Chapter 4 Notes sheet 1 What is matter? n Matter is defined as anything that has mass and volume. n 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=mc 2 matter pure substancesmixtures elementscompounds homogeneous mixtures heterogeneous mixtures Hg element I2I2 HgI 2 compound non-uniform (separate parts) uniform (solution) one type of atomdifferent atoms bonded to each other Classifying Matter H2OH2O Compounds vs. mixtures Compounds vs. mixtures n Compounds are groups of atoms that are held together by chemical bonds. A molecule is the smallest unit of a compound. n The elements within a compound can only be separated by a chemical change, which breaks the bonds in the compound. n Mixtures, however, can be separated by physical methods such as filtration or evaporation, which DO NOT chemically alter the substances. Compounds and Diatomics Compounds and Diatomics n Compounds are groups of two or more different atoms that are held together by chemical bonds. n There are some gases that are bound to themselves as pure elements (diatomic molecules): Hydrogen = H 2 Hydrogen = H 2 Nitrogen = N 2 Oxygen = O 2 Fluorine = F 2 Nitrogen = N 2 Oxygen = O 2 Fluorine = F 2 Chlorine = Cl 2 Chlorine = Cl 2 Bromine = Br 2 Iodine = I 2 How would you classify these? Chemical vs. Physical Properties n Properties of matter are characteristics that can be tested or observed and are used to identify matter. n Physical properties are those which can be observed WITHOUT changing the chemical make-up of the matter. They are observed during physical changes. n 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 n In a physical change, the chemical identity of the substance ___________________. n In a chemical change, the chemical identity of a substance __________. n Another name for a chemical change is a ______________ _________________. doesnt change changes chemical reaction More Examples of Properties n Carbon reacts with oxygen to form carbon dioxide and water. n Water boils at 100C. n Paper burns. n A solution of KNO 3 is colorless. n Iron rusts. n Solid sulfur is dull and yellow. n Gold has a very high density (19.3 g/cm 3 ). Sheet 2 Dalton 1803 Thomson 1897 Rutherford 1909Bohr 1913Present Day Matter is made of atomos of air, fire, wind and earth! 460 BC 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. Daltons Atomic Theory The idea of an indivisible thing that made up all matter was refined by chemist John Dalton in His ideas were: The idea of an indivisible thing that made up all matter was refined by chemist John Dalton in His ideas were: Daltons Law of Multiple Proportions: n Multiple Proportions: Since atoms bond in small, whole number ratios to form compounds, their masses are small whole number ratios. Ex: CO vs. CO 2 Thomsons cathode ray tube and plum pudding model electron beam Ernest Rutherford n While studying radioactive elements, New Zealander Physicist Ernest Rutherford found that radioactive alpha particles deflected when fired at a very thin gold foil. n 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. n This concentration of mass became known as the nucleus. n Link to experiment Link to experiment Link to experiment Millikans oil-drop experiment n Robert Millikans classic oil-drop experiment allowed the charge of a single electron to be determined: 1.60 x C. n 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 proton! Rutherfords gold foil exp. and new atomic model Neils Bohr, a Russian scientist, proposed that the electrons must be in specific energy orbits (called electron shells) around the nucleus. The Bohr Model Lower energy shells closer to the nucleus The Modern Electron Cloud (aka wave-mechanical) Model (not to scale) Chadwicks neutrons Rutherfords space and nucleus Democritus and Daltons atom Bohrs energy levels Thompsons electrons Most atoms ~ 1-5 (Angstroms) = (1x m) Sheet 3 Elements n We currently know of about 110 elements, 92 of which are naturally occurring. n We illustrate an element with an atomic symbol. n The atomic number tells the number of protons and identifies the element. n The mass number is the total number of the protons plus the neutrons. (on the P. T. the average atomic mass of isotopes is given) number of the protons plus the neutrons. (on the P. T. the average atomic mass of isotopes is given) O OXYGEN average atomic mass atomic number Isotope symbols n Isotope symbols give the element symbol, the mass and atomic numbers. From these, the number of protons, neutrons, and electrons can be determined. +2 a charge is shown here for an ion. This tells the number of electrons that have been gained or lost What element is this? What is the mass of this isotope? hey, does this proton make my MASS look big? Ions n An ion is an atom that has gained or lost one or more electrons. n When an atom bonds with another atom, it seeks to gain electrons or lose them. For instance: n Cl has 7 and will gain one electron n Na has 1 and will lose one electron ClCl - NaNa + Positive ions are called cations Negative ions are called anions Practice n How many protons, neutrons, and electrons are present in (a) 27 Al 3+ (b) 79 Se 2- n Write the isotope notation for an ion that contains 20 protons, 21 neutrons, and 18 electrons. n Write the isotope notation for an atom of lead that