chapters 6-7a chemical reactions chapter 6 table of contents 2 6.1 evidence for a chemical reaction...

Chapters 6-7a

Chemical Reactions

Chapter 6

Table of Contents

2

6.1 Evidence for a Chemical Reaction

6.2 Chemical Equations

6.3 Balancing Chemical Equations

7.1 Predicting Whether a Reaction Will Occur

7.2 Reactions in Which a Solid Forms

Section 6.1

Evidence for a Chemical Reaction

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• Chemical reactions often give a visual signal.• But reactions are not always visible.

What are the clues that a chemical change has taken place?

Section 6.1


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Copyright © Cengage Learning. All rights reserved 4

Some Clues That a Chemical Reaction Has Occurred

Section 6.1


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Exercise

What is a clue that a chemical reaction has occurred?

a) The color changes.

b) A solid forms.

c) Bubbles are present.

d) A flame is produced.

Section 6.1


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Exercise

What is a clue that a chemical reaction has occurred?“Colorless hydrochloric acid is added to a red solution of cobalt(II) nitrate, turning the solution blue.”

a) The color changes.

b) A solid forms.



Section 6.1


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Exercise

What is a clue that a chemical reaction has occurred?“A solid forms when a solution of sodium dichromate is added to a solution of lead nitrate.”

a) A gas forms.

b) A solid forms.



Section 6.2

Chemical Equations

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• Chemical reactions involve a rearrangement of the ways atoms are grouped together.

• A chemical equation represents a chemical reaction. Reactants are shown to the left of the arrow. Products are shown to the right of the arrow.

Section 6.2

Chemical Equations

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• In a chemical reaction atoms are not created or destroyed.

• All atoms present in the reactants must be accounted for in the products. Same number of each type of atom on both sides of the

arrow.

Section 6.2

Chemical Equations

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10

Taking 20 kids to the zoo?

What if you came home with only 18 kids?

Parents are funny that way!

What if you came home with 22 kids?

At who’s house would you drop them off?

Section 6.2

Chemical Equations

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11

• Unbalanced Equation:

• Balancing the Equation:

• The balanced equation: CH4 + 2O2 CO2 + 2H2O

Balancing a Chemical Equation

Section 6.2

Chemical Equations

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• Physical states of compounds are often given in a chemical equation. These are sometimes called descriptors.

Physical States

Section 6.2

Chemical Equations

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13

Example

Section 6.2

Chemical Equations

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Exercise

When blue light shines on a mixture of hydrogen and chlorine gas, the elements react explosively to form gaseous hydrochloric acid. What is the unbalanced equation for this process?

a) H2(g) + CH4(g) HCl(g)

b) HCl(g) H(g) + Cl(g)

c) H(g) + Cl(g) HCl(g)

d) H2(g) + Cl2(g) HCl(g)

Section 6.3

Balancing Chemical Equations

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• The principle that lies at the heart of the balancing process is that atoms are conserved in a chemical reaction.

• Atoms are neither created nor destroyed.• The same number of each type of atom is

found among the reactants and among the products.

Section 6.3


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• Chemists determine the identity of the reactants and products of a reaction by experimental observation.

• The identities (formulas) of the compounds must never be changed in balancing a chemical equation.

Section 6.3


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1. Read the description of the chemical reaction. What are the reactants, the products, and their states? Write the appropriate formulas.

Hydrogen gas (H2) and oxygen gas (O2) combine to form liquid water (H2O).

2. Write the unbalanced equation that summarizes the information from step 1.

H2(g) + O2(g) H2O(l)

How to Write and Balance Equations

Section 6.3


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3. Balance the equation by inspection, starting with the most complicated molecule.

Equation is unbalanced by counting the atoms on both sides of the arrow.


Section 6.3


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3. Balance the equation by inspection, starting with the most complicated molecule.

We must balance the equation by adding more molecules of reactants and/or

products.


Section 6.3


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4. Check to see that the coefficients used give the same number of each type of atom on both sides of the arrow. Also check to see that the coefficients used are the smallest integers that give the balanced equation.

The balanced equation is:

2H2(g) + O2(g) 2H2O(l)

or could be:

4H2(g) + 2O2(g) 4H2O(l)

preferred

Section 6.3


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Another Balancing Example:

Section 6.3


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Balancing using the underline method.

