Dr. Streit Week 6 Notes
Dr. Streit Week 6 Notes CHEM 1030 - 003
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This 5 page Class Notes was uploaded by Rachel Ferrell on Friday February 26, 2016. The Class Notes belongs to CHEM 1030 - 003 at Auburn University taught by John D Gorden in Fall 2015. Since its upload, it has received 52 views. For similar materials see Fundamentals Chemistry I in Chemistry at Auburn University.
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Date Created: 02/26/16
Rachel Ferrell CHEM 1030 2/23/16 Chapter 5: Cont. Ionic Compounds and Bonding: • =electrostatic attraction that hold oppositely charged ions together in an ionic compound + -‐ o ex. Na + Cl→Na + Cl →combine to form NaCl or sodium chloride • chemical formula= denotes the elements of the ionic compound and the ratio that they occur in • lattice= a 3D array of cations and anions in an ionic compound o think of like a crystal structure o Lattice energy= amount of energy required to convert a mole of ionic solid to constitute ions in the gas phase § Basically is a measure of how stable the ionic compound is o This formation of ionic bonds releases a large amount of energy o Lattice energy depends on the magnitude of charges and distance between them o High lattice energy→small radius/distance→most stable o High lattice energy→higher charges→most stable § If charges are the same→look at the distance § Is distances are the same→look at the charges § Ex. Arrange MgO, CaO, and SrO by increasing lattice energy 2+ 2+ 2+ • Mg=Mg , Ca=Ca , Sr= • Since same charge, so order by atomic radius • Answer= SrO<CaO<MgO Naming Ions and Ionic Compounds: • Monotomic cation→ add –ion to the element • Monoatomic anion→ add –ide to the element • Some metals can form cations(not anions) of more than one possible charge 2+ o Ex. Fe = Fe(II) o Fe = Fe(III) Formulas for Ionic Compounds: • ionic compounds→always between metals and nonmetals • electrically neutral→sum of charges must be zero o ex. Al , O →Al O 2 3switch charges to be the subscript of the opposite element) • To name ionic compounds: o NaBr→ sodium bromide o FeCl →iro2(II) bromide o CaO→calcium oxide o Mg N → 3 m2gnesium nitride o Fe S →2 3n(III) sulfide o Exceptions § Transition metals can form more than one possible charge; therefore the roman numerals must be used § Always be sure to check that the overall charge is zero Covalent Bonding and Molecules: • Elements with similar properties share electrons to give each atom more stability and a noble gas configuration • Lewis Theory of Bonding o Depicts the bond formation of H 2 as 2 Hs sharing their electrons • Covalent bond= pair of shared electrons • Molecule= combo of two atoms held together by chemical bond o can be an element or a compound • Law of Definite Proportions= different samples of the same compound always have the same mass ratio of elements • Law of Multiple Proportions= if 2 elements can combine with each other to form 2 or more different compounds, then the ratio of masses of 1 element that combines with a fixed mass of the other element can be expressed in small whole numbers o ex. CO 2 can also be CO→ratio of O to C between these molecules is about 2:1 • Diatomic Molecules= contains 2 atoms; can be the same element or different elements o Heteronuclear= elements are different o Homonuclear= elements are the same • Polyatomic moleules= more than 2 atoms Molecular Formulas: • Chemical formula denotes the composition of the substanc2 (H O) • Molecular formula= shows the exact number of atoms of each element in a molecule o using the dot structure • some elements have 2 or more distinct forms known as allotropes o ex. O2 and 3 O • structural formula= shows not only elemental composition, but also the general arrangement of atoms