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Transcript
The Atom Chapter 4 I. History of the Atomic Theory IV. Mass of Atoms A. Democritus A. Atomic Mass B. Aristotle B. Average Atomic Mass C. Lavoisier C. Mole and Molar Mass D. Proust E. Dalton F. Modern Atomic Theory II. History of Atomic Structure V. Radioactive Decay A. Thomson A. Process B. Millikan B. Comparison C. Rutherford C. Types D. Bohr E. Chadwick F. Quantum Atom III. Subatomic Particles A. Comparing Particles B. Atomic Number and Mass Number C. Ions D. Changing Number of Particles E. Nuclear Notation F. Hyphen Notation I. History of the Atomic Theory Remember: a scientific theory explains behaviors and the ‘nature’ of things Theories can be revised when new discoveries are made The theory describing the composition of matter has been revised many times A. Democritus (460-370 BC) 1.Matter is made up of “atoms” that are solid, indivisible and indestructible 2.Atoms constantly move in space 3.Different atoms have different size and shape 4.Changes in matter result from changes in the grouping of atoms 5. Properties of matter result from size, shape and movement B. Aristotle (384-322 BC ) 1. Four kinds of matter a. Fire – Earth – Water – Air 2. One kind of matter can transform into another 3. Rejected idea of the “atom” (idea then ignored for almost 2000 years 4. This theory was more popular and it was easier to accept Aristotle’s Theory of Matter C. Antoine Lavoisier (1770s) Experiment: 2 Sn + O2 2 SnO tin oxygen tin (II) oxide mass before reaction = mass after reaction Law of Conservation of Mass Matter cannot be created or destroyed (in a chemical or physical change) 1788-1799 D. Joseph Proust (1779) Develops Law of Definite Composition- all samples of a specific substance contain the same mass ratio of the same elements a. ex: all samples of CO2 contains 27.3% carbon and 72.7% oxygen b. therefore ‘elements’ are combining in a whole number ratio D. John Dalton (1803) Dalton became a school teacher at the age of 12 (he left school at age 11) Loved meteorology - pioneer in this field Studied works of Democritus, Boyle and Proust Wrote New System of Chemical Philosophy in 1808 Develops Law of Multiple Proportions a. describes the ratio of elements by mass in two different compounds composed of the same elements Example: carbon monoxide vs. carbon dioxide 1 part oxygen : 2 parts oxygen *when compared to the same amount of carbon in each compound Dalton collects data and develops Atomic Theory in 1803 a. Matter is made of small particles-atoms b. Atoms of a given element are identical in size, mass, but differ from those of other elements*. c. Atoms cannot be subdivided or destroyed*. d. Atoms combine in small whole number ratios to form compounds. e. Atoms combine, separate, or rearrange in chemical reactions. * Modified in Modern Atomic Theory F. Modern Atomic Theory (1) 1. All matter is made up of small particles called atoms. 2. Atoms of the same element have the same chemical properties while atoms of different elements have different properties 3. Not all atoms of an element have the same mass, but they all have a definite average mass which is characteristic. (isotopes) F. Modern Atomic Theory (2) 4. 5. Atoms of different elements combine to form compounds and each element in the compound loses its characteristic properties. Atoms cannot be subdivided by chemical or physical changes – only by nuclear changes II. History of the Atomic Structure A. J.J. Thomson (1887) Experiments with cathode ray tubes Voltage source - + Vacuum tube Metal Disks Voltage source - + Voltage source - + Voltage source - + Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end Voltage source By adding an electric field Voltage source + By adding an electric field Voltage source + By adding an electric field Voltage source + By adding an electric field Voltage source + By adding an electric field Voltage source + By adding an electric field Voltage source + By adding an electric field he found that the moving pieces were negative Thomson’s Model of