Exam #3 Study Guide
Exam #3 Study Guide Chem 1010
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This 6 page Study Guide was uploaded by Courtney Burke on Wednesday November 11, 2015. The Study Guide belongs to Chem 1010 at University of Denver taught by Teresa Cowger in Fall 2015. Since its upload, it has received 142 views. For similar materials see General Chemistry 1010 in Science at University of Denver.
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Date Created: 11/11/15
Chapter Exam #3 Study Guide General Chemistry Chapter 4 Acidbase reactions (where an acid reacts with a base) involve water as a reactant or product. + An acid is a substance that produces H ions when dissolved in water. HX H (aq) + X (aq) A base is a substance that produces OH ions when dissolved in water. + MOH M (aq) + OH (aq) Acidic solutions arise when certain covalent Hcontaining molecules dissociate into ions in water. Strong acids AND strong bases dissociate completely into ions when placed in water. Weak acids AND weak bases dissociate very little into ions. Strong acids, such as HNO and 3 SO , a2d w4ak acids, such as HF and H PO , 3 4 have one or more H atoms as part of their structure. 2 Strong bases have either OH or O as part of their structure. Weak bases, such as ammonia (NH ) do not3contain OH ions, but they all have an electron pair on N. The key event in aqueous reactions between a strong acid and a strong base is that + an H ion from the acid and an OH ion from the base form a water molecule. Acidbase reactions occur through the electrostatic attraction of ions and their removal from solution as the product. The ionic compound that results from the reaction of an acid and base is called a salt. Acidbase reactions are metathesis (doubledisplacement) reactions. An acid is a molecule (or ion) that donates a proton. A base is a molecule (or ion) that accepts a proton. + H 3 ion acts as the acid and donates a proton to OH ion, which acts as the base and accepts it. Titration: The known concentration of one solution is used to determine the unknown concentration of another. + Equivalence point: Occurs when the amount (mol) of H ions in the original volume of acid has reacted with the same amount (mol) of OH ions from the buret. End point: Occurs when a tiny excess of OH ions changes the indicator permanently to its basic color. The amount of base needed to reach the end point is the same as the amount needed to reach the equivalence point. Oxidationreduction (redox) reaction: Net movement of electrons from one reactant to another. Ionic compounds: transfer of electrons. Covalent compounds: shift (sharing) of electrons. Oxidation is the loss of electrons. Reduction is the gain of electrons. Example: Formation of MgO Oxidation (electron loss by Mg): Mg Mg + 2e – – 2– Reduction (electron gain by O )2 ½ O +22e O The oxidizing agent is the species doing the oxidizing (causing electron loss). The reducing agent is the species doing the reducing (causing electron gain). In MgO, O 2xidizes Mg by taking electrons that Mg gives up. Mg reduces O 2y providing the electrons that O t2kes, so Mg is the reducing agent, and O 2s the oxidizing agent. The oxidizing agent is reduced, the reducing agent is oxidized. Oxidation numbers: 1. For Group 1(A): O.N. = +1 in all compounds 2. For Group 2(A): O.N. = +2 in all compounds 3. For hydrogen: O.N. = +1 in combination with nonmetals = –1 in combination with metals and boron 4. For fluorine: O.N. = –1 in all compounds 5. For oxygen: O.N. = –1 in all peroxides = –2 in all other compounds (except with F). 6. For Group 7(A): O.N. = –1 in combination with metals, nonmetals (except O), and other halogens lower in the group Transferred electrons are never free because the reducing agent loses electrons and the oxidizing agent gains them simultaneously. Atoms occur as an element on one side of an equation and as part of a compound on the other. In a combination reaction, two or more reactants form a compound. X + Y Z Metal and nonmetal form an ionic compound. The metal is the reducing agent and the nonmetal is the oxidizing agent. Two nonmetals form a covalent compound. Nearly every nonmetal reacts with O t2 form a covalent oxide. In a decomposition redox reaction, a compound forms two or more products, at least one of which is an element. Z X + Y In doubledisplacement (metathesis) reactions, atoms of two compounds exchange places. AB + CD AD + CB In solution, singledisplacement reactions occur when an atom of one element displaces the ion of another. The most reactive metals displace H f2om liquid water. Group 1(A) metals and Ca, Sr, and Ba from Group 2(A) displace H fro2 water. Chapter 6 Thermodynamics: The study of energy and its transformations. Thermochemistry: A branch of thermodynamics that deals with heat in chemical and physical change. System: The part of the universe we are focusing on (in a reaction). Surroundings: Everything else (in a reaction). Internal Energy (E): Sum of all potential and kinetic energy in a system. ∆E is the difference between internal energy after the change (E finaland before the change (E initial Final state minus the initial state: ∆E = E finalE initial productsEreactants A change in the energy of a system must be accompanied by an equal and opposite change in the energy of the surroundings. By releasing some energy in a transfer to the surroundings: E finalE initial∆E < 0 By absorbing some energy in a transfer