CHEM 109 Week 3
CHEM 109 Week 3 CHEM 109
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This 9 page Study Guide was uploaded by Hannah Rapp on Sunday September 11, 2016. The Study Guide belongs to CHEM 109 at University of Nebraska Lincoln taught by Jason Kautz in Fall 2016. Since its upload, it has received 135 views. For similar materials see General Chemistry 1 in Chemistry at University of Nebraska Lincoln.
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Date Created: 09/11/16
STUDY GUIDE FOR EXAM 1 CHAPTER ONE: CHEMISTRY: THE CENTRAL SCIENCE Key words accuracy Kelvin chemical change law chemical property mass compound matter conversion factor mixture density physical change dimensional analysis physical property element precision extensive property qualitative property heterogenous mixture quantitative property homogenous mixture scientiﬁc method hypothesis signiﬁcant ﬁgures intensive property SI unit International System of Units substance theory Key Equations • K = °C + 273.15 • °C = (°F - 32°F) x (5/9) • °F = (9/5) x °C + 32°F • D = (m/v) Section One Summary • Chemistry is the study of matter and the changes matter undergoes • Scientiﬁc Method observations give rise to laws, data gives rise to hypotheses, hypotheses are tested with experiments, and successful hypotheses give rise to scientiﬁc theories Section Two Summary • all matter exists as a substance or a mixture • Substances can be elements or compounds • Mixtures can be homogenous (a solution) or heterogenous (not uniform throughout) • Mixtures can be separated through physical means • Compounds can only be separated using chemical processes • Elements cannot be separated into simpler substances Section Three Summary • the International System of Units (SI units) were developed as a way for scientists around the world to label units and be able to communicate. There are seven base units, such as kilogram for mass and Kelvin for temperature. Density • and volume are derived from these base units. Section Four Summary • Quantitative properties involve numbers while qualitative properties describe the quality, or smell, taste, or look of something. Quantitative: 22 grams • • Qualitative: white, powdery substance • Physical properties can be determined without a chemical change occurring. A Physical change is one that the identity of the matter involved is not changed. • Physical property: boiling point, physical state • Physical change: sublimation, melting, freezing • Chemical properties are determined as a result of a chemical change, where the original substance is altered sue to a chemical reaction. • Chemical change: combustion, baking (baking soda produces carbon dioxide) • Extensive properties are dependent on the amount of substance (mass, volume) while intensive properties are independent of amount (temperature, density) • way to remember: Intensive - In dependent Section Five Summary • measured numbers are inexact. Numbers obtained by counting are exact numbers • Signiﬁcant ﬁgures specify uncertainty a measure number. They must be carried through calculations such that the uncertainty in the ﬁnal answer is reasonable • 450 has 2 sig ﬁgs 45.0 has 3 sig ﬁgs. • • .0004500 has 4 sig ﬁgs • Accuracy deﬁnes how close measured number are to the true value • Precision refers to how close measured numbers are to one another • Accuracy: true value is 45. Measured value were 45.1, 44.8 and 44.9 • Precision: measured values were 44, 44.1, 39.9 • Measurements can be just accurate, just precise, both, or neither Section Six Summary • A conversion factor is a fraction in which the numerator and denominator are the same quantity expressed in different units. • multiplying by a conversion factor is a unit conversion Dimensional analysis is a series of unit conversions used a multistep problem. • Chapter One Example Problems 1. The density of gold in 19.3 g/cm . Find the volume (in cm ) of a gold nugget with a mass of 3.78. 2. Convert 300 seconds to hours. 3. Give an example of a homogeneous mixture and a heterogenous mixture. 4. Give the name of the following elements: C, Li, F, Cu, Mg, As, Cl, N, Ca, K, and Br. 5. Convert 700 km/hour to m/s. 6. Classify the following substances as an element or compound. (a) hydrogen (b) salt (c) gold (d) water 7. What is the difference between mass and weight? 