Honors Chemistry

Arnoldi Honors Chemistry Final Review

Introductory Unit

1. Density = Mass Calculate the density of an object measuring 6.23 cm X Volume 5.00 cm X 3.56 cm and having a mass of 362.1 grams.

The density of an unknown object is 1.32 g/mL. What is the volume, if the mass of the object is 13.0 g?

2. Significant Figures Rules a. All non-zeros ARE significant. b. Zeros that are between significant figures ARE significant. c. Zeros that are AFTER the decimal point AND at the at the END of the number ARE significant. d. Zeros that are placeholders are NOT significant (for example at the end of the number, but before the decimal point, or after the decimal point, but not at the end of the number).

EXAMPLES: How many significant figures are in each of the following?

603 0.00063 6.03 60300 6000. 0.0603

3. Significant Figures Math Rules a. The answer to a X or / problem must be in the least number of significant figures given. b. The answer to a + or – problem must be in the least number of decimal places given.

EXAMPLES: Calculate each of the following using significant figures.

a. (6.30 X 103) X (0.06030 X 10-5) =

b. 60.342 + 11.87 =

c. 18.0 g / 3.2 cm3 =

4. Remember, when using a measurement device, estimate one place further than the device measures.

What is the volume of this liquid?

30 mL __ 20 mL __ 10 mL __ 5. What is being International System American Equivalent Measurement Device measured? Unit Mass Length Time Liquid Volume Temperature

Prefix Abbreviation Relationship to Base Unit Kilo Deca Deci Centi Milli

EXAMPLE: Rank the following from largest to smallest.

10.0 m 216.50 cm 0.0445 km 0.14 dam

EXAMPLE: Subtract the following: 16.5 cL - 2.5 mL.

EXAMPLE: If 2.50 cm = 1.00 inches, how long is 1 foot in centimeters?

EXAMPLE: What is the Kelvin temperature equivalent to 30’C?

6. The density of gasoline is 0.70 g/mL. If a student conducted an experiment and determined the density of gasoline to be 0.67 g/mL, what was her percent error? Classification of Matter

1. Definitions Word Bank: Compound Matter Homogeneous Mixture Element Mixture

a. A pure substance composed of two or more elements chemically combined is called a(n) ______

b. A pure substance that cannot be reduced chemically is called a(n) ______

c. A physical combination of two or more substances is a(n) ______

d. A physical combination of two or more substances in which the properties are the same throughout is a(n)______

e. Anything that has mass and takes up space is ______

EXAMPLES: Identify each of the following as compounds, elements, solutions or mixtures.

a. Air b. NaCl c. KCl (aq) d. Brass e. K

2. More Definitions Word Bank: Chemical Change Physical Change Chemical Property Physical Property

a. A characteristic that differentiates one substance from another, that can be observed without changing the substances is a(n) ______

b. A characteristic that differentiates one substance from another, that cannot be observed without changing the substances is a(n) ______

c. A change that results in a new substance with new properties being formed is a(n) ______

d. A change that merely changes the appearance, but not the identity of the substance is a(n) ______

EXAMPLE: Label each of the following as a physical property, chemical property, physical change, or chemical change.

a. Burning Paper b. Melting Wax c. Dissolving Salt

d. Color e. Melting Point f. Length

3. Complete the following table. Separation Method Property the Separation is Based Examples of this type of Separation Upon

Chromatography

Distillation

Filtration Atomic Theory (One and Two)

1. Definitions Word Bank: Anion Cation Ion Atomic Mass Electrons Isotope

a. Any charged particle is called a(n) ______. A positively charged particle is a(n) ______. A negatively charged particle is a(n) ______These are all formed by losing / gaining ______.

b. An atom that has a set number of protons, but varying numbers of neutrons is a(n) ______.

c. The average mass of all of the isotopes of an element is called ______.

EXAMPLE: Complete the following table. Symbol Protons Electrons Neutrons Mass Atomic Number Number 39K 25 30 56 54 82 S-2 18

EXAMPLE: Given the following information, what is the atomic weight of iron?

