HONORS CHEMISTRY 2 FINAL EXAM

Time to Suffer!!!!!!

6- Properties of Covalent Compounds

6- Properties of Covalent Compounds

• sharing of e-

• occurs between elements that are similar

  • ex) between two nonmetals or metalloids

• never conductive as a solid or as a solution in water

• do not dissolve in water

  • exception: sugar

• shorter bond = higher force of attraction, higher bond E

• longer bond = lower force of attraction, lower bond E

6- Properties of Ionic Compounds

• transferring of e-

• nonmetal takes e- from metal

• forms crystal lattice

• brittle

• high melting/boiling point

• very strong bond

• when solid:

  • ions in fixed position

  • nonconductive

• when dissolved in water:

  • ions move freely

  • conductive

6- Properties of Metallic Bonding

• forms crystal lattice

  • nucleus of a metal

  • delocalized e- (sea of e-)

    • free to move about the entire crystal lattice

  • orbitals overlap

• malleable and ductile

• high melting/boiling point and heat of evaporation

• strong bonds

• conductive as solid

• insoluble

• luster

6- Lewis Structures

• ionic compound- composed of cations (metals) and anions (nonmetals)

• Lewis Structures must be electrically neutral

6- Polarity

• nonpolar molecules:

  • dipoles cancel

  • ends have same partial charges

• polar molecules:

  • dipoles do not cancel

  • ends have opposite partial charges

6- Shapes of Molecules (molecular geometry)

• linear

• bent

• trigonal planar

• trigonal pyramidal

• Tetrahedral

6- Intermolecular Forces

• London dispersion forces (LDF)

  • all molecules have LDF

  • weakest IMF

• dipole-dipole force

  • any polar molecule that doesn’t have H bonding has this

• hydrogen bonding

  • must be polar

  • must have hydrogen bonded to high energy anion (N, O, or F)

  • strongest IMF

7- Naming and Writing Chemical Formulas

• monatomic ions

  • element symbol + charge

  • ex) Na^+

  • ex) O^2-

  • ex) Zn(II) <- like this for all transition metals

• binary ionic compounds

  • cation + anion name with -ide

  • ex) Aluminum Oxide - Al_2 O_3

  • ex) Silver Chloride - AgCl

  • no number prefixes, since cation charge determines number of anions and vice-versa

• binary molecular compounds

  • first element in formula + second element in formula with -ide

  • inorganic prefixes

    • “mono” is understood for first element, but necessary for second

  • ex) dichlorine monoiodide

• ionic compounds with polyatomic ions

  • cation name + polyatomic ion name (anion)

7- Mass, Percent Composition, and Molar Mass Conversions

• percent composition

  • part/whole

  • ex) percent of water in a hydrate:

    molar mass H2O/molar mass of entire formula

• formula mass = molar mass

• Avogadro’s number = 6.022*10^23

• molecules to moles uses Avogadro’s number

  • 6.022*10^23 molecules/1 mol

• formula units to moles uses Avogadro's number

  • 6.022*10^23 formula units/1 mol

7- Empirical and Molecular Formulas

• empirical formula is the smallest whole number ratio of the subscripts in a compound

• molecular formula is the actual formula of the compound

• x(empirical formula) = molecular formula

  • x = molecular formula molar mass/empirical formula molar mass

• to find empirical formula:

  • take grams of all elements in compound and convert to moles (if given percentages, just pretend they are grams, it doesn’t matter)

  • divide all by the smallest amount of moles

  • if there is a _.5, multiply all by 2

  • use these numbers as subscripts for the formula (1s are implied)

8- Evidence of a chemical change

• color change

• gas production

• change in state of matter (formation of a precipitate)

• energy production/consumption

8- Chemical Equations

• reactants on left, products on right

• balanced when the amount of each element is equal on both sides

• must contain states of matter symbols

• coefficients- the only things that change

  • balance atoms that appear only once on each side of the equation first

  • balance polyatomic ions as single units

  • balance H and O last

  • if an atom can be balanced with a coefficient of 1/2, use it, then multiply all coefficients by 2 at the end

8- Predicting Products

• synthesis

  • to or more substances combine to make a compound

  • 1 product

  • A + X → AX

• decomposition

  • a single compound undergoes a reaction that produces two or more simpler substances

  • 1 reactant

  • AX → A + X

• single displacement

  • one element replaces a similar element in a compound

  • USES ACTIVITY SERIES

    • one element can only replace a similar element if that similar element is lower on the activity series (back of periodic table)

  • metal replacement

    • A + BX → AX +B

  • halogen replacement

    • Y +BX → X + BY

• double displacement

  • ions of two compounds trade places in an aqueous solution to form new compounds

  • AX + BY → BX + AY

• combustion

  • a substance combines with oxygen, releasing a large amount of energy in the form of heat or light

  • CxHy (g) + O2 (g) → CO2 (g) + H2O (g) +Energy

  • only contains elements hydrogen, oxygen, and carbon

• remember to balance!!!!!!!!