8. The density of Platinum (Pt) is 21.5 g/cm at 25°C. What is the volume of 78.6 g of Pt at this temperature? 9. Determine the density of the following object in g/cm . A cube with edges of 0.750 m and a mass of 7.56 kg. 10. A student pours 44.3 g of water at 10°C into a beaker containing 113.9 g of water at 10°C. What is the ﬁnal mass, temperature, and density of the combined water? The density of water at 10°C is 1.00 g/mL. 11. Express the following numbers in scientiﬁc notation. (a) 0.00000987 (b) 323 (c) 63,018 (d) 0.045 12. Determine the number of signiﬁcant ﬁgures in the following measurements. (a) 4537 miles (b) 23 mL (c) 70,012 tons (d) 1300 g (e) 0.000000007 cm (f) 0.9 min (g) 30.7 g/mL (h) 4.7 x 1023 atoms 13. Carry out the following operations and express the answer with the correct number of sig ﬁgs. (a) 7310 km/ 5.70 km (b) 3.70 mg - 2.9133 mg (c) 4.51 cm x 3.5555 cm 14. Three students are asked to determine the volume of a certain substance. The results are as follow (A) 87.1, 88.2, 87.6 (B) 86.9, 87.1, 87.2 (C) 87.6, 87.8, 87.9 The true value is 87.0 mL. Discuss the precision and accuracy of each student. CHAPTER TWO: ATOMS, MOLECULES, AND IONS acid Beta (b) ray inorganic compound nucleus Alkali Metal binary compound ion organic compounds Alkaline Earth Metal cation ionic compound oxoacid alkane Chalogens isotope oxoanion allotrope chemical formula Law of Conservation of period Mass alpha particle diatomic molecule Law of Deﬁnite periodic table Proportions Alpha (a) ray electron Law of Multiple polyatomic ion Proportions anion empirical formula mass number polyatomic molecule atom family metal proton atomic ion functional group metalloid radiation atomic mass Gamma (y) rays molecular formula radioactivity atomic mass unit group molecule structural formula atomic number Halogens monoatomic ion transition elements atomic weight hydrate neutron transition metals beta particle hydrocarbon Noble Gases Section One Summary • Dalton’s atomic theory state that all matter is made up of tiny indivisible particles called atoms. Compounds form when atoms of different elements combine in ﬁxed ratios. • According to the Law of Deﬁnite Proportions, any sample in a compound will always contain the same elements in the same mass ratio. • The Law of Multiple Proportions states that if two elements can form more than one compound, the mass ratio of one will be related to the mass ratio of the other by a small, whole number. • The Law of Conservation of Mass state that matter can neither be created no destroyed, but conserved in a chemical reaction. Section Two Summary The atom is the basic unit of an element. Studies with radiation have indicated that atoms • contain subatomic particles, one of which is an electron. • Experiments with radioactivity have shown that some atoms give off different types of radiation. Alpha (a) rays are composed of alpha particles which are actually helium nuclei. • • Beta (b) rays are composed of beta particles which are actually electrons. • Gamma (y) rays are high-energy radiation Most of the mass of an atom is contained in a tiny, dense nucleus. The nucleus contains • positively charged protons and electrically neutral particles called neutrons. The electron occupies a large volume around the nucleus known as the ‘electron cloud.’A neutron is slightly larger than a proton, but each is almost 2000 times larger than an electron Section Three Summary • The atomic number is the number of protons. It determines the identity of the atom. • The mass number is the sum of the protons and neutrons in the nucleus. • Atoms with the same atomic number but different mass numbers are known as isotopes. Section Four Summary • The periodic table arranges the elements in rows (periods) and columns (groups or families). Elements in the same group exhibit similar properties • All elements are either a metal, metalloid, or nonmetal. Section Five Summary • Atomic Mass is the mass of an atom in atomic mass units (amu). • One amu is exactly one-twelfth the mass of a Carbon-12 atom. • The periodic table contains the average atomic mass (sometimes referred to as the atomic weight) of each element. Section Six Summary • An ion is at atom or group of atoms with a net charge. An atomic ion or monoatomic ion consists of just one atom. • A positively charged ion is a cation. It becomes positive by losing an electron. • Na + • A negat-vely charged ion is an anion. It becomes negative by gaining an electron. • Cl • Polyatomic ions are those that contain more than one atom chemically bonded together. • The formulas of ionic compounds are empirical formulas. • Ionic compounds’ names are written as the name of the cation followed by the name of the anion. • Naming polyatomic ions: 1st: name of the element. 