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The Structure of Atoms and Mole Theory

The Structure of Atoms and Mole Theory. Element vs. Compound. Element : is a pure chemical substance made of one type of atom. They are classified as metals, nonmetals or semimetals. For example: Na is a metal in elemental form. Cl forms Cl 2 , a diatomic gas in elemental form.

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The Structure of Atoms and Mole Theory

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  1. The Structure of Atoms and Mole Theory

  2. Element vs. Compound • Element: is a pure chemical substance made of one type of atom. They are classified as metals, nonmetals or semimetals. For example: Na is a metal in elemental form. Cl forms Cl2, a diatomic gas in elemental form. • No charge (neutral state) • Compound: is made up of a combination of two or more elements (more specifically atoms from these elements). For example: An ionic compound is made of a metal and a nonmetal (NaCl). • Charge: Na is 1+ and Cl is 1‒ chemical formula • Chemical formula: 1) Molecular: C12H22O11 (sucrose aka table sugar) 2) Empirical: NaCl (table salt) 2 HgO 2 Hg + O2 chemical equation

  3. Periodic Trends

  4. Metals: shiny, silvery, soft (malleable), good conductors of heat/electricity, react violently in water, all solid (except Hg). • Nonmetals: not silvery (some colored), brittle, poor conductors, some solid, liquid (Br) and gas. • Semimetals (metalloids): Have properties that cross between metals and nonmetals. Partially conduct electricity (rather poor), silvery in appearance. All are solid. B (Boron), Si (Silicon), Ge (Germanium), As (Arsenic), Sb (Antimony), Te (Tellurium), At (Astatine)

  5. Group 1A: known as the alkali metals Li (Lithium), Na (Sodium), K (Potassium), Rb (Rubidium), Cs (Cesium), Fr (Francium) When part of an ionic compound, these cations have a 1+ oxidation state (An electron configuration that is full is preferred)

  6. Group 2A: known as the alkaline earth metals Be (Beryllium), Mg (Magnesium), Ca (Calcium), Sr (Strontium), Ba (Barium), Ra (Radium) When part of an ionic compound, these cations have a 2+ oxidation state (An electron configuration that is full is preferred)

  7. Group 7A: known as the halogens F (Fluorine), Cl (Chlorine), Br (Bromine), I (Iodine), At (Astatine) When part of an ionic compound, these anions have a 1‒ oxidation state (An electron configuration that is full is preferred)

  8. Group 8A: known as the noble gases He (Helium), Ne (Neon), Ar (Argon), Kr (Krypton), Xe (Xenon), Rn (Radon) Are rather inert and only under special circumstances do they form compounds.

  9. Hg(NO3)2(aq) + 2 KI(aq) HgI2(s) + 2 KNO3(aq) Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions. Aqueous solutions of mercury(II) nitrate and potassium iodide will react to form a precipitate of mercury(II) iodide and aqueous potassium iodide. 3.25 g + 3.32 g = 6.57 g 4.55 g + 2.02 g = 6.57 g

  10. Evolution of Atomic Theory https://www.youtube.com/watch?v=UDIprICe9kg Atomic Theory Fill in the Blank Worksheet

  11. Dalton’s Atomic Theory • Elements are made up of tiny particles called atoms. • Each element is characterized by the mass of its atoms. Atoms of the same element have the same mass, but atoms of different elements have different masses. • The chemical combination of elements to make different chemical compounds occurs when atoms bond together in small whole-number ratios. • Chemical reactions only rearrange how atoms are combined in chemical compounds; the atoms themselves don’t change.

  12. The Structure of Atoms: Electrons Cathode-Ray Tubes: J. J. Thomson (1856–1940) proposed that cathode rays must consist of tiny, negatively charged particles, which we now call electrons. Thomson was only able to measure the charge to mass ratio of an electron

  13. The Structure of Atoms: Electrons Was able to measure the mass of an electron and ultimately the charge (from Thomson) Robert Millikan’s Oil Drop Experiment

  14. The Structure of Atoms: Protons and Neutrons Rutherford’s Scattering Experiment

  15. The Structure of Atoms: Protons and Neutrons Atomic Nucleus: When Ernest Rutherford (1871–1937) directed a beam of alpha (α) particles at a thin gold foil, he found that almost all the particles passed through the foil undeflected. A very small number, however (about 1 in every 20,000), were deflected at an angle, and a few actually bounced back toward the particle source. Rutherford explained his results by proposing that a metal atom must be almost entirely empty space and have its mass concentrated in a tiny central core that he called the nucleus.

