Glen Cove High School- Science Department
Chemistry Regents Review

This web site is designed to SUPPLEMENT the instruction that you received in school NOT replace it. These pages cover the topics you will be responsible for in June in a very broad manner. For further clarification, please make sure you go over your notes, old tests and go to speak to your teachers. Remember they are there to help you if you have any questions.

Unit I: Matter and energy

1. Properties of solids - definite shape & volume, fixed atoms; regular geometric pattern

2. Properties of liquids - no definite shape, but definite volume

3. Properties of gases - no definite shape or volume, random particle motion

4. Elements - all atoms have the same ATOMIC #. Can NOT be broken down chemically

5. Mixture - 2 or more elements physically combined. There are different types of mixtures

1. heterogeneous (uneven - lumpy iced tea)
2. homogeneous (evenly mixed à SOLUTION - clear tea)

6. Physical change - no change in the identity of the substance (i.e. gas à liquid à solid)

7. Chemical change - substance changes into new substance with NEW properties (H₂ + O₂ à H₂O: Chemical reaction)

8. For calorie problems: know the following formula àQ = mDt

Q = calories, m = mass, Dt = change in temperature

9. Temperature (Kinetic Energy)- know how to convert from Celsius to Kelvin (+ 273) and back (- 273)

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

10. Fixed points on a thermometer - O^o C - freezing/melting point of H₂O;

100^o C is the boiling/condensation point of H₂0; you need 2 points to create a thermometer

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

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 à

Temp must be in Kelvin

[See diagram #2]

14. STP - standard temp and pressure

1. Temperature à O^o C -or- 273 K [On Table C]
2. Pressure à 760 torr -or- 760 mm Hg -or- 1 ATM [On table A]

15. Density:
 

16. Sublimation à a substance turns directly from a solid to a gas ex. CO_2(s) à CO_2(g); I_2(s) purple crystalsà I_2(g)purple gas

17. Phase change diagrams à

18. Kinetic molecular theory

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

1. No attraction between molecules/atoms
2. Molecules have a negligible volume
3. Collisions are elastic
4. Particle movement is random

Real gases VERY RARELY BEHAVE LIKE IDEAL GASES since

1. There IS an attraction between particle (van der Waals)
2. 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 = 100^o C

22. Vapor pressure - depends on the

1. Temperature of the liquid
2. 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 P_tot = P₁+ P₂ + P₃

Unit II - Atomic structure

1.Parts of the atom à

1. Proton - (+) charged; 1 atomic mass unit
2. Neutron - (+/-) charged; ~1 atomic mass unit
3. Electron - (-) charged; 1/1836 atomic mass unit

2. Nucleon - particles found in the nucleus (protons & neutrons)

3. Nucleus - contains most of the mass of the atom; has a positive charge; The # of protons is called the nuclear charge

4. 1 AMU - the atomic mass unit à based on 1/12 the mass of a carbon 12 atom; on top of the periodic table

5. In a neutral atom the # of protons = the number of electrons.  All the elements on the periodic table have = #’s of protons & electrons as listed. 

6. Atomic # - the # of protons in an atom; used to identify the element

7. Atomic mass = the # of protons + the # of neutrons

8. Isotopes - elements that have the same atomic # but different atomic masses due to a difference in the # of neutrons in the nucleus.

9. To figure out the # of neutrons in an element subtract the atomic # FROM the atomic mass.

₆C¹⁴ has 6 protons, 6 electrons and 8 neutrons

10. Atomic mass is really a weighted average of all of the isotopes that exist in nature for that element. i.e. Carbons atomic mass = 12.011 because there is ₆C¹² and ₆C¹⁴in nature but ₆C¹² is more abundant and therefore skews the average toward 12.