Percent Abundance Isotope Exact Mass 5.82 Iron -54 53.9696 amu 91.66 Iron -55 55.9349 amu 2.19 Iron -57 56.9354 amu 0.33 Iron -58 57.9333 amu

2. Explain the relationship between energy and quantum numbers.

a. 1s vs 2s vs 3s

b. 2s vs 2p 3. Complete the following nuclear equations:

______+ 9Be  12C + 1n 4 6 0

234Pa  234 U + ______91 92

4. A sample of radioactive matter has a half-life of 3.5 days. How much of a 40 g sample will remain after one week.

5. A sample of radioactive matter has a half-life of 16 minutes. How long will it take a 26 g sample of decay to 3.25 g?

6. Give an example of a pair of isotopes.

7. How do we fill the orbital notations for the p and d sublevels?

How many pairs of electrons are in Cobalt?

8. What is the electron configuration of each of the following? What is the shorthand configuration of a and e?

a. Calcium

b. Iron

c. Cl-1

d. K+1

e. Copper 9. Using the basic pattern for completing a Lewis Dot Structure, draw the LDS for each of the substances below. 1 2 5 3 What are the most important exceptions? Sy 8 6 How do you know the number of electrons needed? 4 7

EXAMPLES: Draw the electron dot structure for each of the following:

a. Calcium b. Phosphorus c. Iron d. Cl-1 e. K+1

10. More Definitions Word Bank: Bright-Line Spectrum Continuous Spectrum

a. A spectrum in which one color blends into the next is called a(n) ______

b. A spectrum in which the colors are distinct lines is called a(n) ______

11. How is a bright-line spectrum formed?

12. Protons are (positively / negatively) charged and found in the (nucleus / shells) of an atom. Electrons are

(positively / negatively) charged and found in the (nucleus / orbitals) of the atom. Neutrons (are / are not) charged and are found in the (nucleus / orbitals) of the atom. Periodic Table

1. Groups of the Periodic Table Group Name Location of Valence Shell Ion(s) Formed Reactivity Special Chars. Group e- (oxidation number)

Alkali Column 1 1 +1 Very High Flame Color Metals

Alkaline Earth Column 2 2 +2 High X Metals

Halogen Column 7 7 -1 Very High Salt-Forming Group

Noble Gas Column 8 8 N/A Very Low Inert Gases Group

Transition Middle usually 1 or 2 Most Common +2, Not Very X Metals +3

2. Name an element that will have similar properties to magnesium.

3. What are the valence shell electron configurations of each of the main groups of the periodic table?

Given the following valence shell electron configurations, identify the following elements:

a. 2s1 b. 4s2(3d10)4p5 c. 3s23p2 d. 5s2(4d10)5p4 4. Show the increasing trends for each of the following: electronegativity, ionization energy, atomic size, metallicity. and electron affinity.

EXAMPLES: Label each of the following as a metal, nonmetal or semimetal.

EXAMPLES:

a. Which is bigger, Cl or S?

b. Which has lower electronegativity, Si or C?

c. What atom has higher ionization energy, K or Mg?

d. Which is more active, I or Br?

e. Which is less active, Rb or F?

f. Which is more metallic, K or Ca?

g. Which is smaller, Ca or Ca+2?

h. Which is bigger, Cl or Cl-1?

Bonding 1. An ionic bond forms from the combination of ______or ______It has an electronegativity difference that is ______which means that the two combining elements will be further apart on the ______In fact, the further apart, the more ionic!

A covalent bond forms from the combination of ______, including ______

It has an electronegativity difference that is ______which means that the two combining elements will not be far apart on the periodic table. To determine between a polar covalent bond and a nonpolar covalent bond, you can use electronegativity differences: polar covalent bonds have differences of ______while nonpolar covalent bonds have differences that are ______. It is harder to see this on the periodic table, but in general, nonpolar bonds are either the same element attached to itself, or two elements that are

______on the periodic table.

Metallic bonds form within a PURE sample of a metal. This means that it is only ______type of atom involved.

EXAMPLES: Classify each of the following bonds as ionic, polar-covalent, nonpolar covalent or metallic bonds.

a. Na – Cl

b. C – Cl

c. C – C

d. C – O

e. In Na

f. Na – OH

g. Which pair of ions forms the most ionic bond: K to Cl or Fr to F?

2. When completing Lewis Dot structures for compounds, there are several important things to remember. For example, you can only bond on a side that has ______Also, hydrogen forms only ______bond. All other elements that we combined covalently resulted in ______pairs of electrons. If four pairs of electrons were not present, then ______bonds were required.

When determining molecular shapes, you know that the key is to count the number of pairs of electrons on the central atom. The most common shapes we discussed are shown in the table on the next page.

