Atomic Theory: History of the Atom Atomic Theory: experimental observations that led scientists to postulate the existence of the atom (smallest bit of an element). 1. Law of Conservation of Mass -During a chemical reaction, mass is conserved. (A. Lavoisier 1743-1794) Ex. 71.8 g iron oxide decomposed to 55.8 g iron and 16.0 g oxygen. 2. Law of Definite Proportions - In a given compound, the elements are always combined in the same ratio by mass. (Joseph Proust 1754-1826) Ex. 1.668 g molybdenum disulfide decomposed to 1.000 g Mo and 0.6680 g S. Mass Ratio Mo/S = 5.000 g molybdenum disulfide decomposed to 2.996 g Mo and 2.004 g S. Mass Ratio Mo/S = Dalton s Atomic Theory of Matter John Dalton (1766-1844) proposed this theory to explain the experimental observations given by the laws of conservation of mass and definite proportions. Postulates: Matter consists of tiny particles called atoms. Atoms are indestructible. In chemical reactions, the atoms are rearranged but they do not themselves break apart. Atoms of the same element are identical in mass and other properties. Atoms of different elements differ in mass and other properties. Chemical combination of elements to form compounds occurs. However, in a given compound the atoms of each element are present in a fixed number ratio. 1
Law of Multiple Proportions When two elements can form more than one compound, the mass ratios of the elements in the two compounds occur in small whole number ratios. Example: Consider two different compounds of iron and sulfur. A. Iron(II) disulfide (pyrite or fool s gold ): 5.00 g pyrite decomposed to 2.67 g S and 2.33 g Fe. Mass Ratio S/Fe = B. Iron(II) sulfide: 5.00 g iron(ii) sulfide decomposed to 1.82 g S and 3.18 g Fe. Mass Ratio S/Fe = C. Law of Multiple Proportions: Mass Ratio S/Fe in pyrite:iron(ii) disulfide = In that Dalton s Atomic Theory predicted the law of multiple proportions, this helped to give validity and force acceptance of the theory. Atomic Mass: mass of one atom of an element Atoms are too light to weigh individually. However, scientists could determine the relative mass of one atom of an element to another (i.e. relative atomic mass). Relative Atomic Mass: A. 5.000 g molybdenum disulfide decomposes to 2.996 g Mo and 2.004 g S. Thus, mass ratio Mo/S = B. Also known that in each molybdenum disulfide unit there is one atom of Mo to two atoms of S (MoS 2 ). Thus, atom ratio of Mo/S = C. From mass ratio and atom ratio can find relative atomic mass of Mo:S. 2
Carbon-12 ( 12 C) Atomic Mass Scale Relative atomic masses were not useful until a standard reference point was established. Atomic masses of all elements were referenced to the atomic mass of the most abundant isotope of carbon ( 12 C). Atomic Mass Reference: Carbon-12 or 12 C 1 atom 12 C = 12 amu (exactly) OR 1 amu = 1/12 the mass of an atom of 12 C Designation was arbitrary but gave atomic masses close to whole numbers for most elements. Example: If the relative mass of Mo: 12 C is 7.995, what is the atomic mass of Mo on the 12 C atomic mass scale? Example: If the relative mass of Fe:S is 1.74 and the relative mass of Fe: 12 C is 4.65, what is the atomic mass of S? Example: The relative mass of an unknown element to Zn is 0.245, while the relative mass of Zn: 12 C is 5.45. What is the atomic mass of the unknown element and what is the identity of the unknown element? Atomic masses on 12 C atomic mass scale are shown as non-integer numbers below the elements on the periodic table. But.. Why is atomic mass of carbon given as 12.011 amu instead of as 12 amu? Atomic masses shown on periodic table are average atomic masses taking into account the different isotopes of each element and their percent abundances. Isotopes are atoms of the same element but with a different mass. These isotopes occur in different percentages in nature (percent abundances or isotopic abundances). Thus, the third postulate of Dalton s Atomic Theory (Atoms of the same element are identical in mass) is NOT strictly true. 3
Calculation of Average Atomic Masses Example: It is found that carbon consists of two naturally occurring isotopes ( 12 C and 13 C) with atomic masses and % abundances given below. Calculate the average atomic mass of carbon. Isotope Atomic Mass % Abundance 12 C 12 amu 98.89% 13 C 13.0034 amu 1.11% Example: Naturally occurring boron consists of two isotopes, boron-10 and boron-11. Given the data shown below, calculate the average atomic mass of boron. Isotope Atomic Mass % Abundance 10 B 10.0129 amu 19.8% 11 B 11.0093 amu Calculation of % Abundances From Average Atomic Mass Example: A sample of naturally occurring gallium has an average atomic mass of 69.7 and consists of two isotopes, gallium-69 and gallium-71. Given the information shown below, calculate the % isotopic abundances of the two isotopes. Isotope Atomic Mass 69 Ga 68.9 amu 71 Ga 70.9 amu 4
