Reactions Unit

Chemistry

Chemistry

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Learning Objectives: Reactions

Essential knowledge and skills:

Classify types of chemical reactions as synthesis, decomposition, single replacement, double replacement, neutralization, and/or combustion.

Transform word equations into chemical equations and balance chemical equations. Recognize that there is a natural tendency for systems to move in a direction of randomness (entropy). Recognize equations for redox reactions and neutralization reactions. Determine the number of valence electrons and possible oxidation numbers from an element’s electron configuration.

Essential understandings:

Conservation of matter is represented in balanced chemical equations. A coefficient is a quantity that precedes a reactant or product formula in a chemical equation and indicates the relative number of particles involved in the reaction.

Major types of chemical reactions aresynthesis (A+B → AB)decomposition (BC → B+C)single replacement (A+BC→ B+AC)double replacement (AC+BD → AD+BC)neutralization (HX+MOH → H2O + MX)combustion (CxHy + O2 → CO2 + H2O).

Transition metals can have multiple oxidation states. Reactivity is the tendency of an element to enter into a chemical reaction.

Exothermic and Endothermic Reactions

Almost all chemical and physical reactions involve energy (usually in the form of heat) being released or added.  An exothermic change is a reaction that releases energy.  An endothermic change is one in which the energy must be added for the reaction to occur. For exothermic reactions, energy can be thought of as a product in the reaction. For endothermic changes, energy can be thought of as a reactant in the reaction.

If a chemical reaction occurs at constant pressure, as all of our chemical reactions do we can consider a property called enthalpy.  Enthalpy (H) is the energy (heat) content of a system at constant pressure.  You cannot measure the actual energy or enthalpy of a substance, but you can measure the change in enthalpy for a reaction.  This change is symbolized by ∆Hrxn. 

For exothermic reactions, enthalpy values are always negative, that is the energy of the products is lower than that of the reactants. This is because energy is released as new bonds are formed in the products and this amount of energy is greater than the energy required to break the old bonds in the reactants. 

∆Hrxn =  Hproducts - Hreactants (small # -  BIG#) =  - negative # 

For example: 4Fe   +   3O2     2Fe2O3    ∆Hrxn = -1625 kJ 

For endothermic reactions, enthalpy values are always positive, that is that energy of the products is greater than that of the reactions. This is because the energy released as new bonds are formed in the products is less than the energy required to break the bonds in the reactants.  This energy must be supplied in order for the reaction to occur.  The added energy does not disappear, of course due to the Law of Conservation of Energy. Instead, it becomes stored in the chemical bonds of the products. 

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∆Hrxn =  Hproducts - Hreactants (BIG# - small #) =  positive # 

For example: C +  H2O    CO  + H2      ∆Hrxn = + 113 kJ 

Complete the following chart:

Type of Reaction Sign of ∆Hrxn Which has more energy: reactants or products?

Exothermic  

Endothermic  

Reactivity Series

Entropy (∆S)

What is Entropy3

Entropy is the degree of randomness or disorder in a substance. The symbol for change in entropy is ΔS.

Solids are very ordered and have low entropy. Liquids and aqueous ions have more entropy because they move about more freely, and gases have an even larger amount of entropy. According to the Second Law of Thermodynamics, nature is always proceeding to a state of higher entropy (more disorder). +ΔS means more disorder. –ΔS means more order or less disorder.

Tendency in nature to move toward a state of higher disorder or randomness (Your bedroom gets messy after a couple of days of neglect!)

The Second Law of Thermodynamics states that there is an inherent direction in which processes occur. This direction is towards a state of higher entropy (more disorder.)

For example, an egg falls to the floor and cracks, but it never falls back up and puts itself together.

Several factors can be assessed to determine if a chemical or physical reaction is likely to occur because of increased entropy.

a) Phase Changes:

Solid (great order, low entropy) Liquid (more randomness, higher entropy) Gas (max randomness, highest entropy)

b) Physical Changes:

When a substance is divided in to parts entropy increases

A large crystal is broken in to smaller pieces A solid is dissolved and dissociates

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c) Chemical Changes:

When there are more products than reactants in a chemical reaction, entropy increases. When the products of a reaction are simpler than the reactants entropy is increased. Compounds are much less random than their constituent elements

d) Temperature Changes:

When temperature increases, molecules move faster, increasing disorder and entropy.

Ex) Which direction is this reaction more likely to go in, and why?

2KClO3 (s)2KCl (s) + 3O2 (g) + 875 kJ

This reaction would go to the right because:

1) The Products are simpler increased entropy

2) There are more Products increased entropy

3) Reaction leads to a Gas increased entropy

4) Exothermic Reaction decreased enthalpy

Entropy Worksheet 1

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Determine whether the following reactions show an increase or decrease in entropy and write + or – to indicate this.