Total Pairs of Shared Pairs Unshared Shape Name e- of e- Pairs of e- Linear 1 or 2 1 or 2 0

Trigonal 3 3 0 Planar Bent 3 2 1 Tetrahedral 4 4 0 Trigonal 4 3 1 Pyramidal Bent 4 2 2

We also need to identify the polar / nonpolar nature of the molecule. We know that if there are any unshared pairs of electrons on the central atom, then the molecule must be ______. Otherwise, we look at the symmetry of the bonds. Remember, on the polar bonds, put in the arrows facing the more electronegative element. Then, are the arrows symmetric or asymmetric? (Will the cart move?) Symmetric arrows correlate to

______molecules, while asymmetric arrows correlate to ______molecules.

Finally, we need to know what type of IMFs occur within a sample of a pure compound. If the compound is polar, the IMFs will be ______. A special dipole/dipole is Hydrogen-Bonding which occurs when hydrogen is bonded to ______, as in water. If the compound is nonpolar, the IMFs will be ______or ______

EXAMPLE: What is the Lewis Dot Structure for each of the following? What shape is the molecule for a, c, and d? What IMF occurs in each?

a. PH3 b. C2H2

c. CO2 d. H2S

3. Write a formula for a substance that contains both ionic and covalent bonds. Explain your answer.

4. How can a molecule made of polar bonds be nonpolar?

5. Compare characteristics of ionic and covalent substances. Nomenclature

1. When given a name, and writing the formula for ionic compounds, remember to place the positive ion first and the negative ion second. Then, put in the oxidation numbers. These numbers can come from the periodic table, if it is in group one, two, three, five, six, or seven, they can come from the Roman Numeral, if it is a transition metal, or lead, or tin, or they can be memorized if it is a polyatomic ion. Once the oxidation numbers are put in, crisscross them to become the subscripts on the other ion. Remember to use ( ) if you are putting a number on a polyatomic ion. Lastly, reduce if it is possible.

When given a name, and writing the formula for covalent compounds, remember the prefixes give you the subscripts. You do NOT use oxidation numbers, and you do NOT crisscross. None/mono = 1. Di = 2. Tri = 3. Tetra = 4. Penta =5. Hexa = 6. Hepta = 7. Octa = 8. Deca = 10.

When given the formula, and writing the name for ionic compounds, remember to place the names in the same order as they appear in the formula. Remember, that if a transition metal, tin or lead appears, you will need a Roman Numeral. The Roman Numeral indicates the charge of the metal ion, and while it is sometimes possible to determine this by sauce-apple-cross-criss, using algebra will always get you the correct answer. Also, remember that unless it ends in a polyatomic ion, the ending is always “ide”.

When given the formula, and writing the name for covalent compounds, remember to keep the elements in the same order. Then, add prefixes as needed. (The first element NEVER gets mono). The last element ends in ide. In our naming, you will never have more than two elements getting named using covalent rules! (See prefixes above).

EXAMPLES: Given the formula, write the name. Given the name, write the formula.

a. Oxygen Dichloride g. Fe(ClO4)2

b. Sodium Phosphite h. KNO3

c. Sulfur Trioxide i. SnO

d. Calcium Sulfide j. Ba(NO2)2

e. Potassium Dichromate k. CCl4

f. Mercury (II) Oxide l. Fe2O3

2. Calculate the mass of each of the following.

. a. BaCl2 b. (NH4)2SO4 c. Mn3(PO4)2 4H2O 3. What is the oxidation number of each element in the following compounds and ions?

a. CF4

-2 b. SO3

c. HClO3

4. A compound consists of 38.5 % potassium, 13.85 % nitrogen, and 47.48 % oxygen. What is its empirical formula?

5. What is the percent by mass of sulfur in H2SO4?

. 6. What is the percent of water in H2C2O4 2H2O?

7. How many oxygen atoms are in Cu(ClO3)2?

8. List possible elements for X and Y in the following substance: XY3.

9. A compound has an empirical formula of C1H3. If the Molecular Mass = 30.0 g, what is the molecular formula? Reactions

1. Definitions:

The substances that you begin a reaction with are called the ______.

The substances that you create in the reaction are called the ______.

A solid that forms from the combination of two aqueous substances is called a(n) ______.

When a substance is dissolved in water, its phase is ______, or (__).

EXAMPLE: Identify the reactants, products and phases in the following reaction.