  16. The Structure of Atoms: Protons and Neutrons The mass of the atom is primarily in the nucleus. The charge of the proton is opposite in sign but equal to that of the electron.

  17. Song about Atomic Theory https://www.youtube.com/watch?v=07yDiELe83Y Can you create a song based on what you learned? Create a song using the people of atomic theory and a one liner about their contributions. One that rhymes and will help you remember the pioneers of atomic theory. You can get in groups of 2-3 to help. Be prepared to at least read what you have written. Feel free to use the fill in the blank handout.

  18. Atomic Number and Mass Number Atomic weight is also known as the average atomic mass. The Mass Number comes from this value unless given. Round to whole number.

  19. Some Definitions Atomic Number: number of protons in an atom (different elements have a different number of protons). This number defines the atom. Mass Number: Protons + Neutrons Isotopes: atoms of the same element that have different numbers of neutrons. Ion: atom that has gained or lost 1 or more electrons. Ex: Na+ has a charge of +1 because it has lost 1 electron. Ex: O-2 has a charge of -2 because it has gained 2 electrons.

  20. Determining # of protons, neutrons and electrons. Look up the element in the periodic table. Write its average atomic mass and atomic number down. Round its average atomic mass off to the nearest whole number to make it the mass number. Protons: atomic number Neutrons: Mass number – atomic number Electrons: atomic number – charge (no charge written down, the charge is zero) *if the element has a number after the symbol and a hyphen, use this as the mass number. It is an isotope.

  21. 15 14 N N 7 7 Isotope Notation (If Given) nitrogen-14 or N-14 mass number 7 protons 7 electrons 7 neutrons atomic number nitrogen-15 or N-15 mass number 7 protons 7 electrons 8 neutrons atomic number

  22. Isotope Notation (If Not Given) If you were just given N and asked to determine the number of electrons, protons, and neutrons. Look at periodic table. Top left is atomic number = number of protons = 7 Bottom is average atomic mass. The mass number will be 14. So, 14 = number of protons + number of neutrons 14 = 7 + number of neutrons Number of neutrons = 7 If there is no charge, then the number of protons equals the number of electrons. So, there are 7 electrons.

  23. Isotope Notation (If Not Given) What if you were given N-3 and asked to determine the number of electrons, protons, and neutrons. Look at periodic table. Top left is atomic number = number of protons = 7 Bottom is average atomic mass. The mass number will be 14. So, 14 = number of protons + number of neutrons 14 = 7 + number of neutrons Number of neutrons = 7 A charge of -3 is given (this is an ion). Number of electrons = atomic number – charge Number of electrons = 7 – (-3) = 10

  24. Red phosphorus is an allotrope of phosphorus that can be made by heating white phosphorus to temperatures above 240 oC. The most common isotope of phosphorus can be represented by How many protons,electrons, and neutrons does this isotope have? a) 16, 16, 15 b) 15, 15, 1 c) 15, 15, 16 d) 16, 16, 31 e) 16, 16, 1

  25. The most stable isotope of silver is given by How many protons, neutrons, and electrons does it have? p, n, e a) 107, 47, 47 b) 60, 47, 60 c) 47, 47, 60 d) 47, 60, 47 e) 47, 107, 60

  26. Example Problem #1 1. Determine the number of protons, neutrons and electrons in each. A. Ba B. P-3 C. Na+ D. U-235

  27. Average Atomic Mass The mass number on the periodic table is a reflection of the masses of each type of isotope of an element and the percent of that element present as each isotope. Avg. At. Mass = (%1)(mass1) + (%2)(mass2) + (%3)(mass3)….. Neutrons and protons are most of the mass. The number of protons defines the atom. Isotopes are of the same atom, just differ in number of neutrons. So, a change in mass only really reflects changes in the number of neutrons.

  28. Average Atomic Mass Why is the average atomic mass of the element carbon 12.011 u? carbon-12: 98.89% natural abundance 12 u carbon-13: 1.11% natural abundance 13.0034 u You have to convert percentages to decimal. Divide percent by 100 or move decimal two places to the left. Avg. At. Mass = (0.9889)(12 u) + (0.0111)(13 u) = 11.867 u + 0.144 u = 12.011 u

  29. Example #2 An element (X) has two isotopes. Isotope X-6 (7.5% of 6.015u) and X-7 (92.5% of 7.016u). What is the average atomic mass of element X? What is the element?

  30. Example #3 Chlorine has two isotopes. It is 75.77% of Cl-35 which has a mass of 34.969u and 24.230% of Cl-37 which has as mass of 36.966u. What is the average atomic mass of chlorine?

