Chemistry 1211 Chapter 2 Book & Lecture Notes
Chemistry 1211 Chapter 2 Book & Lecture Notes 1211
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This 14 page Class Notes was uploaded by Caroline Holt on Sunday August 21, 2016. The Class Notes belongs to 1211 at University of Georgia taught by Jay Agarwal in Fall 2016. Since its upload, it has received 16 views. For similar materials see General Chemistry 1 (CHEM 1211) in Chemistry at University of Georgia.
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Date Created: 08/21/16
Chemistry 1211: Agarwal Chapter 2 Book Notes 1 Structure of the Atom 1.1Components of the Atom o Elements are characterized by the number and type of particles of which they are composed o Atoms consist of subatomic particles: protons, electrons, neutrons Proton: carries +1 charge, has a mass of 1.672622E-24 g Neutron: no electrical charge, has a mass of 1.674927E-24 g Electron: carries -1 charge, has a mass of 9.109383E-27 g (relatively massless) o Protons, neutrons are found in the nucleus of the atom o Electrons account for most of the volume of an atom o The mass and charge of an atom affect the physical and chemical properties of the element and the compounds it forms Atomic Mass Unit (u): defined as ½ the mass of a carbon atom that contains 6 protons and 6 neutrons Ion: an atom with an unequal number of protons and electrons; because they aren’t equal, the atom has an overall positive or negative charge Cation: more protons than electrons, positively charged Anion: more electrons than protons, negatively charged 1.2Atomic Number, Mass Number, and Atomic Symbols o Atoms of each element are distinguished by the number of protons in the nucleus Atomic number (Z): equal to the number of protons; identifying number of the element Mass number (A): number of protons and neutrons in the nucleus of an atom Atomic symbol: one or two letter symbol that represents the element along with the atomic number o The number of neutrons is equal to the difference between the mass number and the atomic number 1.3Isotopes & Atomic Weight o All atoms of a specific element have the same number of protons, but some have different numbers of neutrons Isotopes: atoms that have the same number of protons, but a different number of neutrons (same Z, different A) o Isotopes are named using the element name and mass number Atomic weight: average mass of all naturally occurring isotopes of that element, taking into account relative abundance of the isotope Percent Abundance: percentage of the atoms of a natural sample of the pure element represented by a particular isotope, to describe isotope composition for an element o What percentage each isotope of an element makes up of the total number of isotopes of the element o To calculate average atomic weight: ∑ = (mass 1) x (abundance 1) + (mass 2) x (abundance 2) 2 Elements & the Periodic Table 2.1Introduction to the Periodic Table Alkali Metals (1A): all shiny solids that react vigorously with air, water, and halogens Alkaline Earth Metals (2A): react with water to form alkaline solutions Noble Gases (8A): least reactive Main Group Elements: A groups; representative elements Transition Metals: B Groups o 1A: Alkali Metals o 2A: Alkaline Earth Metals o 6A: Chalcogens o 7A: Halogens o 8A: Noble Gases o Period 6 = lanthanides o Period 7 = actinides o Left side of table = metals o Right side of table = non-metals o In between = metalloids, semi-metals o Metals: shiny solids that are ductile, good conductors o Non-metals: dull, brittle solids/gases that do not conduct electricity (Br is only non-metal liquid) o Metalloids: have properties of both non-metals and metals o Many elements exist as molecules consisting of 2 or more atoms of an element o Diatomic Molecules: H 2 N 2 O 2 F 2 Cl2 Br 2 I2 Allotropes: forms of the same element that differ in their physical and chemical properties 3 Covalent Compounds 3.1Introduction to Covalent Compounds Covalent Compounds: consist of atoms of different elements held together by covalent bonds Molecular covalent compounds: each molecule held together by covalent bonds Network covalent compounds: 3-D network of atoms held together by covalent bonds 3.2Representing Covalent Compounds with Molecular and Empirical Formulas Molecular formula: way to represent a molecule; contains the symbol for each element and subscript to identify the number of atoms in the molecule o Chemical formulas always show a whole number ratio of elements Empirical formulas: simplest whole number ratio of elements in the compound o Both molecular covalent and network covalent compounds are represented using empirical formulas Structural formula: shows the linkage of all atoms in the molecule Condensed structural formula: lists the atoms present in groups to indicate connectivity between the atoms o CH 3H 2H C2 O2 3.3Representing Covalent Compounds with Molecular Models Wedge-and-dash model: 2-D representation of a 3-D structure o Key: Ball-and-Stick model: shows atoms as colored spheres connected by sticks that represent covalent bonds Space filling model: interpenetrating spheres represent the relative amount of space occupied by each atom in the molecule o Useful for