1. 2KClO3(s) → 2KCl(s) + 2O2(g) __________

2. H2O(l) → H2O(s) __________

3. N2(g) + 3H2(g) → 2NH3(g) __________

4. NaCl(s) → Na+(aq) + Cl-(aq) __________

5. KCl(s) → KCl(l) __________

6. CO2(s) → CO2(g) __________

7. H+(aq) + C2H3O2-(aq) → HC2H3O2(l) __________

8. C(s) + O2(g) → CO2(g) __________

9. H2(g) + Cl2(g) → 2HCl(g) __________

10. Ag+ + Cl-(aq) → AgCl(s) __________

11. 2N2O5(g) →4NO2(g) + O2(g) __________

12. 2Al(s) + 3I2(s) → 2AlI3(s) __________

13. H+(aq) + OH-(aq) → H2O(l) __________

14. 2NO(g) →N2(g) + O2(g) __________

15. H2O(g) → H2O(l) __________

Entropy Worksheet 2

1. Predict whether the entropy change will be positive or negative for the following:a. H2O (g) H2O (l) S_____b. C6H12O6(s) 2C2H5OH(l) + 2CO2(g) S_____c. 2NH3(g) + CO2(g) H2O(l) + NH2CONH2(aq) S____d. NaCl(s) NaCl(aq) S____e. Cu(s) (100oC) Cu(s) (25oC) S____f. 2NH3(g) N2(g) + 3H2(g) S____

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2. Which of the following reactions will have an increase in entropy? Choose all that apply.a. SO3(g) → 2SO2(g) + O2(g)b. H2O(l) → H2O(s)c. Br2(l) → Br2(g)d. H2O2(l) → H2O(l) + ½ O2(g)

3. Which of the following compounds has the lowest entropy at 25 oC?a. CH3OH(l)b. CO(g)c. MgCO3(s)d. H2O(l)e. H2O(g)

Balancing Chemical Equations

How to balance chemical equations

Tips and Tricks!

A chemical equation describes what happens in a chemical reaction. The equation identifies the reactants (starting materials) and products (resulting substance), the formulas of the participants, the phases of the participants (solid, liquid, gas), and the amount of each substance. Balancing a chemical equation refers to establishing the mathematical relationship between the quantity of reactants and products. The quantities are expressed as grams or moles.

It takes practice to be able to write balanced equations. There are essentially three steps to the process:

1. Write the unbalanced equation. Chemical formulas of reactants are listed on the lefthand side of the equation. Products are listed on the righthand side of the equation. Reactants and products are separated by putting an arrow between them to show the direction of the

reaction. Reactions at equilibrium will have arrows facing both directions.

2. Balance the equation. Apply the Law of Conservation of Mass to get the same number of atoms of every element on each

side of the equation. Tip: Start by balancing an element that appears in only one reactant and product. Once one element is balanced, proceed to balance another, and another, until all elements are

balanced. Balance chemical formulas by placing coefficients in front of them. Do not add subscripts, because

this will change the formulas.

3. Indicate the states of matter of the reactants and products. Use (g) for gaseous substances.

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Use (s) for solids. Use (l) for liquids. Use (aq) for species in solution in water. Write the state of matter immediately following the formula of the substance it describes.

Worked Example Problem Tin oxide is heated with hydrogen gas to form tin metal and water vapor. Write the balanced equation that describes this reaction.

1. Write the unbalanced equation.

SnO2 + H2 → Sn + H2O

2. Balance the equation. Look at the equation and see which elements are not balanced. In this case, there are two oxygen atoms on the lefthand side of the equation and only one on the righthand side. Correct this by putting a coefficient of 2 in front of water:

SnO2 + H2 → Sn + 2 H2O

This puts the hydrogen atoms out of balance. Now there are two hydrogen atoms on the left and four hydrogen atoms on the right. To get four hydrogen atoms on the right, add a coefficient of 2 for the hydrogen gas. Remember, coefficients are multipliers, so if we write 2 H2O it denotes 2x2=4 hydrogen atoms and 2x1=2 oxygen atoms.

SnO2 + 2 H2 → Sn + 2 H2O

The equation is now balanced. Be sure to double-check your math! Each side of the equation has 1 atom of Sn, 2 atoms of O, and 4 atoms of H.

3. Indicate the physical states of the reactants and products. To do this, you need to be familiar with the properties of various compounds or you need to be told what the phases are for the chemicals in the reaction. Oxides are solids, hydrogen forms a diatomic gas, tin is a solid, and the term 'water vapor' indicates that water is in the gas phase: SnO2(s) + 2 H2(g) → Sn(s) + 2 H2O(g)

Balancing Chemical EquationsOne of the most useful devices for communicating information related to chemical changes is the chemical equation. The equation contains both qualitative and quantitative information related to the nature and quantity of the substances involved in the chemical reaction. It may also include the energy change involved.

Atoms are fundamental building blocks of all matter. For the purpose of equation balancing we say that they can be neither created nor destroyed. Thus the number of atoms at the beginning of a reaction (reactants - left side of the equation) must equal the number of atoms at the end of the reaction (products - right side of the equation). Note that the number of atoms on each side of an equation must balance, therefore, the mass (number

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of grams) must balance, but not the number of molecules nor the volume of gases.The subscripts in a correct formula tell the number of atoms in one molecule. The coefficients (numbers

in front of a formula) in a correctly balanced equation tell the number of molecules involved in a reaction.There is a particular order that you can follow in balancing. It is the MINOH method which is very

simple to use by inspection. Where:

M - Metals. Balance metals such as Fe or Na first.I - Ions. Looks for polyatomic ions (such as PO4

3- or SO42- that cross from reactant to

product unchanged. Balance them as a group.

N - Non-metals. Look for Cl or S, these are common ones.O - Oxygen Remember, oxygen by itself is O2

H - Hydrogen. Remember, hydrogen by itself is H2

Often, balancing H and O will involve water on one side or the other. In some cases when balancing. you might want to write water as HOH, instead of H2O. Also, look carefully for elements which occur in only one place on each side of the arrow. These should be balanced before examining elements that are spread over several compounds. Often, either H or O will be spread out over several compounds. This is the one to leave to the last. Remember, you cannot change a subscript to balance the equation, nor can you add in new compounds.

Finally, when an equation is difficult to balance, use the grid method covered in class. Look for repeating numbers in your grid and trying using these numbers as coefficients to balance the equations. Remember, 2-3 and 3-3 combinations are very common due to the polyatomic anions having a 2- or 3- oxidation number and polyvalent cations having a +2 or +3 oxidation number.

