Exam 1 Review Guide
Exam 1 Review Guide CHM 113
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This 5 page Study Guide was uploaded by Andrew Notetaker on Wednesday September 14, 2016. The Study Guide belongs to CHM 113 at Arizona State University taught by Cabirac in Fall 2016. Since its upload, it has received 107 views. For similar materials see General chemistry 1 in Science at Arizona State University.
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Date Created: 09/14/16
Exam 1 Review Wednesday, September 14, 2016 3:48 PM Chapter 1 (Matter and Measurement) The definition of chemistry and the five major branches of chemistry Chemistry is a branch of physical science that studies the composition, structure, properties and change of matter. Five branches are physical, analytical, biochemistry, organic and inorganic. Know the names and symbols for elements 1 – 36 (hydrogen – krypton) Classifications of matter (pure substance, mixture, homogeneous mixture, heterogeneous mixture, element, compound) Pure Substances: Elements- a substance that cannot be made into simpler substances Compounds- two or more elements bonded together Molecules Homogenous Mixtures (solutions) • Constant composition throughout Examples: air, salt water, metal alloys Heterogeneous Mixtures (solutions) Not constant composition throughout Examples: cookies Know the law of constant composition for compounds. All samples of a given chemical compound have the same elemental composition by mass. Ex: oxygen makes up about 8/9 of the mass of water while hydrogen makes up the remaining 1/9 of the mass. The difference between a physical and chemical property and phy/chem changes. Physical properties: Can be observed without changing a substance into another substance • Boiling point, density, mass, volume Chemical Properties: Can only be observed when matter undergoes a chemical change • Iron reacts with oxygen to form iron oxide, sodium reacts with water to give off hydrogen gas Physical changes Changes in matter that do not change the composition of a substance Changes of state, temperature, volume Chemical changes Chemical changes result in new substances The six types of phase changes (freezing, melting, vaporization, condensation, sublimation, vapor deposition) Freezing: liquid to solid, melting: solid to liquid, vaporization: liquid to gas, condensation: gas to liquid, sublimation: solid to gas, deposition: gas to solid The difference between intensive and extensive properties An extensive property is a property that changes when the size of the sample changes. Examples are mass, volume, length, and total charge. An intensive property doesn't change when you take away some of the sample. The 5 fundamental SI units of measurement (kg, m, s, K, mol) Kg- Kilogram m-meter s-second CHM 113 Lecture Page 1 s-second K-Kelvin mol- mole How to convert between Kelvin and Centigrade temperatures; marker temperatures on both scales in reference to water K=273.15+C Kelvin Celsius Melting point 273.15 0 Boiling point 373.2 100 The definitions and common units of volume and density 3 Cm L g kg How to use density as a conversion factor 125 g of x, x has a density of 40 g/l. How many liters of x are there? 125g * 1L/40g= 3.125L The metric prefixes discussed in lecture (red highlighted on slides; Giga pico (p, 1− 12)in Table 1.5) and how to covert between them Giga G 10^9 Mega M 10^6 Kilo K 10^3 Centi C 10^-2 Milli M 10^-3 Micro U 10^-6 Nano N 10^-9 Pico P 10^-12 Significant Figures in numbers and how to apply to calculations Measured value # of SF 2.456 4 1003.2 5 1.0300 6 .0000402 3 230000 2 1.230 x 108 4 When calculating use least number of SF in the given amounts. Ex: 1.00g of CH 28g of H O 4 2 How many sig figs in answer? 2, 28g has two sig figs. The meaning of Precision and Accuracy Accuracy- refers to the proximity of a measurement to the true value of a quantity Precision- refers to the proximity of several measurements together Dimensional Analysis-this could involve one or multiple conversion factors Convert 45 sec to minutes 45 sec * 1 min/ 60 sec Chapter 2 (Atoms, Molecules, and Ions) Dalton’s Atomic Theory and the laws of Conservation of Matter and Constant Composition The theory that atoms are the fundamental building blocks of matter reemerged in the 19th century, from John Dalton. The law of constant composition: In a given compound, the relative numbers and kinds of atoms are constant. The law of conservation of mass: The total mass of materials present after a chemical reaction is the same as the total mass present before the reaction. CHM 113 Lecture Page 2 same as the total mass present before the reaction. The law of multiple proportions: If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers. Groundbreaking experiments that led to the understanding of atomic structure (JJ Thomson’s CRT, Millikan’s Oil Drop expt., discovery of radioactivity, Rutherford’s alpha particle expt.) 8 Thomson using his cathode experiment determined the charge to mass of an electron was 1.76 X 10 coulombs per gram. Once the charge -to- mass ratio was determined, in 1909 Robert Millikan of the University of Chicago measured the charge of an electron using his oil-drop experiment which resulted in a calculation of -19 1.602 X 10 C. He then calculated the mass of an electron to be 9.10 X 10 -2g Rutherford discovered electrons and most of atoms were empty space with his gold-foil experiment Thomson's plum-pudding model explained the atom as filled with electrons, which was short lived. The three elementary particles along with their charges, symbols, and relative masses The structure of the atom The nucleus is made of protons and neutrons, which are equal unless the atom is an isotope. Electrons surround the nucleus. How to use atomic and mass numbers to determine the number of elementary particles Atomic number is number of protons and electrons. Mass number is protons + neutrons. Define the term isotope. An element with a different amount of neutrons than protons How the mass of an element on the periodic table is determined, i.e., determine the average atomic