Learning Commons
` 3 Alkynes
A m i d es
Ar o
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Hydration of Alkenes
T hio
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Salt
Acetal Formation
5 USER GUIDE 7 UNIT I ACIDS & BASES 1.1 – ACIDS & BASE CHEMISTRY 1.2 – GRAM EQUIVALENTS-­‐NORMALITY-­‐TITRATION 1.3 – PH -­‐ POH – [H+] – [-­‐OH] 1.4 – BUFFERS ALKANES 1.5 – NAMING ALKANES 1.6 – FUNCTIONAL GROUPS 1.7 – DRAWING STRUCTURES 1.8 – CONFORMERS & ISOMERS PRACTICE TESTS 1A 1B 9 11 13 15 17 19 21 23 25 27 29 31 39 UNIT II 47 ALKENES, ALKYNES & AROMATICS 2.1 – NAMING ALKENES, ALKYNES, & AROMATICS 2.2 – CIS/TRANS 2.3 – POLYMERS 2.4 – REACTIONS OF ALKENES, ALKYNES, & AROMATICS 49 51 53 55 57 ALCOHOLS, PHENOLS, ETHERS, THIOLS, & HALOGENS 2.5 – NAMING ALCOHOLS, PHENOLS, ETHERS, THIOLS, & HALOGENS 2.6 – REACTIONS OF ALCOHOLS, PHENOLS, ETHERS, THIOLS, & HALOGENS 59 61 63 PRACTICE TESTS 2A 2B 65 73 UNIT III 81 AMINES, ALDEHYDES, KETONES, CARBOXYLIC ACIDS, ESTERS, & AMIDES 3.1 – NAMING AMINES, ALDEHYDES, & KETONES 3.2 – NAMING CARBOXYLIC ACIDS, ESTERS, & AMIDES 3.3 – REACTIONS – TESTS (LUCAS, BENEDICTS, & TOLLENS) 3.4 – REACTIONS OF CARBONYLS 83 85 87 89 91 CARBOHYDRATES 3.5 – CARBOHYDRATE NOMENCLATURE 3.6 – REACTIONS OF CARBOHYDRATES 3.7 – CONVERTING FISCHER TO CYCLIC PRACTICE TESTS 3A 3B 93 95 97 99 TOOLS NOMENCLATURE COMMON FUNCTIONAL GROUPS SIGNIFICANT FIGURES PERIODIC TABLE OF THE ELEMENTS 101 109 117 119 120 121 123 7 Thank you for choosing to use this workbook provided by the PLUS program at Eastern Washington University! The purpose of this book is to provide a medium to which students and tutors can enhance their subject knowledge by focusing on fluency and speed, which are fundamental to success on exams. This workbook is intended for use by a tutor who has successfully completed the course in question, in a 1-­‐on-­‐1 or small group format. In addition, this workbook assumes that the reader has access to a textbook for explanations not provided in this material. This book is intended to be supplemental to a textbook for use by a tutor and a student. NO TE:
Worksheets: Ø
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1 page of problems on a focused topic FRONT = BLANK BACK = ANSWER KEY Print off/copy for student Or work on with student Consult textbook for detailed explanation of topic Practice Exams: These worksheets and
practice exams are
intended to be used with
a knowledgeable tutor
and therefore, the most
problems encountered
are NOT meant to be
easy.
The theory: if a student
Ø 3-­‐4 pages of simulated exam type questions can complete difficult
o There are two types of exams: practice problems with a
§ A – very similar to worksheets tutor, they can excel
§ B – more integrative concepts & analysis independently on a
Ø TREAT IT LIKE A REAL EXAM standard exam level.
o Grab a study carrel or another quite place o No drinks/food, cell phones, textbooks, or computers Ø Allow no more than 50-­‐55 minutes to complete. Ø Answer key immediately follows the blank test Ø Grade exam for students just like a professor would Ø Or grade exam with student – learn from mistakes! In order to get the most out of the practice exams, it is imperative to simulate a test-­‐like environment when taking these exams. Many students unfortunately suffer from test anxiety and much of this anxiety can be avoided by increased knowledge of material, as well as experience with testing environments. Once the student has completed an exam, the tutor may choose to grade the exam during a normal session while the student works on other materials (such as worksheets) and return the graded exam to the student. Alternatively, a tutor could evaluate the exam with the student. It is not advised, however, to simply provide the student with the answer key without a tutor to review the exam also. This is due the simple fact that it is hard to catch our own mistakes. Also, because the practice exam does not directly affect the student’s course grade, the path of least resistance (neglecting to evaluate) is easy to be taken. The point of this book is to help identify problem areas and provide a means to address and remedy them. Happy tutoring!
9 11 Common Proton acceptors (Bases): • NaOH Common Proton • Mg(OH)2 Donators (Acids): • Ca(OH)2 • HClO4 • NH3 • HBr Amount of Titration acid /base • HCl that contains 1 mol of ions • HI M1V1=M2V2 • H3PO4 Equivalents • H2SO4 !"#$%#&'( #!! !" !"! !"#$%&'$
1 Eq = !"#
• CH3COOH Buffers [!! ]
pH = pKa + log!