2 AgNO3 (aq) + Mg(OH)2 (aq) --> Mg(NO3)2 (aq) + 2 AgOH (s)

2. Classify, complete and balance the following equations as is appropriate.

a. MgCl2 + K2CO3  MgCO3 + KCl

b. AlCl3  Al + Cl2

c. C3H4 + O2  CO2 + H2O

d. Li2O 

e. Na + S 

f. Cu + H2SO4 

g. Ca + NiCl2 

h. MgS(aq) + Ba(ClO2)2 (aq) 

i. LiOH (aq) + H2SO4 (aq) --> Moles

1. How many moles are in 3.85 grams of FeSO3?

2. 140 moles of NaH2PO4 is equal to how many formula units?

28 3. How many grams are contained in 5.31 X 10 formula units of Mg(OH)2?

4. How many formula units are in 35.0 g of KCl?

5. What volume of hydrogen gas (careful) at STP is occupied by 14.5 grams?

6. If we have the same number of moles of a substance, regardless of phase, what must be true?

7. What is the mass of one mole of Barium Nitrate? Stoichiometry

1. Calculate the number of grams of zinc chloride that can be prepared from 34.0 grams of zinc given the following balanced equation.

Zn (s) + 2HCl (aq) --> ZnCl2 (aq) + H2 (g)

2. Calculate the number of moles of barium sulfate that can be prepared from 60.0 grams of barium chloride. Use the following balanced equation.

BaCl2 (aq) + Na2SO4 (aq) --> BaSO4 (s) + 2NaCl (aq)

3. A 35.0 g sample of calcium hydroxide reacts with a 54.0 g sample of phosphoric acid according to the following reaction.

3 Ca(OH)2 + 2H3PO4 --> Ca3(PO4)2 + 6H2O

a. What mass of water can be made?

b. What is the percent yield if 15 grams of water are actually produced? 4. How many moles of oxygen are needed to burn 1.65 moles of propane? You must first balance the following equation. Assume STP conditions.

C3H8 (g) + 5O2 (g) --> 3 CO2(g) + 4 H2O (g)

5. 15 Liters of oxygen gas and 10 grams of hydrogen gas react to produce liquid water at STP. What will be in the vessel when the reaction is complete? (use grams for water) Phases of Matter

1. Use the following graph to answer questions a – d.

1.60 __ (A) 1.40 __ (D) 1.20 __

1.00 __

a. What is the normal boiling point of substance “A” ?

b. Which substance evaporates the fastest?

c. Which substance has the stronger IMFs?

d. Which substance has the higher vapor pressure at the same temperature?

2. Complete the following problems using the constants given. Substance Melting Point Boiling Point Solid Cp Liquid Cp Water 0 ‘C 100 ‘C 2.092 J / g’C 4.184 J / g’C

What is the energy change when 52.0 g of water are cooled from 90 ‘C to 25 ‘C?

3. Memorize the following table:

Characteristic Solid Liquid Gas Energy low middle high Density highest middle lowest Volume definite definite indefinite Shape definite indefinite indefinite Fluidity low middle high Particle Movement low middle high Space b/w Particles closest middle furthest Diffusion slowly medium fastest 4. Answer the following questions based upon the graph given.

T (5) E M P Boiling (4) E Point R (3) A T (2) U Melting R Point (1) E (in’C)

HEAT (in joules or calories)

a. What is happening in section (1)? What phase(s) is/are present in section (1)?

b. What is happening in section (2)? What phase(s) is/are present in section (2)?

c. What is happening in section (3)? What phase(s) is/are present in section (3)?

d. What is happening in section (4)? What phase(s) is/are present in section (4)?

e. What is happening in section (5)? What phase(s) is/are present in section (5)?

f. What is the melting point? What is the boiling point?

5. Define Boiling Point.

6. Describe the following phase changes in terms of energy / enthalpy changes.

a. Solid  Liquid

b. Liquid  Solid

c. Liquid  Gas

d. Gas  Liquid

e. Gas  Solid

f. Solid  Gas 7. Why do we heat a liquid in order to cause it to boil? (Explain using the relationship of vapor pressure to temperature).

8. Given the following phase diagram, answer the question below.

1.60 __ (C)

1.40 __ (D) 1.20 __

1.00 __ (B) 0.80 __(A)

On curve BC, what will happen if you increase pressure? Describe using phase change terms.

On curve BC, what will happen if you increase temperature? Describe using phase change terms.

Under what conditions will something on the …

BC line freeze?

