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Chapter 4 Notes sheet 1 What is matter? Matter is defined as anything that has mass and volume. Atoms are defined as the basic units of matter. NOTE: Energy is NOT matter and does not have mass or volume. Energy and matter are related by E=mc2 matter pure substances elements mixtures homogeneous mixtures compounds one type of atom different atoms bonded to each other uniform (solution) heterogeneous mixtures non-uniform (separate parts) Hg I2 HgI2 element element compound Classifying Matter H2O Compounds vs. mixtures Compounds are groups of atoms that are held together by chemical bonds. A molecule is the smallest unit of a compound. The elements within a compound can only be separated by a chemical change, which breaks the bonds in the compound. Mixtures, however, can be separated by physical methods such as filtration or evaporation, which DO NOT chemically alter the substances. Compounds and Diatomics Compounds are groups of two or more different atoms that are held together by chemical bonds. There are some gases that are bound to themselves as pure elements (diatomic molecules): Hydrogen = H2 Nitrogen = N2 Oxygen = O2 Fluorine = F2 Chlorine = Cl2 Bromine = Br2 Iodine = I2 How would you classify these? Chemical vs. Physical Properties Chemical vs. Physical Changes Properties of matter are characteristics that can be tested or observed and are used to identify matter. Physical properties are those which can be observed WITHOUT changing the chemical make-up of the matter. They are observed during physical changes. Chemical properties can only be observed when matter is involved in a chemical reaction, which CHANGES the chemical composition. Chemical changes will create a new substance whereas physical changes do not. Chemical vs. Physical Changes Chemical vs. Physical Changes In a physical change, the chemical identity of doesn’t change the substance ___________________. In a chemical change, the chemical identity of changes a substance __________. Another name for a chemical change is a ______________ chemical _________________. reaction More Examples of Properties Carbon reacts with oxygen to form carbon dioxide and water. Water boils at 100ºC. Paper burns. A solution of KNO3 is colorless. Iron rusts. Solid sulfur is dull and yellow. Gold has a very high density (19.3 g/cm3). Sheet 2 Matter is made of atomos of air, fire, wind and earth! 460 BC Rutherford 1909 Dalton 1803 Bohr 1913 Thomson 1897 Present Day The Law of Conservation of Mass Antoine Lavoisier (1700s) compared the masses of substances before and after a reaction and found that the mass always remained constant. This led to the Law of Conservation of Mass. In other words, matter cannot be created or destroyed, only rearranged in chemical reactions. The Law of Definite Proportions French chemist Joseph Proust The same samples of a pure compound always contain elements in the same mass proportion. Dalton’s Atomic Theory The idea of an indivisible thing that made up all matter was refined by chemist John Dalton in 1803. His ideas were: Daltons Law of Multiple Proportions: Multiple Proportions: Since atoms bond in small, whole number ratios to form compounds, their masses are small whole number ratios. Ex: CO vs. CO2 Thomson’s cathode ray tube and plum pudding model electron beam Ernest Rutherford While studying radioactive elements, New Zealander Physicist Ernest Rutherford found that radioactive alpha particles deflected when fired at a very thin gold foil. This was known as the gold foil experiment, and it suggested that the atom was not a hard sphere as thought, but was mostly space, with a small concentration of mass. This concentration of mass became known as the nucleus. Link to experiment… Millikan’s oil-drop experiment Robert Millikan’s classic oil-drop experiment allowed the charge of a single electron to be determined: 1.60 x 10-19 C. Using these two numbers, we can calculate the mass of an electron: The mass of an electron is about 1/2000 of the mass of a 1.60 x10 C proton! m 9.10 x10 g 19 electron 8 1.76 x10 C / g 28 Rutherford’s gold foil exp. and new atomic model The Bohr Model Neils Bohr, a Russian scientist, proposed that the electrons must be in specific energy “orbits” (called electron shells) around the nucleus. Lower energy shells closer to the nucleus The Modern Electron Cloud (aka wave-mechanical) Model Democritus and Dalton’s atom electron Thompson’s electrons neutron Rutherford’s space and nucleus proton Bohr’s energy levels (not to scale) Chadwick’s neutrons Most atoms ~ 1-5 Å (Angstroms) = (1x10-10m) Sheet 3 Elements 15.9994 We currently know of about 110 elements, 92 of which are naturally occurring. 