Chapter 2: Atoms, Molecules, and Ions
Chapter 2: Atoms, Molecules, and Ions CHEM 1055
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This 9 page Class Notes was uploaded by Katie Rizzolo on Tuesday August 30, 2016. The Class Notes belongs to CHEM 1055 at Virginia Polytechnic Institute and State University taught by Dr. Gordon T. Yee in Fall 2016. Since its upload, it has received 7 views. For similar materials see Honors General Chemistry for Majors in Chemistry at Virginia Polytechnic Institute and State University.
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Date Created: 08/30/16
Monday, August 22, 2016 Chemistry Chapter 2 Notes 2.1 The Early History of Chemistry Greeks were the first to try to explain chemistry – 4 fundamental substances: earth, fire, water, air – Is matter cont. or composed of small indivisible pieces Democritus—atomos (later became atoms) – No experiments, no definitive conclusion Next 2000 years dominated by alchemy – Often “fake” but some impt. discoveries – Discovery of some elements and learned how to prepare mineral acids th Foundations of modern chemistry in 16 century – Development of systematic metallurgy (extraction of metal from ores) Georg Bauer – Medicinal application of minerals Paracelsus Robert Boyle – First “chemist” to perform quantitative experiments – Measure relationship between pressure and volume of gasses – The Sceptical Chemist (1661)—birth of quantitative sciences of physics and chemistry – A substance is an element unless it can be broken down into smaller substances Georg Stahl – Phlogiston flowed out of burning material Joseph Priestly – Discovered oxygen (originally called “dephlogisticated air”) Stahl and Priestly’s observations don’t conflict – What happens does not change, only the why is modified 2.2 Fundamental Chemical Laws 1. Antoine Lavoisier 1. Explained true nature of combustion 2. Law of Conservation of Mass—Mass is neither created nor destroyed 3. Combustion involved oxygen, not phlogiston 4. Life requires oxygen 2. Joseph Proust 1. Law of Definite Proportion—A given compound always contains exactly the same proportion of elements by mass 3. John Dalton 1. Law of Multiple Proportions—When two elements form a series of compounds, the rations of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers 3.3 Dalton’s Atomic Theory A New System of Chemical Philosophy (1808)— Presented theory of atoms 1. Each element is made up of tiny particles called atoms 2. The atoms of a given element are identical; atoms of different elements are different in some fundamental way/s 3. Chemical compounds are formed when atoms combine and said compound has same relative #s and types of atoms 4. Chemical reactions are the reorganization of atoms (how they are bound); the atoms themselves are not changes 1. O to H ratio known in water but couldn’t know the exact mass of each element 2 1. Assumed H was 1.00g and O was 8.00g 2. Did this for many other compounds and constructed first table of atomic masses* 1. Most were later proven incorrect 2. *Atomic Mass (Average)—The weighted average mass of the atoms in a naturally occurring element 3. Joseph Gay-Lussac & Amedeo Avogadro 1. Keys to determining absolute formulas 2. Lussac measured the volumes of gases that reacted with each other 3. Avogadro’s Hypothesis—At the same temperature and pressure, equal volumes of different gases contain the same number of particles 1. 2 volumes of H + 1 volume of O 2 volumes of water vapor can be expressed as 2 mol H + 1 mol O 2 mol of water 4. Most chemists didn’t accept b/c they didn’t atoms would form diatomic molecules 4. For many years there were many ideas and interpretations surrounding elements and their formulas 2.4 Cannizzaro’s Interpretations 1. First International Chemical Congress (1860) 1. Cannizzaro brought forth the ideas that chemists finally began to come to a consensus 1. Compounds contained whole numbers of atoms as Dalton postulated 2. Avogadro’s hypothesis was correct 2. Work was convincing because of the large amount of data he collected 1. Led to universal mass standards 2. Led to approximate vales of the relative atomic masses 2.5 Early Experiments to Characterize the Atom 3 1. JJ Thompson 1. Studied electrical discharges in partially evacuated tubes (cathode-ray tubes) 2. Cathode Rays—The “rays” emanating from the negative electrode (cathode) in a partially evacuated tube; a stream of electrons 1. Ray was produced at the negative electrode and repelled by negative field of the applied electric field 2. Ray is made of a stream of electrons 3. Electron—A negatively charged particle that moved around the nucleus of an atom 1. Charge-to-mass ratio: e=−1.76×10 C/g m 4. Electrons were produced from electrodes made of different metals thus all atoms contained electrons 1. From this the plum pudding model was created 2. Robert Millikan 1. Charged oil drop experiments 1. Was able to calculate the mass of an electron: −31 9.11×10 kg 4 3. Radioactivity (Radioactive Decay)—The spontaneous decomposition of a nucleus to form a different nucleus 1. Alpha, beta, and gamma particle emission 2. Gamma is high-energy “light” 3. Beta is a high-speed electron 4. Alpha has a 2+ charge 1. 