Chemistry 068, Chapter 7

Chemical Equations and Stoichiometry

• Chemical equations are representations of chemical reactions using formulas rather than words.– Correctly written chemical equations must

follow a number of rules.

• Stoichiometry is the quantitative study of the relationships between the amounts of products and reactants in a chemical reaction.– Stoichiometry also involves calculations using

chemical equations.

Evidence of Chemical Reactions

• Many, but not all, chemical reactions result in visible, or at least detectable, changes which indicate that a reaction is taking place.

• Some examples are:– Color change.– Formation of a solid.– Formation of a gas.– Heat absorption or emission.– Light emission.

Writing and Balancing Chemical Equations

• We will look at the reaction of methane and oxygen to form carbon dioxide and water as we go through the rules.

• When writing chemical reactions it is necessary to follow a number of rules.– Reactants are always written on the left,

products are always written on the right.

CH4 O2 CO2 H2O

– Plus signs separate formulas on each side.

CH4 + O2 CO2 + H2O

Writing and Balancing Chemical Equations (Cont’d)

– An arrow () or an equilibrium sign (⇌) separate the products and reactants.

CH4 + O2 CO2 + H2O– Reactions must be consistent with

experimental results (correct phase, molecular formulas rather than empirical formulas).

– Optionally, phase can be represented by writing (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous (dissolved in water) after each formula.

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

Writing and Balancing Chemical Equations (Cont’d)

– Reactions must be balanced (follow conservation of mass) for number and type of atoms but not phase.

• There must be the same number of each type of atom on each side of the equation.

– Coefficients in front of a formula tell you how many there are of that atom/molecule. Ones are excluded.

– Coefficients must be the smallest whole number values possible.

CH4 + 2O2 CO2 + 2H2O

– Reactions must be charge balanced (same overall ionic charge). We will look at this much later.

Writing and Balancing Chemical Equations from Pictures

• Chemists sometimes like to represent chemical information with images of the product and reactant molecules.

• It is possible to write chemical equations from these representations.

Writing and Balancing Chemical Equations from Pictures (Cont’d)

• For example, consider a reaction between A (red) and B (blue) to make products.

• 3A2 and 3B molecules react to form 3A2B molecules. So, we write:

3A2 + 3B 3A2B• Since we always want the smallest possible

coefficients, we finally write:

A2 + B A2B

Balancing Chemical Equations from Pictures Problem

• Write a balanced chemical equation for the reaction of A (red), B (blue), and C (green).

Strategies for Balancing Chemical Equations

• Remember that number and type of atoms is the main thing that must be balanced.– Phase and number of molecules does not

need to be balanced.

• You can only change coefficients (the numbers in front).– You cannot do anything to change the

molecules in the reaction.– You can’t change subscripts.

Strategies for Balancing Chemical Equations (Cont’d)

• Try to work with one element at a time – it’s usually much easier.

• It is often simpler to treat polyatomic ions as single units rather than individual atoms.– This only really works when the polyatomic

ions remain as single units in both reactants and products.

Strategies for Balancing Chemical Equations (Cont’d)

• Check your coefficients.– They must all be whole numbers (no

fractions).– They should be the smallest possible whole

numbers.

• Once you have finished, double check everything.– Equal number of each element, no fractional

coefficients, smallest whole number coefficients.

Balancing Chemical Equation Problems

• Balance the following reactions:– PbO + NH3 Pb + N2 + H2O

– NH3 + HNO3 NH4NO3

– NBr3 + NaOH N2 + NaBr +HBrO

– C5H10 + O2 CO2 + H2O

– Al + Sn(NO3)2 Al(NO3)3 + Sn

Balancing Chemical Equation Problems (Cont’d)

• Balance the following reactions:– FeO + HCl Fe + Cl2 + H2O

– NaOH + H2CO3 Na2CO3 + H2O

– Ba(C2H3O2)2 + (NH4)3PO4 Ba3(PO4)2 + NH4C2H3O2

– C2H6 + O2 CO2 + H2O

– NO + CH4 HCN + H2O + H2

Ions and Solutions

• The basic theory of ionic solutions was established by Svante Arrhenius in 1884.

