General Chemistry II – Flashcards
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Unlock answers| Thermodynamics |
| The science of relationships between heat and other energy. |
| Thermochemistry |
| A branch of thermodynamics that studies specifically chemical reactions. Knowledge can be used practically or to explore science. |
| Energy |
| The potential to move mater. It has many forms including heat, light, chemical, and electrical. It can convert from form to form. |
| Kinetic Energy (Ek) |
| The energy of an object in motion. This depends on the mass and speed of the object. |
| Equation for kinetic energy |
Ek = (mv2)/2
Ek = kinetic energy m = mass v = velocity |
| Joule (J) |
| A unit of energy equal to one kg*m2/s2. Named after James Prescott Joule, a scientist who studied energy. |
| Watt (W) |
| A unit equal to one Joule per second. |
| Kilowatt-hour |
| A unit equal to 3.6 million Joues. This is the unit in which household electrical energy is measured and sold. |
| Calorie (cal) |
| A unit of energy equal to the energy required to heat one gram of water by one degree Celsius. However this can change based on the original temperature of the water. Officially it is equal to 4.184 Joules. |
| Potential Energy (Ep) |
| The energy an object hs based on its physical position. This depends on the object's mass, its height, and the gravitational pull of the planet. |
| Equation for potential energy |
Ep = mgh
Ep = potential energy m = mass g = gravity of the planet. On Earth it is 9.807 m/s2 h = height above a standard level. The standard level is arbitrary; the difference in height is what matters. |
| Internal Energy (U) |
The kinetic and potential energy of the molecules and particles making up a substance. It is made from heat (q) and work (w) added together.
?U = q + w |
| Total Energy (Etot) |
| The sum of the 3 types of energy kinetic potential, and internal. |
| Equation for total energy |
Etot = Ek + Ep + U
Etot = total energy Ek = kinetic energy Ep = potential energy U = internal energy |
| Law of Conservation of Energy |
| A law that states that energy can be converted into different forms, but the total amount of energy in the universe will always stay the same. |
| System |
| The part of the universe that is being studied. |
| Surroundings |
| The universe other than the system. |
| Work (W) |
When energy is transferred into or out of the system and moves an object in the surroundings. When work is being done on the system, w is positive (add energy). WHen work is being done by the system, w is negative (take away energy). |
| Heat (q) |
A form of internal energy. The molecules of a hot gas have higher kinetic energy since they are moving faster. When there is a temperature difference between the system and the surroundings, if the system is not insulated, it is transfered. Heat always flows from where there is a high temperature to where there is a low temperature. When the system absorbs heat q is positive (add energy). When the system releases heat q is negative (take away energy). A system may absorb heat energy from heat or work. |
| Temperature (T) |
| A measure of heat per unit of a substance. |
| Difference equations |
Often the change in an amount is more important than the actual values. What happens in the middle is irrelevent. This general formula is used for U, T, h, and many other things.
?X = Xf - Xi
?X = Difference of X Xf = final value of X Xi = initial value of X |
| State function |
| Any property of a substance that depends only on its current state; its history is irrelevant. |
| The First Law of Thermodynamics |
Energy is always conserved; it cannot be destroyed or created, only transfered between different forms. The change in internal energy is equal to the sum of energy and work. Remember that q and w may be negative.
?U = q + w
?U = change in internal energy q = heat w = work
Using the equation for pressure-volume work, we can replace w with -P?V to get
?U = q -P?V.
If temperature and pressure are fixed, q = ?H, so
?U = ?H - P?V
|
| Heat of reaction |
The heat (q) that a chemical reaction either evolves or absorbs. If the reaction absorbs energy, it is positive. If the reaction evolves energy, it is negative. |
| Exothermic |
| A reaction where q is negative. Heat is evolved. Heat flows from the system to the surroundings. |
| Endothermic |
| A reaction where q is positive. Heat is absorbed. Heat flows from the surroundings to the system. |
| Pressure-volume work (w) |
When the volume of a vessel decreases, pressures increases. Work is done on the system; w is positive. When volume of a vessel increases, pressure decreases. Work is done by the system; w is negative. Pressure-volume depends on outside pressure and the change in volume.
w = -P?V |
| Equation for pressure-volume work |
w = -P?V
w = work P = pressure outside the vessel ?V = change in volume |
| Enthalpy (H) |
A value dependent on internal energy, pressure, and volume. All of these are state functions, so enthalpy is a state function. It is an extensive propery: it depends on the amount of a substance.
