Chemistry Unit 2 Exam

Chapters 4 & 5

Synthesis Chem. Rxn

• Type of chem. rxn

• 2 components → one product

• A + B → AB

Decomposition Chem. Rxn

• Type of chem. rxn

• 1 compound → 2 products

• AB → A + B

Single Replacement Chem. Rxn

• Type of chem. rxn

• Nonmetals switch

• AB + C → B + AC

Double Replacement Chem. Rxn

• Type of chem. rxn

• Metals switch partners

• AB + CD → AD + CB

Solutions

Homogenous mixtures of 2+ pure substances

Solvent

Present in greatest abundance in a solution

All other substances are solutes

Aqueous Solution

When water is the solvent

Dissolving Solutes

• There are attractive forces between solute particles holding them together

• There are attractive forces between solvent molecules

• When mixing solute w/ solvent, there are also attractive forces between solute particles and solvent molecules

• If the attractions between solute and solvent are greater than the other two, the solute will dissolve

Aqueous Solutions and Dissolving Solutes

• All substances dissolve by solvation (surrounding of solute by solvent)

• Ionic compounds dissolve by dissociation, where water surrounds separated ions

• Molecular compounds disperse in water, but mostly remain intact

• Some molecular substances form ions in water when they dissolve

  • Don’t break into individual atoms but stay as a single unit

Limiting Reagent

In most chemical rxns, one substance will be limiting to the reaction.

Idea: if you have 3 bags of flour but 2 eggs, the eggs will limit how many pancakes you can make, while the flour is in excess

Typical Process:

Find mol of product produced with reactant 1

Find mol of product produced with reactant 2

etc….

Which produces a small #mol of product is the limiting reagent.

To find how much excess is actually used, perform stoichiometry from limiting reagent, molar mass, molar ratio, to find actual mol or g of other reactant that was used.

Theoretical, Actual, and & Percent Yield

Theoretical yield is how much of a product you would theoretically have after a chem. rxn

Actual yield is how much of a product you actually produce

Percent yield is the percent of the theoretical yield that was actually produced: (mass actual yield) / (theoretical yield) x 100%

Charge Distribution in Water Molecules

• Uneven distribution is e- within molecule

• Causes oxygen to be partially neg. and hydrogens partially pos.

• Makes water molecules polar

LIKE DISSOLVES LIKE

Non-polar solvents dissolve non-polar solutes

Polar solvents dissolve polar solutes

Electrolytes

• Dissociate: break off into smaller ions or atoms

• Electricity: flow of free e-

• Strong electrolytes: dissociate completely into ions when dissolved in H2O; conducts electricity

  • ex. NaCl

  • equation has single arrow

• Weak electrolytes: dissociates partially when dissolved in H2O; cannot conduct (not enough charged particles)

  • equation represents chem. equilibrium w/ equilibrium arrow

• Non-electrolytes: DOES NOT DISSOCIATE in H2O (may dissolve, not dissociate); does NOT conduct

  • Sugar

• Acids ionize in H2O to varying degrees

  • Completely ionize: strong acids

  • Don’t completely ionize: weak acids

Electrolytes and Solubility

Ionic compounds that are SOLUBLE are STRONG ELECTROLYTES

Ionic compounds that are INSOLUBLE are WEAK ELECTROLYTES

When determining strength of electrolyte, look at solubility rules

Weak and Strong Acids and Bases

Strong Acids:

• HCl

• HBr

• HI

• HNO3

• H2SO4

• HClO3

• HClO4

(SO I BRought NO CLean CLOthes)

SO4, I, Br, NO3, Cl, ClO3, ClO4

H2SO4, HI, HBr, HNO3, HCl, HClO3, HClO4

Strong Bases:

• LiOH

• NaOH

• KOH

• RbOH

• CsOH

• Mg(OH)2

• Ca(OH)2

• Sr(OH)2

• Ba(OH)2

(Like NAthan Knows, CAlifornia is SupeR BAsic)

Li, Na, K, Ca, Sr, Ba

LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

+ RbOH, CsOH, Mg(OH)2

(Main idea: most group 1A and 2A cations)