7300 times the mass of an electron 4. Ernest Rutherford 1. Gold foil experiment to test Thomson’s model 1. Alpha particles should go through foil but many were deflected 2. Deflected particles had “close encounter” with positive center of atom 2. Nuclear Atom—An atom having a dense center of positive charge (the nucleus) w/ electrons moving around the outside 1. Nucleus—The small, dense center of positive charge in an atom 2.6 The Modern View of Atomic Structure: An Introduction −13 1. Atom contains a nucleus with diameter of 10 cm and ele−8rons that move around it at an average distance of 10 cm 2. Nucleus contains protons and neutrons 1. Protons—a positively charged particle in an atomic nucleus 2. Neutrons—a particle in the atomic nucleus with mass virtually equal to the proton’s but with no charge 3. Extremely dense (almost all of the atom’s mass) and a very small portion of the atom 3. # of electrons affects how and atom interacts with others thus the reason atoms of different elements show different chemical behavior 4. Isotopes—Atoms of the same element (the same # of protons) w/ different #s of neutrons. They have identical atomic #s but different mass #s 5 1. Show almost identical chemical properties 5. Atomic Number—The number of protons in the nucleus of an atom 6. Mass Number—The total number of protons and neutrons in the atomic nucleus of an atom 2.7 Molecules and Ions 1. Chemical Bond—The energy that holds two atoms together 1. Covalent Bonding—A type of bonding in which electrons are shared by atoms 2. Many representations of molecules 1. Chemical Formula—The representation of a molecule in which the symbols for the elements are used to indicate the types of atoms present and subscripts are used to show the relative numbers of atoms 2. Structural Formula—The representation of a molecule in which the relative positions of the atoms are shown and the bonds are indicated by lines 3. Space-Filling Model—A model of a molecule showing the relative sizes of the atoms and their relative orientations 4. Ball-and-Stick Model—A molecular model that distorts the sizes of atoms, but shows bond relationship clearly 3.Bonding between ions 1.Ion—An atom or group of ions that has a net positive or negative charge 1.Cation—A positively charged ion 2.Anion—A negatively charged ion 2.Ionic Bonding—The electrostatic attraction between oppositely charged ions 4.Ionic solids are salts are solids consisting of oppositely charged ions 1.Contains simple ions or polyatomic ions 6 1.Polyatomic Ion—An ion containing a number of atoms 2.8 An Introduction to the Periodic Table Periodic Table—A chart showing all the elements arranged in columns with similar chemical properties – Letters in boxes are symbols for elements – # above is atomic number (# of protons) – Most of the elements are metals Metal— Characteristic physical properties—good conductors, malleable, ductile, etc. Tend to lose electrons – The rest are nonmetals Nonmetal— Grouped in top right corner (except H) Lack physical properties that characterize metals Tend to gain electrons Often bond together by covalent bonds Elements in the same column have similar chemical properties – Group/Family—A vertical column of elements having the same valence electron configuration and showing similar properties – Alkali Metals—A group 1A metal – Alkaline Earth Metals—A group 2A metal – Halogens—A group 7A element – Noble Gases—A group 8A element Periods—Horizontal rows of elements 2.9 Naming Simple Compounds 7 Originally had common names (water, sugar, etc.) but too much chaos A system for naming compounds was adopted Binary Ionic Compounds—A two element compound containing a cation and an anion 1. The cation is always named first and the anion second 2. A monatomic cation takes its name from the element 3. A monatomic anion is named by taking the first part of the element’s name and adding –ide 4. For cations with multiple possible charges, the charge must be specified with a roman numeral written in parenthesis after the cation An older system used –ic and –ous to denote the form of the cation if there were only two forms Ionic Compounds with Polyatomic Ions – Polyatomic ions have special names that must be memorized – Oxyanions—Series of that contain an atom of a given element and different numbers of an oxygen atom Smaller # of O ends in –ite, larger ends in –ate If there are more than two, per– and hypo– are used to name the members with the fewest and most oxygen atoms Binary Covalent Compounds—A two element compound containing two nonmetals 1. The first element of the formula is named first, using the full element name 2. The second element is named as if it were an anion 3. Prefixes are used to denote the number of atoms present 8 4. The prefix mono- is never used for naming the first element 5. The final o or a of a prefix is often dropped when the element begins with a vowel Some compounds are always referred to by common names (water, ammonia, etc.) Acids 1. If the anion contains oxygen, the acid is named with the prefix hydro– and the suffix –ic 2. If the anion does not contain oxygen, the acid if formed from the root of the anion with a suffix of –ic or –ous If the anion ends in –ate, the suffix –ic is used If the anion ends in –ite, the suffix –ous is used 9