In a reaction where temperatures and pressure are fixed (but volume may change), q = ?H q = U + PV |
| Equation for enthalpy |
H = U + PV
H = enthaply U = internal energy P = pressure V = volume |
| Enthalpy of reaction (?Hrxn) |
The change in enthalpy in a chemical reaction at a fixed temperature. It can be calciulated from BEs if all the reactants are gases. ?Hrxn = (energy in broken bonds) - (energy in formed bonds) |
QUESTION CARD
|
ANSWER
19.2 cal |
| Open system |
A system in which both mass and energy may leave and/or enter freely. Example: a beaker. |
| Closed system |
| A system where energy can move in or out, but no mass. Example: a sealed flask. |
| Isolated system |
A system where neither mass nor energy can move in or out. Example: an insulated, closed container. |
| State |
| A set of characteristics set by varables such as pressure, volume, temperture, number of moles, and mass. |
| Intensive |
A type of measurement that is independent on the size of the sample. Example: temperature. |
| Extensive |
A type of measurement that depends on the size of the sample. Example: heat. |
| Combustion reaction |
A reaction where a substance reacts with oxygen. The molecules are made more stable, evolving heat and light energy. An activation energy is required. All combustion reactions are exothermic. All combustion reactions are redox reactions. CO2 and H2O are often produced, but not always. The products always include one or more oxides. |
| Calorimetry |
"Heat measure". The science of measuring quantities of heat generated, consumed, or dissipated by a sample. First difference in temperature is measured, then it is converted into the amount of heat. |
| Heat capacity (C) |
| The energy required to raise a substance by 1 degree Celcius (or 1 degree Kelvin). Measured in J/°. |
| Molar heat capacity (C[with line over top] or c) |
The energy needed to raise 1 mole of a substance 1 degree Celcius (or Kelvin). Measured in J/°mol |
| Specific heat capacity |
| The amount of energy needed to raise one gram of a substance 1 degree Celsius (or Kelvin). Measured in J/°g. |
| Equations for heat |
q = C?T q = nc?T q = ms?T
q = heat C = heat capacity ?T = change in temperature n = number of moles c= molar heat capacity m = mass s = specific heat |
| Coffee cup calorimeter |
A calorimeter made out of a coffee cup. It has constant pressure on it (from the atmosphere), so when working with them q = ?H. |
| Bomb calorimeter |
A sealed calorimeter in a tank of water. The temperature of the water is measured to calculate q. The bomb is at constant volume, so pressure may change. Therefore q ? ?H. |
QUESTION CARD 50.0mL of dilute AgNO3 is added to a solution with OH- ions in a coffee cup calorimeter. AgNO3 precipitates and the temperature of the solution rises from 23.78°C to 25.19°C. Assume that the mixture has the same specific heat of water. Calculate qsurr for a mass of 150.0g. Is the reaction exothermic or endothermic? |
ANSWER qsurr = 885 J The reaction is exothermic. |
| The specific heat of water |
| c = 4.184 J/°g |
QUESTION CARD
|
ANSWER
s = 0.251 J/°g |
QUESTION CARD
When ice at 0°C melts to liquid water, it absorbs 334J for every gram of ice. Suppose the heat needed to melt a 35.0g ice cube is absorbed from 0.210kg of water at 21.0°C in a calorimeter. What is the final water temperature? |
ANSWER 6.6°C. *Note this is above freezing point: beware for all party-throwers. |
| Thermochemical equation |
| A chemical equation (complete with phase labels) where the enthalpy change (?H) is given right afterwards in kJ. The coefficients in the equation are equal to moles of the substances. If you multiply the moles by a factor, you must also multiply the enthalpy change by that factor. If you reverse the direction of the equation, you must also switch the sign of the change in enthalpy. |
| Hess's Law |
If equation A is the sum of equations B and C, the change in enthalpy of equation A is equal to the sum of the changes of enthalpy in equations B and C. Because enthalpy is a state function, it doesn't matter what path you take. It is useful for determining the enthalpy changes of reactions that are difficult to do or difficult to measure. |
QUESTION CARD An aqueous solution of hydrogen carbonate (baking soda in water) reacts at constant pressure with hydrochloric acid to produce sodium chloride, water, and carbon dioxide gas. The reaction absorbs 12.7 kJ of heat for each mole of sodim carbonate. Write the thermochemical equation for this reaction. |
ANSWER
|
QUESTION CARD
Using the following equation, give the thermochemical equation for 1 mole of liquid water dissociating into hydrogen and oxygen gases.
2H2(g) + O2(g) --> 2H2O(l); ?H = -572 kJ |
ANSWER
H2O(l) --> H2(g) + (1/2)O2(g); ?H = 286 kJ |
QUESTION CARD N2(g) + 3H2(g) --> 2NH3(g); ?H = -91.8 kJ How much heat is evolved from 9.07e5 g of ammonia? |
| 2.45e6 kJ |
| Standard Conditions (°) |
The conditions of a reaction when it is said to be standard, e.g. standard molar enthalpy. 1 atm (or 1 bar, or 105 Pa) 1 mole 25°C |
| Enthalpy of Fusion (?H?fus) |
| The change in enthalpy of a 1 mole of a substance as it goes from a solid to a liquid (melts). Always positive. Temperature remains constant as it melts. The opposite is enthalpy of freezing. |
| Enthalpy of Freezing |
| The change of enthalpy of 1 mole of a substance as it goes from liquid to solid. Always negative. Temperature remains constant as it freezes. The opposite is enthalpy of fusion. |
| Enthalpy of Vaporization (?H?vap) |
| The change in enthalpy of 1 mole of a substance as it goes from liquid to gas (evaporates). Always positive. Temperture stays constant as it evaporates. The opposite is enthalpy of condensation. |
| Enthalpy of Condensation |
| The change in enthalpy of 1 mole of a substance as it goes from gas to liquid (condenses). Always negative. Temperature stays constant as it condenses. The opposite is enthalpy of vaporization. |
| Enthalpy of Sublimination (?H?sub) |
| The enthalpy change in 1 mole of a substance as it goes from solid to gas. Always positive. |
| Enthalpy of Combustion (?H?comb) |
| The enthalpy change in a reaction where 1 mole of a substance is burned completely in oxygen. Always negative. |
| Enthalpy of atomization |
| Th enthalpy change of 1 mole of a substance when it is dissociated into atoms (not ions, not necesarily each element's standard state). |
| Standard enthalpy of formation (?H?f) |
The enthalpy change in 1 mole of a substance as it is formed from its elements in their standard states in standard conditions. The value for common compounds may be found in the textbook appendix. We can find the standard enthalpy change of any reaction as long as we have these values for all the products and reactants.