K2CrO4 + Fe(NO3)3 Fe2(CrO4)3 + KNO3

Reactants Products

K 2 1CrO4 1 3

Fe 1 2NO3 3 1

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Reaction Classification

Chemical Reactions

Types of chemical reactions

There are 5 general types of reactions and two that are special cases of the following1. Decomposition2. Synthesis3. Double displacement4. Single displacement 5. Combustion

Special cases: Redox and Neutralization

Decomposition

Decomposition = one compound two (or more pieces).

AB A + B

Pieces can be elements or simpler compounds i. Element examples:

1. HgO Hg + O2

2. H2O H2 + O2

3. MgCl2 Mg + Cl2

4. FeS Fe + Sii. Simpler compound examples

1. CaCO3 CaO + CO2

2. Na2CO3 Na2O + CO2

3. KClO3 KCl + O2

4. Ba(ClO3)2 BaCl2 + O2 10

iii. Acids and bases1. (base) Ca(OH)2 CaO + H2O2. (base) NaOH Na2O + H2O3. (acid) HNO3 N2O5 + H2O4. (acid) H3PO4 P2O5 + H2O

Notice how, in every case so far, there is only one substance on the left-hand (reactant) side. This is always the case in a decomposition reaction.

Single Replacement/Displacement

Single displacement, one element replaces another element in a compound. One reactant is always an element. It does not matter if the element is written first or second on the

reactant side. The other reactant will be a compound. Two possibilities:

1. Cations switch.

AX + Y YX + A

Element Y replaced A (in the compound AX) to form a new compound YX and the free element A. Remember that A and Y are both cations (postively-charged ions) in this example.

Examples

1. Cu + AgNO3 Ag + Cu(NO3)2

2. Fe + Cu(NO3)2 Fe(NO3)2 + Cu3. Ca + H2O Ca(OH)2 + H2

4. Zn + HCl ZnCl2 + H2

2. Anions switch places:

A + XY XA + Y

Element A has replaced Y (in the compound XY) to form a new compound XA and the free element Y. Remember that A and Y are both anions (negatively-charged ions) in this example.

Examples

1. Cl2 + NaBr NaCl + Br2

2. Br2 + KI KBr + I2

Double Replacement/Displacement

Double displacement, the cations and anions of two different compounds switch places. Both reactants are compounds, each with a cation part and an anion part. Diatomic elements do not count; they are included in the single displacement category.

AB + XY AY + XB

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A and X are the cations (postively-charged ions) in this example, with B and Y being the anions (negatively-charged ions).

Examples:

1. KOH + H2SO4 K2SO4 + H2O2. FeS + HCl FeCl2 + H2S3. NaCl + H2SO4 Na2SO4 + HCl4. AgNO3 + NaCl AgCl + NaNO3

Synthesis

Synthesis are, at this introductory level, almost always the reverse of a decomposition reaction. Two pieces one, more complex compound. Complex means the product compound has more atoms than the reactant molecules. Usually!! Pieces can be elements or simpler compounds.

A + B AB

Examples: two elements are combining1. Mg + O2 MgO2. H2 + O2 H2O3. K + Cl2 KCl4. Fe + O2 Fe2O3

Examples: two compounds making a more complex compound (or a compound and an element joining together):

1. CaO + CO2 CaCO3

2. Na2O + CO2 Na2CO33. KCl + O2 KClO3

4. BaCl2 + O2 Ba(ClO3)2

Combustion

Combustion, at its most general, can mean the reaction of oxygen gas (O2) with anything. However, we define combustion as the reaction of oxygen with a compound containing carbon and

hydrogen. A common synonym for combustion is burn.

CxHy + O2 CO2 + H2O

Examples: 1. CH4 + O2 CO2 + H2O2. C2H6 + O2 CO2 + H2O3. C6H12O6 + O2 CO2 + H2O4. C2H5OH + O2 CO2 + H2O

Notice that some compounds contain carbon, hydrogen AND oxygen. The products are all the same, in every reaction. Variations include NO2 and SO2 Like this:

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1. C21H24N2O4 + O2 CO2 + H2O + NO2

2. C2H5SH + O2 CO2 + H2O + SO2

There are complexities with combustion as you get deeper into it. o i.e. Not enough O2 CO instead of CO2

Neutralisation

Neutralisation reactions are reactions involving and acid and a base. When an acid and a base mix in the right proportion, they neutralize each other The product is a salt and water – a neutral solution when measured with the pH scale

o Examples: NaOH + HCl NaCl + H2O H2SO4 + 2 KOH K2SO4 + 2 H2O

Neutralization is a special case of double replacement reactions

Redox

Simple definitions of oxidation and reduction are based on the loss/gain of oxygen or the loss/gain of hydrogen. Oxidation is the gain of oxygen or the loss of hydrogen; reduction is the loss of oxygen or the gain of hydrogen. These definitions can only be used when a chemical reaction involves hydrogen and oxygen, and therefore their usefulness is limited.

A more basic and more useful definition of oxidation and reduction is based on the loss/gain of electrons.

OXIDATION IS LOSS OF ELECTRONS

REDUCTION IS GAIN OF ELECTRONS

In reactions involving simple ions, it is usually easy to tell whether electrons are lost or gained, but it is less easy to tell when complex ions or covalent molecules are involved. Oxidation number is a useful concept for helping to decide in these more awkward cases.