Na2O(s) + H2O(l) NaOH(aq)2

CH4(g) + O2(g) CO2(g) + H2O(g)22

Fe(s) + O2(g) Fe2O3(s)2 24 3

LiOH(s) + CO2(g) LiHCO3(s)

KClO3(s) KCl(s) + O2(g)

MnO2

32 2

Section 6.3


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Exercise

Balance the following equation in standard form (lowest multiple integers) and determine the sum of the coefficients?

FeO(s) + O2(g) Fe2O3(s)

a) 3

b) 4

c) 7

d) 14

24

Section 6.3


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Exercise

Which of the following correctly balances the chemical equation given below? There may be more than one correct balanced equation. If a balanced equation is incorrect, explain what is incorrect about it.

CaO + C CaC2 + CO2

I. CaO2 + 3C CaC2 + CO2

II. 2CaO + 5C 2CaC2 + CO2

III. CaO + (2.5)C CaC2 + (0.5)CO2

IV. 4CaO + 10C 4CaC2 + 2CO2

Section 6.3


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Exercise

Of the three that are correct, which one is preferred most (the most accepted convention)? Why?

CaO + C CaC2 + CO2

I. CaO2 + 3C CaC2 + CO2

II. 2CaO + 5C 2CaC2 + CO2

III. CaO + (2.5)C CaC2 + (0.5)CO2

IV. 4CaO + 10C 4CaC2 + 2CO2

Section 6.3


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Concept Check

When balancing a chemical equation, which of the following statements is false?

a) Subscripts in the reactants must be conserved in the products.

b) Coefficients are used to balance the atoms on both sides.

c) When one coefficient is doubled, the rest of the coefficients in the balanced equation must also be doubled.

d) Phases are often shown for each compound but are not critical to balancing an equation.

Section 6.3


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• The number of atoms of each type of element must be the same on both sides of a balanced equation.

• Subscripts must not be changed to balance an equation.

• A balanced equation tells us the ratio of the number of molecules which react and are produced in a chemical reaction.

• Coefficients can be fractions, although they are usually given as lowest integer multiples.

Notice

Section 6.3


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1. Formation of a solid

2. Formation of water

3. Transfer of electrons

4. Formation of a gas

Four Driving Forces Favor Chemical Change

Section 7.1

Section 6.3


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Ukrainian Wolves

Partially charged water molecules act like Ukrainian Wolves in desloving fully charged NaCl ion.

http://www.northland.cc.mn.us/biology/Biolog

y1111/animations/dissolve.swf

Section 7.1

Section 6.3


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• A reaction in which a solid forms is called a precipitation reaction. Solid = precipitate

Precipitation

Section 7.2

Section 6.3


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• The ions separate and move around independently.• Strong electrolyte – each unit of the substance that

dissolves in water produces separated ions.

What Happens When an Ionic Compound Dissolves in Water?

Section 7.2

Section 6.3


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• K2CrO4(aq) + Ba(NO3)2(aq) Products

What Happens When an Ionic Compound Dissolves in Water?

Section 7.2

Section 6.3


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• K2CrO4(aq) + Ba(NO3)2(aq) Products

• The mixed solution contains four types of ions: K+, CrO4

2–, Ba2+, and NO3–.

• Determine the possible products from the ions in the reactants. The possible ion combinations are:

How to Decide What Products Form

Section 7.2

Section 6.3


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• Decide which is most likely to be the yellow solid formed in the reaction.

• K2CrO4(aq) reactant

• Ba(NO3)2(aq) reactant

• The possible combinations are KNO3 and BaCrO4.

KNO3 white solid

BaCrO4 yellow solid

How to Decide What Products Form

Section 7.2

Section 6.3


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Using Solubility Rules

Section 7.2

Section 6.3


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Using Solubility Rules

• Predicting Precipitates Soluble solid Insoluble solid Slightly soluble solid

Section 7.2

Section 6.3


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Let’s Practice Determining Solubility

Which of the following are solublesoluble in water?

Na2CO3 yes

CaCl2 yes

AgCl no

BaSO4 no

(NH4)2S yes

Cu(OH)2 no

Ba(OH)2

Ca3(PO4)2no

yes

Pb(NO3)2 yes

PbCl2 no

Section 7.2

Section 6.3


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1. Write the reactants as they actually exist before any reaction occurs. Remember that when a salt dissolves, its ions separate.

2. Consider the various solids that could form. To do this, simply exchange the anions of the added salts.

3. Use the solubility rules to decide whether a solid forms and, if so, to predict the identity of the solid.

How to Predict Precipitates When Solutions of Two Ionic Compounds Are Mixed

Section 7.2

Section 6.3


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Concept Check

Which of the following ions form compounds with Pb2+ that are generally soluble in water?

a) S2–

b) Cl–

c) NO3–

d) SO42–

e) Na+

Section 7.2

Section 6.3


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Concept Check

A sodium phosphate solution reacts with a lead(II) nitrate solution. What precipitate, if any, will form?

a) Pb3(PO4)2

b) NaNO3

c) Pb(NO3)2

d) No precipitate will form.