2nd: add -ide to the end • KBr is Potassium Bromide • Oxyanions are polyatomic ions that contain one or more oxygen atoms. • Hydrates are compounds whose formulas include a speciﬁc amount of water molecules. Section Seven Summary • A molecule is an electrically neutral group of two or more atoms. Molecules with 2 atoms are diatomic. • N 2s diatomic Nitrogen. • A chemical formula denotes the composition of a substance. A molecular formula speciﬁes the exact number of atoms in a molecule of a compound. • Molecular compounds are names according to a set of rules, including the use of Greek preﬁxes to specify the number of each kind of atom in a molecule. • Binary compounds are those that consist of two elements. • An acid is a substance that generate Hydrogen ions when dissolved in water. Oxyacids are acids based on oxyanions. • • Inorganic compounds generally don't contain Carbon. Organic compounds contain Carbon and Hydrogen, sometimes with other elements. • Hydrocarbons contain ONLY carbon and hydrogen. • The simplest hydrocarbons are alkanes. • A functional group is a group of atoms that determine the chemical properties of an organic compound. Chapter Two Example Problems 1. What is the name for CaSO ? 4 2. What is the correct formula for Nickel (II) perchlorate? 3. State and deﬁne the laws of deﬁnite and multiple proportions. Give an example of each. 4. The elements sulfur and oxygen can form a variety of different compounds. The two most common are SO and SO . Samples of these two compounds decomposed into their 2 3 constituent elements. One produced 1.002 g S for every gram of O and the other produced 0.668 g S for every gram of O. Show that these are consistent with the Law of Multiple Proportions. 5. Calculate the number of neutrons in 239Pu. 6. Determine the mass number of (a) a B atom with 6 neutrons (b) a Mg atom with 13 neutrons (c) a Br atoms with 44 neutrons (d) a Hg atom with 119 neutrons 7. State two difference between a metal and a nonmetal 8. The atomic masses of 203Tl and 205Tl are 202.972320 and 204.974401 amu respectively. Calculate the natural abundance of these two isotopes. the average atomic mass of Thallium is 204.4 amu. Give the number of protons and electrons in each the following ions: (a) Na (b) Ca + 2+ 9. (c) Al+ (d) I (e) F (f) S 2- (g) O 2- 10. Write the formulas for the following ionic compounds (a) sodium chloride (b) iron sulﬁde 2+ (containing the Fe ion) (c) barium ﬂuoride 11. Which of the following are ionic? molecular? SiCl 4, NaBr, BaF , 2Cl , IC4, CsCl, NF 3 12. Name the following compounds: (a) KH PO (b)2HBr 4gas) (c) Li CO (d) PF (2) KC3O 5 (f) Al(OH)3(g) Fe O2(h3 Na O (i)2Ag CO 2 3 13. Write the formula for the following compounds: (a) potassium sulﬁde (b) magnesium phosphate (c) lead (II) carbonate (d) silver perchlorate (e) ammonium sulfate (f) boron trichloride (g) hyrdoiodic acid (h) potassium dihyrdogen phosphate (i) mercury(II) oxide 14. What is the difference between an atom and a molecule? Write empirical formulas for the following: (a) C N (b) P O (c) B H 15. 2 2 4 10 2 6 CHAPTER THREE: STOICHIOMETRY: RATIOS OF COMBUSTION actual yield molar mass aqueous (aq) mole Avogadro’s number molecular mass chemical equation molecular weight combination reaction Percent Composition by Mass combustion percent yield combustion analysis product excess reagent reactant formula mass stoichiometric amount formula weight stoichiometric coefﬁcients limiting reagent theoretical yield Key Equations • percent by mass of an element = n x atomic mass of element x 100 molecular/ formula mass • percent yield = actual yield x 100 theoretical yield Section One Summary • Molecular mass is calculated by summing the masses of all atoms in a molecule. • Molecular weight is another term for molecular mass. • For ionic compounds, we use the analogous terms formula mass and formula weight • All of