Chem12 CPT Review (PART TWO)

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the nature of energy: potential vs. kinetic

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1

the nature of energy: potential vs. kinetic

  • kinetic- energy of an object due to its motion

  • potential- energy of a body or system due to urs position our composition

    • the energy associated with chemical bonds

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2

law of conservation of energy

  • energy is always conserved

  • we cannot create new energy nor can we destroy energy by using it

  • energy can only be converted from 1 form into another

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3

what is the “system”?

  • composed go the reactants & products being studied

  • molecules in the reaction

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4

what are the “surroundings”?

  • all the matter that is not part of the system

    • water, solute, air

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5

open and closed systems (+isolated)

  • open system

    • can exchange both energy and matter with its surroundings

  • closed system

    • can exchange energy, but no matter, with its surroundings

  • isolated system

    • when no energy or matter can be exchanged or moved in or out of the system

    • impossible to have a true isolated system

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6

exothermic reactions

  • when more energy is released from the formation of new bonds than is required to break bonds, the reaction will release energy into the environment

    • products have lower potential energy than the reactant

  • negative q value

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7

endothermic reactions

  • the chemical system absorbs energy from the surroundings and increases its potential energy

  • positive q value

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8

nuclear energy

  • exo

  • fission

    • 1 large molecule splitting into 2

  • fusion

    • 2 smaller molecules bondings together to form 1 larger one

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9

specific heat capacity

  • the quantity of thermal energy required to raise the temp of 1g of substance by 1°C

    • H₂O is 4.18

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10

calorimetry and thermal energy transfer

  • the process if measuring energy changes during physical or chemical changes

    • well-insulated reaction chamber

    • tight fitting cover with insulated holes for a thermometer

    • a mechanism to stir contents within

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assumptions made while measuring calorimetry

  1. any thermal energy released without intent is negligible

  2. any thermal energy absorbed is negligible

  3. all dilute, aq solutions have the same density (1g/mL) & the specific heat capacity (1g/1°C)

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what is q?

  • the total amount of thermal energy absorbed or released in a chemical system

  • depends on 3 factors

    • mass

    • specific heat capacity

    • temp change

  • q=mc∆T

    • qsystem = -qsurroundings

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13

enthalpy change

  • enthalpy: the total amount of thermal energy in a substance (H)

    • enthalpy change → ∆H

    • ∆H = Hproducts - Hreactants

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14

molar enthalpy change

  • ∆H = n∆Hr

    • n= concentration x volume

    • n= mass/volume

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15

standard enthalpies of formation

  • the change in enthalpy that accompanies the formation of 1 mol of a compound from its elements in their standard states

  • ∆H°r = εₙproducts∆H°products - εₙreactants∆H°reactants

    • any singular element or anything not on the table has a value of 0

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16

bond energies

  • the quantity of energy required to break a chemical bond

  • ∆H = εₙxD(bb) - εₙxD(bf)

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factors affecting reaction rates

  1. chemical nature of the reactants (their states)

  2. concentration of the reactants

  3. surface area

  4. temperature of reaction system

  5. catalyst

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18

collision theory and geometry

  • the theory that chemical reactions can only occur if reactants collide with proper orientation and with enough kinetic energy to break reactants bonds and from product bonds

    • the geometry of the molecule is critical to how collision theory works (the proportions and formation)

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19

activation energy

  • the min. energy that reactant molecules must possess for a reaction to be successful

    • it must break the bonds

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20

catalyst theory

  • for any reaction to occur, the kinetic energy of colliding reactant entities must be equal to or greater than the activation energy

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21

reaction rate

  • the change in concentration of a reactant or product per unit time

    • gas volume

    • colour

    • mass

    • pH

    • electrical conductivity

  • rise/run

    • average reaction rate- over a specified time period

    • instantaneous rate of reaction- @ a specific point in time

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22

rate law

  • rate= k[A]ⁿⁿ[B]ⁿ

    • if solving for k, input A & B values from the first set of numbers

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23

NEXT UNIT

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24

dynamic equilibrium

  • a balance between forward and reverse processes occurring @ the same rate

    • forward is left to right

    • reverse is right to left

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25

solubility equilibrium

  • eq is reached when the solute being dissolved is @ the same rate as solute crystallizing out

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phase equilibrium

  • as more molecules of gas disappear (vapour pressure increases) so do the collisions between vapour and liquid

    • eg. water

      • H₂O (l) ⇌ H₂O (g)

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27

chemical reaction equilibrium

eg. CaCO₃ → CaO + CO₂

  • in an open system, CO₂ would escape, preventing the reverse reaction

  • in a closed system, both reactants and products would be present & a reaction would appear to have stopped collision between the reactants and products

    • a state of eq limits the amount of product that can be formed

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ICE charts

  • only for aq and g systems

  • I → initial

    • stands for the initial concentration of reactant or products

  • C → change

    • stands for change in concentration of the reactants or products

  • E → equilibrium

    • stands for the concentration of reactants or products @ equilibrium

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29

equilibrium law

  • when eq concentration are arranged in the following ratio (aA + bB ⇌ cC + dD) the resulting value was the same no matter what the initial concentration were mixed