?H? = ?n?H?f (products) - ?n?H?f (reactants)
|
| Standard state/Reference form |
The state an element is most stable in in standard conditions. Each element has one. Most are a single atom of the element. Exceptions are as follows: H2, N2, O2, P4, S8, Cl2, Br2, I2. Bromine and mercury are the only liquids. Hydrogen, helium, nitrogen, oxygen, fluorine, neon, chlorine, argon, krypton, xenon, and radon are the gasses. Everything else is solid. Some elements have allotropes. One of the allotropes is the standard state. |
| Allotrope |
| Two or more substances which contain only one element, but the atoms are arranged in a different way, producing different substances with different physical properties. One of these formations is the standard state of the element. Example: graphite and diamonds are both pure carbon. Graphite is carbon's standard state. |
| Henri Hess |
| The scientist from University of St. Petersburg who wrote Hess's Law. |
QUESTION CARD What is the enthalpy of reaction (?H) for the formaion of tunsten carbide from its lements in standard state?
W(s) + C(s, graphite) --> WC (s)
2W(s) + 3O2(g) --> WO3(s); ?H = -1685.8 kJ C(s, graphite) + O2(g) --> CO2(g); ?H = -393.5 kJ 2WC(s) + 5O2(g) --> 2WO3(s) + 2CO2(g); ?H = -239.18 kJ |
ANSWER ?H = -40.5 kJ |
| Standard Enthalpy of Reaction (?H°) |
| The enthalpy change of a reaction as it occurs in standard conditions. |
QUESTION CARD
What is the ?H?vap of CS2 in standard conditions?
?H?fCS2(l) = 89.7 kJ ?H?fCS2(g) = 116.9 kJ |
ANSWER ?H?vap = 27.2 kJ |
QUESTION CARD What is the standard enthalpy for this reaction?
4NH3(g) + 5O2(g) --(Pt)--> 4NO(g) + 6H2O(g)
?H?fNH3(g) = -45.9 kJ ?H?fNO(g) = 90.3 kJ ?H?fH2O(g) = -241.8 kJ |
ANSWER ?H? = -906.0 kJ |
| Bond dissociation energy |
| The energy required to break a bond in a molecule. It is a measure of the bond's strength. |
| Bond (Dissociation) Enthalpy/Energy (BDE/BE/?HBD) |
A measure of the average strength of a bond between two specific elements. Obtained from enthalpies of reactions. It will be different in different textbooks since it is an average. Determined by breaking the bonds in a molecule and dividing the energy by the number of bonds. The energy to break a bond is the negative of the energy to form the same bond. It is defined as the ?H of the dissociation of the bond into its atom not ions or molecules. Double bonds are not twice the BE of single bonds, nor are triple bonds thrice the BE of single bonds. However double bonds have bigger BEs than single bonds and triple bonds bigger BEs than double bonds. |
| Covalent bond |
| A bond between atoms where valence electrons are shared. The distance between the atoms is the bond length, which depends on the atoms inovlved. The bigger the bond length, the lower the BE. |
| Ionic bond |
| When two atoms completely transfer valence electrons. This forms ionic solids. An attraction between two oppositely charged ions. It is not an energetically favourable form; energy is required to remove and insert the electron(s) when forming ions. The energy required can be calculated with Coulomb's Law. |
| Ionic solid |
| A solid made from an ionic substance. The similarily charged ions repel and the oppositely charged ions attract. Because of this, they form a uniform formation like a crystal. They have very high melting points and they conduct electricity. When dissolved in water they conduct electricity. |
| Lattice energy |
| The change in energy when an ionic solid is separated into ions in the gas phase. Sometimes it is defined as the energy to put an ionic substance together from gas ions, to go from one or the other simply negate the value. You can calculate it using the Born-Haber Cycle. |
| Electron affinity |
| The energy required to remove an electron from an atom. |
| Born-Haber Cycle |
Invented in 1919 by Max Born and Fritz Haber. A way of calculating lattice energy using Hess's law with these 5 reactions: This example uses NaCl as the ionic solid. 1. Sublimation of the first consituent (Na solid into gas) 2. Ionization of first constituent (Na into Na+) 3. Dissociation of second consituent (Cl2 into Cl) 4. Electron gain of second constituent (Cl to Cl-) 5. Brining the two constituents together (Na+ and Cl- form NaCl) |
| Lattice enthalpy (V) |
The enthalpy of an ionic solid. Depends on the charge of the ions and their bond length. Depends on Coulomb's Law. V = (k*Q1*Q2)/r k = 8.99e9 J*m/C2 (a constant) Q = charges of ions r = bond length in meters |
| Entropy (S) |
The measure of how dispersed the energy of a system is among all the different possible ways that a system can contain energy. Measured in J/K. It is an extensive property (relies on size of system).