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OIL RIG

Balancing Chemical Equations Worksheet

Balance the equations below:

1) ____ N2 + ____ H2 ____ NH3

2) ____ KClO3 ____ KCl + ____ O2

3) ____ NaCl + ____ F2 ____ NaF + ____ Cl2

4) ____ H2 + ____ O2 ____ H2O

5) ____ Pb(OH)2 + ____ HCl ____ H2O + ____ PbCl2

6) ____ AlBr3 + ____ K2SO4 ____ KBr + ____ Al2(SO4)3

7) ____ CH4 + ____ O2 ____ CO2 + ____ H2O

8) ____ C3H8 + ____ O2 ____ CO2 + ____ H2O

9) ____ C8H18 + ____ O2 ____ CO2 + ____ H2O

10) ____ FeCl3 + ____ NaOH ____ Fe(OH)3 + ____NaCl

11) ____ P + ____O2 ____P2O5

12) ____ Na + ____ H2O ____ NaOH + ____H2

13) ____ Ag2O ____ Ag + ____O2

14) ____ S8 + ____O2 ____ SO3

15) ____ CO2 + ____ H2O ____ C6H12O6 + ____O2

16) ____ K + ____ MgBr2 ____ KBr + ____ Mg

17) ____ HCl + ____ CaCO3 ____ CaCl2 + ____H2O + ____ CO2

18) ____ HNO3 + ____ NaHCO3 ____ NaNO3 + ____ H2O + ____ CO2

19) ____ H2O + ____ O2 ____ H2O2

20) ____ NaBr + ____ CaF2 ____ NaF + ____ CaBr2

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21) ____ H2SO4 + ____ NaNO2 ____ HNO2 + ____ Na2SO4

Identify the type of reactionFor the following reactions, indicate whether the following are examples of synthesis, decomposition, combustion, single displacement, double displacement, or acid-base reactions:

1) Na3PO4 + 3 KOH 3 NaOH + K3PO4 _________________________

2) MgCl2 + Li2CO3 MgCO3 + 2 LiCl _________________________

3) C6H12 + 9 O2 6 CO2 + 6 H2O _________________________

4) Pb + FeSO4 PbSO4 + Fe _________________________

5) CaCO3 CaO + CO2 _________________________

6) P4 + 3O2 2 P2O3 _________________________

7) 2 RbNO3 + BeF2 Be(NO3)2 + 2 RbF ________________________

8) 2 AgNO3 + Cu Cu(NO3)2 + 2 Ag ________________________

9) C3H6O + 4 O2 3 CO2 + 3 H2O _________________________

10) 2 C5H5 + Fe Fe(C5H5)2 _________________________

11) SeCl6 + O2 SeO2 + 3Cl2 _________________________

12) 2 MgI2 + Mn(SO3)2 2 MgSO3 + MnI4 _________________________

13) O3 O + O2 _________________________

14) 2 NO2 2 O2 + N2_________________________

Section 2: Practicing equation balancing

1) __ C6H6 + __ O2 __ H2O + __ CO2

2) __ NaI + __ Pb(SO4)2 __ PbI4 + __ Na2SO4

3) __ NH3 + __ O2 __ NO + __ H2O

4) __ Fe(OH)3 __ Fe2O3 + __ H2O

5) __ HNO3 + __ Mg(OH)2 __H2O + __ Mg(NO3)2

6) __ H3PO4 + __ NaBr __ HBr + __ Na3PO4

7) __ C + __ H2 __ C3H8

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8) __ CaO + __ MnI4 __ MnO2 + __ CaI2

9) __ Fe2O3 + __ H2O __ Fe(OH)3

10) __ C2H2 + __ H2 __ C2H6

11) __ VF5 + __ HI __ V2I10 + __ HF

12) __ OsO4 + __ PtCl4 __ PtO2 + __ OsCl8

13) __ CF4 + __ Br2 __ CBr4 + __ F2

14) __ Hg2I2 + __ O2 __ Hg2O + __ I2

15) __ Y(NO3)2 + __ GaPO4 __ YPO4 + __ Ga(NO3)2

Section 3: Predicting the products of chemical reactions

Predict the products of the following reactions:

1) __ Ag + __CuSO4

Type:___________________________

2) __ NaI + __ CaCl2

Type:___________________________

3) __ O2 + __ H2

Type:___________________________

4) __ HNO3 + __ Mn(OH)2

Type:___________________________

5) __ AgNO2 + __ BaSO4

Type:___________________________

6) __ HCN + __ CuSO4

Type:___________________________

7) __ H2O + __ AgI

Type:___________________________

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8) __ HNO3 + __Fe(OH)3

Type:___________________________

9) __ LiBr + __ CoSO3

Type:___________________________

10) __ LiNO3 + __Ag

Type:___________________________

11) __ N2 + __ O2

Type:___________________________

12) __ H2CO3

Type:___________________________

Classifying Reactions

Balance the following equations. Then classify the reactions as synthesis (S), decomposition (D), single replacement (SR), double replacement (DR), or combustion (C). Write the corresponding letter(s) in the blank on the left.