11. Empty space concept - states that atoms are made up of mostly empty space and most of the mass is confined to a very small nucleus. This was proven by the gold foil experiment. [See diagram #5]

12. Bohr's model of the atom - stated that electrons traveled in certain orbits. An absorption of energy will cause electrons to TEMPORARILY jump to higher levels & when the electrons fall back down to lower levels they EMIT this energy in the form of light.

13. Valence electrons - electrons in the outermost energy levels. i.e. ₉F¹⁹ à 1s² 2s² 2p⁵ à has 7 valence electrons SINCE the outer most principle energy level is the 2^nd one. Kernel electrons are the electrons that orbit the nucleus of atom and are NOT considered to be part of the valence shell.

14. Electron dot diagram - uses dots for the valence electrons. [See diagram #6]

15. Orbital diagrams à uses boxes to illustrate the orbit electrons can take around the nucleus. Arrows represent the electrons & two electrons or arrows can fit into each box or orbital. The electrons in the same orbital MUST spin in opposite directions.

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-->

Diagram #7b-->

Unit III - Bonding

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

2. Atoms bonded together to form OCTETS (eight valence electrons are stable s² p⁶ à 8 valence electrons)

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 and from 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

8. Electronegativity - the affinity for electrons. Highest: Fluorine 4.0 [See Table K]

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

11. Helium & Hydrogen need only 2 electrons to fill its outer shell. All the others need 8 electrons.

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.

13. Ions: K⁺ and Cl^- have the same # of electron (18) since formation of ions are caused by the loss or gaining of ELECTRONS.

14. Ionization energy: the amount of energy required to remove the outermost electron from an element. [See Table K]

15. Ionic solids: high melting & boiling point; hard; do not conduct electricity UNLESS dissolved in water -or- in molten form.

16. Metallic solids: mobile electrons, conductors in solids phase, malleable, ductile, only metal that is a liquid at room temp à Hg

17. Molecular solids: held together by van der Waals forces; low melting & boiling points; poor conductors; are soft. ex. Sugar C₆H₁₂O₆

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

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

1. Increase in mass
2. Get closer together
3. 
   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.

23. Formula writing - use the crisscross method. [See diagram #10]

Unit 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
1. Metals - left of stairs
2. Nonmetals - right of the stairs
3. Metalloids - touching the stairs
3. Trends - as you go from left to right across the table in a period
1. Metallic character decreases
2. Atomic radius decreases [See Table P]
3. Ionization energy increases [See Table K]
4. Electronegativity increases [See Table K]
4. As you go down a group
1. Metallic character increases
2. Atomic radius increases [See Table P]
3. Ionization energy decreases [See Table K]
5. Metalloids - have both metal and nonmetal properties. Contact the "staircase".
6. Group IA metal - alkali metals; strongest bases; form +1 ions
7. Group IIA - alkali earth metals; form +2 ions
8. Group O metal - inert or noble gases; generally non-reactive. Kr and Xe can form some bonds in the laboratory.
9. Group VII -halogens - contain elements in ALL three phases. F & Cl are gases, Br is a liquid and I is a solid
10. Elements in the same period fill up the SAME principle energy levels
11. Elements in the same groups have the same # of valence electrons
12. The most active metals are in the lower left corner.
13. The most active nonmetals are in the upper right corner.
14. The MOST active elements for the MOST stable compounds! i.e. RbF
15. Monatomic molecules (one atom) à He, Ne, Ar, Kr, Rn

17. Transition elements -

1. Produce COLORED SOLUTIONS.
2. found in the middle of periodic table
3. emit color in flame test as electrons fall back DOWN from the excited state.
4. 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.

