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Bell Work Prepare for Quiz

Bell Work Prepare for Quiz. Write these in your Bell Work Composition Book. Hg Br Kr K S. Atomic # Atomic weight # of protons # of electrons # of neutrons Group Period. Write these in your Bell Work Composition Book. Hg Br Kr K S. Find the Atomic Weight for MgSO 4.

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Bell Work Prepare for Quiz

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  1. Bell Work Prepare for Quiz

  2. Write these in your Bell Work Composition Book • Hg • Br • Kr • K • S • Atomic # • Atomic weight • # of protons • # of electrons • # of neutrons • Group • Period

  3. Write these in your Bell Work Composition Book • Hg • Br • Kr • K • S

  4. Find the Atomic Weight for MgSO4 • Bell work composition book

  5. Today’s ElementLiquid Metals

  6. Zinc Chemical Properties Physical Properties Periodic Table Importance • Atomic # • Atomic mass • # of Protons • # of Electrons • # of Neutrons • Period • Group

  7. What’s in a Battery? • Modern batteries use a variety of chemicals to power their reactions. Common battery chemistries include: • Zinc-carbon battery: The zinc-carbon chemistry is common in many inexpensive AAA, AA, C and D dry cell batteries. The anode is zinc, the cathode is manganese dioxide, and the electrolyte is ammonium chloride or zinc chloride. • Alkaline battery: This chemistry is also common in AA, C and D dry cell batteries. The cathode is composed of a manganese dioxide mixture, while the anode is a zinc powder. It gets its name from the potassium hydroxide electrolyte, which is an alkaline substance. • Lithium-ion battery (rechargeable): Lithium chemistry is often used in high-performance devices, such as cell phones, digital cameras and even electric cars. A variety of substances are used in lithium batteries, but a common combination is a lithium cobalt oxide cathode and a carbon anode. • Lead-acid battery (rechargeable): This is the chemistry used in a typical car battery. The electrodes are usually made of lead dioxide and metallic lead, while the electrolyte is a sulfuric acid solution.

  8. Water: Separation by Electrolysis Video of Electrolysis: Water to Hydrogen and Oxygen

  9. Atomic Mass of a Compound • H2O Try these • CO2 • C6H12O6 H2 1.01 x 2 = 2.02 O 16 x 1 = 16.00 Add totals 2.02 + 16.00 18.02 1.01 + 1.01 + 16 = 18.02

  10. Practice – Finding Atomic Mass • CO2 C 12.01 x 1 = 12.01 O2 16 x 2 = 32.00 Add totals 12.01 + 32.00 44.01

  11. Practice – Finding Atomic Mass • C6H12O6 C6 12.01 x 6 = 72.06 H12 1.01 x 12 = 12.12 O2 16 x 6 = 96.00 Add totals 72.06 12.12 +96.00 108.18

  12. Percent Composition of Mass for Mixtures • A 6g mixture of sulfur and iron is separated using a magnet. Data Sulfur (S) Iron (Fe) 5g 1g • Calculate the percent composition of S and Fe.

  13. Percent Composition of Mass for Mixtures • A 6g mixture of sulfur and iron is separated using a magnet. Data Sulfur (S) Iron (Fe) 5g 1g • Calculate the percent composition of S and Fe. Part / Whole x 100 = % composition Sulfur: 5g/6g x 100 = Iron : 1g/6g x 100 =

  14. Percent Composition of Mass for Mixtures • A 6g mixture of sulfur and iron is separated using a magnet. Data Sulfur (S) Iron (Fe) 5g 1g • Calculate the percent composition of S and Fe. Part / Whole x 100 = % composition Sulfur: 5g/6g x 100 = 83.33% S Iron : 1g/6g x 100 = 16.66% Fe

  15. Use Percent Composition to find the composition of a compound • Use the periodic table to find the compound’s percent composition of each element. • List the atomic weight of each element in the compound • Note how many of each type of atom is in the compound • Add it all up to get the atomic weight of the whole compound

  16. Atomic Mass of a Compound • H2O Try these • CO2 • C6H12O6 H2 1.01 x 2 = 2.02 O 16 x 1 = 16.00 Add totals 2.02 + 16.00 18.02 1.01 + 1.01 + 16 = 18.02