_____1. ____ Ca(OH)2 + ____ H2SO4 ____ CaSO4 + ____ H2O

_____2. ____ KClO3 ____ KCl + ____ O2

_____3. ____ HCl + ____ NaOH ____ NaCl + ____ H2O

_____4. ____ Mg + ____ HCl ____ MgCl2 + ____ H2

_____5. ____ N2 + ____ O2 ____ N2O5

_____6. ____ Al + ____ O2 ____ Al2O3

_____7. ____ Al + ____ NiBr2 ____ AlBr3 + ____ Ni

_____8. ____ NaCl ____ Na + ____ Cl2

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_____9. ____ CaCl2 + ____ F2 ____ CaF2 + ____ Cl2

_____10. ____ (NH4)2SO4 + ____ Ba(NO3)2 ____ BaSO4 + ____ NH4NO3

_____11. ____ H2(g) + ____ O2(g) ____ H2O (g)

_____12. ____ H2O(l) ____ H2(g) + ____ O2(g)

_____13. ____ Zn (s) + ____ H2SO4(aq) ____ ZnSO4(aq) + ____ H2 (g)

_____14. ____ CO + ____ O2 ____ CO2

_____15. ____ HgO ____ Hg + ____ O2

_____16. ____ KBr + ____ Cl2 ____ KCl + ____ Br2

_____17. ____ CaO + ____ H2O ____ Ca(OH)2

_____18. ____ AgNO3 (aq) + ____ NaCl(aq) ____ AgCl (s) + ____ NaNO3 (aq)

_____19. ____ C4H8 (g) + ____ O2 (g) ____ CO2 (g) + ____ H2O (g)

_____20. ____ H2O2 (l) ____ H2O (g) + ____ O2 (g)

Write and balance equations for the following reactions AND classify them as synthesis (S), decomposition (D), single replacement (SR), double replacement (DR), or combustion (C). Write the corresponding letter(s) in the appropriate blank.

_____21. aluminum nitrate (aq) + sodium hydroxide (aq) aluminum hydroxide (s) + sodium

nitrate (aq)

_____22. sulfur trioxide (g) sulfur dioxide (g) + oxygen (g)

_____23. iron (s) + silver acetate (aq) iron (II) acetate (aq) + ____ silver (s)

_____24. magnesium (s) + oxygen (g) magnesium oxide (s)

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_____25. ethanol (C2H5OH) (l) + oxygen (g) carbon dioxide (g) + water (g)

Solubility Rules

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Solubility Worksheet

Chemical Formula Name Solubility

1. NH4C2H3O2

2. Ba(OH)2

3. Iron (II) Carbonate

4. NaOH

5. RbNO3

6. Cesium Sulfate

7. MgSO4

8. ZnCl2

9. Zinc Hydroxide

10. Zn3(PO4)2

11. AgBr

12. KNO3

13. Al2S3

14. Silver Acetate

15. Sr2CrO4

16. Aluminium Phosphate

17. BaSO4

18. Ca(OH)2

19. BaCO3

20. MgCrO4

21. Lead (II) Chloride

22. NH4CN

23. Silver Iodide

24. Hg2SO4

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25. Lithium Chloride

Net ionic equations

How to write net ionic equations

EXAMPLE: KCl(aq) + Pb(NO3)2(aq)

1. a. Take only one of the first cation(s) and match it with one of the second anion(s). (Write the cation first)

b. Take only one of the second cation(s) and match it with one of the first anion(s). (Write the cation first)

KCl(aq) + Pb(NO3)2(aq) KNO3 +PbCl

2. Correct the formulas of the products based on the charges of the ions.

KCl(aq) + Pb(NO3)2(aq) KNO3 +PbCl2 ◄

3. Balance the equation

2 KCl(aq) + Pb(NO3)2(aq) 2 KNO3 +PbCl2

4. Consult the solubility rules and assign the correct state symbol annotations. This should agree with any observations concerning the formation of a precipitate which gets the symbol (s). If water is formed, water is a molecule; it does not ionize to any significant extent. It is annotated (l).

2 KCl(aq) + Pb(NO3)2(aq) 2 KNO3(aq) ◄ +PbCl2(s) ◄

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5. Write the Total Ionic Equation (T.I.E.). All compounds that are annotated (aq) break up into individual cations and anions in that order.

2 K+(aq) + 2Cl-

(aq) + Pb2+(aq) + 2NO3

- (aq) 2 K+

(aq) +2 NO3-(aq) +PbCl2(s)

6. Eliminate spectator ions. Spectator ions are in the same form on each side of the equation arrow.

2 K+(aq) + 2Cl-

(aq) + Pb2+(aq) + 2NO3

- (aq) 2 K+

(aq) +2 NO3-(aq) +PbCl2(s)

7. Write the Net Ionic Equation (N.I.E.). The convention is to write the cation first followed by the anion on the “reactants” side.

Pb2+(aq) + 2Cl-

(aq) +PbCl2(s)

Note that chemical equations are written using the lowest common coefficients.

If your NIE ended up as 2 H+(aq) + 2 OH-

(aq) 2 H2O(l)

It would become H+(aq) + OH-

(aq) H2O(l)

Net Ionic Equation Worksheet

Write the complete and net ionic equations for the following reactions.

If only the reactants are given, predict the products and balance the equation first.  You also must include the states of matter.

1.  Pb(NO3)2 (aq)   +   2 KCl (aq)         PbCl2 (s)   +  2 KNO3 (aq)

Complete Ionic:

Net Ionic:

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2.  2 HCl (aq)  +  Ba(OH)2 (aq)    BaCl2 (aq)  +  2 H2O (l)

Complete Ionic:

Net Ionic:

3. _____ K3PO4 (aq)  +  _____ Al(NO3)3 (aq)  

Complete Ionic:

Net Ionic:

4.  _____ Cr(NO3)3 (aq)   +  _____ Ba (s)      

Complete Ionic:

Net Ionic:

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5.  beryllium iodide + strontium sulfate -- >

Molecular Equation:

Complete Ionic:

Net Ionic:

6.  zinc + water

Molecular Equation:

Complete Ionic:

Net Ionic:

7.  barium hydroxide + sulfuric acid

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Molecular Equation:

Complete Ionic:

Net Ionic:

One way of accounting for electrons in equations is to use OXIDATION NUMBERS.

e.g. Fe2+ needs to gain two electrons for it to become neutral iron atom therefore its oxidation number is +2.