Weak Acids:

• Not strong acids

Weak Bases:

• No OH group

Empirical Method

Conducting experiments and creating rules based on observations

Solubility of Ionic Compounds

• When an ionic compound dissolves in water, the resulting solution contains component ions in solution, not intact ionic compound

• Some ionic compounds are insoluble, and do not dissolve in water

• Ionic compound in solution:

  • Precipitate is an insoluble ionic compound, which does not dissolve in water

    • Can be in suspension (floating) or collected at bottom

  • Supernate is liquid above collected precipitate at bottom

  • Solution (when no insoluble products present) can be colored, but always clear

Precipitation Reactions

• Reaction that yields an insoluble product (precipitate) when 2 solutions are mixed

• Reactions between aqueous solutions of ionic compounds produce insoluble ionic compounds

• No precipitation means no reaction

Predicting Precipitation Reactions

1. Determine which ions constitute each aqueous reactant

2. Determine formulas of possible products

• exchange ions and balance charges of combined ions to get formula of each product

3. Determine solubility of each product in water

4. If neither products will precipitate, write no reaction after arrow ( → NR)

5. If any possible products are insoluble, write formulas

6. Balance equation

Molecular Equations, CIE, & NIE of Precipitation Reactions

Molecular Equations: the original equation after determining solubility and balancing equation

Complete Ionic Equation: separate all parts into their ions w/ solubility & charges (ex. KI(aq) (ME) = K+(aq) + I-(aq) (CIE))

Net Ionic Equation: only include ions from CIE that are directly involved in producing the precipitate product

• Remove spectator ions (those not involved in formation of the precipitate)

NOTE**: precipitate never written out into its ions since the ions don’t dissociate at all

Aqueous Reactions

• Precipitation Reactions

• Acid-Base Reactions

• Gas Evolution Reactions

• Anion from one reactant combines with cation of another

Acid-Base Reactions

• Also “neutralization reaction”

• Acid reacts with base and they neutralize each other, producing water (or sometimes a weak electrolyte)

• In rxn of an acid & base:

  • H+ from acid combines w/ OH- from base to make water

  • Cation from base combines w/ anion from acid to make a salt

• NIE: H+(aq) + OH-(aq) → H2O(l)

  • (as long as salt that forms is soluble)

Hydronium Ions in Solution

• Acids ionize in water to form H+ ions

  • More precisely, the H+ from the acid molecule is donated to a water molecule to form hydronium ion (H3O)+

    • Most chemists use H+ and H3O+ interchangeably

Polyprotic Acids

• More than 1 ionizable proton and release them sequentially

• First ionizable proton strong, and subsequent ones are weaker

Ex. H2CO3, H3PO4, H2SO4, etc

more than one H in compound

Arrhenius Definition

• Acid: substance that produces H+ in aq solution

• Base: substance that produces OH- in aq solution

Bronsted-Lowry Definition

• Acid: substance that can donate a proton

• Base: substance that can accept a proton

• Water is amphoteric; it acts either as an acid or base

Solution Concentration

• Since solutions are mixtures, composition can vary one sample to another

• We quantify the amount of solute relevant to the solvent, or concentration of solution

• Common way to express solution concentration is with Molarity

  • M= (mol solute)/(L ENTIRE solution)

• Can be used as a conversion factor

Solution Dilution Equation

• Used when stock solutions (concentrated) need to be diluted, meaning more solvent is added

• Amount of solute doesn’t change, just volume

M1V1 = M2V2

Acid-Base Titrations

• Titration: lab procedure where substance in solution of known concentration (titrate) is reacted with another substance of unknown concentration (analyte)

• Equivalence point is point in titration where H+ and OH- from reactants are in their stoichiometric ratio and are completely reacted

• An indicator is a dye whose color depends on acidity or basicity of solution

Gas Evolution Reaction

• Gas is produced, results in bubbling

• Many GE rxns are AB rxns

• Some rxns form a gas directly from ion exchange

• Others form gas through subsequent decomp of one of the ion exchange products into a gas and water

