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Chapter 3: The Atom and the Mole (with nuclear)

Chapter 3: The Atom and the Mole (with nuclear). The investigation and understanding of the atom is what chemistry is all about!. Topics rearranged from your text, pages 62-117.

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Chapter 3: The Atom and the Mole (with nuclear)

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  1. Chapter 3:The Atomand the Mole(with nuclear) The investigation and understanding of the atom is what chemistry is all about! Topics rearranged from your text, pages 62-117. We come here to be philosophers, and I hope you will always remember that whenever a result happens, especially if it be new, you should say, “What is the cause? Why does it occur?” and you will, in the course of time, find out the reason. -Michael Faraday

  2. The Mole • The “mole” represents a number of things….like a dozen. • How many things is a mole? • 6.022137 x 1023… we use 6.02 x1023. • This is Avogadro’s number • named for a lawyer, Amadoe Avogadro, that studied molecular gasses as a hobby. • When you have three moles of atoms, you have (3 x 6.02x1023 =) 1.81x1024 atoms total.

  3. Recall: Parts of the atom(subatomic particles) -1 • Proton charge = +1, mass = 1 • Neutron  charge = 0, mass = 1 • Electron charge = -1, mass = 0 • In a normal, neutral, unreacted atom, the number of electrons equals the number of protons. +1 0 Ions have more or less electrons than protons They have a charge

  4. Atomic History • Greek philosopher Democritus (400BC) • coined the term atomon which means “that which cannot be divided.” • John Dalton (1803) a colorblind chemist. • Among his interests, Dalton was very interested in a scientific explanation for his colorblindness the behavior of gasses. • In his A New System of Chemical Philosophy, Dalton published five principles of matter.

  5. Dalton’s Top Five • All matter is made of indestructible and indivisible atoms. • (atoms are hard, unbreakable, the smallest thing there is) • Atoms of a given element have identical physical and chemical properties. • (all atoms of X will behave the same anywhere) • Different atoms have different properties. • (X behaves differently than Y) • Atoms combine in whole-number ratios to form compounds. • (two H’s and one O = Water (H2O) • Atoms cannot be divided, created or destroyed, • (just rearrangedin chemical reactions).

  6. The Laws: • Constant Composition • Ratios of atoms in a compound is constant for that compound. • Conservation of Mass • Mass is not created or destroyed in a chemical reaction. • Multiple Proportions: • Since atoms bond in small, whole number ratios to form compounds, the ratio of their mass ratios are small whole numbers. hydrogen-oxygen atomic ratio = 2:1 Oxygen-carbon mass ratio = 1.33 x 2 Oxygen-carbon mass ratio = 2.66

  7. Conservation of Mass

  8. Multiple Proportions

  9. The Cathode Ray Tube • The cathode ray tube • A new invention suggested the presence of charges – areas of positive and negative…charge. • This suggested that atoms must be divisible, and Dalton’s theory had to be modified. • J. J. Thomson (1897) • English Physicist proposed that the atom is a sphere of positive charge with small areas of negative charge. • This theory become known as the “plum pudding” model after an English “dessert” of purple bread and raisins. Electrostatics

  10. Millikan’s Oil • Thompson used electrostatics experiments to determine the electron’s charge-to-mass ratio. • Robert Millikan’s (1909) • oil-drop experiment allowed the charge of a single electron to be determined: 1.60 x 10-19 C. • Scientists calculated the mass of an electron to be 1/2000 of the mass of a proton!

  11. Ernest Rutherford • Ernest Rutherford (1910) • New Zealander Physicist, while studying radioactive elements, found that radioactive alpha particles deflected when fired at a very thin gold foil. • The Gold Foil Experiment • the atom was not a hard sphere but • was mostly space, with a small concentration of positively-charged mass (the nucleus). • Link to experiment

  12. The Bohr Model • Niels Bohr • A Danish physicist (and student of Rutherford) rebuilt the model of the atom placing the electrons in energy levels. • Bohr was one of the founders of quantum chemistry: • energy can be taken in and given off in small packets or quanta of specific size. • When a specific amount of energy was added to an atom, an electron could jump into a higher energy level. No more…no less!

  13. Adding the Neutrons • James Chadwick (1932) • British physicist, proved there was too much mass in the nucleus • Suggested the existence of massive, neutral particles in the nucleus. (neutrons)

  14. The Modern Model Dalton’s atom Thompson’s electrons Rutherford’s space and nucleus Bohr’s energy levels Chadwick’s neutrons (not to scale)

  15. Elements 8 O • 112 known elements • 92 of which are naturally occurring. • 93 through 112: transuranium. • Each has an atomic symbol. • Atomic number • is number of protons • Atomic mass • is the total mass of the protons plus the neutrons. OXYGEN 15.9994 Notice that the atomic mass is not a round number, even though protons and neutrons each have a mass of 1. This is due to natural abundance.