these are expressed in amu Section Two Summary • Molecular or formula mass can be used to determine percent composition by mass of a compound. Section Three Summary • A chemical equation is a written representation of a chemical reaction. Chemical species on the left side are reactants. Chemical species on the right side are products. • The physical state of each reactant and product is speciﬁed in parentheses such as (s), (l), (g), and (aq). • Chemical reactions are balanced by stoichiometric coefﬁcients. • Combustion refers to the chemical combination with oxygen. Combustion of hydrocarbons produce carbon dioxide and water. Section Four Summary • a mole is the amount of a substance that contains 6.022 x 10 (Avogadro’s Numbe r) of particles (atoms, molecules, ions, formula units) • Molar Mass is the mass of one mole of a substance, usually expressed in grams. the molar mass of a substance in grams is equal to the atomic, molecular, or formula mass of the substance in amu. • Molar mass and Avogadro’s Number can be used to interconvert among mass, moles, and # of particles. Section Five Summary • Combustion Analysis is used the determine the empirical formula of a compound. The empirical formula can be used to calculate percent composition. • The empirical formula and molar mass can be used to determine the molecular formula. Section Six Summary • A balanced chemical equation can be used to determine how much product will form from given amounts of reactants, how much of one reactant is necessary to react with a given amount of another, or how much reactant is required to produced a speciﬁed amount of product. • Reactants that are combined in exactly the ratio speciﬁed by the balanced equational to to be “combined in stoichiometric amounts.” Section Seven Summary • The limiting reagent is the reactant that is entirely consumed in the reaction. An excess reagent is the reactant this is not completely consumed. The maximum amount of • product is dependent on the limiting reagent. • The theoretical yield of a reaction is the amount of product that will form if all the limiting reagent is consumed by the reaction. • The actual yield is the amount of product actually recovered from the reaction. • Percent yield is a measure of the efﬁciency of the chemical reaction. • Combustion ( a substance is burned in the presence of oxygen), combination (two or more reactant for a single product), and decomposition ( a single reactant forms two products) are three types of chemical reactions that are commonly encountered in chemistry. Chapter Three Example Problems Calcium phosphide Ca P an3 w2ter react to form calcium hydroxide and phosphine (PH ). In a 3 particular experiment, 225.0 g of Ca P 3nd2125.0 g of water are combined. Ca P (s) + 6H O (l) 3Ca(OH) (aq) + 2PH (g) 3 2 2 2 3 1. How much PH can 3e produced? 2. How much Ca(OH) can b2 produced? 3. How much of the excess reactant remains when the reaction is complete? 4. Calculate the molecular mass of the following: (a) C H 6 12 (b) CH C3 (c) SO (d2 BCl 3 (e) C 6 12(f6 H PO 3 4 5. Calculate the percent composition of Chloroform (CHCl ). 3 6. Balance the following equations C + O CO 2 O 3 O 2 Zn + AgCl ZnCl2+ Al NaOH + H SO2 4 Na 2O + 4 O 2 P 4 10H O 2 H 3O 4 H S + HNO 3 2SO 4 NO + 2 O 2 CO 2 KOH K2SO 3 H O 2 7. What is the molar mass of an atom? What are the commonly used units for molar mass? 8. How many atoms are there in 5.10 moles of Sulfur? 9. How many grams of gold (Au) are there in 15.3 moles of Au? 10. Calculate the number of C, H, and O atoms in 1.50 f of glucose (C H 6 12 6 11. Consider the combustion of Carbon monoxide (CO) in oxygen gas 2CO (g) + O (g) 2CO 2 2 Starting with 3.60 moles of CO, calculate the number of moles of CO pro2uced if there is enough oxygen gas to react with all the CO. 12. Ammonia is a principle nitrogen fertilizer. IT is prepared by the reaction between hydrogen and nitrogen: H N 2 3H 2 N 2 3 In a particular reaction, 6.0 moles 2NH were produced. How many moles of H and N were 3 2 2 consumed to produce this? 13. Nitric Oxide (NO) reacts with oxygen gas to form NO 2 2NO (g) + O (2) 2NO 2g) In an experiment, 0.886 moles of NO mixed with 0.503 moles of O Dete2.ine which was the limiting reagent. Calculate the number of moles of NO p2oduced.
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