PRINCIPLES OF CHEMICAL REACTIVITY: CHEMICAL

REACTIONS Chapter 4

Chemical equations

• A chemical equation is the representation of a chemical reaction using the symbols of the elements and the formulas of the compounds

• The reacting species (starting materials or reactants) are listed on the left, and the products of the reaction are listed on the right

Four Things to Know About Equations

One: Reactants are on the left and products are on the right of the arrow.Two: Coefficients are the numbers in front of each formula. If no number is shown, a one is understood.Three: The coefficients tell us how many molecules (moles) of each reactant used and how many molecules (moles) of each product made.Four: The phases of the reactants and products are indicated in parenthesis after the formula.

• gas = (g)• liquid = (l)• solid (s)• aqueous (in a solution) =(aq)

Balanced Chemical Equation

• C6H12O6(s) + 6O2(g) 6CO2(g) + 6H2O(l)

Balancing Chemical Equations

• We balance chemical equations because of the Law of the Conservation of Mass

• We use coefficients in front of the element or compound formula, which are whole numbers that adjust the amounts of the species

• Note that we cannot balance an equation by changing the subscripts of the compound

Balancing coefficients

• The subscripts of the compound are fixed: they cannot be changed in an equation

• The coefficients used should be the smallest whole numbers possible

• The coefficient multiplies every number in the formula (e.g. 2 MgCl2 means that there are 2 Mg atoms and 4 Cl atoms)

• We will balance equations by inspection

Balance the following

• Si2H3 + O2 SiO2 + H2O

• Al(OH)3 + H2SO4 Al2(SO4)3 + H2O

Types of chemical reactions

• Combustion reactions - reaction of an element or a compound with elemental oxygen (O2)

• These reactions are often accompanied by the release of large amounts of heat (and flame)

• Consider: 2C2H6 (g) + 7O2 (g) 4CO2(g) + 6H2O(g)

Combination reactions

• Often called synthesis reactions

• Involves two or more substances combining to form a single substance

• Consider:

• 2Mg(s) + O2 (g) 2MgO(s)

Decomposition reactions

• The reverse of a combination reaction

• One substance breaking apart into two or more substances

• Consider:

• H2CO3 (aq) H2O(l) + CO2 (g)

Reactions in Aqueous Solution

• Single Replacement

• Double Replacement

• Oxidation-Reduction

• All require an understanding of aqueous solutions!

Formation of ions in solution

• When we consider a substance dissolving in another substance, we define the solute as the substance in lesser amount, and the solvent as the substance in greater amount

• Solubility is the relative degree to which a substance dissolves in a solvent

• Soluble substances dissolve to a considerable degree

• Insoluble substances do not dissolve at all or only to a small degree

Dissolution of ionic compounds

• When ionic compounds dissolve in water, they dissociate into separate anions and cations:NaCl(s) Na+

(aq) + Cl-(aq)

• Acids are so named because when they dissolve, the ionize to yield H+

and the associated anionHCl(aq) H+

(aq) + Cl-(aq)

Solubility RulesRule ExceptionSoluble

Group 1 elements, NH4+

NO3 , ClO3

, ClO4

Chlorides, bromides, iodides Ag+, Pb2+, Hg22+

Acetates Ag+, Hg22+

Sulfates Sr2+, Ba2+, Pb2+, Ca2+

Insoluble Compounds

Carbonates, phosphates, oxalates, chromates, sulfides

Group 1, NH4+

Hydroxides, oxides Group 1, Ba2+

Strong Acids

• Strong acids completely dissociate into ions upon dissolution in water

• With strong acids the major components in solution are the hydrogen ion and the anion.

• Examples are HCl, H2SO4

Weak Acids

• Weak acids dissolve, but do not completely dissociate into H+ and anions.

• With weak acids the major component in solution is the weak acid.

• Examples are CH3CO2H, HF

Single replacement reactions

• An element replaces another in a compound

• Zn(s) + CuCl2 (aq) ZnCl2 (aq) + Cu(s)

• This particular form of the equation is termed the molecular equation, where all reactants and products are shown as neutral compounds

Ionic equations

• The total ionic equation writes all of the anions and cations separately

Zn(s) + Cu2+(aq) + 2Cl-

Zn2+(aq) + 2Cl- + Cu(s)

• Ions (such as the Cl- ions) that are in the same state on both sides of the equation are called Spectator ions.

Net ionic equation• The net ionic equation leaves out the

spectator ions

• The complete ionic equationZn(s) + Cu2+

(aq) + 2Cl- Zn2+

(aq) + 2Cl- + Cu(s)

• The net ionic equationZn(s) + Cu2+

(aq) Zn2+(aq) + Cu(s)

Double replacement reactions

• Where the anions and cations in an ionic equation swap places.

