Predicting Different types of Chemical Reactions

PERIODIC TABLE HISTORY Periodic Table (periodic - repeating pattern) Dmitri Mendeleev – Russian chemist,

discoverer of periodic law, (Md) – used existing properties, color, density, boiling point, freezing point, oxide and hydrides formed along with atomic weight to arrange the existing elements in a new periodic table.

Some elements seemed out of order b/c atomic number, not atomic mass determines an elements properties. (Te before I)

Mendeleev had some empty spaces left on his P.T. He predicted the existence and properties of the missing three spaces, which was extremely accurate.

Metals

Metalloids

luster (shiny) malleable (bending) ductile (strung into a wire) good conductors of

electricity – the outer e- are delocalized – can move from atom to atom

lose e- to gain full valence

dull brittle poor conductors – want

electrons gain e- to gain full

valence

some have luster, some don’t some are malleable and ductile, some are notsemiconductors; found in new technology

Metallic character decreases across a period and increases down a family

Nonmetals

SOME THINGS TO MEMORIZE – THE SOONER THE BETTER!

Learn this list of organic names…CH4 - methane

C2H6 - ethane

C3H8 - propane

C4H10 - butane

C5H12 - pentane

C6H14 - hexane

C7H16 - heptane

C8H18 - octane

C9H20 - nonane

C10H22 - decane

C12H22O11 - sugar

SOME THINGS TO MEMORIZE – THE SOONER THE BETTER!

For alcohols: Use the organic base Drop an H, add an OH

methanol (methyl alcohol) Look at methane – CH4

drop an H, add an OHCH3OH - methanol

ethanol (ethyl alcohol)Look at ethane – C2H6

drop an H, add an OHC2H5OH - ethanol

MORE MUST KNOWS!Prefix Table

1 - mono 2 - di 3 - tri4 - tetra5 - penta6 - hexa7 - hepta 8 - octa9 - nona 10 - deca

MUST KNOW CONTINUED!Acid Table

- ide hydro – ic acid

- ite - ous acid

- ate - ic acid

ELECTROLYTES Electrolytes are substances which when dissolved

conduct a current, electrolytes produce ions in solution, it’s called ionic dissociation

  All ionic compounds are electrolytes If dissociation into ions is about 100% complete, it’s a

strong electrolyte. If not, it’s a weak electrolyte.

Most covalent compounds are nonelectrolytes. Exceptions: strong covalent acids and some bases

are electrolytes. They ionize 100% in solution, called covalent ionization.

Acids: HCl, HBr, HI, HNO3, H2SO4, HClO4, HClO3 *if O’s outnumber the H’s by at least two it’s a strong acid

ELECTROLYTES Strong (soluble)• Strong acids• Strong bases• Soluble ionic salts

Weak (slightly soluble)• Weak acids• Weak bases

Non electrolytes (insoluble)• Insoluble ionic salts• Molecules (only nonmetals)

Ex. Gases, liquids

dissociation table (aq)-use for net ionic rxns

STRONG ELECTROLYTES (100% IONIZED)YOU WILL NEED TO MEMORIZE THIS TOO

A. Strong Acids: HCl, HBr, HI, H2SO4, HNO3, HClO4, HClO3

B. Strong Bases: Hydroxides of group IA and II A, Except Be(OH)2 and Mg(OH)2

STRONG ELECTROLYTES (100% IONIZED)YOU WILL NEED TO MEMORIZE THIS TOO

C. Soluble Salts (ionic compounds: metal/nonmetal)

 The following are Always Soluble if they are in an ionic

compoundNO3

-, Group IA metals, NH4+, CH3COO-, ClO4

-, ClO3-

Cl-, Br-, I- ~ Are always soluble except with Ag, Pb, Hg2

2+  

SO42- ~ Is always soluble except with Ag, Pb, Hg2

2+,

Ca, Sr, Ba

SOLUBILITY RULES FOR SALTS IN WATER SOLUTIONS

(another way to look at the Dissociation table)

Rules 1-6 must be applied in order

1. Soluble: All ammonium, NH4, and Group IA metals, Li, Na, K, Rb, Cs, Fr

2. Soluble: All nitrates, NO3; perchlorates, ClO4

-; and chlorates, ClO3- ; and acetates,

CH3COO- Exception: CH3COOH changes b/c it’s a weak acid (a molecule)

When Rules 1 & 2 do not apply3. Insoluble: All silver, Ag+; lead Pb2+; and

mercury (I) salts Hg22+ salts.