  • The product that undergoes decomp is the intermediate

Types of Compounds That Undergo GE Reactions

• Sulfides

  • No intermediate product

  • H2S gas eveolved

• Carbonates/Bicarbonates

  • H2CO3 intermediate

  • CO2 gas evolved

• Sulfites/Bisulfites

  • H2SO3 intermediate

  • SO2 gas evolved

• Ammonium

  • NH4OH intermediate

  • NH3 gas evolved

Redox Reactions

• Electron transferred from one reactant to another

• Many involve the reaction of a substance with oxygen

• Electron transfer does not need to be complete transfer to qualify as a redox rxn

  • Can be unevenly shared

  • Ex. Lower e- density in H and higher e- density in Cl

• Oxidation: loses e-

• Reduction: gains e-

• (OILRIG)

• Oxidized component loses e- and reduced component gains e-

• Oxidized component is reducing agent; reduced component is the oxidizing agent

• Redox rxns are:

  • Any species that changes its oxidation state

  • Any metal + non-metal reacting

  • Sing-replacement rxns

  • MOST synthesis/decomposition rxns

  • NOT double replacement or A-B rxns

Oxidation States

• Imaginary # assigned to elements to determine loss/gain of e-

• Oxidation # of an atom in a compound is the “charge” it would have if all e- were assigned to the atom with the greatest attraction for those e-

• Oxidation states are imaginary charges assigned w/ rules, and written magnification then charge (ex. +1, -1, etc)

  • Note: ion charges are real, measurable charges, and written charge before magnitude (ex. 1+, 1-, etc)

Rules:

1. Free elements have oxidation state of 0 (ex. Cu=0, Cl2=0)

2. Monatomic ions have oxidation state equal to charge (ex. Ca2+=+2, Cl-=-1, etc)

3a. Sum of oxidation states of all atoms in a compound is 0 (ex. H2O, H=+1, O=-2, together=0)

3b. Sum of oxidation states of all atoms in a polyatomic ion equals charge of ion (ex. NO3-, N=+5, O=-2, together=-1)

4a. Group 1 metals have oxidation state +1 in all compounds (ex. Na=+1 in NaCl)

4b. Group 2 metals have oxidation state +2 in all compounds (ex. Mg=+2 in MgCl2)

5. In their compounds, non-metals have oxidation states according to table (higher on top, higher priority):

Nonmetal	Oxidation State	Example
Fluorine	−1	M g F 2 −1 o x state
Hydrogen	+1	H 2 O +1 o x state
Oxygen	−2	C O 2 −2 o x state
Group 7A	−1	C C l 4 −1 o x state
Group 6A	−2	H 2 S −2 o x state
Group 5A	−3	N H 3 −3 ox state

Activity Series

• Top is most reactive and has highest tendency to undergo oxidation

• Decreasing position on list, less reactive

• Half-rxns at top most likely to occur in forward direction

• Half-rxns at bottom most likely to occur in reverse direction

Kinetic Molecular Theory (KMT)

Simplest Model For Behavior of Gases

Postulates of Ideal Gases:

1. Gas particles spread out and so small relative to overall volume that their size and volume and negligible

2. Gas particles are in constant random motion

3. When gas particles hit one another or the walls, collisions are elastic (no KE lost)

4. Assumer there are no attractions between gas particles (just bounce off one another)

5. Average KE (speed) of gas particles is directly proportional to Kelvin temperature

• Not all particles moving at same speed (due to different masses), but avg KE increases as temperature increases

Negligible

RM

Elastic

Proportional

Attraction

KMT-PropANERM

Pressure

• Pressure: force exerted per unit area by gas molecules as they strike surfaces around them

  • Result of many gas molecules exerting forces on surfaces around them is a constant pressure

• Pressure is (Force)/(Area), F/A

• Gas pressure is result of constant movement of gas molecules and their collisions w/ surfaces around them

• Pressure of gas depends on:

  • # gas particles in given volume

  • volume of container

  • average speed of gas particles

• Total pressure exerted by gas depends on several factors including concentration of gas molecules

