Chapter 10 - Chemistry

Energy Types (Chapter 10)

Potential energy: stored energy in chemical bonds

• Correlated to distance between particles and magnitude of charge or attraction

Kinetic energy: movement of particles

• Correlated to temperature

Mechanical energy: total of potential and kinetic energy for an object

Energy Units (Chapter 10)

4 main energy units:

• Joule, J: SI unit of energy

• Kilojoule: 1000 J

• calorie, cal: amount of energy required to raise the temperature of 1 g of water by 1°C, equal to 4.184 J

• Calorie, Cal or kcal: unit of energy used in utrtition, 1 kcal = 1000 cal, 1 kcal = 4.184 kJ

Example:

System & Surroundings (Chapter 10)

System: what is being studied, must be defined by experimenter

Surroundings: the environment around the system taht may exchange matter or energy with it

Types of systems:

• Open: can exchange matter and energy with surroundings

• Closed: can exchange energy with surroundings

• Isolated: cannot exhcnage matter or energy with surroundings

System & Surroundings Examples (Chapter 10)

A pot with boiling water on the stove

• Pot & water are system, stove and air around pot are surroundings, system is open if lid is off, closed if lid is on

A thermos with water and ice

• Water, ice, & thermos are system, table and air around thermos are surroundings, system would be isolated if thermos could insulate perfectly (it can’t)

Work & Heat (Chapter 10)

Work, w: energy resulting from a force acting on an object over a distance

Heat, q: flow of energy that causes a chnage in temperature in an object or its surroundings

The sign of w & q are imporant and must be determined based on the situation

• If work is done on the system (a force is exerted on it), w is positive

• If a system is heated (energy flows to it), q is positive

• If work is done by the system (a force is exterted by it), w is negative

• If a system produces heat (energy flows form it), q is negative

Work & Heat Examples (Chapter 10)

Pot with boiling water on a stove

• The system is heated, +q

A chair is pushed across the room

• A force is exerted on the chair (system), +w

Work & Heat Practice (Chapter 10)

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Law of Conservation of Energy (Chapter 10)

The first law of thermodynamics is the law of conservation of energy

• Energy cannot be created or destroyed, but it can be transferred from one form to another

• The energy of the universe is constant

  • Universe = system + surroundings

Internal Energy (Chapter 10)

Is the all the kinetic and potential energy of a system

• Applies to movement of particles only

• The chnage in internal energy is measured by the sum of heat and work

  • △U = q + w

  • If the internal energy of the system increases, △U is positive

  • If the internal energy of the system decreases, △U is negative

Internal Energy Practice (Chapter 10)

Answer

State & Path Functions (Chapter 10)

All quantitites we’ve discussed are function: heat, work, internal energy

State functions are independent of path; you can start in two different places and end up with the same answer; current state of system is described

• U, internal energy

Path functions are depdent on path; if you start in different places, you will get different asnwers

• q, heat

• w, work

Pressure — Volume Work (Chapter 10)

Pressure — volume work is wokr done on or by a system where in the volume of a gas is increased or decreased agaisnt a constant pressure

• w = -P △ V

• △V = V_final - V_initial

• V is in L, P is in atm

• 1 L•atm = 101.325 J

Generally, systems are setup as gas in a clinder with piston that can move

Pressure — Volume Work Practice (Chapter 10)

Answer

Enthalpy & Heat Exchange (Chapter 10)

Chemists generally do reactions in open systems where in the pressure is constant (atmospheric pressure)

Heat exchange between system and surroundings is of most interest

Enthalpy, H, is teh combination of internal energy and P-V work

• Wokr cancels out of the equation, leaving just q_p - heat flow at constant pressure

• △H = q_p

Enthalpy has the same sign concentions as heat, but different vocabulary

• Endotehrmic: heat absorbed, +△H

• Exothermic: heat released, -△H

Heat Exchange Practice (Chapter 10)

Answer

Specific Heat (Chapter 10)

The specific heat, c, of a substance is the heat required to raise the temperature by 1°C

• Unit: J/g°C

• Difficult to change the temperature of subsatnces with high specific heats

  • Example: H₂O has a very high specific heat of 4.184 J/g°C

The “specific heat equation” is used to calculate how much energu is required to chneg the temperature of agiven amount of specific substance

• q = mc△T

  • m is mass in g

  • c is specific heat in J/g°C

  • △T is final — initial temperature (T_f - T_i) in °C

Specific Heat Practice (Chapter 10)

Answer

Calorimetry (Chapter 10)