When Rule 3 does not apply4. Soluble: All chlorides, Cl-; bromides, Br-; and

iodides I-

When Rules 1, 2, & 4 don’t apply5. Insoluble: All carbonates, CO3

2-; chromates, CrO4

2-; phosphates, PO43-; sulfides, S2-; oxides, O2-;

and hydroxides, OH-.Exceptions: All Group IIA chromates, except,

BaCrO4 are soluble

Group IIA hydroxides, except Mg(OH)2 and Be(OH)2 are soluble. Ca(OH)2 is slightly insoluble.

When Rules 3 & 5 don’t apply6. Soluble: All sulfates, SO4

2-; except BaSO4, CaSO4, and SrSO4

With few exceptions, the solubility rules permit assuming:

A salt with anion 1- charge is soluble. These salts will completely dissociate into ions.

Salts with anions of 2- or 3- charges are insoluble. These salts will precipitate. (Rule 6)

Don’t forget: Infamous seven, diatomic molecules, H2 N2 O2 F2 Cl2 Br2 I2

Crazy cousins: P4, S8

OXIDATION NUMBERS Oxidation numbers (oxidation states) –

the charge that an atom carries

1. All elements in their standard state have a charge of 0.

solid Fe ox. state = 0solid Zn ox. state = 0oxygen gas O2 ox. state = 0

2. Ions of a single atom: oxidation state = charge

Ex. Fe3+ ox. state = +3 Zn2+ ox. state = 2+ O2- ox. state = -2

3. Sum of all oxidation #’s on a compound = 0; sum on a polyatomic ion = overall charge

Ex. H2O

Ca3(PO4)2

4. Fluorine always -1 in compounds

5. IA always +1IIA always +2

IIIA mostly +3

6. Hydrogen is mostly +1 EXCEPT in a metal hydride, H = -1

Ex. NaH

CaH2

7. Oxygen is mostly -2; can be -1, -1/2, +1Ex. Superoxide NaO2 Na1+ O-1/2

TYPES OF REACTIONS: DOUBLE REPLACEMENT-METATHESIS

Double replacement/Metathesis Rxns: (square dance) – Identified by two compounds yielding two different compounds. Cations (+ charged particles) and anions (-

charged particles) change partners. Usually a precipitate is formed. ppt = insoluble

product If water is involved use HOH instead of H2O

Solubility rules need to be known.Ex. AB + CD AD + CB

ppt or molecule is formed

NET IONIC EQUATIONS Net ionic equations - An equation that only

includes the ions and compounds that undergo a chemical change.

AP Reactions must be written using Net Ionic rules

Step 1 – Predict the reaction Step 2 – label each compound as (s), (l), (aq), or

(g) Step 3 – break apart each aqueous compound,

which is the total ionic equation Step 4 – cancel out the spectator ions Step 5 – Rewrite the equation, you now have the

net ionic equation

Use net ionic reactions when applicable (aka. every time you can)

Ex. Sodium chromate + nickel(III) chloride

Rxn: 3Na2CrO4(aq) + 2NiCl3(aq) 6NaCl(aq) + Ni2(CrO4)3(s)

Total ionic: 6 Na++3CrO4

2-+2Ni3++6Cl1-6 Na+ + 6Cl1- + Ni2(CrO4)3(s)

Cancel out the Spectator ions: Na+ and Cl-

Net ionic: 3CrO4

2- + 2Ni3+ Ni2(CrO4)3(s)

