ChemistryCh3AtomicStructure：CH3原子结构化学.doc

Chemistry Ch 3: Atomic StructureObjectives3-1 Define the term atom List the postulates of Daltons atomic theory3-2 Discuss how atomic structure is related to electricity Explain what studies of cathode rays and radioactivity revealed about the atom Discuss Rutherfords alpha-scattering experiment and how it showed the existence of the nucleus3-3 Name and describe the three subatomic particles in an atom Determine the number of protons, neutrons, and electrons in an atom or ion Define isotope and atomic mass3-4 Describe the changes that accompany nuclear reactions Define radioactivity3-1 Early models of the atomWatch video What is the smallest possible particle of a substance? Can you keep dividing a piece of aluminum foil in half indefinitely? In 450 BC, the Greek philosopher Democritus proposed that all matter is composed of tiny, indivisible particles called atomos. Aristotle and other philosophers disagreed, asking what held the particles together. Democritus couldnt answer this so his theory was rejected. Today scientists define an atom as the smallest particle of an element that retains the chemical identity of that element A couple of important principles were discovered in the late 1700s:1. Lavoisiers Conservation of Matter (which states what?)2. Prousts Law of Constant Composition- a given compound always contains the same elements in the same proportions by mass. I.E. Water is always 88.9% O and 11.1% H. In 1803 an English math & physics schoolteacher, John Dalton, wrote down the atomic theory of matter based on the following postulates:1. Each element is composed of extremely small particles called atoms2. All atoms of a given element are identical, but they differ from those of any other element.3. Atoms are neither created nor destroyed in any chemical reaction.4. A given compound always has the same relative numbers and kinds of atoms. There are about 100 different elements- which means about 100 different kinds of atoms These atoms combine to form everything, all matter. Kind of like the letters in the alphabet form words. 3-2 Discovering Atomic StructureStatic Electricity Dalton thought atoms must be hard and round Michael Faraday, an English chemist, (1791-1867) suggested structure of atoms related to electricity. Atoms do contain particles that have electrical charge. The word electricity comes from the Greek word elektron for amber, if amber were rubbed w/ a cloth it would attract dust or other particles Benjamin Franklin (1706-1790) did early experiments w/ electricity. (KITE) o He found two kinds of charges and called them positive and negative. o Particles with like charges repel and opposite charges attractCathode Rays & Electrons A moving stream of electrical charges is called an electrical current In mid 1800s scientists began to investigate movement of electric current through evacuated glass tubes (very little air in them) see figure 3.9 p.97 Cathode Ray Tube- Negative electrode (cathode) and positive electrode (anode) connected to a battery. Line glass w/ fluorescent material (glows in presence of electricity). Energy flowed from the cathode to the anode, called it a cathode ray. By end of 1800s many new discoveries about cathode rays:o They could spin a paddle wheel an their way (suggested it had particles) o Deflected by a magnet in a direction consistent w/ an negative electrical charge (suggested it had a neg. charge) Decided to try and find the mass of 1 particle English physicist J.J. Thomson (1856-1940) started a series of experiments. see transparency o Anode had hole so that cathode ray went through + + - charged plates and through magnetic fieldo He found he could mathematically predict the amount of deflection created. He concluded that a cathode ray is composed of negative particles that come from the cathode. This meant that atoms were not solid, indivisible balls but had a substructure.o He named these negative particles electrons. He could not compute the mass but did figure the ratio of the electric charge to its mass. 1.76 x 10(8) coulombs/gram 1909 American physicist Robert Milliken (1868-1953) measured charge of electron. Used oil droplets to find charge. Found charge of every oil droplet was a multiple of 1.6 x 10(-19) coulomb- so the charge of 1 electron (e-) must be that. Using the charge to mass ratio he found the mass of 1 e- = 9.11 x 10(-28), very, very light! 