has 128 neutrons. n Write the isotope notation for the following: Formula mass (aka molar mass) The formula mass for a compound is the sum of the ______ ______of all elements in the formula. Ex: Calculate the formula mass for Al 2 (SO 4 ) 3 atomic masses Notes 4 Weighted average mass 91.0% 25.2 kg 3.0% 10.0 kg 5.0% 16.1 kg ? % 12.0 kg Natural Abundance - Isotopes n 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. n For instance, 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. Each has a different mass. n In nature, there is a mix of different natural isotopes. We use this mix to calculate average atomic mass. ex: carbon is amu 12 C 13 C 14 C Calculating Average Atomic Mass n 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) n Isotopes are atoms of the same element with different numbers of neutrons. n 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). n Allotropes are forms of the same element, but bonded in different structures. n 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 What particles make up the nucleus? n What is the charge and function of each particle? protons and neutrons proton (+) -determines the identity of the atom neutron (0) helps hold nucleus together #n > #p #n = #p 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? n Nuclear reactions occur when a nucleus becomes unstable. n 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. n Nuclear reactions are DIFFERENT from chemical reactions, because new elements form. What is radioactivity? What is radioactivity? n Radioactivity is the spontaneous emission of radiation from an element to achieve a more stable state. n Uranium was the first radioactive element isolated (by Bequerel), followed by radium and polonium (by Marie Curie and her husband Pierre). n There are no stable isotopes for elements after Bismuth (#83). Types of Radiation NameSymbolsChargeMass Alpha+24 Beta0 gamma00 positron+10 Nuclear Reactions n Nuclear reactions change the composition of an atoms nucleus the element will change!! n Examples of naturally occurring nuclear reactions include alpha and beta decay, and fission and fusion. n 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. n Nuclear reactions can produce enormous amounts of energy as nuclear mass is converted into energy (E=mc 2 ) Radioactive Decay Equations (natural transmutations) n Alpha decay equation: Beta decay equation (n p+e - ) : Beta decay equation (n p+e - ) : n Gamma decay involves energy transitions (electromagnetic waves), no mass is lost! parent isotopedaughter isotope Radioactive Decay Equations (natural transmutations) positron decay equation (n p+e + ) : positron decay equation (n p+e + ) : Electron capture equation ( n+e - p) : Electron capture equation ( n+e - p) : nuclear fission Other nuclear reactions.fusion and fission 2 1 H H 4 2 He n U n Sr Xe n Nuclear Reactor Design Nuclear Accidents Chernobyl Hiroshima and Nagasaki Three Mile Island Products of Nuclear Reactions n In nuclear reactions, unstable nuclei change their number of protons and neutrons. n A DIFFERENT element is created by the reaction, and large amounts of energy are released (E=mc 2 ), much more than in a chemical reaction. n 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? n n Half-life is defined as the time it takes for one half of the original sample to decay. n n Some half-lives are long, some are short (look at Table N) Radioactive decay and half-life concentration Nevada Nuclear Testing -1950s The Radium Girls The Radium Girls were women subjected to radiation exposure at the United States Radium Corporation factory, in Orange, New Jersey around 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 n 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 ( ) n Blame me! I started all of this! In 1926 I described the location and energy of electrons in an atom with my mathematical model. Hi How are compounds different from mixtures? n A mixture is not like a compound in that the parts of a mixture are NOT bonded together as in a pure compound. n Examples of compounds: water (H 2 O), sugar (C 6 H 12 O 6 ), carbon dioxide (CO 2 ) n Examples of mixtures: air, steel, orange juice, Gorp n Mixtures may be homogeneous (evenly distributed in a single phase) or heterogeneous (unevenly distributed in different phases). Daltons Top Five n All matter is made of indestructible and indivisible atoms. (atoms are hard, unbreakable, and the smallest thing there is) n Atoms of agiven element have identical physical and chemical properties. (all atoms of X will behave the same anywhere) n Different atoms have different properties. (X behaves differently than Y) n Atoms combine in whole-number ratios to form compounds. (two Hs and one O = Water (H 2 O) n Atoms cannot be divided, created or destroyed, (just rearranged) in chemical reactions. These two are usually combined Adding the Neutrons n Although theorized by Rutherford, British physicist, James Chadwick, proved the existence of massive, neutral particles. n These particles came to be called neutrons n 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. Bohrs 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. Bohrs Atomic Model n Atoms with electrons in the ground state have 2, 8, 18, and 32 electrons allowed in the first four shells.