1. Auric acetate and sodium dichromate i. List the spectator ions

2. Potassium hydroxide + hydrochloric acidi. What is another name for this type of rxn?

 

 

3. Lithium hydride + hydrobromic acidi. Describe how you know this rxn is taking place.

4. Potassium nitrate + lead(II) acetatei. Illustrate what is happening in the rxn on the molecular level.

Product breakdowns to look for:NH4OH breaks down to NH3 + H2O

H2CO3 breaks down to CO2 + H2O

5. Ammonium chlorate + cesium hydroxide i. How do you know this rxn is taking place.

6. Calcium carbonate and hydrobromic acidi. What would you see happening as this rxn proceeds to completion?

TYPES OF REACTIONS: SINGLE REPLACEMENT/DISPLACEMENT

Single replacement/Displacement – (home wrecker) - Identified by one element and one compound yielding a different element and different compound.

The element trying to break up a compound must be more reactive than the element its replacing in the compound.

(If water is involved use HOH instead of H2O) Perform net ionic rxns whenever possible

REACTIVITY Group IA is the most reactive metals.

Fluorine is the most reactive halogen.

Likes can only replace likes. Metals can only replace other metals

Hydrogen is the only exemption…when hydrogen is replaced it forms H2

Fluorine can replace the other halogens, chlorine can replace everything but fluorine, iodine cannot replace bromine, chlorine, or fluorine…

Ex. Nitric acid and excess iron6HNO3

+ 2Fe 2Fe(NO3)3 + 3H2

Net ionic: 6H+ + 2Fe(s) 2Fe3+ + 3H2(g)

Ex. Nitric acid and platinumHCl + Pt NR

Platinum is not reactive.

All single replacement reactions are also redox reactions. Look at charges/oxidation numbers of the reactants and products.

1. Cesium chloride and fluorine gasi. Describe how you would know chlorine gas was present?

2. Sodium chloride and bromine gasi. Explain why this rxn could not take place.

3. Zinc and hydrofluoric acid i. What would you see happening as this rxn occurs?

TYPES OF REACTIONS: COMBUSTION Combustion – Normally identified by an

organic compound burning in oxygen.

Most combustion reactions Organic compound and oxygen gas

nonmetal oxide and water

A few are a tad more tricky Ionic compound and oxygen gas

metal oxide + nonmetal oxide

Ex. Butane is heated

2C4H10 + 13O2 8CO2 + 10H2O

1. Silicon tetrahydride is burned in oxygen

2. Propanol is burned

3. Barium fluoride undergoes combustion

i. Are combustion rxns exothermic or endothermic?

TYPES OF REACTIONS: SYNTHESIS RXNS

Synthesis rxns (two become one) – identified by 2 or more substances yielding one compound

A + B AB

Use the higher oxidation state on the element if the nonmetal is in excess,

use the lower oxidation state on the element if the nonmetal is limited.

a. Binary synthesis formation: element + element compound.

(beware of valence charges)

Ex. Sodium and fluorine gas combine2Na + F2 2NaF

BINARY SYNTHESIS

BINARY SYNTHESIS

Answer for #1-5 What are the oxidation numbers of each element before and after the rxn?

1. Nickel + excess sulfur

2. Nickel and limited sulfur

3. Sulfur and excess oxygen

 

 

4. Sulfur + limited oxygen

 

5. Silicon and chlorine gas

OXYBASE FORMATIONb. (metallic hydroxide) metal oxide + water oxybase.

Use dissociation rules if the reaction is in solution. (Net ionic)

Ex. SrO + H2O Sr(OH)2

Net ionic: SrO + H2O Sr2+ + 2OH-

OXYBASE FORMATION1. Aluminum oxide and water

OXYBASE FORMATION2. Potassium oxide + water

Compare #1&2, which rxn would make a strong base? Why?

3. Strontium oxide + water  

4. Ferric oxide and water

Compare #3&4, would either produce a weak base? Why?