the Atom a. electrons present (-) b. atom is like plum pudding - bunch of positive stuff (pudding), with the electrons suspended (plums) Calculated the ratio between the charge of the electron and its mass: e/m “Chocolate Chip Cookie” or “Plum Pudding Model” Millikan’s Oil Drop Experiments Robert Milikan (1909) – Oil Drop Experiment – Measured the electrical charge on the electron – Mass can be calculated (Thomson determined the e/m ratio) – Mass is 1/1840 the mass of a hydrogen atom – electron has a mass of 9.11 x 10-28 g B. Millikan’s Oil Drop Experiment (1909) Oil-Drop Experiment: Animation http://wps.prenhall.com/wps/media/object s/602/616761/MillikanOilDropExperiment.h tml http://www.britannica.com/nobel/cap/omil lik001a4.html http://www.daedalon.com/oildrop.html II. History of the Atomic Structure – Summary thus far So, at this point we know: - Atoms are indivisible particles – Electrons are negatively charged – The mass of an electron is very small HOWEVER – Atoms should have a (+) portion to balance the negative part - Electrons are so small that some other particles must account for mass C. Ernest Rutherford (1871-1937) Experiment (1909) Gold Foil Experiment (Expectations) a. Shot alpha particles at atoms of gold b. expected them to pass straight through Lead block Uranium Florescent Screen Gold Foil He thought this would happen: According to Thomson Model He thought the mass of the positive charge was evenly distributed in the atom Here is what he observed: The positive region accounts for deflection Gold Foil Experiment Results a. Most positive alpha particles pass right through b. However, a few were deflected c. Rutherford reasoned that the positive alpha particle was deflected or repelled by a concentration of positive charge Gold Foil Experiment Conclusions a. the atom is mostly empty space b. the atom has a small, dense positive center surrounded by electrons II. History of the Atomic Structure At this point in 1909, we know: – p+ = 1.67 x 10-24 g – e- = 9.11 x 10-28 g – The charges balance! But, – How are the electrons arranged? – There is still mass that is unaccounted for Rutherford Model of the Atom D. Niels Bohr (1913) Electrons orbit nucleus in predictable paths E. Chadwick (1891 – 1974) In 1935 1. Discovers neutron in nucleus 2. Neutron is neutral - does not have a charge n0 3. Mass is 1.67 x 10-24 g slightly greater than the mass of a proton History of the Atomic Theory 1803 1897 1909 1913 1935 Today solid particle electron proton e- orbit nucleus neutron Quantum Atom theory Dalton Thomson Rutherford Bohr Chadwick Schrodinger and others II. History of the Atomic Structure Charges balanced Mass accounted for However – what about the behavior of the electrons? F. The Quantum Atom Theory 1. The atom is mostly empty space 2. Two regions: a. Nucleus- protons and neutrons b. Electron cloud- region where you have a 90% chance of finding an electron III. Subatomic Particles Comparing Particles Relative Actual mass (g) Name Symbol Charge mass Electron e- -1 Proton p+ +1 1amu 1.67 x 10-24 Neutron n0 0 1amu 0 9.11 x 10-28 1.67 x 10-24 III. Subatomic Particles B. Atomic Number and Mass Number 1. Atomic number 1. the number of protons in the nucleus of an atom a. identifies the element b. no two elements have the same atomic number 2. Ex. C is 6, N is 7 and O is 8 carbon nitrogen oxygen B. Atomic Number and Mass Number 2. Mass number a. the number of protons plus neutrons in the nucleus of an atom b. mass number is very close to the mass of an atom in amu (atomic mass units) c. two atoms with the same atomic number but different mass number are called isotopes 1) (mass #) – (atomic #) = #n 0 C. Formation of Ions 1. Electrons and Ions a. For neutral atoms, #e- = #p+ b. When there are more electrons than protons, a negative ion forms (anion) c. When there are less electrons than protons, a positive ion forms (cation) For now, we will work only with neutral atoms C. Formation of Ions Examples of Ions Atom loses electrons Atom gain electrons and form cations and forms anions Cations (+ ions) Anions (- ions) K+ BrCa2+ O2Al3+ N3- C. Formation of Ions From Atoms Na loses an electron and forms a cation Na0 – e- --> Na+ Cl gains an electron and forms an anion Cl0 + e- --> Cl- D. Changing Number of Particles 1. You can never change the number of protons and have the same element 2. If you change the number of neutrons in an atom, you get an isotope 3. If you change the number of electrons in an atom, you get an ion E. Nuclear Notation 1. Nuclear Notation is one method for depicting isotopes of an element 2. contains the symbol of the element, the mass number, and the atomic number Mass number Atomic number X Na 23 11 How many protons? How many neutrons? How many electrons? F. Hyphen Notation 1. Element symbol or name – mass # 2. EXAMPLES a. Fluorine-19 b. C-14 c. U-238 IV. Mass of Atoms A. Atomic Mass 1. Mass of an atom a. too small to measure in grams b. use relative mass (amu) 1) atomic mass unit 2) 1 amu is defined as 1/12 the mass of one C-12 atom B. Average Atomic Mass 1. weighted average mass of all known isotopes a. weighted means that the frequency of an isotope is considered b. mass of each isotope is multiplied by its percent occurrence in nature – then product of all isotopes is added to get the average atomic mass Calculating Average Atomic Mass Process 1. Mulitply % occurrence x mass of isotope 2. Add products for each isotope isotope occurrence isotope occurrence Ex. X- 40 (30.0% ) X-30 (70.0%) (40 x .300) + (30 x .700) = 12 + 21 = 33 amu C. The Mole and Molar Mass 1. measures the amount of substance a. 1 mole = 6.02x1023 (Avogdro’s #) of particles (atoms, molecules, ions, electrons) b. standard – 1mole is the number of atoms in 12g of C-12 isotope 2. Molar mass – mass in grams of one mole (mol) of any substance a. numerically equal to atomic mass in amu b. unit is grams/mol V. Radioactive Decay A. The Process What is radioactive decay? spontaneous nuclear change in which unstable nuclei emit radiation and lose energy radiation – rays and particles emitted by radioactive materials Why do atoms undergo decay? produce a nucleus that is more stable- p/n ratio B. Comparison of alpha, beta, gamma Alpha Form particle Beta particle Gamma ____ electromagnetic radiation Symbol Mass Charge Notation 4 amu +2 no mass -1 no mass none Alpha Particle V. 1. 2. Radioactive Decay C. Types of Decay Beta Decay (neutronproton + electron) a. beta particle (electron) is given off b. atomic number increases by one c. mass number stays the same Alpha Decay a. alpha particle (2p++2n0) is given off b. atomic number decreases by 2 c. mass number decreases by 4 V. Radioactive Decay D. Examples of Decay Beta Decay (n0 p+ + e-) e- released Parent Nuclei Daughter Nuclei Co-60 --------------> Ni-60 + e(z = 27) (z = 28) C-14 ---------------> N-14 + e(z = 6) (z = 7) Beta Decay Beta Decay Beta Decay + beta particle + beta particle + beta particle + beta particle + beta particle Radioactive Decay D. Examples of Decay Alpha Decay (alpha particle (2n0 + 2p+) released) Parent Nuclei Daughter Nuclei Th-232 ----------------> Ra-228 + alpha (z = 90) (z = 88) Ra-226 ---------------> Rn-222 + alpha (z = 88) (z = 86) Alpha Decay Alpha Decay Alpha Decay Alpha Decay + alpha particle + alpha particle + alpha particle + alpha particle + alpha particle Radioactive Decay of Uranium Radioactive Decay of Uranium Radiation Radiation Radiation Mole Calculations Using Conversion Factors 1 mole 6.02 x 1023 molar mass 1 mole PARTICLES <----> MOLES <----> MASS 6.02 x 1023 1 mole 1 mole molar mass Mole Conversions [mass-mole-atoms] Type Equality Used 1. MOLES MASS 2. MASS MOLES 1 mole= molar mass (g) 3. MOLES ATOMS 4. ATOMS MOLES 1 mole = 6.02 x 1023 atoms 5. MASS ATOMS 6. ATOMS MASS 1 mole = 6.02 x 1023 atoms 1 mole = molar mass (g) Mole Calculations (Mass of Helium is 4.00 ) 1. 1.00 mole of helium = 4.00g 2. 2.00 mole of helium = 8.00g 3. 1.00 mole of helium = 6.02 x 1023 atoms 4. 2.00 mole of helium = 1.20 x 1024 atoms 5. 16.0g of helium = 4.00 mol 6. 3.0l x 1023 atoms of helium = .500 moles 7. 8.00g of helium = 1.20 x 1024 atoms.