Chemistry chapters 3 and 4
Chemistry chapters 3 and 4 CHM 160 001
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This 7 page Bundle was uploaded by Devin Mart on Tuesday March 29, 2016. The Bundle belongs to CHM 160 001 at Missouri State University taught by Dr. Richter in Spring 2016. Since its upload, it has received 8 views. For similar materials see General Chemistry 1 in Chemistry at Missouri State University.
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Date Created: 03/29/16
Chapter 3: Molecules, Compounds, and Chemical Equations ● It takes compounds, in all of their diversity, to make life possible. ○ When a balloon filled with H₂ and O₂ is ignited, the two elements react violently to form H₂O. ● The properties of compounds are generally very different from the properties of the elements that compose them. ○ Compounds are composed of elements held together by chemical bonds. ● Ionic bonds formed between metals and nonmetals involve the transfer of electrons from one atom to another. ○ Metals have a tendency to lose electrons, nonmetals have a tendency to gain electrons. ○ The strength of the ionic bond depends on the magnitude of the charges and the distance between the ions. ○ The ionic bond is an electrostatic interaction and does not involve electron sharing. ○ Ionic compounds exist as huge, 3dimensional networks of ions. ● Covalent bonds involve the sharing of electrons between two nonmetals. ○ The electron is shared by both nuclei (both protons) and is therefore more stable. ○ The atoms within all molecular compounds are held together by covalent bonds. ● The quickest and easiest way to represent a compound is with itchemical formula. ○ The more metallic (or more positive) element is usually listed first. ○ An empirical formula simply gives the relative number of atoms in each compound. ■ Ex. Empirical formula HO (Actual formula H₂ O₂ ) ○ A molecular formula gives the actual number of atoms of each element in a molecule of a compound. ■ Ex. Molecular formulas H₂ O₂ ○ A structural formul shows how atoms in a molecule are connected to one another. ■ Ex. HOOH ● Elements may be either atomic or molecular and compounds may be either molecular or ionic. ○ Polyatomic molecules are ions that are composed of two or more atoms covalently bonded to one another. ● Ionic compounds occur throughout the earth’s crust as minerals and are even in the food we eat. ● The formulas of most ionic compounds can be deduced from their constituent elements. ○ Ex. Sodium Chloride ■ Na → Na⁺ + e ■ Cl + e⁻ → Cl⁻ ● Ionic compounds are composed of at least one cation and at least one anion. ● Names are used to identify compounds. ○ Historically names were often given before the formulas for the chemical compounds were known. ○ In ionic compounds name the positive ion first, and then the negative ion. ■ If the cation is a main group element, do not change its name. ● Sodium Na⁺ , Zinc Zn² ■ If the cation is a transition metal, the charge must be listed ● Iron (II) Fe², Iron (III) e³⁺ ■ For simple anions, the ending is changed to ide. ● Fluoride F⁻ , Oxide O ⁻ ○ For polyatomic ions, the name of the polyatomic ion is used whenever it occurs. ■ The name of polyatomic ions usually ends in ate or ite. ■ For polyatomic ions that can have different numbers of oxygens, the one with the higher number of oxygens gets the ate. ● Nitrate NO₃ ⁻, Nitrite NO⁻ ■ If more than two ions exist, prefixes are required. ● Per = highest # of oxygen ● hypo = lowest # of oxygen ● Hydrates contain a specific number of water molecules associated with each formula unit. ● The first step in naming a molecular (or covalent) compound is identifying it as one. ○ Main group or Transition metals. ○ The number of atoms must also be listed, and the ending is changed to ide. ■ sulfur etrafluoide SF₄ ○ Table of prefixes for naming covalent compounds. ■ 1 → Mon ■ 2 → Di ■ 3 → Tri ■ 4 → Tetra ■ 5 → Penta ■ 6 → Hexa ■ 7 → Hepta ■ 8 → Octa ■ 9 → Nona ■ 10 → Deca ● Simple acids are named by adding the prefix ‘hydro’, and the suffix ‘ic’ plus acid. ○ Hydrochloric acid HCl ○ Oxyacids contain hydrogen and an oxyanion (a nonmetal and oxygen). ■ Nitric acid HNO₃ ○ If the oxyanion ends in ‘ite’, the ending is changed to the ‘ous’. ■ Hydrogen nitrite → Nitrous acid ● In