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Chemistry, The Central Science, 11th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 2: Atoms, Molecules, and Ions Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton. Dalton's Postulates Each element is composed of extremely small particles called atoms. Dalton's Postulates All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements. Dalton’s Book: A New System of Chem. Philosophy Dalton's Postulates Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. Dalton's Postulates Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. Law of Constant Composition Joseph Proust (1754–1826) • also known as the law of definite proportions. • It states that the elemental composition of a pure substance never varies. Law of Conservation of Mass Antoine Lavoisier (1754–1826) The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place. Prob. 2.11. A chemist find that 30.82 g of nitrogen react completely with 17.60, 35.20 70.40 or 88.00 g of oxygen to from four different compounds. a) Calculate the mass of oxygen per gram of nitrogen in each compound. b) How is the result explained by Dalton’s atomic theory. Exam Problem. A sample of ascorbic acid (vitamin C) is synthesized in the laboratory. It contains 1.50 g of carbon and 2.00 g of oxygen. Another sample of ascorbic acid isolated from citrus fruits contains 9.22 g of carbon. How many grams of oxygen does it contain? a. 12.3 b. 6.92 c. 2.63 d. 10.5 e. 19.00 The Electron • Streams of negatively charged particles were found to emanate from cathode tubes. • J. J. Thompson is credited with their discovery (1897). The Electron Thompson measured the charge/mass ratio of the electron to be 1.76 108 coulombs/g. Millikan Oil Drop Experiment Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other. Robert Millikan (U. of Chicago) was able to determing the mass of the electron by measuring its charge in this experiment. The mass calculated, me = 9.10 E-28 g. Atom Structure, circa 1900 • The prevailing theory was that of the “plum pudding” model, put forward by Thompson. • It featured a positive sphere of matter with negative electrons imbedded in it. Discovery of the Nucleus, 1911 Ernest Rutherford students (Marsden & Geiger) shot a beam of -particles at a thin sheet of gold foil and observed the pattern caused by their scattering -particles are helium nuclei, 4He2+ The Nuclear Atom Since some particles were deflected at large angles, thus Thompson’s model could not be correct. Rutherford on YouTube The Nuclear Atom • Rutherford postulated a very small, dense nucleus (1.8 fm) with the electrons around the outside of the atom. • Most of the volume of the atom is empty space. Other Subatomic Particles • Protons were discovered by Rutherford in 1919. • Neutrons were discovered by James Chadwick in 1932. quark structure of the neutron Subatomic Particles • Protons and electrons are the only particles that have a charge. • Protons and neutrons have essentially the same mass. • The mass of an electron is so small we ignore it when labeling particles. Symbol 1 p 1 1 0n 0 -1e The neutron to proton ratio is critical in determining the stability of a nucleus. Elements above the belt of stability undergo -decay and elements below the belt of stability undergo positron emission or electron capture. Fusion processes in stars have been shown to form nuclei up to 26 protons and 30 neutrons (5626Fe). Supernova SN 1054 produced the Crab Nebula Supernova produce temperatures (energies) capable of synthesizing elements through Z=94 Radioactivity • Radioactivity is the spontaneous emission of radiation by an atom. • It was first observed by Henri Becquerel in 1896. • Marie and Pierre Curie studied it extensively and isolated the first radioactive elements (polonium and radium). Radioactivity • It is not uncommon for some nuclides of an element to be unstable, or radioactive. • We refer to these as radionuclides. • There are several ways radionuclides can decay into a different nuclide. Radioactivity • Three types of radiation were discovered by in the early 1900s in Ernest Rutherford’s lab: particles particles rays Writing the Symbol of an Isotope A ZX • A is the mass number (p + n) • Z is the atomic number (# protons) • X is the atomic symbol Alpha Decay Alpha decay is the loss of an -particle (a helium nucleus). 4 2 He 238 92 U 234 90 U + 4He 2 Beta Decay Beta decay is the loss of a -particle (a high energy electron). 