q = T?S
It can be based on the positional (pressure) or thermal (heat) of the situation. |
| Spontaneous process |
A reaction that occurs without any external intervention. ?H is negative and ?S is positive: the reaction is always spontaneous.
?Suniverse > 0 ?Ssystem > q/T |
| Non-spontaneous process |
A reaction that requires external intervention in order to occur. Only spontaneous in the reverse direction. ?S < 0 this defies the Second Law of Thermodynamics. ?H is positive and ?S is negative: the reaction will never happen.
?Ssystem > q/T |
| Reversible reaction |
A reaction that can be undone, e.g. melting water; it can be refrozen. Burning paper is irreversible. qrev is the heat supplied reversibly. ?S = qrev/T |
| Second Law of Thermodynamics |
The total entropy of the universe is always increasing.
?Stotal = ?Suniverse = ?Ssystem + ?Ssurroundings
*Note that ?Ssystem can be negative, as long as ?Suniverse is positive its ok. |
| Third Law of Thermodynamics |
| A perfect crystalline substance at absolute zero (T = 0K), entropy is 0. This can never be achieved. |
| Standard molar entropy (S°m) |
| The molar entropy of a substance in standard conditions. Gases are at 1 atm and solutes are at 1M concentration. This increases with melting, vaporization, being dissolved in solvent, increasing mass, increasing number of molecules, or increasing molecular freedom/flexibility. It decreases with freezing, condensation, or when a gas is dissolved in solvent. |
| Entropy of reaction (?Srxn) |
The change in entropy of a reaction. If you have the standard entropies of reactans and products it can be calculated.
?Srxn = ?nS°m(products) - ?nS°m(reactants) |
| Melting and boiling point |
The temperatures at which a substance melts or boils, respectively. At these points the system is at equilibrium.
?S° = (?H°/T) |
| Chemical bond |
| A bond between two atoms in a substance. It takes energy to form bonds and breaking bonds releases energy. |
| Metallic bond |
| When a crystal consists of a regular arrangement of atoms and the valence electrons can move freely about the crystal, holding it together and giving it electrical conductivity. |
| Salt |
| A crystalline solid with a high melting point. Molten salts conduct electricity. |
| Cation |
| A positive ion. It has lost an electron. This state is more stable than the neutrally charged state for some elements. It has an attraction to anions. |
| Anion |
| A negative ion. It has gained an electron. This state is more stable than the neutrally charged state for some elements. It has an attraction to cations. |
| Lewis electron-dot symbols |
| A way to draw ions where the valence electrons are represented by dots around the letter symbol of the element. |
| Coulomb's Law |
Potential energy released from bringing to oppositely charged ions together is found by: E = (KQ1Q2)/r
K = 8.99e9 J*m/C2 Q = charges of the two ions r = distance between the ions
This gives us E for one pair of atoms. Ionic bonds that share more electrons have higher potential energy. And ions with smaller diameter (closer to each other) will have higher potential energy. |
| Coulomb (C) |
| A unit used to measure charge of ions. One electron has a charge of 1.602e-19 C. |
| Avogadro's number |
The number of molecules in a mole.
6.02e23 |
QUESTION CARD CCl4(l) --> CCl4(g); ?H = 39.4 kJ
If one mole of liquid carbon tetrachloride at 25°C has an entropy of 216 J/K. What is the entropy of 1 mole of tetrachloride vapour at 25°C. |
ANSWER
348 J/mol*K |
| Standard entropy/Absolute entropy (S°) |
| The entropy in standard conditions of a substance. 1 atm of pressure and 1 M of aqueous solutes. This value is always positive, except for ions, which may have negative values. |
| J. Willard Gibbs |
| The person who came up with free energy, or Gibbs energy. |
| Gibbs energy/Free energy (G) |
A criterion for spontaneity of a reaction. ?G is a state function (independent on path). G = H - TS ?G = ?H° - T?S° ?G° = ?H° - T?S°
As a reaction progresses entropy decreases and Gibb's energy increases. If a reaction is done so that every single bit of energy is used up, no entropy will be created, however this can never happen in real life.