considering the overall shape of molecule, how molecules interact when they come in contact with one another 3.4Naming Covalent Compounds: o Categorized in 2 ways: Binary Non-Metals Inorganic Acids o Compounds often belong to more than 1 of the categories o Rules for naming Binary Non-Metals: First word = name of 1 element in compound If there is more than 1 atom of the first element, use a prefix Secord word = name of 2 ndelement with -ide ending Always use a prefix to indicate number of atoms CS2= Carbon Disulfide Hydrocarbons: binary non-metal compounds contain ONLY hydrogen and carbon o Class of organic compounds o Named according to the number of +arbon and Hydrogen atoms Inorganic Acids: produce Hydrogen Ion (H ) when dissolved in water o Contain Hydrogen and 1 or more non-metals o Hydrogen is ALWAYS 1 element in compound formula o Same rules as binary non-metal compounds WITHOUT prefixes o Also named using common names: MEMORIZE LIST OF COMMON NAMES o Hydrogen Halides (HF, HCl, HBr, HI) are binary non-metals names when they are gases, but use common names when dissolved in water Oxoacids: group of acids that differ ONLY in number of oxygen atoms o Named according to number of oxygen atoms o -ic ending: acid with more Oxygen atoms o -ous ending: acid with fewer Oxygen atoms 4 Ions & Ionic Compounds 4.1Monoatomic Ions Ionic Compounds: contain ions; carry positive (cation) charge or negative (anion) charge o Ionic bonds held together by strong attractive forces between cations, anions o Most ionic compounds are solids with high melting points Covalent compounds tend to be gases, liquids, and solids with low melting points o Neutral atom has equal number of protons and electrons Monoatomic Ion: when the number of protons and electrons are no longer equal because of gaining/losing electrons o Charge is written with superscript to the right o Cations, anions have physical and chemical properties different from the original element o General Ion Facts: Groups 1A, 2A, and 3A in periodic table form cations with charge equal to group number Groups 5A, 6A, and 7A form anions that have negative charge equal to 8minus the group number Hydrogen forms both cations and anions Transition metals usually form cations with charges from +1 to +3 Many form multiple monoatomic ions Group 4A contains metals and non-metals that behave differently Other than Aluminum, metals in groups 3A, 4A, and 5A form cations with charges that are not predictable o IMPORTANT TO KNOW: Metals form cations Non-metals form anions Noble gases are happy and do not form ions 4.2Polyatomic Ions: Polyatomic ions: groups of covalently bonded atoms that carry OVERALL positive or negative charge o MEMORIZE common polyatomic ions o Most polyat+mic ions are anions NH 4 is the only exception o List of common polyatomic ions: + - NH 4 = Ammonium NO =2Nitrite OH = Hydroxide NO =3Nitrate - - CN = Cyanide ClO = Hypochlorite CH 3O = 2cetate ClO =2Chlorite 2- - SO 3 = Sulfite ClO =3Chlorate SO 42-= Sulfate ClO =4Perchlorate HSO = Hydrogen Sulfate (Bisulfate) CO 2-= Carbonate 2- 3 - S 2 3 = Thiosulfate HCO = 3ydrogen Carbonate (Bicarbonate) 3- 2- PO 4 = Phosphate C 2 4 = Oxalate HPO 42-= Hydrogen Phosphate Cr 2 72-= Dichromate - 2- H 2O =4Dihydrogen Phosphate CrO 4 = Chromate SCN = Thiocyanate MnO = 4ermanganate - OCN = Cyanate 4.3Representing Ionic Compounds with Formulas o Ionic compounds represented by empirical formulas: which show the simplest ratio of cations to anions in the compound o Cation written 1 and then anion o Total canionic positive charge balanced with total anionic negative charge Ionic compounds DO NOT have a charge 4.4Naming Ionic Compounds o Ions and ionic compounds are named by identifying charges on monoatomic and polyatomic ions in the formula o Memorize rules for predicting charges on monoatomic ions o Memorize names, formulas, and charges on polyatomic ions o Rules for naming Ions and Ionic Compounds: Monoatomic Cations: The name of main-group monoatomic cation is the element name followed by the word ion o Na = sodium ion The name of transition metal cation is the element name followed by the cation charge in Roman numerals within parentheses and is also followed by the word ion o Fe2+ = iron (II) ion Monoatomic Anions: The name of a monoatomic ion is the element name changed to include the suffix -ide, followed by the word ion - o Br = bromide ion Polyatomic Ions: The names of the polyatomic ions must be memorized Names are followed by the word ion o NO 3 nitrate ion Ionic Compounds: Consists of the cation name followed by the anion name NOT followed by the word ion, because the compound does not carry an overall charge o Charges of individual polyatomic ions that make up the ionic compound cancel out each other’s charge NO prefixes used to indicate the number of ions present in the formula of an ionic compound 4.5Identifying Covalent & Ionic Compounds I. Covalent: a. Only non-metals b. Use prefixes for naming II. Ionic: a. Contain monoatomic or polyatomic ions b. Contains metals and non-metals or only non-metals c. NEVER uses prefixes d. Sometimes named with cation charge in roman numerals in parentheses
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