Ch 4 notes

Ch 4 notes

Predicting Different types of Chemical Reactions PERIODIC TABLE HISTORY

Periodic Table (periodic - repeating pattern) ______________ 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 ________________, not ________________ determines an elements properties. (Te before I) Mendeleev had some empty spaces left on his

P.T. He __________the existence and properties of the missing three spaces, which was extremely accurate. ____________________

luster (shiny) malleable (bending) ductile (strung into a wire) good conductors of electricity the outer eare 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 dont some are malleable and ductile, some are not semiconductors; found in new technology Metallic character decreases across a period and increases down a family 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 OH

CH3OH - methanol ethanol (ethyl alcohol) Look at ethane C2H6 drop an H, add an OH C2H5OH - ethanol MORE MUST KNOWS! Prefix Table 1 - mono 2 - di 3 - tri

4 - tetra 5 - penta 6 - hexa 7 - hepta 8 - octa 9 - nona 10 - deca MUST KNOW CONTINUED! Acid Table - ide

hydro ic acid - ite - ous acid - ate - ic acid

Dont forget: Infamous seven, diatomic molecules: Crazy cousins: and ELECTROLYTES __________ are substances which when dissolved conduct a current, electrolytes produce ions in solution, its called ionic

dissociation All __________ compounds are electrolytes If dissociation into ions is about 100% complete, its a __________ electrolyte. If not, its considered a __________ ______

(Non-soluble ionic solids are technically a __________ electrolyte we will get to this later)

Most covalent compounds are __________s. __________: 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 Os outnumber the Hs by at least two its a strong acid ELECTROLYTES

Strong (__________) dissociation table (aq) -use for net ionic rxns

Weak (____________________) Strong acids Strong bases Soluble ionic salts Weak acids

Weak bases Non electrolytes (__________) Insoluble ionic salts Molecules (only nonmetals) Ex. Gases, liquids

STRONG ELECTROLYTES (100% IONIZED) YOU WILL NEED TO MEMORIZE THIS TOO A. Strong Acids: Hydrochloric acid, Hydrobromic acid, Hydroiodic acid, Sulfuric acid, Nitric acid, Perchloric acid, Chloric acid B. Strong Bases: Hydroxides of group IA and II A,

sodium hydroxide, lithium hydroxide, barium hydroxide, etc 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 ____________________if they are in an ionic compound

____________________________________________ ________~ Are always soluble except with _________ ______ ~ Is always soluble except with __________ __________ 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 = 0 solid Zn ox. state = 0 oxygen 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 +1 IIA 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, +1 Ex. Superoxide NaO2 Na1+ O-1/2 TYPES OF REACTIONS: 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 H 2O

Solubility rules need to be known. Ex. AB + CD AD + CB 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. Ex. Solutions of Sodium chromate + nickel(III) chloride are mixed 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- Visual Representations of Net Ionic Reactions: Rxn: 3Na2CrO4(aq) + 2NiCl3(aq) 6NaCl(aq) + Ni2(CrO4)3(s) Beaker #1

6 Na++3CrO4 2Beaker #2 2Ni3++6Cl1Combined beaker 6 Na+ + 6Cl1- + Ni2(CrO4)3(s) 1. Auric acetate and sodium dichromate i. List the spectator ions 2. Potassium hydroxide + hydrochloric acid i. What is another name for this type of rxn?

3. Lithium hydride + hydrobromic acid i. Describe how you know this rxn is taking place. 4. Potassium nitrate + lead(II) acetate i. 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 acid i. What would you see happening as this rxn proceeds to completion? TYPES OF REACTIONS:

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 H 2O) Perform net ionic rxns whenever possible REACTIVITY

__________ is the most reactive metals. __________ is the most reactive halogen. Likes can only replace __________.

Metals can only replace other metals Hydrogen is the only exemptionwhen 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 iron

6HNO3 + 2Fe 2Fe(NO3)3 + 3H2 Net ionic: 6H+ + 2Fe(s) 2Fe3+ + 3H2(g) Ex. Nitric acid and platinum HCl + 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 gas

i. Describe how you would know chlorine gas was present? 2. Sodium chloride and bromine gas i. 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 __________ Normally identified by an organic compound burning in oxygen. Most combustion reactions Organic

compound and oxygen gas oxide and water nonmetal 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 __________ __________(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. BINARY SYNTHESIS a. Binary synthesis formation:

element + element compound. (beware of valence charges) Ex. Sodium and fluorine gas combine 2Na + F2 2NaF 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 FORMATION b. (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 FORMATION 1. Aluminum oxide and water

OXYBASE FORMATION 2. Potassium oxide + water Compare #1 and 2, which rxn would make a strong base? Why? 3. Strontium oxide + water 4. Ferric oxide and water Compare #3 and 4, would either produce

a weak base? Why? OXY-ACID FORMATION c. nonmetal oxide + water oxyacid High valence on nonmetal yields ic acid, low valence on the nonmetal yields ous acid. __________ 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+ O2Overall charges: for P2O3 (3+ x2) + (2- x3) = 0 for H3PO3 (1+ x3) + (3+ x1) + (2- x3) = 0 OXY-ACID FORMATION 1. Dinitrogen pentaoxide + water OXY-ACID FORMATION 2. Sulfur dioxide + water If you had the same concentration of the

acids in #1 and 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 and 4, which would have the higher pH?

OXY-SALT FORMATION nonmetal 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 and 2 are not double replacement reactions. TYPES OF REACTIONS: 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 elements

Ex. 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 oxide Ex. 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 + water Ex. 2NaHCO3 Na2O + 2CO2 + H2O h. Metallic nitrite yields metallic oxide + nitrogen monoxide + oxygen gas Ex. 2Mg(NO2)2 2MgO + 4NO + O2 i. Group IA metal nitrate yield Group IA nitrite and oxygen gas

Ex. 2NaNO3 2NaNO2 + O2 j. Any other metallic nitrate yields metal oxide + nitrogen dioxide + oxygen gas Ex. 2Sr(NO3)2 2SrO + 4NO2 + O2 TYPES OF REACTIONS: REDOX REACTIONS __________ __________ __________ 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) __________ __________ : causes something else to be oxidized __________ __________ : causes something else to be reduced

Ex. Feo + O2o Fe2O3 Change in charges Iron 0 to 3+ to 2 Oxygen 0

Oxidizing agent = oxygen reducing agent = iron reduced = oxygen oxidized = iron DISPROPORTIONATION REACTIONS __________ __________ __________ 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+ + MnO41- Fe3+ + Mn2+ +

H2O Ex. 2. Manganese dioxide is added to conc. hydrochloric acid and heated. MnO2 + H1+ + Cl1- Mn2+ + Cl2 + H2O REDOX IN BASIC SOLUTIONS 1. 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 __________ - 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 LIGANDS __________ Ag(NH3)21+ Cu(NH3)42+ Ni(NH3)63+ __________ ______ Al(OH)41Zn(OH)42Cr(OH)63-

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+

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