  16. Natural Abundance - Isotopes • Isotopes: • Each element may have several isotopes • Isotopes differ in the number of neutrons. • Example: • the element carbon has 6 protons, but it could have 5, 6, 7, or 8 neutrons, to form Carbon-11, Carbon-12, Carbon-13, and Carbon-14. • In nature, there is a mix of different natural isotopes. • We use this mix to calculate average atomic mass… 11C, 12C, 13C, 14C

  17. Calculating Average Atomic Mass x • Sum of the products = average atomic mass • Example: • The isotopes of element “Bob” are found below: • Bob-18.0, 25.0% • Bob-19.0, 60.0% • Bob-20.0, 15.0% • What is the average atomic mass of naturally occurring Bob?  x  x  1 amu = 1.66x10-27kg

  18. Review … • Isotopes • atoms of the same _______ • different number of _______ • Ions • atoms of the same _______ • different number of _______ • Allotropes • forms of the same _______ • bonded in different _______ • Quanta / Quantum • Packets of energy of _______ size • Atomic Mass • Is the _______ of all _______ found in nature.

  19. Molar Mass No more AMU: AMU  Molar mass • Molar mass • expressed in grams per mole (g/mol) • mass of one mole of a substance. • link between the atom and the gram. (we can measure) • The average atomic mass of carbon is 12.01. What is the mass of a mole of carbon atoms? • What is the molar mass of Copper, Cu? • What is the molar mass of Nitrogen, N2?

  20. Molar Mass Practice • Determine the Molar Mass of the following elements and compounds: • Ca • Cl2 • CaCl2 • H2O • Ba(OH)2 • FeSO4 • Al2(SO4)3

  21. Mole-to-Mass Conversions • Sodium has an atomic mass of 23 g/mol. How many moles do you have in 115 grams? • How many grams are equal to 3.5 moles of CaCl2? • What is the mass of 0.46 moles of SiO2? Use a T-chart!

  22. Mole-Mass Conversion Practice • Complete the following mole-to-mass conversions: • Mass in grams of 2.25 moles of iron, Fe? • 126 grams Fe • Mass in grams of 0.375 moles of potassium, K? • 14.7 grams K • Number of moles in 5.00 grams of calcium, Ca? • 0.125 moles Ca • Number of moles in 3.60x10-10 grams of gold, Au? • 1.83x10-12 mol Au Use your periodic table to find molar mass

  23. Mole-Atoms Conversions End of Chapter 3 Phew! • Mole = 6.02x1023 things, how many atoms are in: • 3.0 moles of silver, Ag? • 0.010 moles of copper, Cu? • How many moles do you have in: • 2.4x1024 atoms of helium, He? • 3.0x1023 atoms of lithium, Li? • How many moles do you have in 222 grams of copper? • How many atoms in 127.1 grams of copper?

  24. Isotopes – Nuclides - Radioactivity • Nuclides • the nucleus of an isotope • Place the mass above the charge as seen here. • Nuclides undergo decay: • transformation into different nuclides • Balanced nuclear reactions • “Radioactive” • Half Life: time to decay ½ (mass) of a sample mass charge Images from ChemZone

  25. Alpha Decay • Alpha Decay • a helium nucleus is released. • Alpha particles: • move very slowly • because of their size, can be blocked with a few pages of paper or human skin • cause ionization (damaging!) • are positively charged mass charge This is a Nuclear Equation Alpha Decay occurs in all elements with atomic number above 83. Images from ChemZone

  26. Beta Decay • Beta Decay • An electron is ejected from the nucleus • Beta particles • move fast • can penetrate thick low-density materials • but can be blocked with concrete and metals • are negatively charged Beta Decay occurs when a nucleus has a high neutron-proton ratio. Images from ChemZone

  27. Gamma Decay No mass • Gamma Decay • High energy photons (gamma rays) are given off. • Gamma rays • given off as the “spare change” during other radioactive decays…. • extremely penetrating and powerful. Several inches of lead is required to slow these particles down to a stop. • Don’t get included in nuclear equations. No charge Summary of three basic particle decays Images from ChemZone

  28. Nuclear Equations Practice • Complete the following nuclear equations: • Sodium-24 undergoes alpha decay (help on click) • Iodine-131 undergoes beta decay • Tungsten-190 undergoes alpha decay • Uranium-238 undergoes alpha decay, then two beta decays (3 steps)

  29. Nuclear Fission • Nuclear fission: • splitting of large, unstable atoms • releases large amounts of energy • Critical mass (or critical density) • amount of fissionable fuel needed before reaction will begin. • Uncontrolled, nuclear fission proceeds to completion with great speed. • Nuclear Weapons: • two half-spheres of fissionable material are compressed together with conventional explosives, creating the critical mass. • In order to harness nuclear fission to create useable electricity, we slow down the process with control rods… Once fission begins, it is difficult to stop.

  30. Nuclear Fusion • Nuclear Fusion: • Joining of smaller nuclei to form larger nuclei. • Releases far more energy that nuclear fission. • Easier to control than fission. • The sun’s (stars) energy comes from the fusion of hydrogen atoms into helium atoms. • The H-bomb is a fusion weapon. • Fusion: • The power supply of the future? • Why don’t we use it now?

  31. E=mc2 • Einstein: Energy and mass are interchangeable. • E = Energy • m = mass • c = speed of light ( 3 x 108 m/s ) • Very small amounts of mass create large amounts of energy! • We use the formula ΔE= Δmc2 to build new, artificial elements in supercolliders (particle accelerators.) Fermilab cyclotron, Argonne National Laboratory, Chicago

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