9- Yields

• actual yield- the amount of product, affected by human/equipment error attained from a reaction in real-life experiment

• theoretical yield- the supposed amount of product attained from a reaction, based purely on its mol ratio with a known quantity of another reactant or product

• percent yield- way to exemplify the accuracy of an actual yield by comparing it to the theoretical yield

  • percent yield= (actual/theoretical)*100

9- Limiting Reactants

• coefficients indicate mol ratio

• limiting reactant- the reactant whose difference in a have/need table has a negative sign

  • know how to apply have/need table

• excess reactant- the reactant whose difference in a have/need table has a negative sign

  • know how to apply have/need table

10.5, 12, and 13- Solute-Solvent Combinations

• solutions

  • homogenous mixture

  • smallest particle size

    • 0.01-1 nm

  • cannot be separated through filtration

  • particles aren’t large enough to settle out

• colloids

  • appears to be one phase, but is actually heterogenous on the particle level

  • intermediate particle size

    • 1-1000 nm

  • cannot be filtered or settled

  • some separation occurs if left to settle long enough

  • exhibits tyndall effect

    • particles large enough to scatter light (foam, fog, gel, etc.)

  • suspensions

  • visibly heterogenous

  • ex) oil and water mixture

  • can be filtered

  • settles unless consistently agitated

  • largest particle size

    • 1000 nm or greater

10.5, 12, and 13- Rate of Dissolving

• rate of dissolving- how much solute dissolves in solvent over time

• factors that effect ROD:

  • ^temp. = ^ROD

    • ^kinetic energy (particle velocity) = more particle collisions

  • ^surface area = ^ROD

    • ^contact = more particle collisions

  • agitation = ^ROD

    • ^contact = more particle collisions

• solubility curve- graph to model rate of dissolving

  • y-axis: amount of solute

  • x-axis: time

10.5, 12, and 13- Solubility

• the amount of a substance, in grams, required to form a saturated solution with a specific amount of solvent

  • dependent on:

    • nature of solute and solvent

    • temp. of solvent

• solubility equilibrium- solute is recrystallizing at the same rate that the crystal is dissolving

10.5, 12, and 13- Electrolytes/Nonelectrolytes

• electrolytes- when dissolved in water, the resulting solution conducts electricity

  • ionic compounds

  • acids

  • bases

• nonelectrolytes- when dissolved in water, the resulting solution is nonconductive

  • molecular compounds

10.5, 12, and 13- Common Solute-Solvent Combinations

• gas-gas

• gas-liquid

• liquid-liquid

• liquid-solid

• solid-liquid

• solid-solid

10.5, 12, and 13- Miscibility

• like dissolves like

  • polar dissolves in polar

  • nonpolar dissolves in nonpolar

• ONLY APPLICABLE WITH LIQUIDS

10.5, 12, and 13- The Dissolving Process

• 1. dissociation

  • neg. end of water molecule attracted to cation, pos. end of water molecule attracted to anion

  • pulls apart molecule

• 2. hydration

  • water molecules surround anions and cations, oxygen faces cations and hydrogen faces anions

10.5, 12, and 13- Enthalpy of Solution

• enthalpy of solution- the net amount of E absorbed when a specific amount of a solute dissolves

• exothermic

  • neg. value

  • the E of separation of solute and solvent particles is less than the E produced

• endothermic

  • pos. value

  • the E of separation of solute and solvent particles is greater than the E produced

10.5, 12, and 13- Solubility of Gases

• Temperature

  • inversely proportional with solubility

• Pressure

  • directly proportional with solubility (Henry’s Law)

10.5, 12, and 13- Density of Ice

• 0.92 g/ml

10.5, 12, and 13- Net Ionic Equation

• total ionic equation, barring the spectator ions (ions that are identical on both sides of the equation)

10.5, 12, and 13- Molarity

• unit of concentration

• derived unit- mol/liter

• expressed as “M”

17 and 18- Collision Theory

• things that must be obtained in order for 2 reactants to make a product:

  • 1. sufficient energy (different for every reaction)

  • 2. correct orientation

17 and 18- Reaction Mechanisms

• elementary steps- steps that make up a reaction mechanism

• to add elementary steps together, intermediates must be eliminated

  • intermediates- substances that are produced in an early step and consumed in a later step

• slowest step- rate determining step

• even if a reaction is reversible, the reactants are on the left and the products are on the right

  • reversible reaction- a reaction that can be executed both ways, denoted by double-headed arrow

• catalyst- consumed in early step and produced in later step

17 and 18- Energy Profile Diagrams

• y-axis: energy (usually in KJ)