Subatomic Particles: Particles Within the Atom Three key experiments helped to elucidate the structure of the atom and the nature of the subatomic particles. J.J. Thomson s Experiments with Cathode Ray Tubes (1897) Voltage applied across metal plates. Cathode ray obtained independent of metal used to make plates. Cathode ray traveled from negative to positive. Magnet deflected cathode ray. Used amount of deflection and magnetic field strength to calculate charge to mass ratio of cathode ray. Atom consists of parts one of which is the electron. Electron is negatively charged. Charge to mass ratio of e-=1.76x10 8 C/g. R. Milliken s Experiments with Oil Droplets (1909) Milliken watched how fast oil droplets fell: calculated mass of oil droplets. Placed negative charge on oil droplets using X-rays. Applied voltage to plates and suspended oil droplets in midair. From mass of oil droplets and voltage needed to suspend, calculated charge on each oil droplet. Charge always some whole # multiple of 1.60x10-19 C. Fundamental charge on electron: 1.60x10-19 C. From e- charge and Thomson s charge to mass ratio found mass of electron: 9.09x10-28 g. 5
Rutherford s Experiments w/alpha Particles (positive and 7000xheavier than e-) Most of alpha particles went straight through metal foil. 1/20,000 alpha particle deflected at large angles (repelled by something positive). 1/20,000 alpha particle deflected straight back toward source (hitting something massive). Most of atom is empty space. In center of atom there is a massive, positively charged core called the nucleus. Overall Picture of the Atom Atom consists of mostly empty space (wherein reside the electrons) with a small, dense, positively charged core at the center called the nucleus. Ex. Place a pea at center of astrodome (gives relative size of the nucleus to that of the atom). Three Subatomic Particles Particle Rel. Charge Mass Electron -1 9.109x10-28 g (5.486x10-4 amu) Proton +1 1.673x10-24 g (1.007 amu) Neutron 0 1.675x10-24 g (1.009 amu) Protons and neutrons (nucleons) reside within the nucleus. In neutral atom, # protons = # electrons. Isotopes: atoms of the same element but with different mass same # protons (and same atomic number) different # neutrons 6
Subatomic Numbers 1. Atomic Number (Z): gives the # protons in the nucleus a. Given as integer # above the element on periodic table. Ex. C Z=? Ca Z=? b. It is the # protons (or atomic #) that specifies the element. Ex. An element has 15 protons in the nucleus. What element is present? 2. Mass Number (A): gives the # nucleons (#protons plus neutrons) in the nucleus. a. Not given on the periodic table. b. Mass number can be calculated once # neutrons is known. A = # p + #n or. A = Z + #n or.. #n = A Z 3. Examples: Find # of subatomic particles present in boron-10 and boron-11. Ions: Cations and Anions Ions are charged species and are formed by gain or loss of electrons. Cations: positively charged ions formed by loss of electrons # protons > # electrons In ionic compounds, metals tend to form cations. Anions: negatively charged ions Formed by gain of electrons. # electrons > # protons In ionic compounds, non-metals tend to form anions. Example: How many of each type of subatomic particle (# e-, # p, and # n) are present in the following? 27 Al +3 32 S -2 137 Cs +1 An element with A=14, Z=6, and charge=-4. 7
Periodic Table (Mendeleev 1869) Arranged such that elements with similar properties fall within the same group. Periodic Table Terminology Groups: elements that lie within the same column (labeled IA, IIIB, etc.) Periods: elements that lie within the same row (Period 1 7) Metals: all elements to lower left of dividing line (not including H).» Have similar properties: conduct heat & electricity, have luster, are malleable & ductile, and most are solids at RT. Non-metals: elements to upper right of dividing line (includes H).» Have similar properties: poor conductors of heat & electricity, not malleable or ductile (are brittle), and can be s, l, or g at RT. Metalloids: elements bordering dividing line (B, Si, Ge, As, Sb, Te, Po, At)» Have intermediate properties: conduct electricity but not as well as metals. Representative/Main Group Elements: elements in A-Groups Transition Elements: elements in B-Groups Inner Transition Elements: elements in lanthanide and actinide series (elements in rows below main body of periodic table). 8
Alkali Metals: metals in Group IA (Li, Na, K, Rb, Cs)» Have similar properties: form +1 cations in ionic cmpds. & oxides are strongly basic. Alkaline Earth Metals: metals in Group IIA (Be, Mg, Ca, Sr, Ba)» Have similar properties: form +2 cations & oxides are strongly basic but less soluble. Halogens: non-metals in Group VIIA (F, Cl, Br, I, At)» Have similar properties: form -1 anions, exist as diatomic molecules in elemental form at RT (F 2, Cl 2, Br 2, I 2 ), & chlorides have a salty taste. Noble Gases:non-metals in Group VIIIA (He, Ne, Ar, Kr, Xe, Rn)» Have similar properties: not very reactive and do not readily form compounds, & exist as monatomic gases at RT (He(g), Ne(g), Ar(g), etc.) 9