8 We illustrate an element with an OXYGEN atomic symbol. The atomic number tells the number atomic number of protons and identifies the element. The mass number is the total average atomic mass number of the protons plus the neutrons. (on the P. T. the average atomic mass of isotopes is given) O Isotope symbols a charge is shown here for an ion. This +2 tells the number of electrons that have been gained or lost Isotope symbols give the element symbol, the mass and atomic numbers. From these, the number of protons, neutrons, and electrons can be determined. What element is this? What is the mass of this isotope? hey, does this proton make my MASS look big? Ions An ion is an atom that has gained or lost one or more electrons. When an atom bonds with another atom, it seeks to gain electrons or lose them. For instance: Cl has 7 and will gain one electron Na has 1 and will lose one electron Positive ions are called “cations” Cl Cl- Na Na+ Negative ions are called “anions” Practice How many protons, neutrons, and electrons are present in (a) 27Al3+ (b) 79Se2 Write the isotope notation for an ion that contains 20 protons, 21 neutrons, and 18 electrons. Write the isotope notation for an atom of lead that has 128 neutrons. Write the isotope notation for the following: Formula mass (aka molar mass) The formula mass for a compound is atomic masses the sum of the ______ ______of all elements in the formula. Ex: Calculate the formula mass for Al2(SO4)3 Notes 4 Weighted average mass 91.0% 25.2 kg 5.0% 16.1 kg 3.0% 10.0 kg ?% 12.0 kg Natural Abundance - Isotopes There is not just one type of each atom, there are several. When an atom has more or less neutrons than another atom of the same element, we call them isotopes. 12C 13C For instance, the element carbon has 6 protons, but it could have 5, 6, 7, or 8 14C neutrons, to form Carbon-11, Carbon-12, Carbon-13, and Carbon-14. Each has a different mass. In nature, there is a mix of different natural isotopes. We use this mix to calculate average atomic mass. ex: carbon is 12.011 amu Calculating Average Atomic Mass To find average atomic mass, we multiply the relative abundance of an isotope by the mass of the isotope. We then add each of the products for each isotope. Average atomic mass for one mole of atoms Isotopes, Ions, and Allotropes (Oh my) Isotopes are atoms of the same element with different numbers of neutrons. Ions are atoms of the same element with different numbers of electrons. (Ions are easy to create by adding or removing electrons from a neutral atom). Allotropes are forms of the same element, but bonded in different structures. Diamond and pencil graphite are examples of allotropes. They are both pure carbon, but in different structures. Isotopes and allotropes Sheet 5 Nuclear Chemistry and the Band of Stability #n > #p What particles make up the nucleus? #n = #p protons and neutrons What is the charge and function of each particle? proton (+) -determines the identity of the atom neutron (0) –helps hold nucleus together Elements above the band of stability need to decrease their n/p ratio Elements below the band of stability need to increase their n/p ratio Why do nuclear reactions occur? Nuclear reactions occur when a nucleus becomes unstable. Protons and neutrons are attracted to each other by the strong nuclear force. In a stable nucleus, the attraction due to the strong force is greater than the repulsion due to electrostatic force. As elements get heavier, they become more unstable. Extra neutrons must be present to the nucleus (like glue) to increase stability by increasing the strong force. Nuclear reactions are DIFFERENT from chemical reactions, because new elements form. What is radioactivity? Radioactivity is the spontaneous emission of radiation from an element to achieve a more stable state. Uranium was the first radioactive element isolated (by Bequerel), followed by radium and polonium (by Marie Curie and her husband Pierre). There are no stable isotopes for elements after Bismuth (#83). Types of Radiation Name Charge Mass Alpha +2 4 Beta -1 0 0 0 +1 0 gamma positron Symbols γ Nuclear Reactions Nuclear reactions change the composition of an atom’s nucleus –the element will change!! Examples of naturally occurring