• x-axis: course of reaction/reaction progress

• Ea: reactant energy to activated complex

  • activated complex- sufficient energy, peak of graph

  • activated complex lowered by catalyst

• Ea’: product energy to activated complex

• delta E forward = E of products - E of reactants

  • when neg., reaction is exothermic (product E is less than reactant E)

  • when pos., reaction is endothermic (product E is greater than reactant E)

• delta E reverse = same value as delta E forward, opposite sign

17 and 18- Equilibrium

• chemical equilibrium- when the rate of the forward reaction is = to the rate of the reverse reaction

• conc. of products and reactants is unchanged

• it doesn’t matter if the amount of reactants and products are equal, only that the rate of exchange between them is equal

• equilibrium constant:

  • ratio of products and reactants at a given temp. (since there is no conc. change at equilibrium)

  • this ratio is referred to as an equilibrium expression

  • does not include pure solids/liquids

  • uses molarity

  • denoted by K (no units, no sig figs)

  • small K values = reactants favored

  • large K values = products favored

  • K = 1 means reactants and products are equal

17 and 18- LeChatlier’s Principle

• LeChatlier’s principle- if a system at equilibrium is subjected to stress, the equilibrium is shifted in the direction that tends to relieve its stress

• stress: increase conc.-

  • increase reactant- forward, no effect on K

  • increase product- reverse, no effect on K

• stress: decrease conc.-

  • decrease reactant- reverse, no effect on K

  • decrease product- forward, no effect on K

• stress: change in temp.

  • increase temp.- reverse, K gets smaller

  • decrease temp.- forward, K gets bigger

• stress: pressure-

  • increase pressure- forward, no effect on K

  • decrease pressure- reverse, no effect on K

14 and 15- Classifying Acids and Bases

• arrhenius acid

  • produces a hydrogen ion: H+

  • OR

  • produces a hydronium ion: H3O+

  • these are the same thing but proving your claim with a dissociation equation will result in a H+ proof and an ionization equation will result in a H3O+ proof

• arrhenius base

  • produces a OH- ion

  • if ionic: prove with dissociation equation

  • if molecular: prove with ionization equation

• bronsted-lowry

  • acid: proton donor

  • base: proton acceptor

  • proton = H+ or H3O+ ion

• acids and bases are reactants, never products

• acids always have “H” as first listed element in formula

• lewis

  • acid: electron pair acceptor; no unshared pairs on central atom

  • base: electron pair donor; unshared pairs on central atom

• conjugate acid- the product that has received the proton

• conjugate base- the product that has lost the proton

14 and 15- More Acid/Base Definitions

• amphoteric- acts as an acid or a base, dictated by the substance it is reacting with

  • ex) H2O, also written as HOH (has both H+ and OH-)

• nonmetal oxides will react with water to form an acid- can be shown through a simple synthesis equation

  • this is the process for acid rain

• MEMORIZE: 3NO2 (g) + H2O (l) → 2HNO3 (aq) + NO (g)

  • only nonmetal oxide + water equation that won’t just synthesize

14 and 15- Titration

• titration- the controlled addition and measurement of the amount of a solution of known conc. required to completely react with a measured amount of unknown conc.

• equivalence point- moles acid = moles base

  • “end” of titration

  • reached when indicator turns color

14 and 15- Strong vs. Weak Acids and Bases

• strong acids:

  • completely dissociate

  • strong electrolyte (lots of ions in solution)

  • identified by:

    • list of strong acids on periodic table (MEMORIZE!!!)

    • equation has single headed arrow (this is your proof)

• weak acids:

  • slightly dissociate

  • weak electrolyte (few ions in solution)

  • identified by:

    • any acid not on strong acid list

    • equation has double headed arrow (this is your proof)

• strong bases:

  • completely dissociate

  • strong electrolyte (lots of ions in solution)

  • identified by:

    • contains group 1/2 metal and OH- (this is your proof)

• weak bases:

  • slightly dissociate

  • weak electrolyte (few ions in solution)

  • identified by:

    • contains metal not from group 1/2 and OH- (this is your proof)

14 and 15- Equivalence Point

• strong acid and strong base-

  • equivalence point = pH 7

• weak acid and strong base-

  • equivalence point > pH 7

• strong acid and weak base-

  • equivalence point < pH 7

14 and 15- Toolbox Square

14 and 15- Titration Stoichiometry Sig Fig Rules

• of sig figs in conc. (molarity) = # of decimal places in pOH or PH

14 and 15- Weak Acid Stoichiometry

• since weak acids do not completely dissociate, you can’t use a mole ratio

• you must use a RICE table

  • remember: no pure liquids or solids in RICE tables!!!!!!

14 and 15- Endpoint

• the pH indicator is the color that indicates complete titration