CHEMISTRY REGENTS REVIEW—WIDK I. MATTER & ENERGY

The potential energy of a system is considered to be the HEAT of the system.

11. Gas law problems - Combined gas law: (Temp must be in Kelvin)

P1 V1 =P2 V2

T1 T2

12. Boyle's Law - (constant temp) P and V vary inversely à P1 V1 = P2V2 [See diagram #1]

13. Charles Law - (constant P) V and T vary directly

Temp must be in Kelvin  V1 = V2 - or- V1 = T1

T1 T2 V2 T2

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[See diagram #2]

16. Sublimation a substance turns directly from a solid to a gas

ex. CO2(s) CO2(g); I2(s) purple crystals I2(g)purple gas

17. Phase change diagrams à

a. Melting/Boiling [See diagram #3 & 4]

b. Freezing/Condensation [See diagram #3 & 4]

18. Kinetic molecular theory

Ideal gases [How gases should behave but don’t]-

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a. No attraction between molecules/atoms b. Molecules have a negligible volume c. Collisions are elastic d. Particle movement is random

Real gases VERY RARELY BEHAVE LIKE IDEAL GASES since

a. There IS an attraction between particle (van der Waals) b. The volume of particles are NOT negligible, esp. at low temps & high-pressure

since atoms/molecules are close together

***HYDROGEN and HELIUM are the most IDEAL gases.*** Also, Diatomic molecules and nonsymmetrical molecules & noble gases act the most ideal. THE SMALLER THEY ARE THE MORE IDEAL THEY BEHAVE.

19. Heat of fusion - the amount of calories needed to melt one gram of a solid; for H2O it is 80 cal/g [See Reference table A]

20. Heat of vaporization - the amount of calories needed to vaporize one gram of a solid; for H2O it is 540 cal/g [Reference table A]

21. Boiling point - the temp. at which the vapor pressure of a liquid = The atmospheric pressure: for H2O look at Table O. The normal boiling point when the atmospheric pressure = 760 mm Hg = 100o C

22. Vapor pressure - depends on the

a. Temperature of the liquid b. Strength of intermolecular forces (i.e. the stronger the van der Waals forces the stronger

the Intermolecular forces are)

23. Law of partial pressures-Dalton's Law à the sum of all the pressures in a mixture of gases is equal to it's total pressure à Ptot = P1+ P2 + P3

II. ATOMIC STRUCTURE

16. Hund's rule - before an orbital can get a second electron each orbital in that subshell must have at least one in each.

17. Order of filling sublevels: 1s2 2s2 2p6 3s2 3p6 4s2 3d10: WHY? The 4s2 sublevel needs less energy to fill than the 3d10 sublevel.

18. Principle energy levels - [See diagrams #7a and #7b]  

Diagram #7a-->

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Diagram #7b-->

III. BONDING1. When a bond is formed energy is released (exothermic); when a bond is broken energy is absorbed (endothermic)

3. Metals tend to lose electrons and form positive ions.     (Ions formed are smaller than the neutral atoms: Ionic radii < than atomic radii)

4. Nonmetal tend to gain electrons (GER) and form negative ions (ions formed are larger than the neutral atoms: Ionic radii > atomic radii)

5. A chemical bond - results from the simultaneous attraction of electrons by two nuclei

6. Ionic bonds - formed between metal and nonmetal; created by a transfer of electrons; electronegativity difference > 1.7

7. Covalent bond - formed by the sharing of electrons; electronegativity difference < 1.7

9. Exception to 1.7 rule: METAL hydrides are ionic! ex. NaH

10. Diatomic molecules are considered to have NONPOLAR covalent bonding. i.e. N2 à N=N

12. Coordinate covalent bonds - a covalent bond where both of the electrons are donated by one of the elements. [See diagram #8]. Usually found in polyatomic ions.

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18. Network solids: held together by covalent bonds; high melting & boiling points. ; Extremely poor conductors of heat & electricity. i.e. SiO2, diamond - tetrahedral bonding (Cn), graphite (Cn) - hexagonal bonding

19. Van deer Waals forces - attractive forces that exist between ALL particles. They increase when particles à

a. Increase in mass b. Get closer together c.

It's like GRAVITY!

20. Hydrogen bonds - attractive for btw. Molecules that contain hydrogen and atoms of small atomic radius and HIGHELECTRONEGATIVITIES. i.e. H2O and HF. These bonds result in some compounds having higher boiling points than expected.

21. Polar molecules - molecules in which there is a localization of charge that causes part of the molecule to be slightly positively charged [d+]and part of the molecule to be negatively charged[d-]. Tug of war where somebody wins [See diagram #9] These are usually NONsymmetrical molecules ex. H2O, HF, NH3

22. Nonpolar molecule - there may still be localization of charge but there is no NET movement of electrons in any particular direction. This is a tug of war where no one wins.

IV. PERIODIC TABLE

1. Periodic law - states that elements are arranged on the periodic table according to their atomic numbers and chemical properties.