from the surroundings: E final Einitial ∆E > 0 Heat: Thermal energy (symbolized by q) is the energy transferred as a result of a difference in temperature. Work: (symbolized by w) the energy transferred when an object is moved by a force. The total change in a system’s internal energy is the sum of the energy transferred as heat and/or work: ∆E = q + w Energy transferred to the system is positive because the system ends up with more energy. Energy transferred from the system is negative because the system ends up with less energy. Heat flowing out of a system: heat is released so q is negative and ∆E is negative. Heat flowing into a system: heat is absorbed so q is positive and ∆E is positive. Law of Conservation of Energy – First Law of Thermodynamics: The total energy of the universe is constant. ∆E = ∆E – ∆E = 0 universe system surroundings Joule (J): The SI unit of energy 1 J = 1 kg • m /s Calorie (cal): Quantity of energy needed to raise the temperature of 1 g of water by 1 ˚C. 1 cal = 4.184 J or 1 J = 1/4.184 cal = 0.2390 cal British thermal unit: Quantity of energy required to raise the temperature of 1 lb of water by 1 ˚F. 1 BTU = 1055 J ∆E does not depend on how the change takes place, but only on the difference between the final and initial states. Pressurevolume work (PV work): The mechanical work done when the volume of the system changes in the presence of an external pressure (P). w = –P∆V At constant pressure, enthalpy (H) is defined as the internal energy plus the product of the pressure and volume. H = E + PV Change in enthalpy (∆H): The change in internal energy plus the product of the pressure, which is constant, and the change in volume (∆V). ∆H = ∆E + P∆V Exothermic and endothermic process: ∆H = H final HinitialH products reactants Exothermic: Releases heat and results in a decrease in the enthalpy of a system: H productsH reactants∆H < 0 Endothermic: Absorbs hear and results in an increase in the enthalpy of a system: H productsH reactants∆H > 0 q/∆T = constant Heat capacity: [refer to above equation] the quantity of heat required to change its temperature by 1 K. Specific heat capacity: The quantity of heat required to change the temperature of 1 gram of a substance or material by 1 K. Specific heat capacity (c) = q/mass x ∆T Molar heat capacity: The quantity of heat required to change the temperature of 1 mole of a substance by 1 K. Molar heat capacity (C) = q/amount (mol) x ∆T Calorimeter: A device used to measure the heat released (or absorbed) by a physical or chemical process. Finding the specific heat capacity of a solid: csolid H2O x massH2O x ∆T H2O/mass solid ∆T solid A thermochemical equation is a balanced equation that includes the enthalpy change of the reaction (∆H). Hess’s Law: The enthalpy change if an overall process is the sum of the enthalpy changes of its individual steps: ∆H overall∆H 1 ∆H …2+ ∆H n Standard states: o For a gas, the standard state is 1 atm* and ideal behavior o For a substance in an aqueous solution, the standard state is 1 M concentration. o For a pure substance, the standard state is usually the most stable form of the substance at 1 atm and the temperature of interest (the temperature is usually 25 ˚C (298 K). Chapter 18 + Hydronium ion (H O ) 3orms H bonds to several other water molecules. Neutralization: Occurs when an acid and a base react to form H O. 2 + The stronger the acid, the higher [H O 3 is at equilibrium, and the larger the value of K . a Strong acids: HCl, HBr, and HI. 1. Oxoacids in which the number of O atoms equals or exceeds by one the number of ionizable protons, such as HClO, HNO , and H 2O 3 4 2. Carboxylic acids; such as CH COOH 3nd C H COOH. 6 5 Strong bases 1. M O 2r MOH 2. MO or M(OH) 2 Weak bases 1. Ammonia (NH ) 3 2. Amines such as CH CH NH 3 2 2 Water dissociates very little into ions: this process is called autoionization. + – Higher [H O3] lower [OH ] and vice versa. In an acidic solution: [H O3] > [OH ] – + – In a basic solution: [H O3] < [OH ] In a neutral solution: [H O3] = [OH ] – + pH = –log [H O 3 The higher the pH, the lower the [H O ].3An acidic solution has a lower pH (higher [H O3]) than a basic solution. pH of an acidic solution < 7.00 pH of a neutral solution = 7.00 pH of a basic solution > 7.00 pOH = –log [OH ] – Chapter 21 The halfreaction method divides the overall redox reaction into oxidation and reduction halfreactions. Example of the halfreaction method using Cr O (aq) + I (aq) Cr (aq) + I (s):+ 2 7 2 Step #1: Divide the reaction into halfreactions. 2– 3+ Cr 2 7 Cr – I I2 Step #2: Balance atoms and charges in each halfreaction. a. Balance atoms other than O and H. 2– 3+ Cr 2 7 2Cr b. Balance O atoms by adding H O molecu2es. Cr O 2Cr + 7H O 2 7 2 c. Balance H atoms by adding H ions. + 14H + Cr O 2 7Cr + 7H O+ 2 d. Balance the charge by adding electrons. – + 2– 3+ 6e + 14H + Cr O 2 27r + 7H O [reducti2n] Chapter 9 The relative strengths of the bonds in reactants and products determine whether heat is released or absorbed in a chemical reaction. Kinetic energy: Molecules’ movements through space and their rotations and vibrations. Potential energy: Phase changes and changes in the attraction between vibrating atoms. A certain quantity of heat is absorbed (∆H˚ > 0) to break the reactant bonds and form separate atoms. A different quantity of heat is then released (∆H˚ < 0) when the atoms form product bonds.
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