Two key definitions before we begin:

- __thermodynamics__ = the study of energy and its interconversions.

- __1st Law of Thermodynamics__ = the total energy of the universe is constant.

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**The Internal Energy of a System (E)**

- think of the "system" as the "chemical reaction."

* The Internal Energy of a system* - the sum of the kinetic and potential energies of all the "particles" in the system (the reaction).

➞ this internal energy (E) can be __changed__ by a flow of **work (w)**, **heat (q)**, or both:

** ΔE = q + w**

As the image above suggests, let's examine the specifics of heat (q) and work (w)...

... the __signs__ (+ or -) of heat (q) and work (w) are identified from the *point of view of the system*:

A. if the reaction is __endothermic__, heat flows __into__ the system and thus heat (q) is positive (+) ➞ **q > 0**

B. if the reaction is __exothermic__, heat flows __out of__ the system and thus heat (q) is negative (-) ➞ **q < 0**

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C. if the reaction (__system__) does work (w) on the surroundings, energy flows __out of__ the system, so work (w) is negative (-) ➞ **w < 0**

D. if the __surroundings__ do work (w) on the system, energy flows __into__ the system, so work (w) is positive (+) ➞ **w > 0**

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*ex:* Calculate ΔE for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1400 J of work is done on the system.

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**How to Calculate Work (w)**

The "__work__" associated with chemical processes is usually work done * by gases* (through

__expansion__) or work done

*(through*

**to gases**__compression__).

Let's examine 2 CASES (below) and the following relationship between work, external pressure, and volume: **w = - PΔV**

- __external pressure (P)__ = usually a constant in these types of problems, so * w* depends on

**ΔV**__ CASE 1__ ➞ If a gas

__expands__, the volume of the system

__increases__, so the change in volume (ΔV) is positive (+).

Because pressure (P) is always positive (+), then the __work (w)__ must be negative (-).

➞ this corelates well with "C" above because the reaction is doing work * on* the surroundings in order to

__expand__the volume.

-- so, **w < 0** (-) because energy is flowing __out of__ the system.

__ CASE 2__ ➞ If a gas gets

__compressed__, the volume of the system

__decreases__, so the change in volume (ΔV) is negative (-).

Because pressure (P) is always positive (+), then according to "w = -PΔV", the __work (w)__ must be positive (+).

➞ this corelates well with "D" above because the surroundings are doing work * on* the system in order to

__compress__the volume.

-- so, **w > 0** (+) because energy is flowing __into__ the system.

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*ex:* Calculate the work associated with the expansion of a gas from 46L to 64L at a constant external pressure of 15atm.

This is an "expansion" (CASE 1).

Don't think of the (-) minus __sign__ as a "negative number," but instead think of the (-) sign as a __direction__.

➞ since the gas expands, __work__ is negative *(case 1)*.

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*ex:* A balloon is being inflated by heating the air inside it. The volume of the balloon increases from 4.00 x 10

^{6}L to 4.50 x 10

^{6}L by the addition of 1.3 x 10

^{8}J of energy as heat.

Assuming that the balloon expands against a constant pressure of 1.0atm, calculate ΔE for this process. (*note*: 1 L^{.}atm = 101.3 J )

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Next, we'll talk about how **"heat" (q)** is __measured__ experimentally using a concept called Calorimetry.

So, stick around for my next post on **SECTION 6 - Thermochemistry**,

where we'll discuss Calorimetry and Specific Heat Capacity...