OXY-ACID FORMATIONc. nonmetal oxide + water oxyacid

High valence on nonmetal yields “ic” acid, low valence on the nonmetal yields “ous” acid.

Do not change the oxidation state (charge) on the nonmetal

Ex. P2O3 + 3H2O 2H3PO3

(P keeps the charge of 3+ on both sides of )

P3+ O2- H1+ P3+ O2- Overall charges: for P2O3 (3+ x2) + (2- x3) = 0

for H3PO3 (1+ x3) + (3+ x1) + (2- x3) = 0

OXY-ACID FORMATION1. Dinitrogen pentaoxide + water

OXY-ACID FORMATION2. Sulfur dioxide + water

If you had the same concentration of the acids in #1 & 2, which would have the lower pH?

3. Diphosphorus pentaoxide + water  

4. Dinitrogen trioxide + water

If you had the same concentration of each acid in #3&4, which would have the higher pH?

OXY-SALT FORMATIONnonmetal oxide + metal oxide an oxysalt

metal w/oxide radical (polyatiomic ion)

Ex. P2O5 + 3K2O 2K3PO4

(same charge on P, 5+)

~ for the nonmetal:higher charge = ate ion, lower charge = ite

ion

1. Barium oxide and dinitrogen trioxide

2. Diarsenic pentaoxide + aluminum oxide

Describe why reactions #1&2 are not double replacement reactions.

TYPES OF REACTIONS: DECOMPOSITION RXNS:Decomposition rxns Identified by 1

compound yielding two or more simpler substances. AB A + B

-accomplished by heating or electrolysis, starts with ONE reactant only.

a. Binary compound decomposition: 1 compound yields two elementsEx. Sodium fluoride undergoes electrolysis 2NaF 2Na + F2

1. Calcium nitride undergoes electrolysis

i. Describe what would you see happening as this reaction takes place.

2. Diphoshphorus pentasulfide is heated

i. What are the oxidation numbers of each element before and after the rxn?

MORE DECOMPOSITION b. oxy-base decomp: an oxybase yields a

metallic oxide and water (use net ionic if appropriate)Ex. Ca(OH)2 CaO + H2O

1. Sodium hydroxide is decomposed in extreme heat

 

c. oxy-acid decomp: an oxy-acid yields a nonmetal oxide and water

(keep charges the same on both sides)Ex. 2H3PO3 P2O3 + 3H2O

1. Chloric acid undergoes electrolysis

d. oxy-salt decomp: oxysalt yields a nonmetal oxide and a metallic oxide

Ex. 2K3PO4 P2O5 + 3K2O

1. Lithium arsenate decomposes

 

e. Metal carbonates yield carbon dioxide and a metallic oxideEx. MgCO3 CO2 + MgO

1. During electrolysis Cobalt(III) carbonate

f. Metallic chlorates yield metallic chlorides and oxygen gas.Ex. 2LiClO3 2LiCl + 3O2

1. Lead (IV) chlorate is heated

DECOMP EXCEPTIONS – YOU DO NOT HAVE TO KNOW THESE!!!

g. Metal bicarbonates yield metallic oxide + carbon dioxide + waterEx. 2NaHCO3 Na2O + 2CO2 + H2O

h. Metallic nitrite yields metallic oxide + nitrogen monoxide + oxygen gasEx. 2Mg(NO2)2 2MgO + 4NO + O2

i. Group IA metal nitrate yield Group IA nitrite and oxygen gasEx. 2NaNO3 2NaNO2 + O2

j. Any other metallic nitrate yields metal oxide + nitrogen dioxide + oxygen gasEx. 2Sr(NO3)2 2SrO + 4NO2 + O2 

TYPES OF REACTIONS: REDOX REACTIONS:

reduction/oxidation rxns – change of oxidation states on various elements. Learn the list of important oxidizers and reducers. (Usually an acidified solution is a helpful hint that redox is happening)

oilrig: oxidized loses electrons, reduction gains electrons

Redox rxns:~ something is reduced (charge goes down)~ something is oxidized (charge goes up)

