Energy in chemical physical changes
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Energy in Chemical & Physical Changes. Thermochemistry. Study of changes that accompany chemical reactions and phase changes The Universe is considered to be made of 2 parts: 1. System: part that contains the reaction or process 2. Surroundings: everything else. ENERGY.

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Energy in Chemical & Physical Changes

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Energy in chemical physical changes

Energy in Chemical & Physical Changes


Thermochemistry

Thermochemistry

  • Study of changes that accompany chemical reactions and phase changes

  • The Universe is considered to be made of 2 parts:

    1. System: part that contains the reaction or process

    2. Surroundings: everything else


Energy

ENERGY

  • defined as the ability to do work or transfer heat energy.

    2 types of energy

  • Potential Energy (PE): Energy at rest due to the position of an object; chemical potential energy is the energy stored in a substance’s bonds.

  • 2. Kinetic energy (KE): Energy of the motion of particles in a substance and is directly proportional to temperature. As temperature increases, KE also increases.


Law of conservation

Law of Conservation

  • Law of Conservation of Energy states that energy is neither created nor destroyed, just changed in form

    C8H18 + O2 H2O + CO2 + Energy

  • Stored PE converts to 25% work and 75% heat

    (ENERGY)


Exothermic reactions

Exothermic Reactions

  • HOT PACK

  • An exothermic reaction is when the system releases energy; heat flows out of a reaction and the surroundings get warmer. They have a NEGATIVEH.

  • H products < H reactants

    4Fe + 3 O2 2 Fe2O3 + 1625 kJ OR

    4Fe + 3 O2 2 Fe2O3H = - 1625 kJ


Endothermic reactions

Endothermic Reactions

  • COLD PACK

  • An endothermic reaction is when the system absorbs energy; heat flows into a reaction and the surroundings get cooler. They have a POSITIVEH

  • H products > Hreactants

    27kJ + NH4NO3(s) NH4(aq)+1+NO3(aq)-1 OR

    NH4NO3(s) NH4(aq)+1+ NO3(aq)-1H = + 27 kJ


Reaction co ordinates

Reaction Co-ordinates


What is the difference between temperature heat

What is the difference between Temperature & Heat?

Temperature

  • Instrument: thermometer

  • Units: Celsius, Fahrenheit, Kelvin

  • Definition:

  • A measure of the average kinetic energy of the molecules in a substance

  • A measure of the motions of the molecules

  • A measure of how hot or cold something is


What is the difference between temperature heat1

What is the difference between Temperature & Heat?

Heat

  • Instrument: calorimeter

  • Units: calories, joules

  • Definition:

  • The total amount of energy in a substance.

  • A form of energy that is transferred between objects because one is warner than the other.

  • Heat transfer is always from hot to cold

  • Depends on 3 things:

    1. amount of substance (mass)

    2. Temperature change

    3. type of material (specific heat)


Units of heat energy

Units of Heat Energy

  • A calorie is defined as the amount of heat needed to raise the temperature of 1 g of water by 1 C

    1 cal= 4.184 J

  • Food “Calories” are kilocalories.

    1kcal = 1000 calories.


Temperature heat

Temperature ≠ Heat

Greater Thermal Energy


Specific heat

Specific Heat

  • Amount of heat required to raise the temperature of 1 g of a substance by 1 C

  • Different substances have different specific heats.

    Water has a specific heat of 4.184 J/gC.

    Iron(Fe) has a specific heat of .449 J/gC.

    Gold (Au) has a specific heat of .129 J/gC.

  • The higher the specific heat the more energy it takes to change its temperature.


Calculating heat

Calculating Heat

c= specific heat

q = heat in joules or galories

m= mass

T = change in temperature = Tf – Ti

c= q_

mT


Example

Example

  • A 155 g sample of an unknown substance was heated from 25.0 C to 40.0 C. The substance absorbed 5696 J of energy. What is the specific heat?


Example1

Example

  • How much heat is needed to change the temperature of 12.0 g of silver with a specific heat of 0.057 cal/g°C from 25.0°C to 83.0 °C?


Measuring heat in a calorimeter

Measuring Heat in a Calorimeter

  • A coffee cup calorimeter measures heat at constant pressure; works on the premise that the amount of heat released in a reaction(-q) or physical change is equal to the amount of heat absorbed by the water(+q) - q = +q 

  • Rearrange the specific heat equation:

    q = m x c x T


Example2

Example

  • A piece of unknown metal with mass 17.19 g is heated to an initial temperature of 92.50 °C and dropped into 25.00 g of water (with an initial temperature of 24.50 °C) in a calorimeter. The final temperature of the system is 30.05°C. What is the specific heat of the metal? Specific heat of water = 4.184 J/g°C


Example3

Example

  • A 32.07 gram sample of vanadium was heated to 75.00 °C (its initial temperature). It was then dumped into a calorimeter. The initial temperature of the calorimeter’s water was 22.50 °C. After the metal was allowed to release all its heat to the calorimeter’s water, 26.30 °C was the final temperature. What mass of distilled water was in the calorimeter?

  • Specific heat of vanadium = .4886 J/gC Specific heat of water = 4.184 J/g°C


Measuring heat during phases changes

Measuring Heat during Phases Changes


Heat of fusion solidification

Heat of Fusion/Solidification

  • Heat of fusion (Hfus ) is the heat energy required to melt one gram of a solid at its melting point

    For water, Hfus = 334 J/g

    q = Hfus x mass

  • Heat of solidification (Hsolid ) is the heat energy lost when one gram of a liquid freezes to a solid at its freezing point

    For water, Hsolid = -334 J/g

    q = Hsolid x mass


Heat of vaporization condensation

Heat of Vaporization/Condensation

  • Heat of vaporization (Hvap) is the heat to vaporize one gram of a liquid at its normal boiling point

    For water, Hvap= 2260 J/g

    q = Hvap x mass

  • Heat of condensation (Hcond ) is the heat energy released when one gram of a liquid forms from its vapor

    For water, Hcond = -2260 J/g

    q = Hcond x mass


Example4

Example

  • How much heat is needed to melt 500.0g of ice at 0 C?


Example5

Example

  • How much heat is evolved when 1255 g of water condenses to a liquid at 100°C?


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