0 −1 131 53 I or 131 54 0 −1 e Xe + 0e −1 Positron Emission Some nuclei decay by emitting a positron, a particle that has the same mass as but an opposite charge to that of an electron. 0 1 11 6 C e 11 5 B + 0e 1 Electron Capture (K-Capture) Addition of an electron to a proton in the nucleus is known as electron capture or Kcapture. – The result of this process is that a proton is transformed into a neutron. 1 1 p + 0e −1 1 0 n Gamma Emission This is the loss of a -ray, which is highenergy radiation that almost always accompanies the loss of a nuclear particle. 0 0 Writing the Symbol of an Isotope A ZX • A is the mass number (p + n) • Z is the atomic number • X is the atomic symbol Symbols of Elements Elements are symbolized by one or two letters. For this class you should know the names and symbols for the first 36 elements. Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z) Atomic Mass The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom. Atomic Mass Atomic and molecular masses can be measured with great accuracy with a Mass Spectrometer. Isotopes • Isotopes are atoms of the same element with different masses. • Isotopes have different numbers of neutrons. 11 C 6 12 C 6 13 C 6 14 C 6 Oxygen isotopes and decay modes Name Symbol Mass (amu) Oxygen-13 13 13.02810 O 8 Oxygen-14 14 15 O 14.008595 16 % Natural Abundance e 8.9 ms ---- 0 e 70.6 s --- e 122 s --- 1 O 15.003065 8 Oxygen-16 0 Half-Life 1 8 Oxygen-15 Decay Mode 0 1 O 15.994915 --- --- 99.76 O 16.999131 --- --- 0.04 O 17.999160 --- --- 0.20 O 19.003577 0 e 26.9 s --- e 13.5 s --- e ≈3 s --- 8 Oxygen-17 17 8 Oxygen-18 18 8 Oxygen-19 19 8 Oxygen-20 20 -1 O 20.004075 8 Oxygen-21 21 8 0 -1 O 21.008730 0 -1 Average Mass • Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. • Average mass is calculated from the isotopes of an element weighted by their relative abundances. Example Neon is found in three isotopes in nature. Isotope Mass (amu) Natural abundance (%) Neon-20 19.9924 90.4838 Neon-21 20.99395 0.2696 Neon-22 21.9914 9.2465 Average atomic mass of neon = (19.9924 x 0.904838) + (20.99395 x 0.002696) + (21.9914 x 0.092465) = 20.1797 amu Another Isotope Problem… Using the following table of abundances and masses of the three naturally occurring argon isotopes, calculate the exact mass of 40Ar. Symbol Exact Mass (amu) Natural Abundance (%) 36Ar 35.96755 0.337 38Ar 37.96272 0.063 40Ar ? 99.60 Ave. 39.948 Ans: 39.96 amu Mendeleev’s Periodic Table • Dmitrii I. Mendeleev (1869) arranged elements in the periodic table by their chemical and physical properties. • He left open spaces in his periodic table to account for elements not yet discovered. Periodic Table • It is a systematic catalog of the elements. • Elements are arranged in order of atomic number , Z, (rather than atomic weights). • For this class, you should know the names and symbols for elements 1-36 Periodic Table • The rows on the periodic chart are called Periods. • Columns are Groups. • Elements in the same group have similar chemical properties. Groups These five groups are known by their historic group names. Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H). Periodic Table Metalloids (simiconductors) border the stairstep line (with the exception of Al, Po, and At). Periodic Table Metals are on the left side of the chart. Exam Question: Potassium is a __________ and chlorine is a __________. A) metal, nonmetal B) metal, metal C) metal, metalloid D) metalloid, nonmetal E) nonmetal, metal Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. Chemical Formulas Molecular compounds are composed of molecules and almost always contain only nonmetals, i.e. elements on the right side of the periodic table. The bonding between non-metals is more covalent (sharing of electrons) Diatomic Molecules These seven elements occur naturally as molecules containing two atoms. Types of Formulas • Empirical (experimental) formulas give the lowest whole-number ratio of atoms of each element in a compound. • Molecular formulas give the exact number of atoms of each element in a compound. Types of Formulas • Structural formulas show the order in which atoms are bonded. • Perspective drawings also show the three-dimensional array of atoms in a compound. • Space-filling model represent the boundaries of electron density. Ions • When atoms lose or gain electrons, they become ions. – Cations are positive and are formed by elements on the left side of the periodic chart. – Anions are negative and are formed by elements on the right side of the periodic chart. Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals by transfer of charge (electrons). Exam Question When a metal and a nonmetal react, the __________ tends to lose electrons and the __________ tends to gain electrons. A) metal, metal B) nonmetal, nonmetal C) metal, nonmetal D) nonmetal, metal E) None of the above, these elements share electrons. Writing Formulas • When compounds are electrically neutral, one can determine the formula of an ionic compound this way: – The charge on the cation becomes the subscript on the anion. – The charge on the anion becomes the subscript on the cation. – If these subscripts are not in the lowest wholenumber ratio, divide them by the greatest common factor. Exam Question What is the formula of the compound formed between strontium ions and nitrogen ions? A) SrN B) Sr3N2 C) Sr2N3 D) SrN2 E) SrN3 Common Cations (Table 2.4)* *Commit to Memory Common Anions* *Memorize…the anions as well. Note that many common anions are polyatomic. Inorganic Nomenclature (i.e. naming inorganic compounds) • Write the name of the cation. • If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion (see previous, simply write the name of the polyatomic ion. • If the cation can have more than one possible charge (Fe2+ vs. Fe3+) write the charge as a Roman numeral in parentheses. Patterns in Oxyanion Nomenclature • When there are two oxyanions involving the same element: – The one with fewer oxygens ends in -ite. • NO2− : nitrite; SO32− : sulfite – The one with more oxygens ends in -ate. • NO3− : nitrate; SO42− : sulfate Exam Question: The correct formula of iron(III) bromide is __________. A) FeBr2 B) FeBr3 C) FeBr D) Fe2Br2 E) Fe3Br What is the systematic name for Ca(NO3)2? Exam Question: The correct name for CaH2 is __________. A) hydrocalcium B) calcium dihydride C) calcium hydroxide D) calcium dihydroxide E) calcium hydride What is the systematic name for (NH4)2SO4? Exam Question: 50) The correct name for CCl4 is __________. A) carbon chloride B) carbon tetrachlorate C) carbon perchlorate D) carbon tetrachloride E) carbon chlorate What is the systematic name for Co(OH)2? NaH2PO4? Write a chemical formula for cobalt(II) hydrogen carbonate. Patterns in Oxyanion Nomenclature • The one with the second fewest oxygens ends in -ite. – ClO2− : chlorite • The one with the second most oxygens ends in -ate. – ClO3− : chlorate Patterns in Oxyanion Nomenclature • The one with the fewest oxygens has the prefix hypoand ends in -ite. – ClO− : hypochlorite • The one with the most oxygens has the prefix per- and ends in -ate. – ClO4− : perchlorate Acid Nomenclature • If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- . – HCl: hydrochloric acid – HBr: hydrobromic acid – HI: hydroiodic acid Acid Nomenclature • If the anion in the acid ends in -ite, change the ending to -ous acid. – HClO: hypochlorous acid – HClO2: chlorous acid Acid Nomenclature • If the anion in the acid ends in -ate, change the ending to -ic acid. – HClO3: chloric acid – HClO4: perchloric acid Exam Question: The correct name for H2SO3 is __________. A) sulfuric acid B) sulfurous acid C) hydrosulfuric acid D) hydrosulfic acid E) sulfur hydroxide Nomenclature of Binary Molecular Compounds • The less electronegative atom is usually listed first. • A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however) . Nomenclature of Binary Compounds • The ending on the more electronegative element is changed to -ide. – CO2: carbon dioxide – CCl4: carbon tetrachloride Nomenclature of Binary Compounds • If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. N2O5: dinitrogen pentoxide not…dinitrogen pentaoxide Exam Question: The name of PCl3 is __________. A) potassium chloride B) phosphorus trichloride C) phosphorous(III) chloride D) monophosphorous trichloride E) trichloro potassium Nomenclature of Organic Compounds • Organic chemistry is the study of carbon compounds. • Organic chemistry has its own system of nomenclature (i.e. naming). Nomenclature of Organic Compounds The simplest hydrocarbons (compounds containing only carbon and hydrogen) are alkanes. Alkanes contain only single bonds and are often referred to as “saturated” Nomenclature of Organic Compounds The first part of the names above correspond to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.). Nomenclature of Organic Compounds • When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane. • The ending denotes the type of compound. – An alcohol ends in -ol.