nuclear reactions include alpha and beta decay, and fission and fusion. Some nuclei can become unstable by artificial transmutation, where a nucleus is bombarded (or shot) with a particle that creates instability and causes radioactive decay. Nuclear reactions can produce enormous amounts of energy as nuclear mass is converted into energy (E=mc2) Radioactive Decay Equations (natural transmutations) Alpha decay equation: parent isotope daughter isotope Beta decay equation (n→p+e-): Gamma decay involves energy transitions (electromagnetic waves), no mass is lost! Radioactive Decay Equations (natural transmutations) positron decay equation (n→p+e+): Electron capture equation (n+e-→p): Other nuclear reactions….fusion and fission 2 3 H→ 4 He + 1 n H + 1 1 2 0 1 n→ 90 nuclear 143 fission 1 n U + Sr + Xe + 3 92 0 38 54 0 235 Nuclear Reactor Design Nuclear Accidents Hiroshima and Nagasaki Chernobyl Three Mile Island Products of Nuclear Reactions In nuclear reactions, unstable nuclei change their number of protons and neutrons. A DIFFERENT element is created by the reaction, and large amounts of energy are released (E=mc2), much more than in a chemical reaction. Nuclear reactions result in the production of new, more stable nuclei. Unstable nuclei will continue to decay until a stable isotope is produced. Radioactive Decay Series A decay series ends with a stable isotope. Sheet 6 How long does radioactive material last? Half-life is defined as the time it takes for one half of the original sample to decay. Some half-lives are long, some are short… (look at Table N) concentration Radioactive decay and half-life Nevada Nuclear Testing -1950’s The Radium Girls The Radium Girls were women subjected to radiation exposure at the United States Radium Corporation factory, in Orange, New Jersey around 1917. These women painted the dials of watch faces with luminous paint and developed cancer (typically of the mouth and jaw) as a result. For fun, the Radium Girls painted their nails, teeth and faces with the deadly paint produced at the factory, sometimes to surprise their boyfriends when the lights went out. They mixed glue, water and radium powder, and then used camel hair brushes to apply the glowing paint onto dial numbers. The going rate, for painting 250 dials a day, was about a penny and a half per dial. The brushes would lose shape after a few strokes, so the U.S. Radium supervisors encouraged their workers to point the brushes with their lips, or use their tongues to keep them sharp. THE END It is disconcerting to reflect on the number of students we have flunked in chemistry for not knowing what we later found to be untrue. – Quoted in Robert L. Weber, Science With a Smile (1992) Erwin Schrodinger (1887-1961) Hi Blame me! I started all of this! In 1926 I described the location and energy of electrons in an atom with my mathematical model. How are compounds different from mixtures? A mixture is not like a compound in that the parts of a mixture are NOT bonded together as in a pure compound. Examples of compounds: water (H2O), sugar (C6H12O6), carbon dioxide (CO2)… Examples of mixtures: air, steel, orange juice, “Gorp” … Mixtures may be homogeneous (evenly distributed in a single phase) or heterogeneous (unevenly distributed in different phases). Dalton’s Top Five All matter is made of indestructible and indivisible atoms. (atoms are hard, unbreakable, and the smallest thing there is) Atoms of agiven element have identical physical and chemical properties. (all atoms of X will These two are behave the same anywhere) usually combined 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 rearranged) in chemical reactions. Adding the Neutrons Although theorized by Rutherford, British physicist, James Chadwick, proved the existence of massive, neutral particles. These particles came to be called neutrons… and their discovery in 1932 opened the door for more in depth investigations into radioactive materials and gave the WWII Allies the ability to enrich and purify fissionable uranium, a necessity in the production of nuclear weapons. Bohr’s Atomic Model Only certain orbits are stable, and their energies are related to their radii. An electron that absorbs energy becomes “excited” and jumps from the “ground state” to a higher energy orbit. Light is emitted when the electron returns to ground state. Bohr’s Atomic Model Atoms with electrons in the ground state have 2, 8, 18, and 32 electrons allowed in the first four shells.