Dr. Streit Exam 2 Study Guide
Dr. Streit Exam 2 Study Guide CHEM 1030 - 003
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This 5 page Study Guide was uploaded by Rachel Ferrell on Sunday March 6, 2016. The Study Guide belongs to CHEM 1030 - 003 at Auburn University taught by John D Gorden in Fall 2015. Since its upload, it has received 142 views. For similar materials see Fundamentals Chemistry I in Chemistry at Auburn University.
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Date Created: 03/06/16
Rachel Ferrell CHEM 1030 Exam 2 Study Guide: Chapter 4 Vocab: Ionization energy= the minimum energy required to remove an electron from an atom o Increases from left to right o Exceptions: decreases Group 2A-3A and Group 5A-6A Electron affinity= energy released when an atom accepts an electron o Increases from left to right o Exceptions: decreases from Groups 1A-2A and Group 4A-5A Q1xQ2 Columb’s Law= F α d2 Isoelectrons= when an atom gets the same electron configuration as noble gases due to gaining or losing electrons Ionic radius= the radius of a cation or an anion o Cation radius<atomic radius o Anion radius> atomic radius Chapter 5: Vocab: Ionic compound= electrostatic attraction that holds oppositely charged ions together Lattice= 3D array of cations and anions in an ionic compound Lattice energy=measure of how stable the ionic compound is o High lattice energysmall radius/distancemost stable o High lattice energyhigher chargesmost stable 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 between the 2 compounds can be represented by a whole number o ex. CO c2n also be CO ratio of O to C between these molecules is about 2:1 Diatomic molecules= contain 2 atoms; can be the same element or different elements Heteonuclear= 2 molecules are different Homonuclear= elements are the same Polyatomic molecules= more than 2 atoms Molecular formula= exact number of atoms in each element Empirical formula= the simplest chemical formula Structural formula= shows elemental composition and the general arrangement of atoms in space Hydrate= a compound that has a specific number of water molecules associated with its solid structure (ex. CuSO x 5H O) 4 2 Anhydrous= when hydrate is heated and the water molecules are driven off; often can cause the substance to change color Molecular mass= sum of atomic masses in a molecule of atoms; only for molecular compounds Formula mass= same as molecular mass but for ionic compounds (found the same way) Percent composition by mass= nxatomic massof element x 100 molecular∨ formulamassof anelement Chapter 5: Concepts to Know: Naming Compounds o Ionic Comstunds: between metal ndd nonmetal; neutral 1 element is same, 2 element ends in –ide for d-block elements roman numeral represents charge (except for Ag , Zn , Al )+ ex. iron (II) bromide Oxoanions Add –ate to the base 1 more than baseper…ate 1 less than base-ite 2 less than basehypo…ite o Molecular compounds: between 2 nonmetals Use greek prefixes Ex. nitrogen trifluoride o Acids: contains enough hydrogen to cancel out the charge of the other element Add hydro- to beginning Add –ic to the end Ex. HClhydrochloric acid o Oxoacids: contains enough hydrogen to cancel the charge of the polyatomic ion( that contains oxygen) Base (-ate) change to –ic acid 1 less than base(-ite)change to –ous acid 2 less than base(hypo-)change to hypo…ous acid 1 more than base(per-)change to per…ic acid o Organic Compounds: usually is just carbon and hydrogen Alkanes=simplest hydrocarbons: 1 carbon= methane 2 carbon=ethane 3 carbon= propane 4 carbon=butane 5 carbon= pentane Memorize polyatomic ions: (don’t forget charges!!) Common name compounds: o NH a4monia o PH phosphine 3 o H S2hydrogen sulfide Chapter 6: Vocab: The Octet Rule= says atoms with lose/gain/share electrons in order to get a full valence shell Bond length= the distance between nuclei of 2 covalently bonded atoms; shows strength of chemical bond o Shorter bond length= stronger bond o Double bonds are shorter/stronger than single bonds, etc. Formal charge= valence - (lines+dots); must equal the overall charge of the compound Chapter 6: Concepts to know: Electronegativity= the ability of an atom to draw in electrons o Increases from left to right/bottom to top o exceptionsd-block elements Types of bonds that can form: o 1) pure covalent bond=neutral atom; equally shared electrons; usually same element (ex. O ) 2 o 2) polar covalent bond= partially charged atom; unequally shared electrons (M δ +X δ -) o 3) ionic bond= oppositely charger ions; held together by electrostatic attraction(M X )- How to know what type of bond o Look at the difference it atoms electronegativity o If the atoms electronegativity differs by: Less than 0.5nonpolar 0.5-2.0polar covalent greater than 2.0ionic Dipole moment= the measure of the polarity of a bond; also indicates direction of electron shift o Dipole moment = Q x r Q= charge R= radius (distance) Dipole moment al-30s positive; expressed in debye units (D) 1D= 3.336x10 cxm charge of separation represented by δ (+ or -) Exceptions to the Octet Rule: o Group 3A can have less than 8 valence o Coordinate covalent bond1 atom donate both electrons o Lewis basedonates electron o Lewis acidaccepts electron o Free radical= molecule with odd number of electrons o Period 3 and beyond can have more than 8 valence Practice Problems: 1. Arrange MgO, CaO, and SrO by increasing lattice energy a. Answer= SrO<CaO<MgO 2. Write the empirical formula a. C 6 126H O 2 b. N 2N O 2 3. Name the compound or write the formula a. CuSO =4copper(I) sulfate b. NH 4O =3ammonium nitrate c. KClO 3 potassium chlorate d. CuCl= copper(I) chloride e. Mn S = manganese (III) sulfide 2 3 f. FeCl 2 iron (II) chloride g. Nickel(II) sulfate= NiSO 3 h. HIO= hypoiodous acid i. HBrO =4perbromic acid j. Lithium hydrogen phosphate= Li HPO 2 4 k. HBr= hydrobromic acid 4. Practice drawing lewis structures 5. Practice working with the Dipole moment equation 6. Don’t foget to memorize the elements and their symbols
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