chapter 13 chemical reactions

Chapter 13:Chemical Reactions

Principles of Science II

This lecture will help you understand:

• Chemical Equations

• Energy and Chemical Reactions• Reaction Rates • Catalysts

• Acids Donate Protons, Bases Accept Them

• Relative Strengths of Acids and Bases

• Acid, Basic, and Neutral Solutions • Acidic Rain and Basic Oceans • Losing and Gaining Electrons • Harnessing the Energy of Flowing Electrons

• Electrolysis

• Corrosion and Combustion

Reactants Products

Chemical Equations

• During a chemical reaction, one or more new compounds are formed as a result of the rearrangement of atoms.

Chemical Equations

• Law of mass conservation: No atoms are gained or lost during any reaction.

• The number of times atoms appear before the arrow must be equal to the number of times atoms appear after the arrow.

• One of the most important principles of chemistry is the law of mass conservation, which states that matter is neither created nor destroyed during a chemical reaction.

– Instead, atoms simply change partners to form new materials.

Chemical Equations

• The number and type of atoms before a reaction, therefore, are always the same as the number and type of atoms after the reaction.

• To reflect this fact, the chemical equation must be BALANCED.

Balancing guidelines

1. Balance one element at a time.

2. If you incidentally unbalance an element, leave it alone.

3. Make successive passes.

Al2O3 C CO2+ Al +

(not balanced)

Chemical Equations

4 32 3

Al Al

OOO C

Al

C OOAl

C OO

C OOAl Al

OOO

C

C

Al

Al

Al2O3 C CO2+ Al +

Chemical Equations

Balanced

Energy and Chemical Reactions

• Bond energy is the amount of energy required to pull two bonded atoms apart, which is the same as the amount of energy released when they are brought together.

• In an EXOTHERMIC reaction there is a net release of energy.

2 H2 O2 2 H2O+


The reaction is EXOTHERMIC.

Energy absorbed:

Energy released:

–486 kJ/molNet energy of reaction:

–1856 kJ/mol

+1370 kJ/mol


Reaction progress

En

ergy

This reaction is exothermic.


• In an ENDOTHERMIC reaction there is a net absorption of energy.

+1444 kJ/mol

–1262 kJ/mol

Energy absorbed:

Energy released:

Net energy of reaction:

N2 O22 NO+


+182 kJ/mol

This reaction is ENDOTHERMIC.


Reaction progress

Ene

rgy

This reaction is endothermic.

Reaction Rates

• The reaction rate is how quickly the concentration of reactants decreases and how quickly the concentration of products increases.

• The rate of a reaction is dependent on the collisions among molecules.

• Increasing the concentration increases the number of collisions per second and therefore increases the rate of the reaction.

Reaction Rates

Reaction Rates

• Not all collisions lead to products. For example, at ambient temperatures, N2 and O2 molecules do not have sufficient energy to form nitrogen monoxide.

• So, where might this reaction occur?

Reaction Rates

• Activation energy, Ea, is the minimum energy required to overcome the initial breaking of bonds in reactants.

Reaction progress

Ene

rgy

Reaction Rates

Reaction Rates

• Reactants must be moving fast enough (have sufficient kinetic energy) to overcome the energy of activation.

Reaction Rates

• Reaction rates are affected by:– Concentration– Temperature– Catalysts

Premise:

Reactant molecules have to make physical contact with each other in order to transform into products.

Reaction Rates

• The more concentrated a sample of nitrogen and oxygen, the greater the likelihood that N2 and O2 molecules will collide and form nitrogen monoxide.

Reaction Rates

• Slow-moving molecules may collide without enough force to break the bonds. In this case, they cannot react to form product molecules.

Catalysts

• A catalyst is a substance that lowers the activation energy of a chemical reaction, which allows for the reaction to proceed at a faster rate.

Reaction progress

Ene

rgy

OzoneO3 Oxygen

O2

2 O3 3 O2

Cl

Reaction progress

Ene

rgy

OzoneO3 Oxygen

O2

2 O3 3 O2

Catalysts

CCl3F CCl2F2

Catalysts

• Chlorofluorocarbons were once commonly used as refrigerants, until it was recognized that these compounds are a major source of chlorine atoms in the stratosphere.