?G is always negative in spontaneous reactions; free energy decreases. A reaction is at equilibrium when ?G is as small as possible. |
| Enthalpy driven |
| When ?S = 0. A spontaneous process such as a ball rolling down a hill. |
| Entropy driven |
| When ?H = 0. A spontaneous process such as two inert gases mixing. |
| Reaction quotient (Q) |
In the the reaction aA + bB --> cC + dD
Qc = ([C]c[D]d)/([A]a[B]b) Qp = (PCcPDd)/(PAaPBb)
At equilibrium Q = K. If Q > K, the reaction will proceed towards the reactants. If Q < K, the reaction will proceed towards the products. This number is dimensionless. |
| Van't Hoff Equation |
| ln(K2/K1) = (?H°/R)((T1-T2)/T1T2) |
| Exergonic |
| A reaction where Gibbs energy is released. ?G = negative. |
| Endergonic |
| A reaction where Gibbs energy is absorbed. |
| ATP |
| A molecule in cells that is high in Gibbs energy. It is made from metabolizing glucose. It is used in reactions in the cell which are non-spontaneous, and use the Gibbs energy in the ATP molecule in order to proceed. |
| Equilibrium constant (K) |
Expresses the partial pressures (in atm) or the concentrations (in M) of things in equilibrium. Q is this value when not at equilibrium.
aA + bB -> cC + dD
K = ([A]a[B]b)/([C]c[D]d) K = (PAaPBb)/(PCcPDd) |
| Oxidation-reduction (redox) reaction |
| A reaction involving the transfer of an electron from one species to another. |
| Oxidation number |
The hypothetical charge of an atom if it were on its own. Rules for assigning oxidation numbers: 1. Lone atoms or atoms in single-element molecules have oxidation number of 0. 2. Oxidation number of an ion is equal to the charge. 3. Except in peroxides (H and O only in the compound), oxygen has an oxidation number of -2. 4. Except in binary compounds with metals (when it is -1), hydrogen has an oxidation number of +1. 5. The sum of the oxidation numbers should equal the overall charge of the molecule. |
| Oxidized |
| The term to describe an atom whose oxidation number has gone up in a reaction. It has accepted electrons. |
| Reduced |
| The term to describe an atom whose oxidation number has gone down in a reaction. It has accepted electrons. |
| Half reaction |
| One of two parts of a redox reaction. There is an oxidation half reaction and a reductioin half reaction. |
| Oxidizing agent |
| The species that oxidizes another species and is itself reduced. It accepts electrons. The higher the standard electrode potential, the stronger the oxidizing agent. |
| Reducing agent |
| The species which reduces another species and is itself oxidized. It donates electrons. The lower the standard electrode potential, the stronger the reducing agent. |
| Combination reaction |
| A reaction where two substances combine to form a new substance. Most, but not all, are redox reactions. |
| Decomposition reaction |
| A reactio where a compound decomposes to give two or more substances. Most, but not all, are redox reactions. |
| Displacement/Single-replacement reaction |
| A reaction where an element reacts with a compound, displacing another element from it, replacing it. A redox reaction. Whether or not a displacement reaction will occur depends on the activity series. |
| Activity series |
| A list of elements in the order of activity. Will tell you whether or not a displacement reaction will occur. |
| Half reaction method |
| A method of balancing a redox reaction where you split up the reaction into its half-reactions, balance those, then add it back together. |
| Alessandro Volta |
| The person who made the first battery around year 1800. He made a volatile stack of metal disks separated by paper disks soaked in salt water. Any two metals can be used to make a battery like this. |
| John Frederick Daniell |
| The person who made a popular battery in 1836, using zinc and copper. Today batteries like this are found in many electronics. |
| Voltaic/Galvanic cell |
| An electrochemical cell which uses spontaneous redox to make an electrical current which can be used as an energy source. Chemical energy transformed into electrical energy. It is two half-cells connected by a salt bridge. Generates electric current as electrons move from anode to cathode. |
| Electrochemical cell |
| A system consisting of electrodes dipped in electrolytes, giving a chemical reaction that uses or generates electric current. |
| Electrolytic cell |
| An elecrochemical cell where an external energy source drives an otherwise non-spontaneous reaction. Electrical energy is transformed into chemical energy. |
| Half-cell |
| A portion of an electochemical cell where the half-reaction takes place. A metal strip dipped into a solution of its metal ionl. There are anodes and cathodes. |
| Salt bridge |
| A tube of electrolyte gel that connects two half-reactions in a voltaic cell. If the two reactions were to come into contact with each other, the battery would shut down. Maintains electroneutrality, and allows ion migration. Can be made from agar and KCl in a gel form in a U-tube. Or a porous seaparator. |
| Anode |
| The half-reaction in a voltaic cell where oxidation occurs. Electrons are donated. It has a negative sign. In notation it appears on the left. The - terminal. |
| Cathode |
| The half-reaction in a voltaic cell where reduction occurs. Electrons are accepted. Has a positive sign. In notation it appears on the right. The + terminal. |
| Voltaic Cell Notation |
| The anode is written on the left and the cathode on the right. The salt bridge is denoted by "||". The phase boundary is denoted by "|". If there are multiple ions in one solution, separate them by a comma. To fully specify the equation, add the concentrations of each ion in parentheses. If one of the reactants is a gas, platinum serves as the surface for the half-reaction to take place. It is a catalyst. |
QUESTION CARD Write the cell reactions for a) Tl(s)|Tl+(aq)||Sn2+(aq)|Sn(s) b) Zn(s)|Zn2+(aq)||Fe3+(aq),Fe2+(aq)|Pt |
ANSWER a) 2Tl(s) + Sn2+(aq) --> 2Tl+(aq) + Sn(s) b) Zn(s) + 2 Fe3+(aq) --> Zn2+(aq) + 2Fe2+(aq) |
| Potential Difference |
The difference in electrical potential of to half-reactions. The energy produced by the cell is based on it. It is measured in volts.
