2 Ions - charged particles formed by the gain or loss of electrons.

183. (c) 2006, Mark Rosengarten Positive Ion – Formed when an atom, usually a metal, loses 1 or more electrons Negative Ion – Formed when an atom, usually a nonmetal, gains 1 or more electrons. Ionic bond – bond that forms when a metal transfers valence electrons to a nonmetal. Covalent bond – bond that forms when nonmetals share valence electrons Metallic Bond – Bond that forms between metal atoms such as in copper wire – Described as “positve ions in a sea of mobile electrons” Ionic Compound - made of metal and nonmetal ions. Molecular Compound - made of nonmetal atoms bonded to form a distinct particle called a molecule. REDOX Reaction – Short for oxidation-reduction - driven by the loss (oxidation) and gain (reduction) of electrons.

184. (c) 2006, Mark Rosengarten Oxidation – loss of electrons – oxidation number increases Reduction – gain of electrons – oxidation number decreases Precipitate – compound that forms as a result of a double replacement reaction which is insoluble in water Intermolecular Attractive Forces(IMAF) – force of attraction between molecules such as hydrogen bonding, dipole-dipole, etc… Hydrogen Bond – A special type of dipole-dipole attraction that occurs when hydrogen is bonded to N, O or F. Gram Formula Mass - sum of atomic masses of all elements in the compound – equal to the mass of one mole of a compound Catalyst – speeds up a chemical reaction by lowering the activation energy. Activation Energy – amount of energy needed to start a reaction Heat of Reaction(  H) – amount of heat absorbed or released during a chemical reaction

185. (c) 2006, Mark Rosengarten Chemical Equilibrium – When the rate of the forward and reverse reactions are equal Solubility - the maximum quantity of solute that can be dissolved in a given quantity of solvent at a given temperature Arrhenius Acid - molecules that dissolve in water to produce H+ or H 3 O + (hydronium) as the only positively charged ion in solution. Arrhenius Base - molecules that dissolve in water to produce OH- (hydroxide) as the only negatively charged ion in solution. Bronsted-Lowry Acid – proton (H+) donor Bronsted-Lowry Base – proton (H+) acceptor Voltaic cell - Produce electrical current using a spontaneous redox reaction – used to make batteries Electrolytic cell - Use electricity to force a nonspontaneous redox reaction to take place. Anode – electrode at which oxidation occurs Cathode – electrode at which reduction occurs

186. (c) 2006, Mark Rosengarten Salt bridge – allows for the movement of ions Hydrocarbons - Molecules made of Hydrogen and Carbon Alkanes – saturated hydrocarbons with only single bonds between carbon atoms Alkenes – unsaturated hydrocarbons with at least one double bond between carbon atoms Alkynes – unsaturated hydrocarbons with at least one triple bond between carbon atoms Esterification - reaction between an alcohol and organic acid which produces an ester and water Fermentation – reaction of a sugar with an enzyme that produces alcohol and CO 2 Polymerization – process of joining many small molecules(monomers) to make a large molecule(polymer). Saponification – A fat or oil reacts with a strong base and produces a soap

187. (c) 2006, Mark Rosengarten Isomer – compounds that have the same chemical formula but different structures