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Unit 2: Atomic Structure Democritus ATOMS Democritus (4th century BC) 1. IDEA OF THE ATOM a. Matter consists of tiny, indivisible particles (ATOMOS) b. Atoms of elements differ from one another in size and shape - no experiments to support this! ATOMIC THEORY – John Dalton - 1803 a. All matter is made up of tiny indivisible particles called atoms 1. Based on EXPERIMENTAL evidence ATOM (modern definition) Smallest particle of an element that still retains the properties of that element John Dalton 1766-1844 Atomic Theory of Matter based on the following postulates: 1) Each element is composed of indivisible particles called atoms. 2) All atoms of a given element are identical but they differ from those of any other element. 3) Atoms are neither created nor destroyed in any chemical reaction. 4) A given compound always has the same relative numbers and kinds of atoms. SEEING ATOMS Law of Definite (Constant) Composition A compound always has the same composition by mass. CO2 : 27% C CO: 43% C 73% O 57% O COMPOSITION OF THE ATOM 1.ELECTRON: J.J. THOMSON – 1897 a. Cathode rays are composed of small, negatively charged particles – ELECTRONS b. These particles are present in ALL atoms. JJ Thompson •Cathode Ray Tube (CRT) •Discovered the electron Millikan c. MILLIKAN – Oil drop experiment 1. Determined the SIZE of charge on the electron (-1). This allowed calculation of the electron’s mass, 9.1 x 10-28 g or ~ 1/1840 mass of H atom. Models of the atom GOLD FOIL EXPERIMENT 2. MODEL OF ATOM a. Gold foil experiment- Rutherford (1911) 1. Bombarded thin gold foil with alpha particles (+ charge) 2. Most went straight through: “Most of the atom is empty space.” 3. A few were strongly deflected or actually BOUNCED BACK GOLD FOIL EXPERIMENT 4. Conclusion Discovered two main parts of the atom: a. Atoms have a very small, positively charged center containing almost all of the mass of the atom (1/10,000 diameter of the atom) – the NUCLEUS b. Negatively charged electrons occupy the space around the nucleus c. Positive charge on the nucleus = negative charge outside the nucleus From Thomson to Rutherford MODERN CONCEPT: Electrons occupy “cloud” outside the nucleus Orbit Electron Cloud Progression of the Atom 5. OTHER PARTICLES: a. PROTON: CHARGE: +1 MASS: 1.673 X 10-23 g or 1 amu LOCATION: nucleus b. NEUTRON 1. CHARGE: 0 2. MASS: 1.675 X 10-23 g or 1 amu 3. LOCATION: nucleus 4. Discovered by Chadwick (1932) Remember for the ELECTRON 1. CHARGE: -1 2. MASS: 9.1 X 10-28 g or 0 amu 3. LOCATION: electron cloud ATOMIC MASS of an element a. average mass of an atom of an element, compared to a standard b. Standard: Carbon-12 (12.000) c. Atomic mass unit (amu): 1/12 mass of a carbon-12 atom d. Atomic mass of an element on the periodic table is an AVERAGE of the naturally occurring isotopes Isotopes of Hydrogen Hydrogen – 1 (protium) 1p 1 H 1e 1 0n Hydrogen – 2 (deuterium) 1p 2 H 1e 1 1n in “heavy water” Hydrogen – 3 (tritium) 1p 3 H 1e 1 2n Radioactive Note: Use the mass number on the periodic table, unless it is specifically given. Isotopes of Carbon C-12 C-14 or or 12 6 C 6p 6e 6n 14 6 C 6p 6e 8n ATOMIC MASS 1. There are two isotopes of copper, Cu-63 and Cu-65. Which is more abundant? ATOMIC MASS 2. Calculate the atomic mass of Chlorine, given that 