Rules for Assigning Oxidation Numbers

This is a prioritized list. If two rules contradict each other, follow the rule that appears higher on the list.

1. The atoms in pure elements are assigned an oxidation number of zero.

2. Monatomic ions are assigned an oxidation number equal to their charge.

3. For atoms in covalent molecules and polyatomic ions:

a. The sum of all the oxidation numbers of the atoms in a covalent molecule must equal zero. The sum of all the oxidation numbers of the atoms in a polyatomic ion must equal the charge on the ion.

b. Fluorine is assigned an oxidation number of –1.

c. Oxygen is assigned an oxidation number of –2 (an exception to this is when oxygen occurs as the peroxide ion, O2

-2, where it is assigned an oxidation number of –1).

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Oxidation numberThe oxidation number of an atom shows the number of electrons which it has lost or gained as a result of forming a compound

d. Hydrogen is assigned an oxidation number of +1 apart from metal hydrides which have a value of -1

Examples

1. The oxidation number of S in H2SO4

H2 S O4

2 x +1 ? 4 x -2 = 0+2 ? -8 = 0+2 +6 -8 = 0

s = +6

2. The oxidation number of S in S2O82-

S2 O4

? 8 x -2 = -2? -16 = -2+14 -16 = -2S = +7

3. The oxidation number of Cl in NaClO3.

Na Cl O3

+1 ? 3 x -2 = 0+1 ? -6 = 0+1 +5 -6 = 0

Cl = +5

4. The oxidation number of Mn in MnO4-

Mn O4

? 4 x -2 = -1? -8 = -1+7 -8 = -1Mn = +7

How to assign oxidation numbers/states

Another tutorial on oxidation numbers/statesOxidation Numbers Worksheet

Rules for Assigning Oxidation Numbers

1. The oxidation number of any uncombined element is 0.

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2. The oxidation number of a monatomic ion equals the charge on the ion.3. The more-electronegative element in a binary compound is assigned the number equal to the charge it

would have if it were an ion.4. The oxidation number of fluorine in a compound is always -1.5. Oxygen has an oxidation number of -2 unless it is combined with F (when it is +2), or it is in a peroxide

(such as H2O2 or Na2O2), when it is -1.6. The oxidation state of hydrogen in most of its compounds is +1 unless it is combined with a metal, in

which case it is -1.7. In compounds, the elements of groups 1 and 2 as well as aluminum have oxidation numbers of +1, +2,

and +3 respectively.8. The sum of the oxidation numbers of all atoms in a neutral compound is 0.9. The sum of the oxidation numbers of all atoms in a polyatomic ion equals the charge of the ion.

Directions: Use the Rules for Assigning Oxidation Numbers to determine the oxidation number assigned to each element in each of the given chemical formulas.

Formula Element and Oxidation Number Formula Element and Oxidation Number

1. Cl2 Cl 16. Na2O2 Na O2. Cl- Cl 17. SiO2 Si O3. Na Na 18. CaCl2 Ca Cl4. Na+ Na 19. PO4

3- P O5. O2 O 20. MnO2 Mn O6. N2 N 21. FeO Fe O7. Al+3 Al 22. Fe2O3 Fe O8. H2O H O 23. H2O2 H O9. NO3

- N O 24. CaO Ca O10. NO2 N O 25. H2S H S11. Cr2O7

2- Cr O 26. H2SO4 H S O12. KCl K Cl 27. NH4Cl N H Cl13. NH3 N H 28. K3PO4 K P O14. CaH2 Ca H 29. HNO3 H N O15. SO4

2- S O 30. KNO2 K N O

Answer Key1. Cl:0 7. Al:+3 13. N:-3 H:

+119. P:+5

O:-225. H:+1 S:-2

2. Cl:-1 8. H:+1 O:-2

14. Ca:+2 H:-1

20. Mn:+4 O:-2

26. H:+1 S:+6 O:-2

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3. Na:0 9. N:+5 O:-2

15. S:+6 O:-2

21. Fe:+2 O:-2

27. N:-3 H:+1 Cl:-1

4. Na:+1 10. N:+4 O:-2

16. Na:+1 O:-1

22. Fe:+3 O:-2

28. K:+1 P:+5 O:-2

5. O:0 11. Cr:+6 O:-2

17. Si:+4 O:-2

23. H:+1 O:-1

29. H:+1 N:+5 O:-2

6. N:0 12. K:+1 Cl:-1

18. Ca:+2 Cl:-1

24. Ca:+2 O:-2

30. K:+1 N:+3 O:-2

Redox

The term REDOX stands for REDUCTION-OXIDATION.

Oxidation can be defined as gain of oxygen or loss of hydrogen. Reduction can be defined as loss of oxygen or gain of hydrogen.

The most important definition is given in terms of electrons.

OXIDATION is LOSS of ELECTRONS REDUCTION is GAIN of ELECTRONS

Using oxidation numbers it is possible to decide whether redox has occurred.

Increase in oxidation number is oxidation. Decrease in oxidation number is reduction.

Redox Reactions

When magnesium is placed into a solution of copper (II) sulfate, a reaction occurs which in simple terms is called a “displacement reaction”.

Chemical equation: Mg + CuSO4 MgSO4 + Cu Ionic equation: Mg(s) + Cu2+

(aq) Mg2+(aq) + Cu(s)

The copper in this reaction is taking electrons from the magnesium.