Unit V - Stoichiometry and mathematics in chemistry

1. Mole = 22.4 liters at S.T.P. & contains 6.02 X 10²³ molecules
2.  

3. Chart C - gives you densities of some gases
4. Avogadro's Law - equal volumes of gases contain equal # of molecules
5. Volume - volume problems à set up a ratio. [See diagram # 11]

7. 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. 
    [See Table d & E]
15. Solubility of a gas increase as temperature decreases and pressure increases. Think of when soda goes flat (CO₂ escapes)
16. Boiling point elevation - for every mole of substance dissolved in solution the boiling point increase by .520. [See chart A in reference tables]
17. Freezing point depression - for every mole of substance dissolved in solution the freezing point decreases by 1.860. [See chart A in reference tables] 

    

18. 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]
19. How do you know when a substance is an electrolyte? If it is ionically bonded it is an electrolyte. i.e. NaCl (salt) or HCl (acid) or NaOH (base)
20. Molecular formula - the actual # of atoms in the covalently bonded molecule. i.e. C₆H₁₂O
21. Empirical formula - shows the simplest ratio of atoms in a molecule.
22. 
    i.e. C₆H₁₂O₆ à CH₂O
23. 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.
24. 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.
25. Percentage comp. - Total mass of the element in the compound x 100 = total mass of the compound
26. Percent error - (Good for group 12 questions!)

Unit 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]





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
6. If the heat is listed on the right side (with the products) the reactions is exothermic.
7. If the heat is listed on the left side (with the products) the reactions is endothermic.
8. Factors effecting the reactions rate
1. Catalyst - speeds up the reaction by reducing the activation energy needed to start a reaction. A catalyst does NOT effect the heat of reaction or the potential energy of the products or the reactants.
2. Increasing the concentration of one of the substances à shifts the equilibrium away from the increase to the other side of the reaction while decreasing the concentration of ALL of the other compounds on the side of the increase.
3. increase in temperature à shifts the equilibrium away from the heat. Favors the endothermic reaction.
4. Increase in pressure à shifts the equilibrium to the side with the least number of moles.
5. Increase in surface area à increases the reaction rate in both directions {like pounding it into a powder]
9. Entropy (DS)- the randomness of a system. If (DS) is + then there is an increase in entropy or Randomness and if (DS) = - then there is a decrease.
10. Order of increasing entropy: solidsà liquidsà gas
11. 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

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



15. Solubility product equation - K_sp = Dissociated ions ONLY [Ions are charged particles; +/-]
16. When K_sp is large that means that the reaction favors the dissociated. More dissolved in.
17. When K_sp 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 K_a instead.
19. When K_a is large that means that the reaction favors the dissociated. This is a strong ACID.
20. When K_a is small that means that the reaction favors the non dissociated part of the equation. This is a weak ACID. [See Table L]

Unit VII - Acids and Bases

1. Electrolyte - a compound that breaks into ions in solution or when melted. Usually ionically bonded.
2. Non-electrolyte - a compound that does not break into ions in solution or when melted. Covalently bonded
3. Arhennius theory of
1. Acid à gives of a H⁺ ion, as the ONLY positive ion
2. Base à gives off an OH^- ion
4. Bronsted-Lowry Theory
1. Acid = proton donor (losses H⁺ )
2. Base = proton acceptor (gains H⁺)
5. Salt - a metal combined with a nonmetal [ex. NaCl, Na is the metal & Cl is the nonmetal]

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

1. Turns blue litmus red
2. pH less than 7.0
3. Reacts with metals (below H on chart N) to form salt and H₂ gas
4. Taste sour
5. Reacts with base to form salt and water (neutralization)
6. The more they ionize, the better they conduct electricity
7. They contain more H+ (H₃O+) than (OH-)

8. Traits of bases

1. Turns red litmus blue, pink in phenolthalein
2. pH greater than 7.0
3. Reacts with acids - neutralization
4. Taste bitter
5. Feel slippery
6. The more they ionize, the better they conduct electricity
7. They contain more OH- than H+

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



10. pH scale [See diagram #17]



11. Neutralization

1. Acid + Base à salt + water
2. H⁺ + OH^- à H₂O (net reaction)

12. In neutralization, moles of acid and moles of base must be equal.
13. Formula for titration (neutralization)

14. List of conjugate acid-base pairs on Chart L - Strongest acid =largest K_{a and} Weakest acid = smallest K_a
15. Amphoteric or amphiprotic - substance can behave as an acid or a base. Found on both sides of Chart L
16. 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
17. Hydrolysis reaction (opposite of a neutralization reaction) Salt + water --> acid + base [See diagram #18]



18. 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.