Electrical work = Charge*potential difference |
| Volt |
| A unit of measurement used for electricity. The voltage is at its maximum when current is zero. The difference between maximum voltages of half-reactions is cell potential. |
| Faraday constant (F) |
The magnitude of a charge in one mole of electrons. Equal to 9.6485e4 C/moles of e-.
|
| Cell potential/Electromotive force (emf)(Ecell) |
The difference between the maximum voltages of the two half-reactions, or oxidation and reduction potentials. Measured with an electric digital voltometer. If the number is negative, the reaction is non-spontaneous.
wmax = -nFEcell Ecell = Ecathode - Eanode
n = number of electrons transfered F = Faraday constant |
| Electric digital voltmeter |
| Gives the voltage with a negative/positive sign so you can tell which is the anode and which is the cathode. |
QUESTION CARD
The cell potential of this voltaic cell is 0.650 V. Calculate the maximum electrical work of this cell when 0.500 g of H2 is consumed.
Hg2+(aq) + H2(g) --> 2Hg(l) + 2H+(aq) |
ANSWER
-3.09e4 J |
| Oxidation potential |
The potential of the oxidation half-reaction.
Oxidation potenital = -Reduction potential of reverse reaction. |
| Reduction/Electrode potential (E) |
The potential of the reduction half-reaction. The ability of a species to act as an oxidizing agent. There is a list of reduction potentials, very useful when doing calculations. An intensive property (doesn't depend on the amount of a species).
Oxidation potenital = -Reduction potential of reverse reaction. |
| Standard Cell Potential (E°cell) |
| The cell potential of a voltaic cell under state conditions: solutions at 1M, gasses at 1 atm, temperature at 25 C. |
| Standard Electrode Potential (E°) |
| The electrode potential under standard-state conditions: solutions at 1M, gasses at 1 atm, temperature at 25 C. Calculated using the standard hydrogen electrode as equal to zero. A table of standard electrode potentials can be used to calculate the cell potential of any cell. |
| Standard Hydrogen Electrode |
| Its standard electrode potential is arbitrarily held at zero volts, so we can calculate other potentials by comparing them to this. |
| Nernst equation |
Ecell = E°cell - (RT/nF) ln Q or Ecell = E?cell - (0.0592/n) log Q Note: the second one only works at 25?C. |
| Glass electrode |
| An ion-selective electrode that may be used to replace a hydrogen electrode. |
| Ion-selective electrode |
| An electrode that is sensitive to one ion. Can be used to measure the concentration of non-ion substances. |
| Zinc-carbon/Leclanche |
A battery used in flashlights, radios, and other portable appliances. Voltage is aound 1.5V, but decreases over time. Not good in cold weather. The anode is zinc. The cathode is a graphite rod surrounded by a paste made of manganese dioxide, ammonium, zinc chloride, and black carbon. Zn(s) --> Zn2+(aq) + 2e- 2NH4+(aq) + 2MnO2(s) + 2e- --> Mn2O3(s) + H2O(l) + 2NH3(aq) |
| Alkaline dry cell |
Similar to the Leclanche, but potassium hydroxide replaces ammonium chloride. This makes it perform beter in cold weather.
Zn(s) + 2OH-(aq) --> Zn(OH)2(s) + 2e- 2MnO2(s) + H2O(l) + 2e- --> Mn2O3(s) + 2OH-(aq) |
| Lithium-iodine battery |
| A true dry battery (no paste). The anode is lithium metal. The cathode is an I2 complex. They are separated by a thin layer of lithium iodine crystal. Low voltage, but high resistance. Used in pacemakers. Lasts ten years. |
| Lead storage cell |
Electrodes are lead alloy grids. The anode is packed with spongy lead. The cathode is packed with lead (IV) oxide. They are soaked in aqueous sulfuric acid. Delivers 2V (adding more electrodes increases this number in factors of 2). Can be recharged by an external current. More water needs to be added, unless it is maintenance-free.
Pb(s) + HSO4-(aq) --> PbSO4(s) + H+(aq) + 2e- PbO2(s) + 3H+(aq) + HSO4-(aq) + 2e- --> PbSO4(s) + 2H2O(l) |
| Maintenance-free lead storage cell |
| A lead storage cell that has calcium metal in its lead alloy grids. This prevents it from using up water, so no water needs to be added when recharging. It is sealed shut. |
| Nickel-cadmium cell/Nicad cell |
A common storage battery. Anode is cadmium. Cathode is hydrated nickel oxide on nickel. The electrode is potassium hydroxide. Used in calculators, portable power tools, shavers, and tooth-brushes.