The copper gains electrons (ox. no. has decreased) - it is REDUCEDThe magnesium loses electrons (increase in ox. no.) - it is OXIDISED

So this is a REDOX reaction.

Whenever one substance gains an electron another substance must lose an electron, so reduction and oxidation always go together.

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Oxidising and reducing reagents

How to predict oxdising and reducing agents

An oxidising agent causes another material to become oxidised. In the above example of adding magnesium to copper sulfate, the magnesium is oxidised. Since the copper ions in the copper sulfate cause this oxidation, they are the oxidising agent. In the same way the Mg causes the reduction of copper ions so it is the reducing agent.

Mg(s) + Cu2+(aq) Mg2+

(aq) + Cu(s)

In this example the oxidising agent (copper ions) is reduced and the reducing agent (magnesium) is oxidised.

This always happens with redox reactions: in a redox reaction the oxidising agent is reduced and the reducing agent is oxidised.

REDUCING AGENT + MATERIAL

Oxidation number and redox reactions

When a redox reaction occurs an electron transfer takes place and so the oxidation numbers of the substances involved changes.Consider the following reaction: 2HOBr + 2H+ + 2I- Br2 + I2 + 2H2O

Reactants ProductsSpecies Ox. No. Species Ox. No.H in HOBr +1 Br in Br2 0O in HOBr -2 I in I2 0Br in HOBr +1 H in H2O +1H+ +1 O in H2O -2I- -1

The table shows us that the oxidation number of Br goes from +1 to 0, so it is reduced. The iodine goes from -1 to 0, so this is oxidised.

Another example 3NaOCl 2NaCl + NaClO3

Reactants Products

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electrons

The reducing agent loses electrons and so is oxidised.

oxidising agentreducing agent

Species Oxid’n No Species Oxid’n NoNa in NaOCl +1 Na in NaCl +1O in NaOCl -2 Na in NaClO3 +1Cl in NaOCl +1 Cl in NaCl -1

Cl in NaClO3 +5O in NaClO3 -2

In this reaction the Cl in NaOCl is oxidised in one reaction to +5 and in another reaction is reduced to -1. Such an occurrence is called disproportionation.

Oxidation Reduction Worksheet 1

1. Determine the oxidation number of each atom in the following substances

a. NF3 N F

b. K2CO3 K C O

c. NO3- N_________ O__________

d. HIO4 H I O

2. For the following balanced redox reaction answer the following questions

2 Fe2+(aq) + H2O2(aq) 2Fe3+(aq) + 2 OH-(aq)

a. What is the oxidation state of oxygen in H2O2?

b. What is the element that is oxidized?

c. What is the element that is reduced?

d. What is the oxidizing agent?

e. What is the reducing agent?

3. For the following balanced redox reaction answer the following questions

4NaOH(aq) + Ca(OH)2(aq) + C(s) + 4ClO2(g) 4NaClO2(aq) + CaCO3(s) + 3H2O(l)

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Disproportionation takes place a particular species undergoes simultaneous oxidation and reduction.

a. What is the oxidation state of Cl in ClO2(g)?

b. What is the oxidation state of C in C(s)? ______

c. What is the element that is oxidized?

d. What is the element that is reduced?

e. What is the oxidizing agent?

f. What is the reducing agent?

4. For the following balanced redox reaction answer the following questions

16 HCl(aq) + 5 SnCl2(aq) + 2 KMnO4(aq) 2 MnCl2(aq) 5 SnCl4(aq) + 8 H2O(l) + 2 KCl(aq)

a. What is the oxidation state of Mn in KMnO4(aq)?

b. What is the oxidation state of Cl in SnCl2(aq)?

c. What is the element that is oxidized?

d. What is the element that is reduced?

e. What is the oxidizing agent?

f. What is the reducing agent?

g. How many electrons are transferred in the reaction as it is balanced?

5. Determine which element is oxidized and which is reduced when lithium reacts with nitrogen to form lithium nitride.

6 Li(s) + N2(g) 2 Li3N(s)

a. element oxidized:

b. element reduced:

6. Determine which atom is oxidized and which is reduced in the following reaction

Sr(s) + 2 H2O(l) Sr2+(aq) + 2 OH-(aq) + H2(g)

a. element oxidized:31

b. element reduced:

Oxidation/Reduction Worksheet 2

1) 4Fe (s) + 3 O2 (g) → 2 Fe2O3(s)

__________was oxidized __________was the oxidizing agent

__________was reduced _________ was the reducing agent

2) Cu(s) + AgNO3(aq) → Ag(s) + CuNO3(aq)

__________was reduced _________ was the reducing agent

__________was oxidized __________was the oxidizing agent

3) 2Na(s) + Cl2(g) → 2NaCl(s)

__________was reduced and was the __________________agent

__________was oxidized and was the __________________agent

4) 2HNO3(aq) + 6HI(aq) → 2NO(g) + 3I2(s) + 4 HOH(l)

__________was oxidized _________ was the reducing agent

__________was reduced __________was the oxidizing agent

Oxidation/reduction Worksheet 3

1. Assign oxidation numbers to chlorine in each of the following chemicals.

HCl(aq), Cl2(g), NaClO(s), Cl–(aq), HClO3(aq),

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ClO3–(aq), KClO2(s), ClO2(g), HClO4(aq)

2. Assign oxidation numbers to manganese in each of the following chemicals.

MnO2(s), KMnO4(s), Mn(s), MnO42–(aq), MnCl2(s),

Mn2O7(s) Mn2+(aq)

For the following reaction equations, use oxidation numbers to identify the oxidation(the atom oxidized) and the reduction (the atom reduced).