• AgNO3 (aq) + NaCl(aq) AgCl(s) + NaNO3 (aq)

What causes double replacements?

• There has to be some process to cause the reaction to take place– Precipitation of a solid– Evolution of a gas– Formation of a Molecular Substance– Neutralization reaction between an acid and a

base

• If one of these processes does not occur, no reaction takes place

Double replacement reactions

• Precipitation– AgNO3 (aq) + NaCl(aq) AgCl(s)

+ NaNO3

(aq)

• Gas evolution– 2HCl(aq) + Na2CO3 (aq) 2NaCl(aq)

+ CO2 (g)

Double replacement reactions

• Neutralization– an acid is a proton donor– a base is a hydroxide donor– proton plus hydroxide yields water

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

Net Ionic Equations

• In a net ionic equation the spectator ions are omitted.

• Spectator ions are those that do not change in the reaction.

• Net ionic equations are only used for reactions occurring in solution between ionic compounds – single and double replacement reactions – as well as oxidation-reduction

Write the Net Ionic Equation

• AgNO3 (aq) + NaCl(aq) AgCl(s) + NaNO3 (aq)

• 2HCl(aq) + Na2CO3 (aq) 2NaCl(aq) + CO2 (g)

• HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Oxidation – Reduction

• Oxidation• loss of at least one electron during a

reaction..– Ni(s) + H+(aq) Ni2+(aq) + H2(g)

• Reduction • gain of at least one electron during a reaction.

– In above example, H+ gains an electron to become reduced.

Oxidation-Reduction Cont.

• Every reaction must have an oxidation and reduction.

• Metals react with acids to form salts and hydrogen gas.– Cu(s) + 2HNO3(aq) Cu(NO3)2(aq) + H2(g)

• Metals also oxidized with salts– Fe(s) + Cu(NO3)2(aq) Fe(NO3)2(aq) + Cu(s)

Oxidation Number

• Oxidation number (state): • The charge that an atom in a

substance or monatomic ion has or would have if all the bonding electrons were assigned to the most electronegative element.

Ca in CaO +2

Ca2+(aq) +2

Cl(aq) 1

Cr in CrO3 +6

Fe in Fe2O3 +3

Cr in K2Cr2O7 +6

Assigning Oxidation Numbers

– Elemental form: 0– Monatomic ions: charge of ion– Oxygen: 2, except in H2O2 and other peroxides.– Hydrogen: +1, except with metal hydrides when it is

1.– Halogens: 1 (except when bound to oxygen or a

halide above it)– Alkali and alkaline earth metal ions have a charge of

+1 and +2, respectively.– Compounds and ions: sum of the charges on the

atoms in a compound add up to 0 and to the ion charge in the ion.

Displacement Reactions: Activity series of the elements

• A relative reactivity scale allows us to predict if reaction will occur when two substances are mixed together.

E.g. Copper ions in solution are reduced to the metal when an iron nail is placed in the solution.

– Cu2+(aq) + Fe(s) Fe2+(aq) + Cu(s) Iron displaces copper.

– Fe2+(aq) + Cu(s) NR copper will not displace iron.

• Iron more reactive than copper.E.g. Predict which reaction will occur when:

– Li is mixed with K+ and – Li+ is mixed with K.

E.g. In which of the following mixtures will reaction occur:

– Li+ + Mg– Al + Mn2+ – Fe + Cd2+ – Cr + Zn2+

Li

Reacts vigorously with acids to give H2

Reacts with H2O to give H2

K

Ba

Ca

Na

Mg Reacts with acids to give H2

Reacts slowly with H2O to give H2

; more vigorous with steam

Al

Zn

Cr

Fe

Cd

Etc.

Balancing: Oxidation-Number Method

• Determine oxidation # for each atom- both sides of equation.

• Determine change in oxidation state for each atom.

• Left side: make loss of electrons = gain.

• Balance other side.

• Insert coefficients for atoms that don't change oxidation state.

Example

FeS(s) + CaC2(s) + CaO(s) Fe(s) + CO(g) + CaS(s)

Balancing: Half-Reaction Method

Write unbalanced half reactions for the oxidation and the reduction

• Balance the number of elements except O and H for each.

• Balance O's with H2O to the deficient side. • Balance H's with H+ to the hydrogen deficient side

– Acidic: add H+

– Basic: add H2O to the deficient side and OH to the other side.

• Balance charge by adding e to the side that needs it. • Multiply each half-reaction by integers to make electrons

cancel.• Add the two half-reactions and simplify.

Half-Reaction Acidic

Zn(s) + VO2+(aq) Zn2+(aq) + V3+(aq).

Half Reaction Basic

Ag(s) + HS(aq) + CrO42(aq) Ag2S(s) +

Cr(OH)3(s).