BD line boil? Gas Laws

1. What are the five assumptions of an Ideal Gas? Under what conditions is a gas most likely to behave ideally?

2. Standard temperature is ______or ______. Standard pressure is ______, ______, ______, ______or ______.

3. Memorize the following table: Law Conditions Compared Relationship Equation What is held constant?

Avogadro's Principle V, n direct none T, P

Boyle’s P, V Inverse Po Vo = Pn Vn T, n

Charles’ V, T Direct Tn Vo = To Vn P, V

Gay – Lussac’s T, P Direct Pn To = Po Tn V, n

4. Problems

a. A gas has a volume of 401 mL when the pressure is 1.0 atm, and the temperature is 25’C. If temperature is unchanged, but volume is decreased to 250 mL, what will the pressure be?

b. A gas has a volume of 5.01 L at 25’C. If pressure doesn’t change, what volume will the gas occupy at 10’C?

c. A gas is contained in a bottle. It has a pressure of 2.0 atm at 32 ‘C. At what temperature, in degrees Celsius, will the pressure be 1.5 atm? d. What is the molar mass of a 0.750 g sample of a gas that occupies 500 mL at 35 ‘C and 575.0 mmHg?

e. A 30.0 L sample of a gas exerts a pressure of 250 mmHg at 10’C. what volume will the gas have at 310 mmHg and 20’C?

f. At what pressure will 0.750 moles of hydrogen occupy 14.5 L at 0’C?

g. Three of the primary components of air are carbon dioxide, nitrogen and oxygen. In a sample containing a mixture of these gases at one atmosphere of pressure, the partial pressures of carbon dioxide and nitrogen are given as 0.285 mmHg and 593.525 mmHg, respectively. What is the partial pressure of oxygen?

5. If moles and temperature are kept constant, what happens to volume if pressure is cut in half? What happens if that pressure is then held constant, while temperature (in Kelvin) is doubled? Solubility

1. Use the solubility graph from your book (or old homework packet) to answer the following questions.

a. Name a substance that has an endothermic Heat of Solution.

b. Name a substance that has an exothermic Heat of Solution.

2. Complete each of the following calculations.

a. What is the molarity of a solution consisting of 200.0 g HCl in 982 mL of water?

b. What is the percent by mass of a solution containing 112 g of MgCl2 in 335 mL of water?

3. Complete the following paragraph about solubility factors.

Three factors are important in determining solubility. The first is the phrase “Like Dissolves Like”. This is referring to the arrangement of electrons in particles. We know that (think ionic, polar and nonpolar)

______. The universal solvent is ______because it can dissolve both ______and ______. Pressure also affects solubility, but only for ______. In gases, as pressure increases, the solubility of the gas in a liquid solvent ______. Finally, temperature affects solubility. For gas solutes in liquid solvents, as temperature ______, solubility decreases. For solid solutes in liquid solvents, changing temperature can either increase solubility or decrease solubility. The specific case is dependent on ______. If the heat of solution is positive, then as temperature increases, solubility

______. If the heat of solution is negative, then as temperature ______, solubility increases. Most heat of solutions for solids are positive, therefore for most solid solutes, ______temperature increases solubility.

4. What is the net ionic equation for the following reaction?

2HNO3 (aq) + Mg(OH)2 (aq) --> Mg(NO3)2 (aq) + 2H2O (l) 5. Give examples of substances that can dissolve in each of the following: (explain your reasoning)

a. Water

b. Alcohol (polar solvent)

c. Toluene (a nonpolar solvent)

6. I want to dissolve most solids in water. What should I do to temperature and pressure?

7. I want to dissolve a gas in water. What should I do to temperature and pressure?

8. What affect does adding a solute have upon the boiling point and freezing point of a substance? Acids and Bases

1. Combining a solution of known concentration to a solution of unknown concentration, in order to determine the unknown concentration is ______.

2. Convert Names to Formulas and Formulas to Names as needed.

a. H2CO3 a. Hydrobromic Acid b. KOH b. Magnesium Hydroxide

c. Mn(OH)2 c. Perchloric Acid

d. HFO2 d. Cobalt (II) Hydroxide

e. HI e. Nitrous Acid

3. The five properties of an acid are:

4. The five properties of a base are:

5. Complete the following equations. Label each as Ionization, Dissociation, Single Replacement, or Neutralization.

a. H2SO4(aq) + NH4OH (aq)

b. KOH (s) + H2O (l)

c. K (s) + H3PO4(aq)

d. Fe(OH)3 (s) + H2O(l)

e. H2SO4 (l) + H2O (l) 6. Answer each of the following problems.

