Reaction Energy
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Reaction Energy

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Reaction Energy




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Slide1

Reaction Energy

Chapter 16 page 500

Slide2

A. ThermoChemistry

Introduction

Every chemical reaction causes a change in energy

Endothermic or exothermic

Thermochemistry – study of the transfer of energy as heat that happens during chemical reaction and physical changes

Slide3

2. Temperature and Heat are Related

a. Calorimeter – tool which measures energy absorbed or

released as heat in a chemical or physical change

b. Temperature – measure of the average kinetic energy

of the particles

1) Typically use °C or K

2) K = 273.15 + °C

3) Measured by energy transfer in joules (J)

Slide4

c. Joule 1) SI unit of heat as well as all other forms of energy 2) Amount of heat needed to raise the temp of 1 g of water 0.239°C 3) Derived from units for force and length 4) N x m or d. Heat 1) Energy transferred between samples of matter because of a difference in their temperatures 2) Always moves from higher temperature to lower temperature until equilibrium is reached (Only heat moves, never cold)

 

Slide5

3) Unit is either joules or calories

a) A

c

alorie – amount of heat it takes to raise the temperature

of 1 gram of water 1°C

b) A dietary

C

alorie = 1000 calories

c) 1 Cal = 1 kcal = 1000 calories

d) 1

cal

= 4.184 J

Slide6

3. Energy transfers changes with the reaction

a. Energy transfer as heat depends on:

1) Mass of material

2) Type of material

3) Size of the temperature change

4) Example: 1 gram of iron heated to 100°C and then

cooled to 50°C transfers 22.5 J in a calorimeter. Silver

under the same conditions only transfers 11.8 J

Slide7

b. Specific heat 1) Amount of energy required to raise the temperature of one gram of substance by one Celsius degree (1°C) or one kelvin (1 K) 2) Unit: J/(g·K) or J/(g·°C) 3) Fig 1.2 (pg 503) 4) Water 4.18 J/(g·K) a) Very high for a common substance b) Allows for temperature regulation

Substance

Specific Heat J/(

g·K

)

Water (l)

4.18

Water (s)

2.06

Water (g)

1.87

Ammonia (g)

2.09

Ethanol (g)

2.44

Aluminum (s)

0.897

Iron (s)

0.449

Lead (s)

0.129

Slide8

5) cp= a) cp = specific heat (p is for pressure, b/c its done under a constant temperature) b) q = energy lost or gained i. Endothermic = positive value ii. Exothermic = negative value c) m = mass of sample d) ∆T = change in temperature6) Energy gained or lost: q = cp x m x ∆T

 

Slide9

Calorimetry

Two kinds

https://

www.youtube.com/watch?v=EAgbknIDKNo

Slide10

7) Examples:

a) A 4.0 g sample of glass was heated from 274 K to 314 K, a

temperature increase of 40 K, and was found to have

absorbed 32 J of energy as heat. What is the specific heat

of this type of glass?

b) How much energy will the same glass sample gain when it is

heated from 314 K to 344 K?

Slide11

c) Determine the specific heat of a material if a 35 g sample

absorbed 96 J as it was heated from 293 K to 313 K.

Answer: 0.14 J/(

g•K

)

d) If 980 kJ of energy are added to 6.2 L of water at 291 K,

what will the final temperature of the water be?

Answer: 329 K

Slide12

e) A piece of copper with a mass of 95.4 g increases in

temperature from 25.0°C to 48.0°C when the copper

absorbs 849 J of heat. What is the specific heat of

copper?

Answer: 0.387 J/(g·°C)

Slide13

4. Heat energy transferred

a. Enthalpy Change

1) Amount of energy absorbed by a system as heat

during a process at constant pressure.

2) ∆

H

– means “change in enthalpy”

3) ∆

H

=

H

products

H

reactants

4)

q

=

Δ

H

Slide14

b. Enthalpy of Reaction

1

) Quantity of energy transferred as heat during a chemical

reaction

2) Difference b/w stored energy of reactants and products

3) AKA: “heat of reaction”

Slide15

c. Enthalpy of Reaction – Exothermic Reactions

1) Chemical equations don’t typically show energy release

2) Thermochemical equation

a) Equation that includes the quantity of energy released or

absorbed as heat during the reaction

b) 2 H

2

(g) + O

2

(g)

2 H

2

O + 483.6 kJ

3) Reactants have larger enthalpy than products

Slide16

d. Enthalpy of Reaction – Endothermic Reactions

1) Products have larger enthalpy than reactants

2) 2 H

2

O (g) + 483.6 kJ

2 H

2

(g) + O

2

(g)

Slide17

e. Thermochemical Equations

1) Typically not written with energy as product or reactant

2) Written with

Δ

H

a)

Δ

H

is negative = exothermic (system loses energy)

b)

Δ

H

is positive = endothermic (system gains energy)

3)

2 H

2

(g) + O

2

(g)

2

H

2

O (g)

Δ

H

= -483.6 kJ

4) 2 H

2

O (g)

2 H

2

(g) + O

2

(g

)

Δ

H

= 483.6 kJ

Slide18

5) Endothermic or Exothermic?

a)

CaO

+ H

2

O

Ca(OH)

2

+ 65.2 kJ

b) 2 NaHCO

3

+ 129 kJ

Na

2

CO

3

+ H

2

O + CO

2

c) What is

Δ

H

for a) and b)?

a)

Δ

H

= -65.2 kJ

b)

Δ

H

= 129 kJ

Slide19

5. Enthalpy of Reaction

Examples

a. 25.0 mL of

HCl

and 25.0 mL of

NaOH

, both at 25.0°C

are mixed in a calorimeter. The final temperature is

32.0°C. Calculate

Δ

H in kJ.

Is this reaction exothermic or endothermic? Negative or positive value?

Slide20

What we know

Δ

H = - (m x

c

q

x

Δ

T)

C

q

of water = 4.18 J/(g·°C)

M = 50.0 g

Δ

T = 7.0°C

Δ

H = -(50.0g) 4.18 J/(g·°C) (7.0°C)

Slide21