In a calorimetry experiment, two susbatnces of different temperratures are put in contant with each other so the trnasfer of heat can be invetiagted

• Metaal in water or mixxing solutions

• Consatnt-prssure calorimetry - a chemical reaction atle spalce at sonatnt pressure or aolute is diddovled in as olvent

• Constant-volume calorimetry - a chemical reaction takes place at constant volume in a a bomb calorimeter

Metal in Water Calorimetry (Chapter 10)

General Process:

• Metal is heated, typically in hot water

• Room temperature water is added to calorimeter

• Metal is added to colrimeter and temperature taken when thermal equilibrium is established

Example:

Metal in Water Calorimetry Practice Pt.1 (Chapter 10)

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Metal in Water Calorimetry Practice Pt.2 (Chapter 10)

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Constant-Pressure Calorimetry (Chapter 10)

The heat of solution, q_soln is what is measured in a calorimetry experiment wherein a substance dissolves

Most dissolving processes are either endothermic or exothermic, so a change in temperature can be measured as the solute dissolves in the solvent

• The temperatures recorded will be of the solvent (initially) and teh solution (finally)

Since the dissolving process itself cannot have its temperature measured, it can be derived from that of the solution

• q_soln = -q_rxn (rxn = reaction)

• Could also use q_p for heat of reaction (P for constant-pressure)

Constant-Pressure Calorimetry Practice Pt.1 (Chapter 10)

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Constant-Pressure Calorimetry Practice Pt.2 (Chapter 10)

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Constant-Volume Calorimetry (Chapter 10)

A bomb calorimeter is used to measure the heat of combustion of a reaction

Pressure is varibale and volume constant

The total heat capacity of the system includes the calorimeter itself, along with the water inside

• You are given a value for this, C_cal, which applies to any process in that calorimeter with specific amoint of water

Calculate constant-volume heat of reaction using:

• q_v = -C_cal △T

• q_comb can be used instead of q_v

Example:

Constant-Volume Calorimetry Practice (Chapter 10)

Answer

Enthalpy as a Conversion Factor (Chapter 10)

When given the value for the entahlpy (heat released or gained) for a chemical reaction, it cna be related to the moles of each substance in that reaction

• Example: formation of liquid water:

Enthalpy as a Conversion Factor Practice (Chapter 10)

Answer

Hess’s Law (Chapter 10)

Hess’s law states that the enthalpies of 2 or more chemical reaction that can add up to an overall reaction, can also be added up to get the overall enthalpy

• Similar to a system of equation in math where you manipulate one equation to cancel out a variable in both

• The goal is to get the equations listed as steps to add up to be teh overall reaction by having certain substances cancel out. Cancellation occurs by one reaction having a subatnce as a reactant and another having it as a product

Hess’s Law Rules (Chapter 10)

Rules for chemical euqation manipulation:

• You can;t chnage the chemicals in any way, just the coefficients

• You can multiply or divide the coefficients as needed, but all coefficients in the equation must beaffected as well as the enthaply change

• You can reverse a reaction by making the products the reactants and vice versam but the enthalpy chnage sign needs to be reversed as well

• Substances can aptially cancel out if there are, for example, 2 mol A as a reactant and 1 mol A as a product in two of the reactions. This would result in 1 mol A as a reactant and no mol A as a product - like subtraction.

Hess’s Law Practice Pt. 1 (Chapter 10)

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Hess’s Law Practice Pt. 2 (Chapter 10)

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Enthalpy of Formation & Reaction (Chapter 10)

Enthalpy of formation, △H_f, is the amount of energy associated with forming 1 mol of a compounds from its elements

If there is a degree sign in the symbol (△H_f°), that means that this energy is associated with the standard state of the lement or teh state in which the susbatnce is most stabke at 25°C)

Can use enthalpies of formation to find the nehalpy of reaction, △H_rxn°

Enthalpy of Formation & Reaction Practice Pt. 1 (Chapter 10)

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Enthalpy of Formation & Reaction Practice Pt. 2 (Chapter 10)

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Bond Length & Strength (Chapter 10)

Bond length and strength are correlated

• Shorter bonds are stronger, resulting in less reactive molecules

• Longer bonds are weaker, resulting in more reactive molecules

• Single bonds are longest, followed by double and the triple

Bond enthalpy is the enrgy associated with breaking a chemical bind in 1 mol of gas molecules

• Bond breaking requires energy - endothermic

• Bond making releases eegry - exothermic

Bond Enthalpies Example (Chapter 10)

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Bond Enthalpies Practice (Chapter 10)

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