Unit VIII - Redox and electrochemistry

1. Know the rules for determining the oxidation states.
2. Sum of the oxidation states in a neutral atom must always equal ZERO.
3. Oxidation - loss of electrons causes the oxidation # to increase (LEO)
4. Reduction - gaining of electrons causes the oxidation # to decrease.(GER)
5. To have a Redox reaction there must be a change in oxidation # and you CANNOT have oxidation without having reduction.
6. Spectator ions - ions that are not involved in being reduced or oxidized.
7. Chart N- has ONLY reduction reactions, in order to change them into oxidation reactions you must flip them and change to sign of the E^o value. The strongest reducers are on the TOP of the chart and the strongest oxidizers are on the BOTTOM of the chart.
8. Only metals below H₂ will react with acids to produce Hydrogen gas.
9. Hydrogen is used as the standard on which the entire table is based.
10. To calculate the E^o 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
11. If E^o is + then the reaction is spontaneous.
12. If the E^o is - then the reaction is non-spontaneous.
13. If the E^o = 0 then the system is at equilibrium
14. 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.



15. Electrolytic cell - need a battery to get going, Anode is (+) electrode & the cathode is the (-) electrode.
16. 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]



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



18. The substance being reduced is considered to be the oxidizing agent.
19. The substance being oxidized is considered to be the reducing agent.
20. RED CAT; AN OX
21. To figure out the Eo value for a redox reaction you should use the following reaction E^o_total = E^o_reduced - E^o_oxidized

Unit IX - Organic Chemistry

1. Hydrocarbons - contain only hydrogen and carbon
2. Homologous series - successive members differ by -CH₂ groups.
3. Alkanes - C_nH_2n+2 contain ALL single bonds - saturated compounds, ending is - ane.
4. Alkenes - C_nH_2n contain one double bond - unsaturated compounds ending are - ene.
5. Alkynes - C_nH_2n-2 contain one triple bond - unsaturated compounds ending is - yne.
6. Benzene series - C_nH_2n-6 ring structure, toluene is related, [See diagram # 22]



7. Naming compounds - [See diagram #23]



8. Alkyl radials (side groups) regular prefixes but they end in -yl.
9. 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.
10. Properties of organic molecules - non-electrolytes, low boiling points. & melting points ., insoluble in polar solvents (like water), react slowly & are molecular in structure.
11. Isomers - have the same chemical formula but a different structural formula, which means that they behave differently
12. General formulas:

a.Alcohol's à R –OH

b.Organic acids à R-COOH

c.Ester à R₁ - COO - R₂

d.ketones à R₁ – CO – R₂ [the oxygen is double bonded to the carbon]

e.ethers à > R₁ – O – R₂ [the oxygen connects two carbon chains]

f.aldehydes R – COH [the oxygen is double bonded to the carbon]

13. Types of alcohol's
1. Monohydroxl - one OH group
2. Dihydroxl - two OH groups
3. Trihydroxyl - three OH groups. (Glycerol) [See diagram #24]
14. 
15. Primary Alcohol's - OH group bonded to a carbon atom that is in turn bonded to ONE (or none) other carbon atom.
16. Secondary Alcohol's - OH group bonded to a carbon atom that is in turn bonded to TWO other carbon atoms.
17. Tertiary Alcohol's - OH group bonded to a carbon atom that is in turn bonded to THREE other carbon atoms.
18. 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.

Copyright © 2001 Glen Cove School District
Dosoris Lane
Glen Cove, NY  11542
Last Updated: 02/10/2006 10:23 AM