Cd(s) + 2OH-(aq) --> Cd(OH)2(s) + 2e- NiOOH(s) + H2O(l) + e- --> Ni(OH)2(s) + OH-(aq) |
| Nickel metal hydride (NiMH) |
A common replacement for a nicad cell. Use metal hydrides as anodes instead of cadmium; less toxic. Can undergo more discharge-recharge cucles. Used in electrid hybrid cars. |
| Lithium-ion cell |
| A common replacement for nicad cells. The anode is carbon. The cathode is lithium cobalt dioxide or lithium manganese. Used in consumer electronics. Lightweight for the amount of power they supply. Can be cycled through discharge-recharge many times before failure. |
| Fuel cell |
A battery, but has a continuous supply of fuel. Hydrogen gas is the anode. The H+ ions enter the PEM and electrons enter the external circuit. The H+ ions react with oxygen at the cathode, which is taking up electrons from the external circuit. Water is formed. Potential of 0.7. It can run at low temperatures. Used in space, emergency systems, cars, and buses.
H2(g) --> 2H+(aq) + 2e- O2(g) + 4H+(aq) + 4e- --> 2H2O(l)
Other versions can use hydrocarbons or methanol as fuel. |
| Using pipes as a fuel cell to prevent rusting |
A drop of water sitting on iron is a tiny voltaic cell. Water and air are the the cathode. The iron is the anode. O2(g) + 2H2O(l) + 4e- --> 4OH-(aq) Fe(s) --> Fe2+(aq) + 2e-
To prevent rusting of iron pipes, a magnesium anode can attach to the pipe, and then the iron can't act as an anode as well, greatly reducing rust. |
| PEM |
| Proton-exchange membrane. Found in fuel cells. |
| Oxidating |
| Losing electrons. The overall charge goes up. Occurs at the anode. Connects to - terminal. |
| Reduction |
| Gaining electrons. Overall charge goes down. Occurrs at the cathod, attached to + terminal. |
| Voltage reading |
| The reading on the voltometer. It should be positive if set up correctly (anode at - terminal, cathode at + terminal) |
QUESTION CARD
Zn(s) + 2Ag+(aq) --> Zn2+(aq) + 2Ag(s) |
ANSWER
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QUESTION CARD
Zn(s) + Cl2(g) --> Zn2+(aq) + 2Cl-(aq) |
ANSWER
|
QUESTION CARD
Zn(s) + Cu2+(aq) <--> Zn2+(aq) + Cu(s) |
ANSWER
|
QUESTION CARD
|
ANSWER
Ecell = 1.22 V |
| Electrolysis |
| The reaction that occurs in an electrolytic cell. It is non-spontaneous, but driven by the external electrical force. |
| Downs cell |
| An electrolytic cell filled with molten NaCl. Sodium metal forms at the cathode and chlorine gas forms at the anode. These must be kept away from ech other or they will react. This is used to make solid sodium metal. |
| Reduction reaction of water |
| 2H2O(l) + 2e- --> H2(g) + 2OH-(aq) |
| Oxidation reaction of water |
| 2H2O(l) --> O2(g) + 4H+(aq) + 4e- |
| Reduction reaction of sulfuric acid |
| 2H+(aq) + 2e- --> H2(g) |
| Oxidation reaction of sulfuric acid |
| 2SO42-(aq) --> S2O82-(aq) + 2e- |
| Overvoltage |
| The extra voltage needed for an electrolytic cell in real life over the calculated voltage. |
| Chor-alkali membrane cell |
| A cell with aqueous NaCl where the anode and cathode are separated by a plastic membrane. Only cations can pass through. |
| Chor-alkali mercury cell |
| A cell with aqueous NaCl where the electrodes are a mercury amalgam. |
| Amalgam |
| An alloy containing mercury. |
| Electrogalvanizing |
| Coating a metal with a thin layer of zing. Done by using it as an electrode in a zinc salt. |
| Coloumb (C) |
| A measure of charge equal to amperes*seconds. |
| Ampere (A) |
| The SI unit of current. |
| Silver/silver chloride and SCE electrodes |
| A common way to measure Cl- concentration. The cathode is a silver/silver chloride electrode, and the anode is a Saturatd Calomel Electrode (SCE). |
| pH meter |
| An ion selective electrode with a membrane permeable to only H+ ions. Inside the membrane there is a HCl solution of fixed concentration. |
| Dental Voltaic Cell |
| Accidentally biting down on an amalgam dental filling causes a jolt of pain. This is because the foil acts as an anode and teh saliva acts as a cathode, causing a short circuit. |
| Chemical kinetics |
| The study of reaction rates, and how they change under various conditions. |
| Factors effecting reaction rate |
1. Concentration of reactants 2. Concentration of catalyst 3. Temperature 4. Surface area, if a solid reactants or catalyst is involved. |
| Catalyst |
| A chemical that increases the rate of reaction without itself being consumed in the overall reaction. It decreases the Ea. It is written above the arrow in a chemical equation. |
| Reaction Rate |
The increase in molar concentration of products (or decrease in reactants) per unit of time. Measured in mol/L*s. Determined by removing samples at different points in a reaction, or monitoring a physical property of the reaction (volume, temp, colour).