a. The pH of a solution is 1.8. What is the pOH, [H+], and [OH-]? Is this an acid or a base?

b. What is the pH, pOH, [H+], and [OH-] if 0.00001 moles of calcium hydroxide completely dissociate in 8.50 L of water?

c. How many mL of 0.20 N H2SO4 would be required to titrate 50.0 mL of 0.10 N NaOH?

d. 35 mL of 1.0 N KOH is required to titrate a 14 mL sample of an unknown acid. What is the normality of the acid?

e. What volume of a 1.00 Molar solution of H2SO4 is required to titrate 35 mL of a 2.50 M solution of

NH4OH?

f. I have 25 mL of a 2.50 M solution of KOH. I want 30 mL of a 1.50 M solution. How much of my KOH will I need? 7. In each of the following equations, identify the Bronsted /Lowry Acid and Bronsted/Lowry Base pairs.

+ - a. HI + H2O H3O + F

-1 -1 b. H2PO4 + H2O H3PO4 + OH

-2 8. What is the conjugate acid and base of H2O? HPO3

9. What is the relationship between Ka and strength of an acid?

10. What is the relationship between molarity and strength of an acid? Molarity and concentration of an acid? Kinetic Theory

1. You need to know how substances react. Answer the following questions for the mechanism below.

(1) Br2  2Br Fast

(2) Br + H2  HBr + H Slow

(3) H + Br2  HBr + Br Fast

H2 (g) + Br2 (g)  2HBr (g)

a. ______Which step is the rate-determining step?

b. ______What happens to the rate if [Br2] is increased?

c. ______What happens to rate if Pressure is increased?

d. ______What is H?

2. What is a rate determining step?

3. Explain HOW a catalyst works using ideas from the thermodynamics unit. Thermodynamics

1. You need to be able to interpret a potential energy diagram.

Energy __ (joules) 40 __

Reactants  Products Time

EXAMPLE: a. Identify each of these terms on following graph.

Reactants

Products

Activation Energy

Activated Complex

b. What is the value of the EA? ______

c. What is the value of the Hrxn? ______

d. Is this reaction endo- or exothermic? ______

2. Calculate the Heat of Reaction using Hess Law (the short way): Don’t forget to balance! The heats of formation are as follows:

CH4 = -74.8 kj/mol CHCl3 = -132 kj/mol HCl = -92.0 kj/mol

CH4(g) + Cl2(g)  CHCl3 (l) + HCl (g)

o 3. The Hf for NaCl is – 411.0 kJ/mole. What is the change in enthalpy when 11.50 g of Sodium reacts with an excess of Chlorine? Equilibrium

1. You need to know how to write equilibrium expressions.

EXAMPLE: Write the equilibrium expression for the following reactions.

a. 2 CO (g) + O2 (g)  2CO2 (g) + 167 kJ

b. N2(g) + 3H2(g)  2NH3(g)

2. You need to be able to interpret equilibrium equations in terms of Le Chatelier’s Principle.

EXAMPLE: 2 CO (g) + O2 (g)  2CO2 (g) + 167 kJ

a. Fill in the table for each of the reaction above.

Stress Equilibrium Shift Change in K Value What happens to [CO2]?

Pressure Decreases

[CO] Increases

Temperature Decreases

b. To produce the most carbon dioxide, would you want each of the following to be increased, decreased, or left the same?

(i) Pressure ______

(ii) Temperature ______

(iii) [O2] ______3. For the dissociation of Iron (II) Chlorate, (endothermic to dissolve)

a. Write the dissolving equation.

b. Fill in the blank for each situation:

What can I do to __ to make more product?

[Fe+2]

[Fe(ClO3)2]

Pressure

Temperature

If I add Sodium Chlorate to this system, what will happen?

4. For the equilibrium system described by the equation: 2 SO2(g) + O2(g)  2 SO3(g), at a particular temperature, the equilibrium concentrations are 0.75 mol/L , 0.30 mol/L and 0.15 mol/L, respectively. What is the equilibrium constant for the reaction? Electrochemistry

Given the following reaction:

HCl(aq) + Mg(s)  MgCl2(aq) + H2(g)

a. Who is oxidized?

b. Who is reduced?

c. Who is the oxidizing agent?

d. Who is the reducing agent? Organic Chemistry

1. What are the three functional groups we discussed? What is the key structural / formula feature for each?

2. What are the three products of all (complete) organic combustion reactions?

3. How does boiling point relate to the number of carbon atoms in a hydrocarbon?