o shows polarity, etc. Empirical Formula: • molecular substances can be represented using empirical formulas= whole number ratio of elements in a compound o the simplest chemical formula o ex. Molecular formula= N 2 4 o Empirical formula= NH 2 • empirical and molecular formulas can often be the 2same (H O) • ex. Problems: o C H 6 →12 O6 2 o C H 5 →5HN5 o N O→2 O 2 Naming Molecular Compounds: • Binary molecular compound→just 2 elements (two nonmetals) • How to name o 1) Name the first element in formula o 2)Name the 2 element, change ending to –ide • ex. HCl→hydrogen chloride • Greek prefixes used o Mono=1 o Di=2 o Tri=3 o Etc. • Rule= don’t use mono-‐ for the first element, only for the second o Ex. Carbon monoxide not monocarbon monoxide • Ex. Problems: o 1)NF →nit3ogen trifluoride o 2) N O → dinitrogen tetroxide 2 4 o 3) Sulfur tetrafluoride→ SF 4 o 4) Tetraphosphorus decasulfide→ P S 4 10 Compounds Containing Hydrogen: • usually don’t conform to normal naming system used for molecular compounds • many compounds with Hydrogen have common names: o B H =2ib6rane o SiH = Silane 6 o NH = 4Ammonia o PH = 3 Phosphine o H O= 2 Water o H S= 2 Hydrogen sulfide • Acid= produces H ions when dissolved in water o ex. HCl is an acid when dissolved in water→ o Naming Acids: § 1) Remove –gen from hydrogen § 2)change –ide ending on 2 element to –ic o ex. Hydrochloric acid o a compound must contain at least one ionizable hydrogen atom to be an acid when dissolved § ionizable hydrogen is one that separates from the molecule when dissolved in water + and becomes an H ion Organic Compounds: • also do not follow the normal naming system for molecular compounds • organic molecule= has carbon and hydrogen, sometimes other atoms • alkanes=simplest hydrocarbons o 1 carbon= methane o 2 carbon= ethane o 3 carbon= propane o 4 carbon= butane o 5 carbon= pentane • many organic compounds have functional groups o often determine a molecules reactivity o Alcohol= -‐OH o Aldehyde= -‐CHO o Carboxyl(acid)= -‐COOH o Amine(base)= -‐NH 2 Covalent Bonding in Ionic Species: • Polyatomic ions= combination of 2 or more atoms, held together by covalent bond • Compounds with polyatomic ions→follows same rules as naming ionic compounds o Molecule must be neutral overall o Ex. Ca and P4 →Ca (PO3) 4 2 • Common polyatomic ions below: MEMORIZE • • Ex. Problems: o 1) Fe (SO ) → Fe , SO → iron(III) sulfate 2 4 3 4 o 2) Al (OH) → 3 Al , OH → aluminum hydroxide o 3) Hg O→ 2 Hg , O → mercury (I) oxide • Oxoanions o =polyatomic ions that contain 1 or more oxygen atoms and one atom (central atom) of another element o Naming Oxoanions: § 1) Start with oxoanions that end in –ate and as reference point § 2) Ion with one more O= per…ate ion • ClO =3hlorate, ClO = 4erchlorate ion § 3) Ion with one less O= -‐ite ion • ClO =2 chlorite ion § 4) Ion with 2 less O= hypo…ite ion • ClO= hypochlorite ion o Must memorize common –ate oxoanions § Chlorate= ClO 3 -‐ -‐ § Nitrate= NO 3 § Phosphate= PO 4 3-‐ § Sulfate= SO -‐ 4 § Bromate= BrO 3-‐ • Oxoacids + o Polyatomic ions that have 1 or more oxygen atom plus central atom, AND can release H ion when dissolved in water o Naming oxoacids: § 1) An acid with -‐ate→change to –ic acid • HClO =ch3oric acid § 2) An acid with -‐ite→change to –ous acid • HClO = 2 chlorus acid § 3)Prefixes in oxoanion names stay the same for corresponding acids • HClO = 4 perchloric acid • HClO= hypochloric acid o Oxoacids can be monoprotic or polyprotic § Monoprotic= only has 1 ionizable H atom § Polyprotic= has more than 1 ionizable H atom § Names of anions indicate the number of remaining hydrogen • H 3O = 4 phosphoric acid • H 2O = 4 dihydrogen phosphate ion 2-‐ • HPO = 4 hydrogen phosphate ion • PO =4 phosphate ion
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