CHM2045 Exam I Study Guide
CHM2045 Exam I Study Guide CHM 2045
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This 8 page Study Guide was uploaded by Daniel Donovan on Friday September 16, 2016. The Study Guide belongs to CHM 2045 at University of Florida taught by Dr. Martina Sumner in Fall 2016. Since its upload, it has received 14 views. For similar materials see Freshman Chemistry 1 in Chemistry at University of Florida.
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Date Created: 09/16/16
CHM2045 Exam I Study Guide Ch 1 Important things you should know… Distinguish between physical and chemical properties and change See the pictures of molecular arrangements of the three phases Be able to define kinetic and potential energy as well as relationship Explain the phlogiston theory of combustion Summarize the scientific approach Know the units of mass, time, length, temperature, volume Determine the number of sig figs Significant figure rules: Multiplying/Dividing: the number of significant figures will the same as the number with the least number of significant figures Ex. 5.36 * 12.547 * 12.3 /(1.2) : the number of sig figs would be 2 since the lowest number, 1.2 has only 2 significant figures. Adding and Subtracting: the number of significant figures will be the same as the least number of decimal places. Ex. 12.001 + 12.4 : Number of sig figs = 3 because the least number of decimal points is 1. Ch 2: The Components of Matter Elements, Compounds, and Mixtures Elements and compounds are pure substances. Elements – consist of only one type of atom and may exist as individual atoms or as molecules (two or more atoms bonded together). Compounds – two or more different elements that are chemically bonded together. Mixtures – impure substances that consist of pure substances physically mixed together in no specific ratio. May either be homogenous ( no boundaries visible) or heterogenous (boundaries are visible). Section 2.2/2.3 Atomic view of matter and Dalton’s Atomic Theory Law of Mass Conservation – mass is always conserved in chemical transformations. Law of Constant Composition – All molecules of water have the same chemical formula and if each kind of atom has on the average the same mass. Mass Fraction – the mass of an element divided by the total mass of the compound. Dalton’s Theory: 1. Each element is made up of tiny particles called atoms. 2. Atoms of one element cannot be converted into atoms of another element. Instead they can rearranged by chemical reactions. 3. Atoms of an element are identical, and are different from atoms of other elements. 4. Compounds form when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. Avogadro’s Proposal: Equal volumes of gases contain equal number of gas particles. Hydrogen and chlorine exist as polyatomic molecules that split apart and recombine to form gaseous hydrogen chloride molecules. 2.4 Observations that led to the discovery of the Nucleus. JJ Thompson concluded that the rays appeared when negatively charged particles collided with the gas particles in the tube. o He called these particles electrons. Robert Millikan measured the charge of the electron by measuring the effect of an electric field on the rate charged oil drops fell under gravity. Ernest Rutherford established that an atom contains a tiny, positively charged center called the atomic nucleus. o He bombarded a piece of thin gold foil with alpha particles. The Discovery of Neutrons James Chadwick discovered the electron. This particle is present in the nucleus with the protons and accounts for the remaining mass of the atom. Atom Definitions: Atomic Number (Z): The number of protons in the nucleus of an atom. Z= #protons Mass Number (A): The total number of protons and neutrons in an atom. A = #protons + #neutrons A – Z = (number of protons + number of neutrons) – protons = neutrons = N Atomic Symbol: A symbol for each element based on its English, Latin, or greek name. Isotopes: Atoms with the same number of protons (same element), but a different amount of neutrons. Atomic mass unit (amu) : ½ the mass of a carbon -12 atom. The mass of a carbon-2 atom is defined as 12 atomic mass units. Atomic mass : The average of the masses of an element’s naturally occurring isotopes weighted according to their abundance. 2.6 The Periodic Table Elements are classified as metals, nonmetals, or metalloids depending on their physical and chemical properties. 