75.77% of its atoms are Cl-35 (34.97 amu) and 24.23% are Cl-37 (36.97 amu). ATOMIC MASS 3. The atomic mass of aluminum is 26.98. The atomic mass of nitrogen is 14.01. An aluminum atom is how many times heavier than a nitrogen atom? ATOMIC MASS 4. Element Z has two naturally occurring isotopes, Z-203 (AM of 202.97) and Z-205 (AM of 204.97). If the average atomic mass is 204.41, in what percentages does each of these isotopes occur? ATOMS FORM MOLECULES! MOLECULE A group of two or more atoms held together by covalent bonds - definition a. Basic unit of covalent compounds Ex. H2O, CO2, C6H12O6 b. Building blocks of SOME elements H2, O2, N2, Cl2, Br2, I2, F2 DIATOMIC ELEMENTS MOLECULAR FORMULA Shows the number of atoms of each element in one MOLECULE of a substance Ex. C6H12O6 6 C atoms 12 H atoms 6 O atoms 24 total atoms ATOMS FORM IONIC COMPOUNDS! IONS An atom or group of atoms with a positive or negative charge a. Positive ions Form when an atom LOSES electrons Na → Na+ + 1 e 11 p 11p 11 e 10e (+11 -10) = +1 or 23Na+ 11 Negative ions Form when an atom GAINS electrons O + 2 e → O28p 8p 8e 10e (+8 -10 = -2) May be represented as 16O28 IONIC COMPOUNDS a. Made up of ions held together by STRONG electrical attractions b. No discrete molecules (No NaCl molecule) EMPIRICAL FORMULAS (simplest formula) a. Show only the LOWEST RATIO of atoms of each element in the compound b. Represent ALL ionic compounds WRITING EMPIRICAL FORMULAS PRINCIPLE: Compounds are electrically neutral Ba2+O2- = BaO Ca2+F1- = CaF2 Al3+O2- = Al2O3 Sn4+O2- = SnO2 PRACTICE Write correct formulas for the following: Li1+ and O2Sr2+ and S2Be2+ and Br1Al3+ and S2K1+ and I1- MOLECULAR MASS(FORMULA MASS) How heavy a molecule is compared to Carbon – 12 a. How? b. Find the molecular mass of C4H10 (butane). 4 C + 10 H = 4(12.01) + 10( 1.01) = 58.14 amu c. Compare the relative masses of Benzene (C6H6) and butane. 6C+6H= 6(12.01) + 6(1.01) = 78.12 amu 78.12 = 1.344 58.14 A C6H6 molecule is 1.344 x heavier than a C4H10 molecule. Determine the formula masses of: Ba(NO3)2 CaSO4 * 2 H2O 3. Al2(SO4)3 4. CaCl2 * 10 H2O THE MOLE 1 MOLE equals 6.022 x 1023 particles. 1 mole Cu contains 6.022 x 1023 atoms. 1 mole CO2 contains 6.022 x 1023 molecules. A mole has the mass in grams which is numerically equal to the molecular mass (formula mass) of the substance. This formula mass is the sum of the atomic masses of the atoms in the substance. 63.55 1 mole Cu atoms weighs __________ grams. 1 mole CO2 molecules weighs __________ grams. 44.01 1 mole Ba(NO3)2 weighs ___________ grams. The MASS IN GRAMS OF 1 MOLE is: 1 gram-formula mass (GFM) 1 gram-molecular mass (GMM) OR 1 gram-atomic mass (GAM) MOLAR MASS MOLE – GRAM CONVERSIONS 1. Find the mass in grams of 1.20 moles of propane, C3H8. MOLE – GRAM CONVERSIONS 2. Calculate the number of moles in 62.8 grams of K2O, potassium oxide. 3. What is the mass of 0.785 mole of C2H5OH, ethanol? 4. How many moles are present in 300. g of Ca(NO3)2, calcium nitrate? 5. Given: 2.35 grams of propanol, C3H7OH. How many moles of propanol are in the sample? How many propanol molecules? How many carbon atoms? 6. How many molecules are present in 100. grams of hexane, C6H6? How many Hydrogen atoms? How many Carbon atoms? 6. How many molecules are present in 100. grams of hexane, C6H14? How many Hydrogen atoms?