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Semester Review

EVERYTHING YOU NEED TO KNOW!. Semester Review . Types of Research. Basic Research- done to increase knowledge Applied Research- done to solve a problem Technological Development- done to improve quality of life

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Semester Review

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  1. EVERYTHING YOU NEED TO KNOW! Semester Review

  2. Types of Research • Basic Research- done to increase knowledge • Applied Research- done to solve a problem • Technological Development- done to improve quality of life • technology – application of knowledge (usually scientific) for practical purposes

  3. Types of Properties • Physical- characteristics that can be observed or measured without changing the identity of a substance • Chemical- relates to a substances ability to undergo changes that transform it into a different substance • Easiest to see when a chemical is reacting

  4. Important SI Base Units

  5. Prefixes • Prefixes are added to the base unit names to represent quantities smaller or larger

  6. Volume • amount of space occupied by object • SI: m3 = m x m x m • We will use mL (cm3) more than m3 • non-SI: 1 liter = 1000cm3 = 1000mL

  7. Density • ratio of mass to volume • SI: • characteristic property of substance (doesn’t change with amount ) because as volume increases, mass also increases • density usually decreases as T increases • exception: ice is less dense than liquid water so it floats

  8. Density Example • Solve the following problem: What is the Density of an object with a mass of 25g and a volume of 50 ml?

  9. Example • What is 298 K in Fahrenheit?

  10. Using Sig Figs in Calculations • Adding/Subtracting: • end with the least number of decimal places

  11. Adding and Subtracting Example • Solve the following problem: 275.25 cm + 25 = ?

  12. Using Sig Figs in Calculations • Multiplying/Dividing: • end with the least number of sig figs

  13. Multiplying and Dividing Example • Solve the following problem & report to correct number of significant figures: (2.07 x 107)(4.2 x 109) =

  14. Dalton’s Atomic Theory • All mass is made of atoms • Atoms of same element have the same size, mass, and properties • Atoms can’t be subdivided, created or destroyed • Atoms of different element combine in whole number ratios to make compounds • In chemical reactions, atoms are combined, separated, and rearranged.

  15. Discovery of Electron • J.J. Thomson (English, 1897) led the experiments that proved the cathode rays were composed of negative particles • found ratio of charge to mass of this particle • since the ratio stayed constant for any metal that contained it, it must be the same in all of the metals • named this particle the ELECTRON • Robert Millikan (American, 1909) proved Thomson’s work (definitively determined the mass and charge of the electron

  16. Discovery of Nucleus • Rutherford discovered the nucleus by shooting alpha particles (have positive charge) at a very thin piece of gold foil • he predicted that the particles would go right through the foil at some small angle

  17. atoms of the same element with different numbers of neutrons • most elements exist as a mixture of isotopes • What do the Carbon isotopes below have in common? What is different about them?

  18. The mole • 6.022x1023 is called Avogadro’s Number in honor of all of his contributions to chemistry • can be used as a conversion factor between a number of things and mole

  19. Conversion Factors # Atoms Grams Moles Use Molar Mass: grams per mole Use Avog.’s Number: atoms per mole

  20. Mole Example #1 • 14 g of Carbon = ___moles = ___ atoms

  21. Mole Example #2 • How many mols are in 105g of sodium chloride?

  22. Wave Calculations • Wavelength (λ) - distance between two peaks . Measured in meters • Frequency (v) - number of peaks that pass a point each second. • Hz = Hertz = s-1 • c = λ v  • where c = 3.0 x 108 m/s

  23. Electrons start in lowest possible level - ground state. • Absorb energy - become excited and shift upward. • Dropping back down - emits photons (packets of energies equal to the previously absorbed energy). • Hydrogen Emission Spectrum

  24. Highest Frequency and Lowest Frequency

  25. Nucleus: • contains protons and neutrons • takes up very little space • Electron Cloud: • contains electrons • takes up most of space

  26. One s orbital • spherical

  27. Three p orbitals

  28. Order for Filling Sublevels

  29. Rules for Arrangements • Aufbau Principle- an electron occupies the lowest-energy orbital that can receive it • Beginning in the 3rd energy level, the energies of the sublevels begin to overlap

  30. Rules for Arrangements • Pauli Exclusion Principle- no two electrons in the same atom can occupy the same place at the same time • Hund’s Rule- orbitals of a sublevel are each occupied by one electron before any orbital is occupied by a second • all unpaired electrons must have the same spin

  31. Electron Configuration Examples Refer to the Electron Configuration Video! “Electron Configuration Part 1” http://www.youtube.com/user/BryanMossChemistry?feature=mhum#p/u/1/xsEm8yNN0Hg

  32. Mendeleev • Russian, Dmitri Mendeleev • when he arranged them by atomic mass, he found similar properties at certain intervals • published the first periodic table in 1869 • left empty spaces where he predicted undiscovered elements should be • confirmed his predictions and persuaded other chemists

  33. Moseley • In 1911, Henry Moseley (English) found that the pattern worked best if arranged by number of protons • Our current periodic tables use this method or arrangement

  34. Noble Gases • Contain a octet- 8 electrons in its outer main energy level • Makes them inert or very stable so they do not react easily with other elements

  35. Periodic Groups • Refer to the “Periodic Families Ipod Lab” Video. • http://www.youtube.com/user/BryanMossChemistry?feature=mhum#p/u/33/DMT6nP4oeYc

  36. Ionization Energy • Increases across a period • Decreases down a column

  37. Ionization Energy

  38. Electron Affinity Electron Affinity – the energy change when an electron is added to a gaseous neutral atom • exothermic (-) A + e-  A-+ energy

  39. Electron Affinity

  40. Electronegativity • applies when an atom is in a compound NOT alone • Electronegativity – measure of how strongly an atom attracts electrons when it is in a compound • Fluorine (the most electronegative element) is assigned a 4.0 and then all the others were determined by comparison

  41. Electronegativity

  42. Atomic Radii • Decrease across a period • Increase down a column • Smallest Helium • Largest Francium

  43. Atomic/Ionic Radii

  44. Specific Ions • Cation – positive ion caused by an atom losing electrons (metals only) • Anion – negative ion caused by an atom gaining electrons (non metals only) • Metalloids can be either cations or anions

  45. Naming Compounds • Refer to the “Naming Compounds Practice Problems” • http://www.youtube.com/user/BryanMossChemistry?feature=mhum#p/u/4/7NvRYVDIOCA

  46. Formula Mass • Refer to the “Average Atomic Mass Gram Formula Mass” Video • http://www.youtube.com/user/BryanMossChemistry?feature=mhum#p/u/44/9bcp-aRnuak

  47. Chemical Bonds • atoms rarely exist alone • when atoms are bonded together, they have less potential energy and are more stable • What is potential energy? • chemical bond – mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

  48. Ionic vs. Covalent

  49. Ionic vs. Covalent

  50. Octet Rule • representative elements can “fill” their outer energy level by sharing electrons in covalent bonds • Octet Rule- a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing or sharing electrons • Duet Rule- applies to H and He

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