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Outline for Review

Outline for Review. 1) The Atom (Nuclear, Electron Config) 2) Matter (Phases, Types, Changes) 3) Bonding (Periodic Table, Ionic, Covalent) 4) Compounds (Formulas, Reactions, IMAF’s) 5) Math of Chemistry (Formula Mass, Gas Laws, Neutralization, etc.)

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Outline for Review

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

  2. The Atom 1) Nucleons 2) Isotopes 3) Natural Radioactivity 4) Half-Life 5) Nuclear Power 6) Electron Configuation 7) Development of the Atomic Model (c) 2006, Mark Rosengarten

  3. Nucleons • Protons: +1 each, determines identity of element, mass of 1 amu, determined using atomic number, nuclear charge • Neutrons: no charge, determines identity of isotope of an element, 1 amu, determined using mass number - atomic number (amu = atomic mass unit) • 3216S and 3316S 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. (c) 2006, Mark Rosengarten

  4. Isotopes • 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. • Rounding off the atomic mass on the Periodic Table will tell you what the most common isotope of that element is. (c) 2006, Mark Rosengarten

  5. Weight-Average Atomic Mass • 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) (c) 2006, Mark Rosengarten

  6. Most Common Isotope • 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: • Co Ag • S Pb • 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. (c) 2006, Mark Rosengarten

  7. Natural Radioactivity • Alpha Decay • Beta Decay • Positron Decay • Gamma Decay • 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. (c) 2006, Mark Rosengarten

  8. Alpha Decay • The nucleus ejects two protons and two neutrons. The atomic mass decreases by 4, the atomic number decreases by 2. • 23892U  (c) 2006, Mark Rosengarten

  9. Beta Decay • 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. • 146C  (c) 2006, Mark Rosengarten

  10. Positron Decay • 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. • 5326Fe  (c) 2006, Mark Rosengarten

  11. Gamma Decay • 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. (c) 2006, Mark Rosengarten

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

  13. Half-Life • 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 backwards in time • Radioactive Dating (c) 2006, Mark Rosengarten

  14. Going Forwards in Time • 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 (c) 2006, Mark Rosengarten

  15. Going Backwards in Time • 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 (c) 2006, Mark Rosengarten

  16. Radioactive Dating • 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 (c) 2006, Mark Rosengarten

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

  18. Artificial Transmutation • 4020Ca + _____ -----> 4019K + 11H • 9642Mo + 21H -----> 10n + _____ • Nuclide + Bullet --> New Element + Fragment(s) • The masses and atomic numbers must add up to be the same on both sides of the arrow. (c) 2006, Mark Rosengarten

  19. Particle Accelerators • 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. (c) 2006, Mark Rosengarten

  20. Nuclear Fission • 23592U + 10n 9236Kr + 14156Ba + 3 10n + 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. (c) 2006, Mark Rosengarten

  21. Chain Reaction (c) 2006, Mark Rosengarten

  22. Nuclear Fusion • 21H + 21H 42He + 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. (c) 2006, Mark Rosengarten

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

  24. Basic Configuration • 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) (c) 2006, Mark Rosengarten

  25. Valence Electrons • The valence electrons are responsible for all chemical bonding. • 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. (c) 2006, Mark Rosengarten

  26. Electron-Dot Diagrams • 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 (c) 2006, Mark Rosengarten

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

  28. Excited vs. Ground State • 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 (c) 2006, Mark Rosengarten

  29. What Is Light? • 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. • EXAMPLE SPECTRUM (c) 2006, Mark Rosengarten

  30. EXAMPLE SPECTRUM 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. (c) 2006, Mark Rosengarten

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

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

  33. Rutherford Model • 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 shot at a thin sheet of gold foil: most go through (empty space). Some deflect or bounce off (small + charged nucleus). (c) 2006, Mark Rosengarten

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

  35. Quantum-Mechanical Model • Electron energy levels are wave functions. • Electrons are found in orbitals, regions of space where an electron is most likely to be found. • 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. (c) 2006, Mark Rosengarten

  36. Matter 1) Properties ofPhases 2) Types of Matter 3) Phase Changes (c) 2006, Mark Rosengarten

  37. Properties of Phases • Solids: Crystal lattice (regular geometric pattern), vibration motion only • Liquids: particles flow past each other but are still attracted to each other. • 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) (c) 2006, Mark Rosengarten

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

  39. Liquids 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. (c) 2006, Mark Rosengarten

  40. Gases 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 1023) gas molecules. Since there is space between gas molecules, gases are affected by changes in pressure. (c) 2006, Mark Rosengarten

  41. Types of Matter • Substances (Homogeneous) • Elements (cannot be decomposed by chemical change): Al, Ne, O, Br, H • Compounds (can be decomposed by chemical change): NaCl, Cu(ClO3)2, KBr, H2O, C2H6 • Mixtures • Homogeneous: Solutions (solvent + solute) • Heterogeneous: soil, Italian dressing, etc. (c) 2006, Mark Rosengarten

  42. Elements • 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. (c) 2006, Mark Rosengarten

  43. Compounds • Lead has two charges listed, +2 and +4. This is a sample of lead (II) chloride (PbCl2). 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 (PbCl4). Notice how both samples of lead compounds have consistent composition throughout? Compounds are homogeneous! (c) 2006, Mark Rosengarten

  44. Mixtures • 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 PbCl2 and PbCl4 formula units. Again, they are in no particular ratio to each other and can be separated without chemical change. (c) 2006, Mark Rosengarten

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

  46. Phase Change Types (c) 2006, Mark Rosengarten

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

  48. Heat of Phase Change • How many joules would it take to melt 100. g of H2O (s) at 0oC? • q=mHf = (100. g)(334 J/g) = 33400 J • How many joules would it take to boil 100. g of H2O (l) at 100oC? • q=mHv = (100.g)(2260 J/g) = 226000 J (c) 2006, Mark Rosengarten

  49. Evaporation • 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. • 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). (c) 2006, Mark Rosengarten

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

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