  17. Practice – Percent Composition H2O part / whole x 100 = % composition % composition of H % composition of O

  18. Practice – Percent Composition CO2 part / whole x 100 = % composition % composition of C % composition of O

  19. Practice – Percent Composition C6H12O6 part / whole x 100 = % composition % composition of C % composition of H % composition of O

  20. Law of Conservation of Mass • Mass is neither created nor destroyed in any process. It is conserved. Mass reactants = Mass products 2H2O + electricity yields 2H2 + O2

  21. Isotopes • The atomic weight found on the periodic table is based on the average weight of all the isotopes of the element • Isotope – atoms of the same element with the same number of protons but different numbers of neutrons • M&M activity

  22. Writing Isotopes

  23. Reading Isotopes Mass number - the sum of the protons and neutrons

  24. Isotopes of Hydrogen

  25. Write Isotopes for Iron

  26. More isotopes Argon 36, Argon 37…

  27. M&Mium Isotope Activity

  28. http://www.chem.memphis.edu/bridson/FundChem/T07a1100.htm

  29. S.I.Units • http://2012books.lardbucket.org/books/general-chemistry-principles-patterns-and-applications-v1.0/section_05.html#averill_1.0-ch01_s09_s01_s02_t02

  30. http://chemwiki.ucdavis.edu/Physical_Chemistry/Atomic_Theory/The_Mole_and_Avogadro's_Constanthttp://chemwiki.ucdavis.edu/Physical_Chemistry/Atomic_Theory/The_Mole_and_Avogadro's_Constant

  31. Measurements and Calculations Where to Round Song

  32. Steps in the Scientific Method • 1. Observations • - quantitative • - qualitative • 2. Formulating hypotheses • - possible explanation for the observation • 3. Performing experiments • - gathering new information to decide whether the hypothesis is valid

  33. Outcomes Over the Long-Term • Theory (Model) • - A set of tested hypotheses that give an overall explanation of some natural phenomenon. • Natural Law • - The same observation applies to many different systems • - Example - Law of Conservation of Mass

  34. Law vs. Theory • A law summarizes what happens • A theory (model) is an attempt to explain why it happens.

  35. Nature of Measurement Measurement - quantitative observation consisting of 2 parts • Part 1 - number • Part 2 - scale (unit) • Examples: • 20grams • 6.63 x 10-34Joule seconds

  36. The Fundamental SI Units(le Système International, SI)International System of Unitsa system of measurement units agreed on by scientists to aid in the comparison of results worldwide.

  37. SI Units

  38. Metric PrefixesCommon to Chemistry

  39. Metric Prefixes and Conversion Examples

  40. Uncertainty in Measurement • A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.

  41. Why Is there Uncertainty? • Measurements are performed with instruments • No instrument can read to an infinite number of decimal places Which of these balances has the greatest uncertainty in measurement?

  42. Precision and Accuracy • Accuracyrefers to the agreement of a particular value with the truevalue. • Precisionrefers to the degree of agreement among several measurements made in the same manner. Precise but not accurate Precise AND accurate Neither accurate nor precise

  43. Types of Error • Random Error(Indeterminate Error) - measurement has an equal probability of being high or low. • Systematic Error(Determinate Error) - Occurs in the same directioneach time (high or low), often resulting from poor technique or incorrect calibration.

  44. Rules for Counting Significant Figures - Details • Nonzero integersalways count as significant figures. • 3456has • 4sig figs.

  45. Rules for Counting Significant Figures - Details • Zeros • -Leading zeros do not count as significant figures. • 0.0486 has • 3 sig figs.

  46. Rules for Counting Significant Figures - Details • Zeros • -Captive zeros always count as significant figures. • 16.07 has • 4 sig figs.

  47. Rules for Counting Significant Figures - Details • Zeros • Trailing zerosare significant only if the number contains a decimal point. • 9.300 has • 4 sig figs.

  48. Rules for Counting Significant Figures - Details • Exact numbershave an infinite number of significant figures. • 1 inch = 2.54cm, exactly

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