Section 7.2

Section 6.3


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Concept Check

Consider a solution with the following ions present:

When all are allowed to react (and there is plenty available of each), how many different solids will form? List them.

Five different solids will form.

PbCl2, PbSO4, Pb3(PO4)2, AgCl, Ag3PO4

- 2+ + + - 2- 3-3 4 4NO , Pb , K , Ag , Cl , SO , PO

Section 7.2

Section 6.3


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Chapters 6-7b

Chemical Reactions

Vv

Section 6.3


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7.3 Describing Reactions in Aqueous Solutions

7.4 Reactions That Form Water: Acids and Bases

7.5 Reactions of Metals with Nonmetals (Oxidation–Reduction)

7.6 Ways to Classify Reactions

7.7 Other Ways to Classify Reactions

Section 6.3


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Net Ionic Equations

AgNO3(aq) + NaCl(aq)

1. Divorce Ag++NO3- + Na++ Cl-

2. Change Partners Ag++NO3- + Na++ Cl- AgCl + NaNO3

3. Soluble? Ag++NO3- + Na++ Cl- AgCl(s) + Na+ + NO3

-

4. Cross out Spectator Ions

Ag++ Cl- AgCl(s) 5. Balance

Ag++ Cl- AgCl(s)

Total Ionic Equation

Molecular Equation

Net Ionic Equation

AgCl(s) + NaNO3(aq)

Section 7.2

Section 6.3


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Pb(NO3)2(aq) + NaI(aq)

1. Divorce Pb+2+NO3- + Na++ l-

2. Change Partners Pb+2+NO3- + Na++ l- Pbl2 + NaNO3

3. Soluble? Pb+2+NO3- + Na++ l - Pbl2(s) + Na+ NO3

-


Pb+2+ 2 l- Pbl2(s) 5. Balance

Pb+2+ l- Pbl2(s)


Molecular Equation

Net Ionic Equation

Pbl2(s) + NaNO3(aq)

Section 7.2

Net Ionic Equations

Section 6.3


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BaCl2(aq) + Na2SO4(aq)

1. Divorce Ba+2 + Cl- + Na++ SO4-2

2. Change Partners Ba+2 + Cl- + Na++ SO4-2 BaSO4 + NaCl

3. Soluble? Ba+2 + Cl- + Na++ SO4-2 BaSO4(s)+ Na+ + Cl-


Ba+2+ SO4-2 BaSO4(s)5. Balance

Ba+2+ SO4-2 BaSO4(s)


Molecular Equation

Net Ionic Equation

BaSO4(s) + NaCl(aq)

Section 7.2

Net Ionic Equations

Section 6.3


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1. Molecular Equation Shows the complete formulas of all

reactants and products. It does not give a very clear picture of

what actually occurs in solution.

Types of Equations for Reactions in Aqueous Solutions

Section 7.3

Section 6.3


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2. Complete Ionic Equation All strong electrolytes are shown as ions.

Notice: K+ and NO3– ions are present in solution

both before and after the reaction.


Section 7.3

Section 6.3


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2. Complete Ionic Equation Spectator ions – ions which do not

participate directly in a reaction in solution.


Section 7.3

Section 6.3


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3. Net Ionic Equation Only those components of the solution

that undergo a change.

Notice: Spectator ions are not shown in the net ionic equation.


Section 7.3

Section 6.3


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Concept Check

Write the correct molecular equation, complete ionic equation, and net ionic equation for the reaction between cobalt(II) chloride and sodium hydroxide.

Molecular Equation:

CoCl2(aq) + 2NaOH(aq) Co(OH)2(s) + 2NaCl(aq)

Complete Ionic Equation:

Co2+(aq) + 2Cl(aq) + 2Na+(aq) + 2OH(aq)

Co(OH)2(s) + 2Na+(aq) + 2Cl(aq)

Net Ionic Equation:

Co2+(aq) + 2OH(aq) Co(OH)2(s)

Section 7.3

Section 6.3


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Types of Reactions

1. Combination (one product) A + B C

2. Decompostion (one reactant) A B + C

3. Single Replacement A + BC AC + B

4. Double Replacement AB + CD AD(s) + CB(l)

5. Acid-Base (Neutralization) HA + BOH H2O + BA

6. Combustion- Organic + O2 CO2 + H2O

7. No Reaction- both products are (aq).

Section 6.3


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• A strong acid is one in which virtually every molecule dissociates (ionizes) in water to an H+ ion and an anion.