• He proposed that certain substances split into freely moving ions when in solution and that those ions allow for electrical conduction.

Electrolytes and Nonelectrolytes

• Nonelectrolytes – substances that dissolve in water to produce a poorly or non conducting solution.

• Electrolytes – substances that dissolve in water to produce a conducting solution. They are usually ionic compounds.– Strong electrolytes completely dissociate into ions.– Weak electrolytes only partially dissociate into ions.– Same idea as strong acids/bases or weak

acids/bases.

Solubility Rules

• The solubility rules are a set of qualitative rules used to determine weather or not ions are soluble in water.

• Soluble ions are normally strong electrolytes.

• They arise from the “like dissolves like” rule.

• There are still quantitative limits to these rules.

Solubility Rules (Cont’d)Soluble

• All group IA and ammonium – Li+, Na+, K+, NH4

+.• All acetates and nitrates –

C2H3O2-, NO3

-.• Most chlorides, bromides, and

iodides – Cl-, Br-, I-.– Exceptions: Silver [AgCl, Ag

Br, Ag I], Mercury [Hg2Cl2, HgBr2, Hg2Br2, HgI2, Hg2I2], Lead [PbCl2, PbBr2, PbI2].

• Most sulfates – SO42-.

– Exceptions: Calcium [CaSO4], Strontium [SrSO4], Barium [BaSO4], Silver [AgSO4], Lead [PbSO4].

Insoluble• Most carbonates – CO3

2-.– Exceptions: Group IA and

ammonium carbonates.• Most phosphates – PO4

3-.– Exceptions: Group IA and

ammonium phosphates.• Most sulfides – S2-.

– Exceptions: Group IA, IIA, and ammonium sulfides.

• Most hydroxides – OH-.– Exceptions: Group IA

hydroxides, Calcium [Ca(OH)2], Strontium [Sr(OH) 2], Barium [Ba(OH) 2].

Precipitation Reactions

• Uses the solubility rules to predict weather or not any precipitate(s) form.

• The insoluble compound(s) will form solid precipitate(s).

• Simplest form is a double replacement reaction where ions from the two mixed solutions switch places.

Precipitation Reaction Problems

• Predict the identity of any precipitate(s) that forms during the following reactions.– NaBr + AgNO3 ?

– NaOH + CaCl2 ?

– CaSO4 + BaS ?

Molecular, Complete Ionic, and Net Ionic Equations

• Three different kinds of equations can be written to describe ionic solution phase reactions.

• Molecular – no ions shown.– This is the way we have been writing equations up

until now.• Complete Ionic – all aqueous molecules written

as ions, including spectator ions.– Spectator ions – ions in the chemical equation that

don’t take part in the reaction – they are both products and reactants.

• Net Ionic – spectator ions removed from the equation.

Molecular, Complete Ionic, and Net Ionic Equation Examples

• Mg(s) + 2HBr(aq) MgBr2(aq) + H2 (g)

– Molecular• Mg(s) + 2H+(aq) + 2Br-(aq)Mg2+(aq) + 2Br-(aq) + H2(g)

– Complete Ionic

• Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)

– Net Ionic

• Note that all of these equations can be used to represent the same reaction.

Writing and Balancing Complete Ionic and Net Ionic Equations

• Complete ionic and net ionic equations are written and balanced in much the same way as molecular equations.

• The key difference is in how charge must be balanced as well as number/type of atoms.– The total charge of all ions added up must be

the same on both the reactants and products side of the equation.

– This usually, but not always, means a total charge of zero on both sides.

Writing and Balancing Complete Ionic and Net Ionic Equations (Cont’d)

• It is also important to establish rules for what compounds do, and do not, split up into ions in water.

• The following kinds of compounds split up:– All soluble salts.– All strong acids.– All strong bases.

• All other compounds are written as molecules rather than ions.