2.7 Ionic and Covalent Bonding Terms: Ionic compounds: form when an atom of one element transfers electrons to atoms of another element. (Metals and Nonmetals) Covalent compounds: form when atoms share electrons Ion: a charged particle Cation: a positively charged particle (paws-itive) Anion: a negatively charged particles Monoatomic ion: an ion derived from a single atom Polyatomic ions: two or more atoms bonded covalently and having a net positive or negative charge, i.e : NH4, SO 4- Types of Formulas Empirical formula: shows the simplest whole number ratio of atoms of each element in a compound. Molecular Formula: shows the actual number of atoms of each elements in the molecule. Structural formula: shows the actual number of atoms of each element in a molecule an the bonds between them. M 2.8 Naming Compounds Ionic Compounds Consider NaCl o The cation is named first followed by the anion o The name of the cation (sodium) is the same as the name of the neutral metal. o The name of the anion (chloride) takes the root of the nonmetal name and adds the suffix “-ide Ionic Compounds Containing Polyatomic Ions Polyatomic ions stay together as a charged unit. If there are more than one of the polyatomic ion, the ion is written in parentheses with the subscript outside o Ex. AgNO 3nd Ca (3O )4 2 Naming Oxoanions: If a polyatomic ion contains one or more oxygen atoms they are called oxoanions. A. With two oxoanions in the same family use the suffixes: a. “-ate” for the one with more O atoms b. “-ite” for the one with less O atoms B. With four oxoanions in the family, “per” (more than) and “hypo-“ (less than) is added to the names in addition to the previous 2. EX. ClO4is perchlorate - ClO3 is chlorate ClO2 is chlorite - ClO is hypochlorite. Naming Acids An acid is a molecule with one or more H+ ions associated to be an anion. o If the anion does not contain oxygen, the acid is name with the prefix “hydro” and the suffix “ic” Ex. HCL = hydrochloric acid o If the anion contains oxygen, the names are similar to those of the oxoanions, except: Ate becomes “ite” and “-ite” becomes “ous” Binary Covalent Compounds Ex CO 2 The first element is the one on the lower group number in the periodic table. If both elements are in the same group, the element with the higher period number is named first. The second element is named with the suffix “-ide” Greek prefixes indicate the number of atoms of each element in the compound. o Ex. PC3 = phosphorus trichloride o Cl2O7= dichlorine heptaxide Alkanes Special type of binary covalent compounds that contain only carbon and hydrogen Molecular Masses from Chemical Formulas Molecular mass = sum of atomic masses Molecular mass of NH = 31 * atomic mass of N) + (3 * atomic mass of H) = (14.01 amu) + (3 * 1.008 amu) = 17.03 amu 2.9 Classification of Mixtures Heterogeneous mixtures - one or more visible boundaries between between the components Homogeneous – no visible boundaries because the components are mixed evenly. Also called a solution. Solutions in water are called aqueous solutions. Ch 3 Stoichiometry 3.1 The Mole Atomic mass units allow us to know only the relative atomic masses of the elements. The mole is the SI unit for amount Relative atomic masses: 1. How many grams of iron would contain the same number of iron atoms as the same number of carbon atoms contained in 1.000 g of Carbon? Solution: (55.85 g Fe)/(12.01 g C) * (1.000 g C) = 4.650 g Fe Avogadro’s Number: 1 mole = 602,200,000,000,000,000,000,000 Avogadro’s number = 6.022 * 10^23 entities The Atomic Point of View of Mass Total Mass = mass per unit population * population Total mass = molar mass * amount m = M * n Molar mass = M = m/n The common units for M, Molar mass g/(mole) atoms (often written as g/mol) Converting form mass to atoms: First, calculate the mass of the compound or element Next, convert the mass to moles Finally, convert moles to atoms. o Ex. You weigh a nickel and find that it weighs 5.0012 grams. How many nickel atoms are in the nickel. Nickel is 25% nickel and 75% copper o Solution: 0.2500 * 5.0012 = 1.250 g 1.250 g /(58.69 g/mol) = 0.02130 moles 0.02130 moles *(1 amu / 6.022 * 10^23 atoms) = 1.283 * 10 ^22 atoms Converting from mass to moles and molecules First calculate the molar mass Use the molar mass to convert from grams to moles. Finaly, convert the moles to molecules. 