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Honors Chemistry Chapter 4 –Atomic Structure (Student’s edition) Chapter 4 problem set: Useful Diagrams: 4.1 36, 40, 50, 51, 59, 65, 80, 81 4.2, 4.4, 4.5, 4.7, 4.10 and tables 4.1, 4.3 Defining the Atom Atoms: are the . Atoms: the smallest piece of an element that still retains the of that element. Some Historical Background: Democritus’s Atomic Theory: In approximately 460 BC, Democritus (Greek) coins the term “atom” (means ). Before that matter was thought to be one piece - called the theory of matter. Democritus creates the theory of matter. Solid Sphere In the 18th century , evidence appears to support the idea of atoms. Law of Conservation of Mass: Antoine Lavoisier (French) - 1770’s (see Ch2) Law of Definite Proportions: In 1799, Joseph Proust (French) states: “The of masses of chemicals in reactions are always the Examples: 8g O + 1 g H yields ____ g H2O 16g O + 2 g H yields ____ g H2O .” Law of multiple proportions: In 1803, John Dalton - English school teacher states: “The mass of one element combines with masses of other elements in simple, Examples: 2H + 1O yields 2H + 2O yields ratios.” Dalton’s Atomic Theory : Dalton put together the laws of conservation of mass, definite proportion, and multiple proportion to create his own atomic theory. *1. *2. 3. 4. #1 is #2 is Sizing up the Atom: Individual atoms are observable with instruments such as a microscope. 4.2 Structure of the Nuclear Atom Subatomic Particles: Atoms are made up of Electron: charge, 1/1837 amu (.0005) In the 1870’s, English physicist William Crookes studied the behavior of gases in tubes. Crookes tubes - forerunner of picture tubes in e- e- e- e- Crookes’ theory was that some kind of radiation or particles were traveling from the cathode across the tube. He named them 20 years later, J.J. Thomson (English) repeated those experiments and devised new ones. In 1897, JJ Thomas discovered the Thomson used a variety of materials, so he figured cathode ray particles must be to all atoms. J.J. Thomson gets credit for discovering the electron. Plum Pudding Model + - + Thomson and Milliken (oil drop experiment) worked together (their data, not them) to discover the and of the electron Electron charge: Electron mass: this is the smallest charge ever detected this weight is pretty insignificant Proton: In 1886, Eugen Goldstein found evidence for the charge and a mass of amu. . It has Neutron: In 1932, James Chadwick confirmed the existence of the slightly more than amu. NIB - Quarks: They are made up ups +2/3 charge and . It is . There are 6 types: downs - 1/3 charge so.... 2 ups, 1 down = proton charge 2 downs, 1 up = neutron charge Particle accelerators: miles long with propel particles along the chamber. The particles into each other at high speeds. This results in nuclear . Other simpler particles: . The Atomic Nucleus: The Rutherford Gold Foil Experiment: In 1911, Rutherford (New Zealand) … - The Experiment: towards a thin sheet of the screen. Concluded: particles from (in the lead box) were released foil. Most of the particles went through and were seen on alpha particles bounced back. 1– 2– 3– 4– Analogy: if an atom is the size of the Eagle’s stadium, then the nucleus is the size of a tennis ball floating in the middle of the stadium. + - Shortcomings of the Rutherford Model: According to gravity, electrons should move towards the nucleus eventually - they don’t. So.... more work needs to be done to understand the structure of the atom. Technology and Society: In 1931, Ernst Ruska and Max Knoll built the first electron microscope. It uses an electron beam and “lenses” that consist of magnetic or electric fields. Objects can be magnified over 100,000 times and projected on to a monitor. Biochemistry, Microelectronics, and Biology all use this technology. 4.3 Distinguishing Among Atoms Atomic Number: Nucleons: particles that make up the . Proton and Neutrons make up most of the Protons: 1 amu, Atomic # (Z): Always a # of element is of atoms. , positive charge, determines of the atom. number, number on the periodic table. in the nucleus, also indicates the # of electrons if the charged Neutron: neutral, determines the of the atom, mass is slightly more than 1 amu Electrons: not a nucleon, negative charge, nucleus, determine of an atom Mass Number: Originally it was thought that all atoms of the same element had the Hydrogen was observed to have the mass (Dalton) mass (assigned a weight of “1”). Original periodic table listed elements in order of their atomic not true today (see Iodine). . This is Mass number: Represented by the letter “A” . It is the sum of the protons and neutrons in a nucleus. This number is rounded from atomic mass due to the fact that there are isotopes. # neutrons = Example - # of neutrons in Li = Isotopes: atoms of the same element with different numbers of . because they have different Hydrogen Isotopes Protons Neutrons Electrons Mass % Abundance Protium Deuterium Tritium (artificial and radioactive) 1 1 amu 99.85% 1 2 amu .15% 1 3 amu 0% Some isotopes occur naturally. Most isotopes are produced artificially. Counting protons, electrons, and neutrons: Mass # = Neutrons = Protons = Atomic # = Electrons = Isotope 40 +1 K + Charge indicates the of an electron - Charge indicates the of an electron Protons Electrons 12 Neutrons Atomic # Mass # 40 12 53 36 10 19 C S-2 Na+1 Br 14 Atomic Mass: Mass of Cl thought to be 35.5 times that of hydrogen. Today we know this isn’t true. It’s the weighted average of 2 isotopes: 75% 35 Cl and 25% 37 Cl Analogy for weighted averages: If your homework grade is 80.0 and your test grade is 95.0, then what is your average? Note: homework is worth 50.0% of your grade. If your homework grade is 80.0 and your test grade is 95.0, then what is your average? Note: homework is worth 20.0% of your grade. Average atomic mass = [(%)(mass of 1st isotope)] + [(%)(mass of 2nd isotope)]...... Sample problem: find the average atomic mass of B B11 = 80.20% B10 = 19.80% Sample problem: find the %’s of 2 isotopes of Carbon given the following information: average atomic mass = 12.0111 isotope 1 = 12 C , isotope 2 = 13 C History lesson - originally H was the basis of all atomic masses and was given the mass of 1.0. Later, chemists changed the standard to oxygen being 16.000 (which left H = 1.008). In 1961, chemists agreed that 12C is the standard upon which all other masses are based. 1/12 of the mass of 1 atom of 12C = 1 amu Another example: Antimony consists of two naturally occurring isotopes: Antimony-121 (57.2500% abundance and an exact mass of 120.9038) and some other isotope. Calculate the percent and mass of the other isotope. Periodic Table Preview: see chapter 6 NIB – Technology and Society - The mass spectrometer - - - - the mass spec (chemist lingo) is used to detect, analyze, and identify unknown chemicals samples are vaporized, bombarded with electrons (in order to create + ions [positively charged particles due to a loss of 1 electron]), and placed in electrical and magnetic fields. Due to differences in mass ( # of neutrons) the paths of the molecules curve based on their individual mass. Heavier particles curve less. This change in curvature causes the particles to land on different places on a detector. The mass spectrometer was invented in 1912. By 1922 it was discovered that there were 300 naturally occurring isotopes that existed out of the 92 elements known at that time. Used for identifying components of mixtures, analyzing pollution, and dating works of art.