Writing and Balancing Complete Ionic and Net Ionic Equations (Cont’d)

• To write a complete ionic or net ion equation you should follow the following steps:

1. Write and balance the molecular equations as normal.

HCl + NaOH NaCl + H2O

2. Split compounds into ions as appropriate, and double check charge.

1. At this point you are done if all you are writing is a complete ionic equation.

H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O• If writing a net ionic equation, cancel out any

spectator ions.H+ + OH- H2O

Writing and Balancing Complete Ionic and Net Ionic Equation Problems

• Write complete and net ionic equations for each of the following reactions:– MgCl2 + AgNO3 Mg(NO3)2 + AgCl

– H2SO4 + KOH H2O + K2SO4

– NaNO3 + KCl KNO3 + NaCl

Acid-Base Reactions

• Arrhenius Acid – A substance that produces hydrogen ions (H+) when dissolved in water.

Ex: HCl(aq)H+(aq)+Cl-(aq)

• Arrhenius Base – A substance that produces hydroxide ions (OH-) when dissolved in water.

Ex: NaOH(aq)Na+(aq)+OH-(aq)

• Bronsted-Lowry Acid – A proton donor in a proton transfer reaction.

• Bronsted-Lowry Base – The proton acceptor in a proton transfer reaction.

Ex:

NH3(aq)+H2O(aq)NH4+(aq)+OH-(aq)

Acid-Base Reactions (Cont’d)

• Strong Acid/Base – Completely ionizes in water.

• Weak Acid/Base – Only partially ionizes in water.

• Strong/Weak only refers to ionization, not how harmful or how reactive it is.

• Concentration is a much better indicator of how dangerous an acid/base is.

Strong Acids and Bases• Acids – HClO4, H2SO4,

HI, HBr, HCl, HNO3.

• Bases – LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2.

Acid-Base Reactions (Cont’d)

• Acid-Base Indicator – A dye used to distinguish between acidic and basic solutions by color.

• Neutralization – A reaction of an acid and a base that produces a salt and possibly water.HA(aq) + BOH(aq) H2O(l) + AB(aq)

Ex: HCl(aq) + NaOH(aq) H2O(l) + NaCl(aq)

• Polyprotic Acid – An acid that contains more than one acidic hydrogen.

Acid-Base Problems

• Write molecular, ionic, and net ionic equations for the following neutralization of Barium hydroxide by hydrobromic acid.

Acid-Base Problems (Cont’d)

• Write molecular, ionic, and net ionic reaction equations for the neutralization of each acidic hydrogen in sulfuric acid by sodium hydroxide; and for the overall (1 step) reaction.

Acid-Base Gas Formation Reactions

• Reactions with gas formation – certain salts give off gas when they react with acids.– Most notably carbonates (CO3

2-), sulfites (SO3

2-), sulfides (S2-), and ammoniums (NH4+).

– They produce CO2, SO2, H2S, and NH3 respectively.

Acid-Base Gas Formation Reactions (Cont’d)

• The reactions of sulfides are a simple one step metathesis reaction resulting in a salt and a gaseous product.

• Example:

Na2S(aq)+2Hl(aq) 2NaI(aq)+H2S(g)

Acid-Base Gas Formation Reactions (Cont’d)

• The reactions for carbonates and sulfites are very similar.– The reaction takes place in two steps. The first step similar to a

metathesis reaction, which produces a salt and an unstable intermediate product.

– The intermediate then decomposes into water and a gas.

• Examples:

K2CO3(aq)+2HBr(aq) 2KBr(aq)+H2CO3(aq)

2KBr(aq)+H2CO3(aq)2KBr(aq)+H2O(l)+CO2(g)

Li2SO3(aq)+2HCl(aq) 2LiCl(aq)+H2SO3(aq)

2LiCl(aq)+H2SO3(aq)2LiCl(aq)+H2O(l)+SO2(g)

Acid-Base Gas Formation Reactions (Cont’d)

• Like the reactions for carbonates and sulfites, ammonium reactions are two step..– The reaction takes place in two steps. The first step similar to a

metathesis reaction, which produces a salt and an unstable intermediate product.