Catalysts

Catalysts

Concentration of chlorinemonoxide, ClO

Concentration of ozone, O3

Catalysts• Chemists have been able to harness the power

of catalysts for numerous beneficial purposes.

Acids Donate Protons, Bases Accept Them

• An acid is a chemical that donates a hydrogen ion, H+.

• A base is a chemical that accepts a hydrogen ion, H+.

H3O+Cl–H2OHCl + +


(acid) (base) donor

acceptor

NH4+OH– +H2O NH3

donor

acceptor

+


(acid) (base)

NH4+OH– +H2O NH3+


(acid) (base) (acid) (base)

A salt is an

Relative Strengths of Acids and Bases

• Strong acids and bases ionize completely in water.

Relative Strengths of Acids and Bases

• Weak acids and bases do not ionize completely in water.

Acidic, Basic, and Neutral Solutions

• Water can behave as an acid or a base.


• In pure water, for every one hydronium ion, H3O+, formed, there is one hydroxide ion, OH–, formed.

• So, in pure water, [H3O+] = [OH–] = 0.0000001 M = 10–7 M.


• Add hydronium ions, H3O+, and the solution is "acidic."

• Add hydroxide ions, OH–, and the solution is "basic."


• pH is a measure of the concentration of hydronium ions, H3O+:

pH = –log [H3O+]

• For pure water:pH = –log (10–7)pH = –(– 7)pH = 7


• The "log" of a number is simply the power to which 10 is raised. The log of 103, for example, is 3.

Quiz Time

What is the log of 102?

Log 102 = 2

(the power to which 10 is raised)


103 = 1000

102 = 100

101 = 10

100 = 1

10–1 = 0.1

10–2 = 0.01

10–3 = 0.001


pH = –log [H3O+]

For acidic water pH < 7; for example:

pH = –log (10–5)

pH = –(–5)

pH = 5


pH = –log [H3O+]

For basic water pH > 7; for example:

pH = –log (10–9)

pH = –(–9)

pH = 9

Acidic Rain and Basic Oceans

• Acid rain has a pH lower than 5.

2 SO2(g) + O2(g) SO3(g)

SO3(g) + H2O(l) H2SO4(aq)

• SO2 is released from burning coal and oil.


• CO2 levels in the atmosphere are rising.

CO2(g) + H2O(l) H2CO3(aq)

Losing and Gaining Electrons

• Acid–base reactions involve the transfer of a proton.

• Oxidation–reduction reactions involve the transfer of an electron.


• Oxidation is the loss of an electron.

2 Na 2 Na+ + 2 e–

• Reduction is the gain of an electron.

Cl2 + 2 e– 2 Cl–


2 Na + Cl2 2 Na+ + 2 Cl–

Harnessing the Energy of Flowing Electrons

• Electric currents are generated by oxidation– reduction reactions.

• Oxidation–reduction reactions are used in batteries and fuel cells.


• Ions must be able to flow in order to generate a current.


• A salt bridge allows this to happen.



• A battery is a self-contained voltaic cell.


• The positive electrode is the cathode, where reduction occurs.

• The negative electrode is the anode, where oxidation occurs.


• Several types of batteries:– Dry cell– Alkaline– Rechargeable

• NiMH• Lithium


• Fuel cells convert the chemical energy of a fuel into electric energy.

Electrolysis

• Electrolysis uses electric energy to produce chemical change.

• Examples:– Recharging your car battery– Purifying metal ores

Electrolysis

Corrosion and Combustion

• Corrosion is the process in which a metal deteriorates through oxidation–reduction reactions.

• Corrosion can be prevented by coating the metal with zinc, which oxidizes first.

Corrosion and Combustion

• Combustion is an oxidation–reduction reaction between a nonmetallic material, such as wood, and oxygen.

chapter 13 chemical reactions

Science

kjmolnet energy of reaction

chemical reactionsbond

chemical reactionsthe

endothermic reaction

exothermic reaction

sufficient energy

energy of activation

minimum energy