mass or relative abundance given information about existing isotopes. Atomic weight = SUM [(Isotope mass) X (fractional isotope abundance)] All isotopes of the element The layout of the modern periodic table. Names and locations of the 4 groups listed in class: 1(1A), 2(2A), 17(7A), and 18(8A). Alkali metals (1st column from the left), alkali earth metals (2nd column from the left), halogens (17th column from the left) , noble gases (18th column from the left). The location of metals, metalloids, and nonmetals on the periodic table. Metals are the left side, metalloids are a staircase dividing metals from nonmetals, nonmetals are on right side Basic properties of metals, nonmetals, metalloids, and noble gases – at this point, just know how these differ by tendency to gain/lose electrons. Metals- lose electrons Metalloids-Both gain and lose electrons CHM 113 Lecture Page 3 Metalloids-Both gain and lose electrons Nonmetals- gain electrons Chemical formulas of the elements including the ones that exist naturally as diatomic molecules A molecule made up of two different atoms is called a diatomic molecule. Compounds composed of molecules contain more than one type of atom and are called molecular compounds. CH fo4 example. Chemical formulas that indicate the actual numbers of atoms in a molecule are called molecular formulas. Chemical formulas that give only the relative number of atoms of each type in a molecule are called empirical formulas. H 2 2> HO Characteristics of an ionic bond and what elements form them Ionic compounds are generally combos of metals and nonmetals. Molecular compounds are generally composed of nonmetals only as in H O2 These two ions are not covalently bound, but because opposite charges attract they associate with an ionic bond(Metallic ion and non-metal ion) Predicting the chemical formula of a compound formed by two ions Na + Cl = NaCl Characteristics of a covalent bond and what elements form them Bond between two non-metals, electrons are shared Be able to assign oxidation number (ox states) to elements in compounds (rules provided) MgCl 2 Mg Cl2 -1 How to name ionic and binary molecular compounds (binary ionic, binary ionic with multiple oxidation states, ionic with polyatomic anions, and binary molecular) Ionic Compounds Names of ionic compounds consist of the cation name followed by the anion name. Names and Formulas Binary Molecular Compounds The procedure used for naming binary (two-element) molecular compounds are similar to naming ionic compounds 1. The name of the element farther left is written first, (unless the compound contains oxygen and chlorine, bromine or iodine [any element except fluorine]) 2. If both are in the same group the one closer to the bottom is named first. 3. The name of the second element is given an -ide ending 4. Greek prefixes indicate the number of atoms in each element. (di-,tri-….) The chemical formulas and names of acids, and bases (including which acids/bases are strong/weak) – know the definitions of acids and bases. Names and Formulas of Acids Acids containing anions whose names end in -ide are named by changing the -ide to -ic adding the prefix hydro- to this anion name , and then following with the word acid. HCl- hydrochloric acid Acids containing anions whose names end in-ate or -ite are named by changing -ate to -ic and -ite to - ous and then adding the word acid. Perchlorate Perchloric acid Chlorate Chloric acid Chlorite Chlorous acid Hypochlorite Hypochlorous acid Chapter 3 (Chemical Reactions & Reaction Stoichiometry) How to balance a chemical equation and why it is necessary Ensure both sides have equal amounts of atoms. ONLY CHANGE COEFFICIENTS General Guidelines for balancing equations 1. Write correct formulas for reactants and products 2. Can only adjust coefficients, not subscripts 3. Start with more complex molecules 4. Balance polyatomic ions as a single unit 5. Check to verify coefficients Be able to identify the 3 types of reactions discussed in 3.2: Combination, Decomposition, and Combustion CHM 113 Lecture Page 4 Combination: two molecules combine to form one Decomposition: One molecule reacts to form two Combustion: a hydrocarbon reacts with water to form carbon dioxide and water The meaning of the mole concept and why it is used A mole is a unit of measurement, it is used for very large or very small numbers. 1 mole= 6.02*10^23 Using the mole and molar masses in calculations How to calculate the % composition(% by mass) of an element in a compound and use % composition data to determine mole/numbersof atoms. Calculate empirical formula form %composition data Assume 100g of compound 30.45% Nitrogen and 69.56% O -30.45g N -60.56g O Convert to moles Divide each by the lowest number of mol (2.173 mol) 2.173 mol N/2.173 = 1 mol 4.348 mol O/ 2.173 = 2 mol The relationship between an empirical and molecular formula Empirical formula is the simplified molecular formula How to determine the molecular formula of a compound from the empirical formula and the molecular mass Use molecular weight to determine how much of each atom is in the molecular formula. Stoichiometry calculations Review conversions The meaning and significance of a limiting reactant; be able to identify the limiting reactant. Limiting reactant is the molecule that runs out first in the reactants, limiting the amount of product produced. Easily calculate limiting reactant by calculating moles of each reactant and divide by the balanced equation coefficient. The smaller number is the limiting reactant. How to solve stoichiometry problems with limiting and excess reactants Be able to convert between moles and grams. The difference between theoretical and actual yields Theoretical Yields Theoretical yield is what's calculated from the limiting reactant. Actual yield is what's produced when the reaction is run. How to calculate theoretical and % yields Limiting reactant-> moles of limiting reactant-> moles of product using mole ratio -> theoretical yield of product Percent yield= Actual/Theoretical x 100% CHM 113 Lecture Page 5
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