1896- Henri Bequerel accidentally discovered radioactivity when he placed a sample of uranium on photographic film. Radioactivity is the spontaneous emission of radiation from an element. Early 1900s New Zealand Scientist Ernest Rutherford showed two kinds of radiation, alpha & beta. Gamma was discovered later. JJ Thomson concluded that since atoms have electrons with a negative charge but have a neutral overall charge they must contain particles with a positive charge. Plum pudding model. Rutherfords gold foil experiment showed that a beam was scattered went it went through thin gold foil thus proving that the particles were deflecting off the nucleus, a small concentrated core at the center. This contradicted the plum pudding model. Finally he came up with the nuclear model of the atom (see transparency) Read last paragraph page 1023-3 Modern Atomic Theory Parts of an atomo Nucleus- protons (positive) and neutrons (neutral), most of mass. Neutrons have slightly more mass than protons.o Electrons- (negative charge) move in space around the nucleus, small mass takes about 2000 to = 1 proton. Negative charge attracted to positive charge of nucleus (protons) Rutherfords (actually Bohrs) atom- saw an atom as a miniature solar system. Not correct they do not orbit in a well defined path, do not know exactly where they are. Instead electrons are more accurately depicted as indistinct clouds around the nucleus. See transparency For convenience sake (although not completely accurate, see chart page 104):o Charge of proton = +1, actually +1.062 x 10(-19). Mass = 1amu (atomic mass unit), actually 1.673 x 10(-24) g.o Charge of neutron = 0. Mass = 1 amu, actually 1.675 x 10(-24)o Charge of electron = -1, actually -1.062 x 10(-19). Mass = 0, actually 9.109 x 10(-28) g. Atomic number- discovered by Henry Moseley (1887-1915) a student of Rutherford, it is the # of protons in an atom. In periodic table the atomic number is written above the element, so the identity of the atom comes from the # of protons in the nucleus. Since an atom under normal conditions is neutral, its # of electrons = # of protons. So the atomic # also reflects the # of electrons. (Talk through sample problem 1, practice prob. 1, 2, alt. pract. 1,2) Ions- when an atom gains or loses and electron it then has either a positive or negative net electrical charge. The charge of the ion is the difference between the # of protons and the # of electrons. Ionic charge = # protons - # electrons. Example:1. Na loses one electron (e-), so 11-10 = +12. Cl gains one e-, so 17-18 = -13. Ca loses 2, so 20-18 = +2 Ions are written- chemical symbol with the charge in superscript. I.E. the +/- can be before or after the number. If the charge is +/- 1 the 1 is omitted. I.E. (Do sample prob. 2, Practice 3 & 4, Alt sample 2, Alt sample 3 & 4) Isotopes- every element has a set # of protons, however they may have different #s of neutrons. I.E. All Cl have 17 protons but some 18 neutrons, others have 20. Atoms w/ same # protons but different # of neutrons are called isotopes. Generally one form is the most common. I.E. the most common form of H has 0 neutrons. H w/ one neutron called deuterium (rare) H w/ two neutrons called tritium (very rare) In nature elements are found as a mixture of isotopes. These mixtures usually stay in the same ratios- no matter where you find them. How different are isotopes? They have different masses because where does the mass of an atom come from? Protons + neutrons. If different # of neutrons then different masses. Cl- Mass # 37 atomic # 17. Since mass = proton + neutron then 37- 17 = 20 neutrons. Talk through sample problem 3 Do practice 5 & 6. Alt. sample prob3 & alt. practice 5 & 6 Mass of an atom Since the mass of an atom is so incredibly small it is impractical to use grams Chemists use another unit called Atomic Mass Unit (amu) 1 amu = 1/12 (mass of a C-12 atom) = 1.66 x 10(-24)g atomic mass (average atomic mass or atomic weight) = the average mass of an elements atoms Homework 3-3 Practice, 3-3 Review3-4 Changes in the nucleus Chemical Reactions- atoms interact through outer electrons Changes in the nucleus are called nuclear reactions. These change the composition of an atoms nucleusNuclear Stability Almost all atoms in nature are stable- not radioactive. A very few are naturally radioactive. Why are some stable and some not? It has to do with the # of protons and neutrons- not all combinations are stable Nucleus is protons and neutrons packed very closely together. But since protons have a + charge they repel each other. Why dont they fly apart? Strong Nuclear Force - a force that overcomes the electric repulsion between protons. It is a force that is only significant between subatomic particles. Neutrons have no charge, so experience no electrical repulsion. However they do have Strong Nuclear Force, which helps hold them together with the protons. The presence of the neutrons adds a net attractive force to the inside of the nucleus. Think of neutrons as the glue that holds together the nucleus.o Neutrons- dont repel they attracto Protons- both repel/attract other protons and attract neutrons. Why? All stable nuclei between atomic numbers 1-20 have about equal #s of protons and neutrons Beyond 20 need increasingly more neutrons to hold nucleus together When atomic # exceeds 83 no # of neutrons can adequately hold nucleus together. We say that these elements above atomic # 83 are radioactive. Examples? Too few neutrons makes nucleus unstable. Any isotope containing very different # (too many or too few) neutrons will be radioactive. Too many neutrons will emit beta particles. Types of Radioactive Decay Radioactive elements emit many different kinds of radiation. The three most common are:1. alpha 2. beta3. gamma Alpha ( ) radiation- stream of high energy alpha particles con