    • K= [C]c[D]d**/**[A]a[B]b

  • forward and reverse reactions are simply reciprocals of each other

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magnitude of K

  • K > 1

    • reaction processes towards completion

  • K = 1

    • reactant and products are = @ eq

  • K < 1

    • very small amount of product(s) form

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limitations of eq constants and % reaction values

  • the value of the eq constant (K) depends on temp

  • % reaction values are dependent on temp and concentration

  • % reaction values & eq constants give no info about he rate of reaction

    • only a measure of eq position in the reaction

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homo and hetero equilibria

  • homogeneous

    • systems in which the reactants and products ae all in the same phase

  • heterogeneous

    • systems in which the reactants and products are in different phases

  • DO NOT INCLUDE SOLIDS OR LIQUIDS IN FINAL EQ EQUATIONS

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33

le chatelier’s principle

  • when a chemical system @ eq is disturbed by a change, the system adjusts in a way that opposes the change

  • changes cause an eq to shift when either more product will produce, or more will revert back to reactant

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le chatelier’s: concentration

  • shifts to the right to consume some of the product to “reuse”

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35

le chatelier’s: temperature

  • if the system is cool, the eq shifts towards the direction of heat production

  • if heat is added, the eq shifts towards the side that absorbs heat

    • always towards the heat

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le chatelier’s: gas volume changes

  • if pressure is increased, eq will shift to the side which possesses the least amount of moles

  • if pressure decreases, eq will shift to the side that possess the most amount of moles

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changes that do not affect eq shifts

  1. adding catalysts

  2. adding inert gas

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38

the reaction quotient (Q)

  • used when both reactants and products are present in an initial reaction, & it becomes difficult to determine which direction the reaction will go before reaching eq

    • we place all initial reactant and product concentrations into the equilibrium law equation to generate a trial value (Q)

  • Q=K → system is @ eq

  • Q>K → system must shift to the left to reach eq

    • product:reactant ratio is too high

  • Q<K → system must shift to the right to reach eq

    • product:reactant ratio is too low

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39

hundred’s rule

  • if [initial]/K > 1, we can ignore the x expression

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40

Ksp

  • usually used for substances with very low solubilities

  • values are very small

    • when evaluating, you take the Ksp value and leave it on the left side. get the exponents from the compounds and change them to [#]x. then evaluate and simplify

<ul><li><p>usually used for substances with very low solubilities</p></li><li><p>values are very small</p><ul><li><p>when evaluating, you take the Ksp value and leave it on the left side. get the exponents from the compounds and change them to [#]x. then evaluate and simplify</p></li></ul></li></ul>
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41

precipitate prediction (Q and Ksp)

  • to predict whether an precipitate will form, we compare Ksp value for the salt of the ions to the trial ion product (Q)

    • product, Q > Ksp → supersaturated, precipitate

    • product, Q = Ksp → saturated, no precipitate

    • product, Q < Ksp → undersaturated, no precipitate

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42

the common ion effect

  • a shift in eq when dissolving into a solution when a substance containing a common ion product is present in the eq system

    • must use an ICE chart

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43

Arrhenius acids and bases

  • an acid is a substance that produces H ions when it is dissolved in water

  • a base is a substance that produces OH ions when it is dissolved in water

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problems with Arrhenius (acid & base)

  • any substance with H ions can form acids, and any substance with OH ions can form bases

  • all acid-base reactions occur in aq solutions

  • proposed all salts must be neutral, but this is false

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45

Bronsted-Lowry Theory

  • acid is a H ion donor

  • base is a H ion acceptor

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46

conjugate acid-base pairs

  • conjugate acid- an acid formed when a base accepts a H ion from an acid

  • conjugate base- a base formed when an acid donates a H ion to a base

  • conjugate acid-base pairing- 2 substances related to each other by donating and accepting single H ions

    • does not need to be aq

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amphiprotic substance

  • amphiprotic substance- able to donate or accept a H ion, thus acts as both a Bronsted-Lowry acid, and Bronsted-Lowry base

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48

acid ionization constant, Ka

  • used for the actions in which an acid, HA (aq), reacts with water

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49

strong and weak acids

  • strong acids- an acid that ionized almost 100% in water, producing H ions

  • weak acids- an acid that only partially ionizes in water, producing H₂O ions

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50

strong and weak bases

  • strong base- a compound that dissociates completely un water, producing OH ions

  • weak base- a base that undergoes an eq with water to produce hydroxide ions

  • base ionization constant (Kb) - eq constant for the ionization os a base

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51

autoionization of water

  • acts as an acid or base in the same reaction

    • ion-product constant (the reaction that occurs to make water have these properties)

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52

pH

  • pH- the negative logarithm of the concentration of H ions in an aq solution

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53

pOH

  • the negative logarithm of the concentration of OH ions in an aq solution

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