Average rate = ?[A]/?t |
| Instantaneous Rate |
| The average rate of a very short period of time. |
| Rate of decomposition |
| The rate at which a reactant disappears. |
| Rate of formation |
| The rate at which a product appears. |
QUESTION CARD
How is the rate of formation of NO2F related to the formation of F2 in this reaction?
2NO2(g) + F2(g) --> 2NO2F(g) |
ANSWER CARD
|
| Spectrometer |
| A tool used to measure light absorption in order to determine reaction rate. |
| Manometer |
| An instrument used to measure pressure to determine rate of reaction. |
| Rate Law |
An equation that tells the rate of reaction in comparison to concentration of reactants and catalysts, raised to various powers.
Rate = k[A]m[B]n |
| Rate constant (k) |
A factor in the rate law. A constant at fixed temperature. It depends on 1. Z: the collision frequency 2. f: fraction of collisions having energy higher than Ea. 3. p: the fraction of collisions with the molecules properly oriented. |
| Reaction order |
| The number to which components in a rate law are raised. If a species is not present in a rate law, it is zero order. |
| Zero order |
| When a species is not present in the rate law. The reaction will occur no matter what the concentration of the species is, as long as there is some of the species. |
| Isomerization |
| A reaction where the molecule changes structure, but doesn't break up or form new substances. |
| Integrated Rate Law |
| A function of rate law over temperature. |
| First Order Rate Law |
| Integrated Rate Law: ln[A]t / [A]0 = -kt |
| Second Order Rate Law |
| Integrated Rate Law: 1 / [A]t =kt + (1 / [A]0) |
| Zero Order Rate Laws |
| Integrated Rate Law: [A]t = -kt + [A]0 |
| Half life (t1/2) |
| The amount of time it takes for a reactant to decrease to one half of its original amount. This is always the same, no matter what the original amount is. |
| Collision theory |
| For a reaction to ocur, the reactant molecules must collide with more energy than the activatioin energy, and be in the correct orientation. |
| Activation Energy (Ea) |
| The minimum energy molecules need to have in order to collide and succesfully react. |
| Transistion-state theory |
| When two molecules collide in a molecular reaction, they form an activated complex. |
| Activated complex |
| A high-energy, unstable grouping of atoms that forms when two molecules collide. If they are in the right orientation and have enough energy, they will form the products of the reaction. If not, they will go back to being reactants. |
| Arrhenius Equation |
Expresses the dependence of rate constant on temperature. Written by Svante Arrhenius.
k = Ae -Ed/RT
Where A is the frequency factor. |
| Elementary Reaction |
| A theoretical step reaction within a reaction. A single molecular event. |
| Reaction mechanisms |
| A series of elementary reactions that add up to a full reaction. They can only be theorized, you can never prove them. However, the mechanism has to correlate with the experimental rate law. |
| Reaction Intermediate |
| Species that aren't in the reactants or the products of a reaction, but are made during one of the elementary reactions, then consumed by another elementary reaction. |
| Molecularity |
| The number of molecules on the reactant side of an elementary reaction. Indicates how many molecules have to collide simultaneously for the reaction to occur. The larger the number, the smaller the likelihood they will collide. |
| Unimolecular reaction |
| An elementary reaction with a molecularity of 1. There are no collisions, instead the molecule splits up. The reaction is usually a decomposition reaction. |
| Biomolecular Reaction |
| An elementary reaction with a molecularity of 2. Two molecules collide. This is the most common molecularity. |
| Termolecular Reaction |
| An elementary reaction with a molecularity of 3. Three molecules collide. |
| Rate-determining step |
| The slowest elementary reaction in a reaction mechanism. It determines the overall rate of reaction. |
| Enzyme |
| A catalyst that is a protein. Found in living cells and organisms. It is usually very specific to its substrates. It has an active ste. |
| Substrate |
| A chemical upon which an enzyme reacts. Enzymes have very specific substrates. The substrate binds to the enzyme's active site. |
| Active site |
| A region on an enzyme to which the substrat binds. This is where the reaction is catalyzed. |
| Homogeneous catalyst |
| A catalyst in the same phase state as the reacting species. |
| Heterogeneous catalyst |
| A catalyst in a different phase state as the reacting species. It is usually a solid. |
| Adsorption |
| The attraction of molecules to a surface. |
| Physical adsorption |
| Adsorption from the weak intrmolecular forces. |
| Chemisorption |
| The binding of a spcies to a catalyst surface by chemical bonding forces. |
| Catalytic hydrogenation |
| When H2 is added to a compound to break a double bond, using platinum or nickel as the catalyst. This is how they hydrogenate vegetable oil to make solid margarine. |