Arrhenius Acids and Bases

Section 7.4

Section 6.3


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Strong Acids Behave as Strong Electrolytes

Section 7.4

Section 6.3


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• A strong base is a metal hydroxide that is completely soluble in water, giving separate OH ions and cations. Most common examples: NaOH and KOH


Section 7.4

Section 6.3


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• The products of the reaction of a strong acid and a strong base are water and a salt. Salt Ionic compound

• Net ionic equation H+(aq) + OH−(aq) H2O(l)

• Reaction of H+ and OH− is called an acid-base reaction. H+ acidic ion OH− basic ion


Section 7.4

Section 6.3


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1. The common strong acids are aqueous solutions of HCl, HNO3, and H2SO4.

2. A strong acid is a substance that completely dissociates (ionizes) in water (into H+ ions and anions).

3. A strong base is a metal hydroxide compound that is very soluble in water (and dissociates into OH– ions and cations).

Summary of Strong Acids and Strong Bases

Section 7.4

Section 6.3


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4. The net ionic equation for the reaction of a strong acid and a strong base is always the same: it shows the production of water.

5. In the reaction of a strong acid and a strong base, one product is always water and the other is always an ionic compound called a salt, which remains dissolved in the water. This salt can be obtained as a solid by evaporating the water.


Section 7.4

Section 6.3


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6. The reaction of H+ and OH– is often called an acid-base reaction, where H+ is the acidic ion and OH– is the basic ion.


Section 7.4

Section 6.3


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Concept Check

The net ionic equation for the reaction of HNO3 and LiOH is

a) H+ + NO3– + LiOH → H2O + LiNO3

b) HNO3 + LiOH → H2O + LiNO3

c) H+ + OH– → H2O

d) Li+ + NO3– → LiNO3

Section 7.4

Section 6.3


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• Reactions between metals and nonmetals involve a transfer of electrons from the metal to the nonmetal.

• A reaction that involves a transfer of electrons. 2Mg(s) + O2(g) 2MgO(s)

Oxidation–Reduction Reaction

Section 7.5

Section 6.3


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Concept Check

Which of the following best describes what is happening in the following representation of an oxidation–reduction reaction:

a) Metal Al gains 3 e– and O2 – in Fe2O3 loses these 3e–.

b) Metal Al gains 3 e– and Fe3+ in Fe2O3 loses these 3e–.

c) Metal Al loses 3 e– and O2 – in Fe2O3 gains these 3e–.

d) Metal Al loses 3 e– and Fe3+ in Fe2O3 gains these 3e–.

Section 7.4Section 7.5

Section 6.3


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1. A metal–nonmetal reaction can always be assumed to be an oxidation–reduction reaction, which involves electron transfer.

2. Two nonmetals can also undergo an oxidation–reduction reaction. At this point we can recognize these cases only by looking for O2 as a reactant or product. When two nonmetals react, the compound formed is not ionic.

Characteristics of Oxidation–Reduction Reactions


Section 6.3


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• Formation of a solid when two solutions are mixed.

• Notice this is also a double–displacement reaction. AB + CD AD + CB

Precipitation Reaction


Section 6.3


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• Involves an H+ ion that ends up in the product water. H+(aq) + OH−(aq) H2O(l)

HCl(aq) + KOH(aq) H2O(l) + KCl(aq)

Acid–Base Reaction


Section 6.3


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• Transfer of electrons 2Li(s) + F2(g) 2LiF(s)

Oxidation–Reduction Reaction


Section 6.3


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• Oxidation–reduction reaction• Single–replacement reaction

A + BC B + AC

Formation of a Gas


Section 6.3


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• Involve oxygen and produce energy (heat) so rapidly that a flame results. CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

Special class of oxidation–reduction reactions.

Combustion Reactions


Section 6.3


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• A compound forms from simpler materials. C(s) + O2(g) CO2(g)


Synthesis (Combination) Reactions


Only one product!

Section 6.3


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• Occurs when a compound is broken down into simpler substances. 2H2O(l) 2H2(g) + O2(g)


Decomposition Reactions


Only one reactant!

Section 6.3


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Summary


chapters 6-7a chemical reactions chapter 6 table of contents 2 6.1 evidence for a chemical reaction...

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