chemical reactions, matter is neither created nor destroyed. ○ A chemical reaction is composed of reactants and products. ■ Reactants on the left, Products on the right. ○ A reaction written in paper is a prediction and does not have to happen. ■ 2H₂ + O₂ → 2H₂ O ● The Law of Conservation of mass states that matter can neither be created nor destroyed, but can be rearranged. ○ In a balanced chemical reaction, the law of conservation of matter is observed. ■ A balanced chemical equation obeys the law of conservation of matter. ○ Number of atoms on the left equals the number of atoms on the right and you can’t change their identity. ■ 2Al + Fe₂ O₃ → Al₂ O₃ + 2Fe ● Balancing Equations. ○ If an element is present in just one compound on each side, balance it first. ○ Balance anything that exists as a free element last. ○ Balance polyatomic ions as a unit. ○ Check when done same number of atoms, and same total charge on both sides. ● Always conserved. ○ Identity of atoms in reactants = Identity of atoms in products. ○ Number of atoms in reactants = Number of atoms in products. ○ Mass of all reactants = Mass of all products. ● May change. ○ Number of molecules in reactants vs. Number of molecules in products. ○ Physical states (solid, liquid, or gas) of reactants vs. physical state of products. ● Historically, organic compounds were derived from living material. ○ Organic compounds primarily contain carbon and hydrogen. ○ Compounds containing only carbon and hydrogen are called hydrocarbons. ■ Functionalized hydrocarbons are hydrocarbons containing a ‘functional group’. ● Historically, inorganic compounds were obtained from nonliving matter. Chapter 4: Chemical Quantities and Aqueous Reactions ● The Greenhouse Effect is the warming of the planet when heat is trapped by gases in the atmosphere and returned to the planet. ○ Sunlight passes through atmosphere and warms Earth’s surface. ○ Some of the heat radiated from the Earth’s surface is trapped by greenhouse gases. ■ Common examples of the greenhouse effect include ‘greenhouses’ and automobiles. ○ Atmospheric gases trap and return a major portion of the heat radiating from the Earth. ■ Average temperature without the greenhouse effect would be 0℉ . ● Global warming is the popular term used to describe the increase in average global temperatures. ○ In the last 120 years the average temperature has increased between 0.4 and 0.8℃. ● Experimental evidence implicates carbon dioxide from humanrelated sources as the cause of recent global warming. ○ The vast majority of scientists agree that global warming is occurring and that humans are contributing to it. ■ Global warming skeptics disagree with the overwhelming consensus, and state that other factors are responsible for global warming. ● Stoichiometry involves taking a balanced chemical equation and using it to make predictions. ○ The coefficients in a chemical reaction specify the relative amounts in moles of each of the substances involved in the reaction. ● The amount of product that can be made in a chemical reaction is limited by the amount of reactants one starts with. ○ The limiting reactant (or reagent) is the reactant that makes the least amount of product. ● The general strategy for solving limiting reactant problems is: ○ Write down the reaction from the word problem. ○ Balance the equation. ■ If you assume the smaller mass reacting is the limiting reagent you’re wrong. ○ Calculate the yield of one product from both reactants. ● Solutions contain two components, a solute and a solvent. ○ An aqueous solution contains water as the solvent. ■ The human body is an aqueous solution. ○ Water is called the “universal solvent”. ○ The amount of solute in a solution is variable. ● Solution concentration is usually measured in units of Molarity. ○ Molarity = amount of solute (moles) / volume of solution (L) ● Partspermillion (ppm) is the relative amounts of the particles of various compounds in the mixture. ● A dilute solution contains less solute than a concentration solution. ○ The dilution equation simply says that the number of moles that you take from the concentrated solution, equals the number of moles in the dilute solution. ■ M₁ V₁ = M₂ V₂ ● Ionic bonds formed between metals and nonmetals involve the transfer of