2. Elements are classified in 3 categories a. Metals - left of stairs b. Nonmetals - right of the stairs

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c. Metalloids - touching the stairs 3. Trends - as you go from left to right across the table in a period

a. Metallic character decreases b. Atomic radius decreases [See Table P] c. Ionization energy increases [See Table K] d. Electronegativity increases [See Table K]

4. As you go down a group a. Metallic character increases b. Atomic radius increases [See Table P] c. Ionization energy decreases [See Table K]

5. Group IA metal - alkali metals; strongest bases; form +1 ions 6. Group IIA - alkali earth metals; form +2 ions 7. Group O metal - inert or noble gases; generally non-reactive. Kr and Xe can form some

bonds in the laboratory. 8. Group VII -halogens - contain elements in ALL three phases. F & Cl are gases, Br is a

liquid and I is a solid 9. Elements in the same period fill up the SAME principle energy levels 10. The most active metals are in the lower left corner. 11. The most active nonmetals are in the upper right corner. 12. The MOST active elements for the MOST stable compounds! i.e. RbF

17. Transition elements -

a. Produce COLORED SOLUTIONS. b. found in the middle of periodic table c. emit color in flame test as electrons fall back DOWN from the excited state. d. lose both s & d electrons & therefore have multiple oxidation states

17. Van der Waals forces increase as you go down a group since the size of the atom increase. This causes the boiling and melting points to increases as well. Remember this when you get to ORGANIC chemistry.

18. Atomic radius decreases as you go across a period since there is an increase of nuclear charge (# of protons) which pulls the electrons in closer thereby shrinking the size of the atom.

V. STOICHOMETRY AND MATHEMATICS IN CHEMISTRY7. Solution - homogeneous mixture (evenly mixed) 8. Unsaturated solution - holds less solute than the maximum 9. Saturated - holds the exact amount of solute the solvent can hold 10. Super-saturated - holds more than the maximum amount of solute 11. Concentrated solution - holds a large amount of solute 12. Dilute solution - holds a little amount of solute 13. Solubility of a solid- (ability to dissolve) generally increases as temperature increases. 14. Solubility of a gas increase as temperature decreases and pressure increases. Think of

when soda goes flat (CO2 escapes)

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15. Boiling point elevation - for every mole of substance dissolved in solution the boiling point increase by .520. [See chart A in reference tables]

16. Freezing point depression - for every mole of substance dissolved in solution the freezing point decreases by 1.860. [See chart A in reference tables] 

17. When figuring out boiling point elevation and freezing point depression keep in mind that electrolytes (molecules that split into ions) create more moles in solution than the would seem to. [See diagram #12]

18. Finding the empirical formula from percentages. 1. Divide the percentages by the atomic masses (see periodic tables) 2. Divide the resulting numbers by the smallest result and this gives you your ratio

for the empirical formula. 19. Finding the molecular formula from percentages. You MUST be given the total mass to

do this 1. Divide the percentages by the atomic masses (see periodic tables) 2. Divide the resulting numbers by the smallest result and this gives you your ratio

for the empirical formula. 3. Figure out what the empirical formulas mass is and see how many times it goes

in to your total mass.

VI. KINETICS AND EQUILIBRIUM

1. Heat of reaction (DH)- the difference between the potential energy of the reactants and the products

2.(does NOT change with the addition of a catalyst)

3. Diagrams of exothermic and endothermic reactions. [See diagrams # 13 & 14]

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4. Exothermic reactions à release energy, (DH) = -, products formed are MORE stable compounds than the reactants

5. Endothermic reactions à absorb energy, (DH) = +, products formed are LESS stable compounds than the reactants

1. Gibb's equation DG = DH - T DS states whether or not a reaction occurs spontaneously or not. If DG is negative the reaction will occur spontaneously and if DG is positive the reaction will occur nonspontaneously. When DG = O the system is at equilibrium

11.

Equilibrium constant equation: Keq =  ProductsReactants

12. When Keq is large that means that the reaction favors the products. [[See bottom of Table M]]

13. When Keq is small that means that the reaction favors the reactants. 14. Remember that the coefficients in front of the compounds become the exponents in the

equilibrium constant equation. [See diagram #15]

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15. Solubility product equation - Ksp = Dissociated ions ONLY [Ions are charged particles; +/-] 16. When Ksp is large that means that the reaction favors the dissociated. More dissolved in. 17. When Ksp is small that means that the reaction favors the non-dissociated part of the

equation. 18. Ionization constant for acids- same as solubility product constant but you use Ka

instead. 19. When Ka is large that means that the reaction favors the dissociated. This is a strong

ACID. 20. When Ka is small that means that the reaction favors the non dissociated part of the

equation. This is a weak ACID. [See Table L]

VII. IONIC SOLIDS1. Bronsted-Lowry Theory

a. Acid = proton donor (losses H+ ) b. Base = proton acceptor (gains H+)