3. AsO33–(aq) + IO3

–(aq) AsO43–(aq) + I–(aq)

4. CuO(s) + NH3(g) N2(g) + H2O(l) + Cu(s)

5. MnO4–(aq) + H2Se(g) + H+(aq) Se(s) + Mn2+(aq) + H2O(l)

Polyatomic Ions

AsO43- arsenate

BO33- borate

B4O72- tetraborate

BrO3-1 bromate

CHO2- formate

C2H3O2- acetate

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C2O42- oxalate

C4H4O62- tartrate

ClO- hypochlorite

ClO2- chlorite

ClO3- chlorate

ClO4- perchlorate

CN- cyanide

CO32- carbonate

CrO42- chromate

Cr2O72- dichromate

Fe(CN)63- ferricyanide

Fe(CN)64- ferrocyanide

HCO3- hydrogen

carbonate (bicarbonate)

Hg22+ mercury (I)

H3O+ hydronium

H2PO4- dihydrogen

phosphate

HPO42- hydrogen

phosphate

HSO3- hydrogen sulfite

HSO4- hydrogen sulfate

I3- triiodide

IO3- iodate

MnO4- permanganate

MnO42- manganate

MoO42- molybdate

N3- azide

NH4+ ammonium

NO2- nitrite

NO3- nitrate

O22- peroxide

OCN- cyanate

OH- hydroxide

PO33- phosphite

PO43- phosphate

P2O74- pyrophosphate

SCN- thiocyanate

SeO42- selenate

SiF62- hexafluorosilicate

SiO32- silicate

SO32- sulfite

SO42- sulfate

S2O32- thiosulfate

Reactions Worksheet

Write balanced chemical equations from the following word equations

1. Potassium nitrate Potassium nitrite + oxygen (gas)

2. Zinc + hydrochloric acid Zinc chloride + hydrogen (gas)

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3. Potassium Chlorate Potassium Chloride + Oxygen (gas)

4. Ammonium Nitrate Nitrogen (gas) + oxygen (gas) + Water

5. Calcium Oxide + hydrochloric acid Calcium Chloride + Water

6. Ammonia + Oxygen Nitrogen Monoxide + Water

7. Iron (III) Oxide + Carbon Monoxide Iron + Carbon Dioxide

8. Calcium Oxide + diphosphorous pentoxide Calcium Phosphate

9. Aluminium Hydroxide + acetic acid Aluminium Acetate + Water

10. Aluminium Hydroxide + Cupric Chloride Aluminium Chloride + Copper (II) hydroxide

11. Iron + Silver Acetate Iron (III) Acetate + Silver

12. Bromine + Calcium Iodide Calcium Bromide + Iodine

13. Sodium Hydroxide + Sulfuric Acid Sodium Sulfate + Water

14. Lithium + Water Lithium Hydroxide + Hydrogen

15. Magnesium + Oxygen Magnesium Oxide

16. Mercury (II) Oxide Mercury + Oxygen

17. Fluorine + Potassium Chloride Potassium Fluoride + Chlorine

18. Oxygen + Iron Ferrous oxide

19. Calcium Carbonate Calcium Oxide + Carbon Dioxide

20. Aluminium Oxide Aluminium + Oxygen

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21. Magnesium Oxide + Carbon Dioxide Magnesium Carbonate

22. Copper + Sulfuric Acid Cupric Sulfate + Sulfur Dioxide + Water

23. Calcium Hydroxide + Phosphoric Acid Calcium Phosphate + Water

24. Magnesium Nitrate + Sulfuric Acid Magnesium Sulfate + Nitric Acid

25. Potassium Carbonate + Barium Chloride Potassium Chloride + Barium Carbonate

26. Aluminium Chloride + Sulfuric Acid Aluminium Sulfate + Hydrogen Chloride

27. Cadmium Phosphate + Ammonium Sulfate Cadmium Sulfide + Ammonium Phosphate

28. Manganese (IV) Oxide + Hydrochloric Acid Manganese (II) Chloride + Water + Chlorine

29. Magnesium hydroxide + ammonium phosphate magnesium phosphate + ammonium hydroxide

30. Ferric bromide + ammonium sulfide ferric sulfide + ammonium bromide

31. Calcium oxide + diphosphorous pentoxide calcium phosphate

32. Magnesium chloride + silver nitrate magnesium nitrate + silver chloride

33. Sodium carbonate + sulfuric acid sodium sulfate + carbon dioxide + water

34. Aluminium hydroxide + acetic acid aluminium acetate + water

35. Plumbous nitrate + copper (II) sulfate plumbous sulfate + copper (II) nitrate

36. Aluminium + cupric chloride aluminium chloride + copper

Reaction Predictions36

Label each equation according to its reaction type, then predict the products and balance the equations

1. Aqueous silver nitrate reacts with aqueous calcium chloride.

2. _____ Cl2 (g) + _____ NaI (aq)

3. Solid zinc reacts with aqueous copper (II) nitrate.

4. Predict the reactant isotope that yields iridium-181 during alpha decay.

5. Solid magnesium is placed in a beaker of hydrochloric acid.

6. _____ P4 (s) + _____ O2 (g)

7. _____ Ca(OH)2 (aq) + _____ HgCl2 (aq)

8. Barium chlorate is heated.

9. _____ CaBr2 (aq) + _____ KOH (aq)

10. _____ NH4OH (aq)

11. Benzene (C6H6) is ignited in the presence of oxygen gas.

12. Nitrogen gas reacts with oxygen gas.

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13. _____ Al2(SO4)3 (aq) + _____ Ca(OH)2 (aq)

14.

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