3.2 Determining Empirical and Molecular Formulas %C = 52.1, %H = 13.1, %O = 34.7 Assume that this is out of a 100.0 gram sample 52.1 g C / (12.01 g/mol) = 4.34 mol C (13.1 g H)/(1.008 g/mol) = 13.0 mol H (34.7 g O) / (16.000 g/mol) = 2.17 mol O Preliminary equation = C 4.3413.02.17 Next divide each number of moles by the smallest number C: 4.34 / 2.17 = 2.00; H: 13.0/2.17 = 5.99; O: 2.17/2.17 = 1.000 The empirical formula is C H O 2 6 M = (2 mol C * 12 g/mol) + (6 mol H * 1 g/mol) + (1 mol O * 16 g/mol) = 46 g / mol Molecular formula is the same since the masses are about the same. Limiting Reactant: When one reactant is used up and the reaction stops. The limiting reactant is simply the reactant that runs out first. How to find: o Write the balanced equation o Calculate moles of each reactant o Using the coefficients from the balanced equation, determine how many moles one substance would react with another, o Calculate the total amount produced by the limiting reactant Theoretical, Actual, and Percent Yield %Yield = (actual yield)/(theoretical yield) * 100 Actual yield – the amount of product actually attained Theoretical yield – the amount of product formed based on the limiting reactant 3.5 Solution Stoichiometry Molarity = moles of solute / liters of solution Ex. What is the molarity of a 750 ML solution containing 14.8 g of Silver nitrate 14.8 g /(169.9 g/mol) = 0.08711 mole Molarity = 0.08711 mole/ 0.750 L = 0.116 M Ch 4 The Major Classes of Chemical Reactions 4.2 Aqueous Ionic Reactions : Precipitation and Acid-Base Reactions Precipitation reactions When 2 aqueous ionic com+ounds are -ixed, a precipitate, (s), may form o AgNO (3) -> Ag (aq) + NO 3aq) Total ionic equations – show all the ions, even if they are not involved in the reaction. Net Ionic equation – eliminate spectator ions and only record ions involved in precipitation Molecular Equations show all reacts and products as unassociated compounds. Usually occurs as a double-displacement reaction. Solubility Rules 1. Salts containing Group I elements are soluble (Li+, Na+, K+, Cs+, Rb+). Exceptions to this rule are rare. Salts containing the ammonium ion (NH4+) are also soluble. 2. Salts containing nitrate ion (NO3-) are generally soluble. 3. Salts containing Cl -, Br -, I - are generally soluble. Important exceptions to this rule are halide salts of Ag+, Pb2+, and (Hg2)2+. Thus, AgCl, PbBr2, and Hg2Cl2 are all insoluble. 4. Most silver salts are insoluble. AgNO3 and Ag(C2H3O2) are common soluble salts of silver; virtually anything else is insoluble. 5. Most sulfate salts are soluble. Important exceptions to this rule include BaSO4, PbSO4, Ag2SO4 and SrSO4 . 6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not soluble. 7. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS, ZnS, Ag2S are all insoluble. Arsenic, antimony, bismuth, and lead sulfides are also insoluble. 8. Carbonates are frequently insoluble. Group II carbonates (Ca, Sr, and Ba) are insoluble. Some other insoluble carbonates include FeCO3 and PbCO3. 9. Chromates are frequently insoluble. Examples: PbCrO4, BaCrO4 10. Phosphates are frequently insoluble. Examples: Ca3(PO4)2, Ag3PO4 11. Fluorides are frequently insoluble. Examples: BaF2, MgF2 PbF2. (From the Chemistry: A Molecular Approach Textbook) Acid-Base Reactions: o Acid – substance that produces protons (H+ ions) when dissolved in water, o Base substance that produces hydroxide (OH-) ions when dissolved in water. Neutralization Reaction o Reaction of an acid that produces water and a salt (on evaporation) in a neutralization reaction. o The net ionic equation shows an acid base as a kind of “proton transfer process) 4.3 Oxidation-Redux Reactions Rules for Assigning Oxidation Numbers o An atom in its elemental form has an oxidation number of 0 o A monoatomic ion has an oxidation number equal to its charge o The sum of the oxidation numbers for the atoms in a compound equals 0. o Some elements have the same oxidation number in almost all compounds Rules for Balancing Redox Reactions o Assign oxidation numbers o Identify the oxidizes and reduced atoms (atoms whose oxidation number changes) o Make the number of electrons lost equal to the number of electrons gained by multiplying by appropriate factors. o Complete the balancing by Inspection. Key: Color - definitions Color – equations Color – important people
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