– The intermediate then decomposes into water and a gas.

• Example:

NH4Cl(aq)+KOH(aq) KCl(aq)+NH4OH(aq)

NH4OH(aq)+KCl(aq) H2O(l)+KCl(aq)+NH3(g)

Acid-Base Gas Formation Reactions Problems

• Write a reaction for the production of carbon dioxide when calcium carbonate reacts with hydrochloric acid.

• If 12.0g of calcium carbonate are reacted, what mass of carbon dioxide is produced?

Oxidation – Reduction Reactions

• Oxidation-reduction reactions (also called redox) involve electron transfer from one species to another, or in which atoms change oxidation number (state).

• For example, consider these molecular and net ionic equations:Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s)Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)

• Iron and copper switch charged states. Iron goes from neutral to 2+ while copper goes from 2+ to neutral. Sulfate does not change state.

Combustion Reactions

• Combustion reactions are reactions of a substance with oxygen which include a release of heat and flame.– Typically, oxygen is from the air.

• Example reactions:

2C2H2 + 5O2 4CO2 + 2H2O

CS2 + 3O2 CO2 + 2SO2

Predicting Products of a Combustion

• Carbon and hydrogen atoms commonly form the following compounds as products during a combustion:– C forms CO2

– H forms H2O– O forms either CO2 or H2O

• It is more difficult to predict products formed from other elements.– S forms SO, SO2, SO3

– N forms NO, NO2, NO3

– It is only possible to determine which is formed experimentally.

Predicting Products of a Combustion (Cont’d)

• For the moment we will only deal with compounds of carbon, hydrogen, and oxygen.– Each reactant carbon forms one CO2.

– Every two reactant hydrogens for one H2O.

– Ignore oxygens – we assume they are all used up to make water or carbon dioxide.

• You can then determine the number of oxygens needed as reactants by chemical balancing.

Predicting Products of a Combustion Problems

• Predict the products of the combustion of ethanol C2H6O and write the balanced equation for the reaction.

Predicting Products of a Combustion Problems (Cont’d)

• Predict the products of the combustion of benzene C6H6 and write the balanced equation for the reaction.

Predicting Products of a Combustion Problems (Cont’d)

• Predict the products of the combustion of methane CH4 and write the balanced equation for the reaction.

Types of Chemical Reactions

• There are many types of chemical reactions and a single course could not possible look at all of them.

• The following one method to classify chemical reactions.

• We will look at 4 classes of chemical reactions in this chapter.

• They are:– Synthesis reactions.– Decomposition reactions.– Single replacement reactions.– Double replacement reactions.

Synthesis Reactions

• Synthesis, sometimes called combination, reactions are reactions which form a single product from two or more reactants.– Reverse of decomposition reactions.

• Example reactions:

2NO2 + H2O2 2HNO3

Ni + S NiS

Decomposition Reactions

• Decomposition reactions are reactions where a single reactant breaks down into two or more products.– Reverse of synthesis reactions.

• Example reactions:

2H2O 2H2 + O2

Al2(CO3)3 Al2O3 + 3CO2

Single Replacement Reactions

• Single replacement reactions are reactions where one element in a compound is replaced by another (single) element.– Single elements but can be multiple atoms of

the same element.

• Example reactions:

Fe + CuSO4 Cu + FeSO4

Mg + Ni(NO3)2 Ni + Mg(NO3)2

Double Replacement Reactions

• Double replacement reactions are reactions where two compounds switch parts with one another to create two new compounds.– Again, can be multiples of the same part.

• Example reactions:

NaF + HCl NaCl + HF

AgNO3 + HCl AgCl + HNO3

Types of Reactions Problems

• Determine the type of chemical reaction for each of the following:

• SO3 + H2O H2SO4

• Na2CO3 + Ca(OH)2 2NaOH + CaCO3

• Fe + 2CuNO3 2Cu + Fe(NO3)2

• 2SO2 + O2 2SO3

• K2CO3 K2O + CO2