  31. Example #4 Silicon has three isotopes shown in the table below: Isotope Mass (u) Abundance(%) Si-28 27.98 92.21% Si-29 28.92 4.70% Si-30 29.97 3.09% What is the average atomic mass of silicon?

  32. What the heck is a mole?

  33. Nope, not a little animal. A mole is a word that means a number. The word “dozen” means 12 The word “mole” means 6.022 x 1023 (this is called Avogadro’s number) You can have a mole of atoms, a mole of molecules or a mole of pennies (that would be awesome!!)

  34. Mole concept The mass number on the periodic table represents the mass of one mole of each element. 1 mole C = 12.01g = 6.022 x 1023 atoms 1 mole of Fe = 55.85g = 6.022 x 1023 atoms etc.

  35. Converting between grams-moles and atoms Write down the number, the unit, and the chemical that is given. Make a bracket with a line. Place the unit you currently have on the bottom, the unit you need to switch to on the top. In the brackets: next the unit “mole” place a 1, next to the unit “grams” place the number on the periodic table, next to the unit “atoms” place Avogadro’s number. Put the number on the left in the calculator first. Multiply by the top number, divide by the bottom number.

  36. Example #5 A. How many moles are there in a 4,000,000atom sample of copper? B. How many grams are in a 0.037mole sample of iron? C. How many atoms of sodium are in a 3.08g sample of sodium? D. How many grams of aluminum will a 3 x 1020 atom same weigh? E. How many moles are in a 4g sample of potassium?

  37. Molar Mass or Molecular Weight There are about 118 elements (so far) but there are millions of compounds. To perform grams mole conversions with compounds we have to calculate the mass of one mole. Subscript of the first element (atomic mass of the first element) + subscript of the second element (atomic mass of the second element)….. If there is a poly atomic ion with parenthesis and another subscript you need to multiply the subscripts.

  38. Example #6 Determine the molar mass of C6H12O6. 6(12.01) + 12(1.01) + 6(16) = 180.18g/mol Round the atomic masses to two places past the decimal.

  39. Example #7 Determine the molecular weight of Cr(NO3)3. 1 Cr, 3x1 = 3 N, 3 x 3 = 9 O 1(52) + 3(14.01) + 9(16) = 238.03g/mol

  40. Example #8 (let’s try some together) Find the molar mass (molecular weight) for: • NaOH • CaCO3 • AlPO4 • C12H22O11 • (NH4)2SO4 • K2C2O4

  41. Gram-mole-molecule conversions Write down the number the unit and the chemical that is given. Make a bracket with a line. Place the unit you currently have on the bottom, the unit you need to switch to on the top. Inside the Bracket: next the unit “mole” place a 1, next to the unit “grams” place the number from finding molar mass, next to the unit “molecule” place Avogadro’s number. Put the number on the left in the calculator first. Multiply by the top number, divide by the bottom number.

  42. Example #9 A. How many moles of NaCl are there in a 75g sample? B. How many molecules of Al2O3 are present in a 0.0059mole sample? C. How many molecules of Fe2(CO3)3 are there in a 0.045g sample? D. What mass will 1,000,000 molecules of (NH4)2SO4 have? E. What mass will a 0.59mole sample of K2C2O4 have?

  43. Percent composition Determine the molecular weight. Place the total mass of one element over the molecular weight x 100%. Do this for each element. Percent composition is based on mass.

  44. Example #10 Determine the percent composition of Fe2(CO3)3. Determine the percent composition of C2H6. Determine the percent composition of C12H22O11.

  45. Empirical & Molecular Formulas Empirical: The formula with the simplest whole number ratio of elements in the compound. Molecular: The formula with the exact numbers of each type of atom present in the molecule. CH2O C6H12O6 empirical molecular

  46. Finding the empirical formula If percents are given, assume you have the same number of grams. Do a gram mole conversion for each element present in the molecule. Identify the element with the smallest number of moles. Divide all the moles by that number. The answer will give you the subscripts for each element.

  47. What if you don’t get a whole number when you divide? You decide what fraction the decimal represents. Multiply all the numbers by the denominator of the fraction. 0.5 = ½ 0.2 = 1/5 0.33 = 1/3 0.4 = 2/5 0.67 = 2/3 0.6 = 3/5 0.25 = ¼ 0.8 = 4/5 0.75 = 3/4 if the decimal is <0.2 or >0.8 just round off to a whole number.

  48. Example #11 A compound is analyzed and determined to be 27.37% sodium, 1.20% hydrogen, 14.30% carbon and 57.14% oxygen. Determine the empirical formula for this compound.

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