6. Organic compounds- begins with C. i.e. C6H12O6 - usually NOT electrolytes. Except organic acids [functional group –COOH]

7. Traits of Acids

a. Turns blue litmus red b. pH less than 7.0 c. Reacts with metals (below H on chart N) to form salt and H2 gas d. Taste sour e. Reacts with base to form salt and water (neutralization) f. The more they ionize, the better they conduct electricity g. They contain more H+ (H3O+) than (OH-)

8. Traits of bases

a. Turns red litmus blue, pink in phenolthalein b. pH greater than 7.0

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c. Reacts with acids - neutralization d. Taste bitter e. Feel slippery f. The more they ionize, the better they conduct electricity g. They contain more OH- than H+

9. Ionization Constant of water (Chart M) = Kwà [H+] x [OH-] = 1 x 10-14. Use this to figure out pH. [See diagram #16]

10. pH scale [See diagram #17]

14. Amphoteric or amphiprotic - substance can behave as an acid or a base. Found on both sides of Chart L

15. Finding pH of salt solutions: Find pH of sodium carbonate (Na2Cl3) Solution: NaOH is a strong base; HCl is a strong acid so the pH of the resulting solution will be ~7

16. Hydrolysis reaction (opposite of a neutralization reaction) Salt + water --> acid + base [See diagram #18]

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17. It should be noted that Group IA and IIA are strong bases when combined with OH; Bases [OH combined with a metal] get weaker as you move across the periodic table from left to right.

VIII. REDOX AND ELECTROCHEMISTRY1. Sum of the oxidation states in a neutral atom must always equal ZERO. 2. Oxidation - loss of electrons causes the oxidation # to increase (LEO) 3. Reduction - gaining of electrons causes the oxidation # to decrease.(GER) 4. Spectator ions - ions that are not involved in being reduced or oxidized. 5. Chart N- has ONLY reduction reactions, in order to change them into oxidation reactions

you must flip them and change to sign of the Eo value. The strongest reducers are on the TOP of the chart and the strongest oxidizers are on the BOTTOM of the chart.

6. Only metals below H2 will react with acids to produce Hydrogen gas. 7. Hydrogen is used as the standard on which the entire table is based. 8. To calculate the Eo of a cell first determines which one of your elements is the substance

being reduced and which one is being oxidized. Flip the sign on the element being oxidized and add them up

9. If Eo is + then the reaction is spontaneous. 10. If the Eo is - then the reaction is non-spontaneous. 11. If the Eo = 0 then the system is at equilibrium 12. Electrochemical cell - [See diagram # 19], spontaneous, electrons flow to better

reducer, salt bridge allows for the migration of ions in BOTH directions to sustain the reaction. Cathode is (+) electrode & the anode is the (-) electrode.

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13. Electrolytic cell - need a battery to get going, Anode is (+) electrode & the cathode is the (-) electrode.

14. Electroplating - plating occurs at the reduction or negative electrode. Car bumpers can be coated with protective metal in this manner. Mass increases at the site of plating and decreases at the oxidation or positive electrode. [See diagram # 20]

15. Balancing Redox equations - balance with respect to charge and mass. [See Diagram # 21]

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16. The substance being reduced is considered to be the oxidizing agent. 17. The substance being oxidized is considered to be the reducing agent. 18. RED CAT; AN OX 19. To figure out the Eo value for a redox reaction you should use the following reaction Eo

total

= Eoreduced - Eo

oxidized

IX. ORGANIC CHEMISTRY1. Benzene series - CnH2n-6 ring structure, toluene is related, [See diagram # 22]

1. Alkyl radials (side groups) regular prefixes but they end in -yl. 2. As the molecular mass of each of these homologous series increase so to do their boiling

points and melting points due to an increase in the van der Waals forces.

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3. Properties of organic molecules - non-electrolytes, low boiling points. & melting points ., insoluble in polar solvents (like water), react slowly & are molecular in structure.

4. Isomers - have the same chemical formula but a different structural formula, which means that they behave differently

5. Primary Alcohol's - OH group bonded to a carbon atom that is in turn bonded to ONE (or none) other carbon atom.

6. Secondary Alcohol's - OH group bonded to a carbon atom that is in turn bonded to TWO other carbon atoms.

7. Tertiary Alcohol's - OH group bonded to a carbon atom that is in turn bonded to THREE other carbon atoms.

8. Organic reactions to know 1. Addition - adds a pair of halogens to an unsaturated hydrocarbon. One product. 2. Substitution - adds a halogen to a saturated hydrocarbon. Two Products. 3. Esterfication - acid + alcohol à ester + water 4. Fermentation - C6H12O6à alcohol + carbon dioxide 5. Saponification - fat + base à soap + glycerol 6. Polymerization - n (C2H4) à (C2H4)n 7. Combustion - hydrocarbon + oxygen à carbon dioxide + water 8. Cracking --> the separation of a polymer.

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