  • Higher the concentration, higher the pressure

  • As volume increases, concentration decreases (#particles is same, but concentration is decreasing), this results in less collisions and therefore lower pressure

  • Fewer gas particles, lower force per unit area, and so lower pressure too

    • Higher density of particles, higher pressure; lower density, lower pressure

• Gas pressure and temp. impact how gases behave

Common Pressure Units

Atmospheres - - - - - - - - - - atm - - - - - 1atm (avg air pressure at sea level)

Pounds per square inch - - - psi - - - - - - 14.7psi

Torr - - - - - - - - - - - - - - - mmHg/torr - - 760torr or 760mmHg

Inches of mercury - - - - - - - inHg - - - - - 29.92inHg

Pascal (1N/m²) - - - - - - - - - (Pa) - - - - - - 101,325Pa

Ideal Gas Law

PV=nRT

P= pressure (atm)

V= volume (L)

n= # moles (mol)

R= gas constant; 0.08206 (L x atm)/(K x mol)

• OR 8.314 (J)/(mol x K)

• OR 8.314 (L x kPa)/(mol x K)

T= temperature (K)

Can derive simple gas laws (which dictate behavior of gas if 2 variables change and others remain constant)

Simple Gas Laws

• Describe relationships between pairs of properties involved with gases

  • Properties interrelated

• Boyle’s Law: P1V1 = P2V2

• Charles’s Law: V1/T1 = V2/T2

• Gay-Lussac’s Law: P1/T1 = P2/T2

• Combined Gas Law: (P1V1)/T1 = (P2V2)/T2

• Avogadro’s Law: V1/n1 = V2/n2

Standard Temperature and Pressure (STP)

1 atm

273K

ANY gas at STP, molar volume is 22.4L (1mol=22.4L)

Molar volume: V=(nRT)/P; and w/ STP, n=1mol, T=273, and P=1atm, meaning V=273

Density of a gas

Tells you how “heavy” a gas is

d=(PM)/(RT)

M=(dRT)/P

Dalton’s Law of Partial Pressure

Ptotal = Pa + Pb + Pc +…

Pa = X(mole fraction)a + Ptotal

Pn = nn(RT/V)

Average Kinetic Energy

KE=(1/2)mv²

Avg KE = (3/2)RT

R= 8.314 (J/Kxmol)

Average KE is only reliant upon temperature, nothing else

Since there are different masses depending on the type of gas, the only way to keep the same avg KE is if lighter particles move faster while slower particles move slower. This allows gases to have the same avg KE. This can be seen through examples of the equation KE=(1/2)mv², because as m grows, v is smaller.

Root-Mean-Square Velocity

Avg speed of gas particles; depends on molar mass

At same T, all gases have same avg KE, but this is because they have different speeds.

Lighter molecules will move faster while slower particles move slower

Urms= sqrt(3RT/M)

Mean Free Path

Molecules in a gas move in a straight line until colliding

MFP is the average distance a molecule travels between collisions

So, as P increases, meaning the V decreases, MFP also decreases

Graham’s Law of Effusion

Diffusion: movement of gas molecules from an area of high concentration to an area of low concentration

Effusion: when a collection of molecules escape through a small hole in a vacuum, meaning the molecules are still traveling to areas of lower concentrations, but the process to getting there takes longer since molecules must escape through a small hole or a membrane first (meaning smaller particles escape quicker since they’re faster and smaller)

Law states that lighter particles move faster, and so they will escape quicker than heavier particles during effusion

r1/r2 = sqrt(M2/M1)

Real Gases and Ideal Behaviors

Real gases do not behave ideally under low T and high P

• Smaller volume, so particles are closer together (KMT states particles so small relative to overall volume that they’re size and volume is negligible; this breaks down when particles are in small spaces (they take up more space))

• Under low T, avg KE and Urms decrease, meaning particles move much slower. Since real gases can have polar molecules with intermolecular forces, this situation gives particles a better chance of attracting to one another since they move slower and will be close by for a longer period of time (KMT states that we should assume molecules do not have any intermolecular forces between them, but this breaks down if polar gas molecules are nearby and attract each other)