electrons from one atom to another. ● Covalent bonds involve the sharing of electrons between two nonmetals. ○ Salts are ionic compounds, water is a covalent/molecular compound. ● The ability of an atom to attract electrons to itself in a chemical bond is called electronegativit. ● A polar bond is intermediate between a pure covalent bond and an ionic bond. ○ Water is a polar covalent molecule. ● When a salt is put into water, the cations are attracted to the negative end of water, and the anions to the positive end. ○ The attraction between water molecules and the ions of sodium chloride allows NaCl to dissolve in the water. ■ Partial charges on sugar molecules and water molecules result in attractions between them. ○ Ionic compounds dissociate into ions when they dissolve, molecular compounds do not. ● Substances that will dissociate in solution to form ions are called electrolytes. ○ When ions (charged particles) are in aqueous solutions, the solutions are able to conduct electricity. ■ Sodium chloride in water is an electrolyte solution it conducts electricity. ■ Sugar in water is a nonelectrolyte solution it does not conduct electricity. ○ Electrolyte drinks have ions in them. ○ Electrolyte tests are routinely used by medical personnel to assess health. ■ The human body is (in part) an electrolyte solution. ● Acids are molecular/covalent compounds that ionize when they dissolve in water. ○ Weak acids do not completely dissociate in water. ● A compound is soluble if it dissolves in water and insoluble if it does not. ○ Silver nitrate is very soluble in water and is a strong electrolyte. ● Solubility is very important for living systems and the environment. ● In a precipitation reaction, a precipitate forms upon mixing two solutions. ○ 2KI (aq) + Pb(NO₃ )₂ (aq) → PbI₂ (s) + 2KNO₃ (aq) ○ Precipitation reactions do not always occur when aqueous solutions are mixed. ■ KI (aq) + NaCl (aq) → No Reaction ● Molecular equations show the complete neutral formulas for each compound in the reaction as if they existed as molecules. ○ Na₂ CO₃ (aq) + 2HCl (aq) → 2NaCl (aq) + (g) +HO (l) ● Complete ionic equations list individually all of the ions present in a chemical reaction. ○ 2Na⁺ (aq) + CO² (aq) + 2H (aq) 2Cl⁻ (aq) →2Na⁺ (aq) 2Cl⁻ (aq) + CO (g) + HO (l) ○ Spectator ions appear unchanged on both sides of the reaction and do not participate in the reaction. ● A net ionic equation shows only the species that change during a chemical reaction. ○ CO₃ ²⁻ (aq) + 2 (aq) → CO₂ (g) + HO (l) ● Acidbase (neutralization) reactions occur when an acid reacts with a base to form water. ○ HCl + NaOH → H₂ O + NaCl ○ An acid produces H⁺ in aqueous solution and a base produc in aqueous solution. ■ Both H⁺ and H₃O⁺ are used interchangeably to designate an acid. ■ Common acids include those found in lemons, limes, vinegar, Vitamin C, aspirin, the human stomach, etc. ■ Bases are also found in many household products including cleaners and antacids. ○ An acidbase neutralization forms a salt plus water. ● Oxidationreduction (redox) reactions involve the transfer of electrons from one reactant to the other. ○ 2Mg + O₂ → 2MgO ○ Step 1: Assign the oxidation numbers for each atom in the reaction. ■ Oxidation numbers are the charge an atom would have if the compound were ionic. ■ Oxidation number is also known as the oxidation state of an atom or ion. ■ It’s often necessary to look at the other species in a compound to determine an atom’s oxidation number. ■ Neutral elements have an oxidation number of zero. ■ The more electronegative element is assigned the negative charge. ■ Group 1 and 2 elements have +1 and +2 charges. ■ Oxygen usually has a 2 charge. ■ The sum of the oxidation states must equal the charge on the molecule. ■ Elements in the same group usually have the same oxidation number. ■ The largest possible positive oxidation number is equal to the group number. ■ The largest negative oxidation number possible is equal to the group number minus 8. ● Al = 3 8 = 5 ○ Step 2: Determine what is beinoxidized. ■ Oxidation = loss of electrons. ● Mg → Mg²⁺ + 2e ○ Step 3: Determine what is beineduced. ■ Reduction = gain of electrons. ● O₂ + 4e⁻ → 2O²⁻ ■ OIL RIG ○ Combustion reactions are redox reactions.
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