Oxidizing agent: causes something else to be oxidized

Reducing agent: causes something else to be reduced

Ex. Feo + O2o Fe2O3

Change in charges Iron 0 to 3+ Oxygen 0 to 2-

Oxidizing agent = oxygen reducing agent = iron

reduced = oxygen oxidized = iron

DISPROPORTIONATION REACTIONS

Disproportionation reaction is a redox rxn in which the same element is oxidized and reduced.Ex. Cl2 + H2O HCl + HClO

RULES FOR BALANCING REDOX EQNS.

1. Verify redox rxn. Predict products and cancel out spectator ions

2.Split into half rxns, 1 for oxidation, 1 for reduction

3.Balance half rxns for mass, get # of atoms equal. ~ you can add H+ or H2O to balance out hydrogen or oxygen

4. Balance half rxns for charge by adding e- to a side

5. Make e- equal in half rxns by

multiplying coefficients

6. Add half rxns – cancel anything you can

7. If basic, add OH-1 to H+1 side

Hints: Become familiar with important reducers and oxidizers on your list of things to memorize. Most redox rxns will say “acidified solution” or “added acid”

BALANCING REDOX EXAMPLES

Ex. 1. A soln of iron(II) nitrate is added to an acidified soln of potassium permanganate.Fe2+ + H+ + MnO4

1- Fe3+ + Mn2+ + H2O

Ex. 2. Manganese dioxide is added to conc. hydrochloric acid and heated.MnO2 + H1+ + Cl1- Mn2+ + Cl2 + H2O

REDOX IN BASIC SOLUTIONS1. In a basic soln, sodium hypochlorite

and lithium chromite, LiCrO2 react to produce sodium chromate and lithium chloride

2. Potassium permanganate is titrated with cesium nitrite in a basic soln.

SUBATOMIC PARTICLES Counting Subatomic particles X = symbol A = atomic mass Z= atomic

# Ex.

Isotopes: atoms of the same element that have different masses due to the different number of neutrons.

DEFINITIONS OF ACIDS AND BASES

1. Arrhenius Theory: Acid ~ substance that contains hydrogen and produces H+ in aqueous solutions. Base ~ substance that contains OH and produces hydroxide ions in aqueous solutions.

Ex. Acid – HCl Base - NaOH

2. Bronsted-Lowry Theory: Acid ~ a species that acts as a proton donor. Base ~ a species that acts as a proton acceptor

Ex. Acid – NH4+ Base – F-

3. Lewis Theory: Acid ~ a substance that accepts a share in an electron pair to form a coordinate covalent bond. Base ~ a substance that makes available a share in an electron pair to form a coordinate covalent bond.

Ex. Acid – BCl3 Base – NH3

 

MOLECULAR VIEW Ex. of a Lewis acid rxn: Gases Boron

trichloride and ammonia are mixed.BCl3 + NH3 BCl3NH3 or Cl3B-NH3

1. Phosphorus trifluoride + boron triiodide

LIGANDS Ligands - also called complex ions,

coordination chemistry

Ligands (Lewis bases) are bonded to a central atom that is usually a transition metal ion (Lewis acids). Most frequently occurring ligands are NH3 and OH-1

The number of ligands attached to a central metal ion is usually twice the oxidation # of the central metal.

Ex. Fe3+ + 6CN1- Fe(CN)63-

TYPES OF LIGANDSAmmonia Excess hydroxideAg(NH3)2

1+ Al(OH)41-

Cu(NH3)42+ Zn(OH)4

2-

Ni(NH3)63+ Cr(OH)6

3-

The breakup of a complex ion is achieved by adding an acid. The products are the metal ion and the ligand reacting with the H1+ ions.

Ex. Tetraammine copper(II) ions are reacted with nitric acid (ammine = NH3)

Cu(NH3)42+ + 4H+ Cu2+ + 4NH4

+