! Gram Equivalents 1 g-­‐Eq = !"#$% !"## !" !"#$ !" !"#$
Potential #H+ or OH− produced
[!"]
Be able to recognize conjugate acid/base pairs Normality 1 N = Ex: Acetic Acid – Conj Acid Acetate – Conj Base ! !"
L of soln
= M x (#H ! or OH ! ) pH = -­‐log[H+] pH + pOH = 14 . [H+][OH-­‐]=10-­‐14 10-­‐pOH = [OH-­‐] Log Rules! pOH = -­‐log[OH-­‐] log(x) = y à x = 10y log(AB) = log(A) + log(B) !
log!!! = log(A) – log (B) 10-­‐pH = [H+] 0 -­‐-­‐ ACIDIC -­‐-­‐ 7 -­‐-­‐ BASIC -­‐-­‐ 14 Predicting Reaction Directions Rxn will always favor the formation of the weaker acid & base Stronger Stronger Weaker Weaker Acid Base …….…. Acid Base 13 1.1 – Worksheet
Con jugate Acid/Base Pairs
Please complete the following table by providing the correct names and formulas of the following acids or bases. Name Acid Conjugate Base Name HI H2PO4 Hydrocyanic Acid Water HNO3 Ammonia Carbonic Acid –
HCO3
Name Acetic Acid Hydrobromic Acid Conjugate Acid HCl HClO4 Base –OH –NO
2 2–SO
4 Name Phosphate Ion In the following equations, label the correct Brønsted-­‐Lowery acid/base. Then, label the conjugate acid/base. HClO3 + H2O Mg(OH)2 + +H3O H2CO3 + –HO HI + NH3 –ClO + +H O 3
3
2+Mg + H
–HCO
2O 3 + H2O +NH + –I 4
1.1 – Worksheet – KEY
Con jugate Acid/Base Pairs
Please complete the following table by providing the correct names and formulas of the following acids or bases. Name Acid Conjugate Base Name –I Hydroiodoic Acid HI Iodide Ion –
Dihydrogen Phosphate Phosphoric Acid H3PO4 H2PO4 –
Hydrocyanic Acid HCN CN Cyanide Ion +
Hydronium Ion H3O H2O Water –
Nitric Acid HNO3 NO3 Nitrate Ion +
Ammonium NH4 NH3 Ammonia –
Carbonic Acid H2CO3 HCO3 Bicarbonate Ion –
2–
Bicarbonate Ion HCO3
CO3 Carbonate Ion Name Water Hydrogen Phosphate Acetic Acid Nitrous Acid Hydrochloric Acid Hydrogen Sulfate Hydrobromic Acid Perchloric Acid Conjugate Acid H2O 2–HPO 4
CH3COOH HNO2 HCl –HSO
4 HBr HClO4 Base –OH 3–PO 4
CH3COO– –NO
2 –Cl 2–SO
4 –Br –ClO
4 Name Hydroxide Ion Phosphate Ion Acetate Nitrite Ion Chloride Ion Sulfate Ion Bromide Ion Perchlorate Ion In the following equations, label the correct Brønsted-­‐Lowery acid/base. Then, label the conjugate acid/base. HClO3 + H2O Acid
Base
Mg(OH)2 + +H3O Base
Acid
H2CO3 + –HO Acid
Base
HI + NH3 Acid
Base
–ClO + +H O 3
3
Conj. Base
Conj. Acid
2+Mg + H
Conj. Acid
–HCO
2O Conj. Base
3 + H2O Conj. Base
Conj. Acid
+NH + –I 4
Conj. Acid
Conj. Base
15 1.2 – Worksheet
Gram Equ ivalents/Normality/ Titration
1. How many equivalents are in 4.2 g of acetic acid? 2. In a solution with 0.45 Eq of H2SO4 how many grams of acid are in the solution? 6. A 25 mL solution of 0.2 M HI is titrated with 0.30 M NaOH. How many milliliters of NaOH are used? 7. A solution containing 45 mL of NH3 is titrated with 50 mL of 0.60 M HBr. What was the original concentration of NH3? 3. What is the gram equivalent of H3PO4? 8. Equal volumes of acid and base are used to titrate an acidic solution of 0.50 M. What is the concentration of the base used? 4. A 225 mL solution contains 7.3 g of H3PO4. a) What is the molarity of this solution? b) What is the normality? c) What is the concentration of this solution in milliequivalents per liter? 5. A 400 mL solution has a concentration of 0.56 M H2SO4, how many equivalents are in this solution? 9. A 0.35 M sample of H2S is titrated with 0.065 L of 0.15 M NaOH. a) What is the initial volume of the acid? b) What are the equivalents per liter? c) What are the microequivalents per liter? 10. 35 mL of titrant (KOH) is used neutralize a 75 mL solution containing 4.5 grams of H2SO4. How many grams of KOH were used? 1.2 – Worksheet - KEY
Gram Equ ivalents/Normality/Titration
1. How many equivalents are in 4.2 g of acetic acid? ! !"# !! !""#
! !" x = # of Eq à 4.2 g !" ! !! !!""# ∙ ! !"# = x !
x = 0.07 Eq 2. In a solution with 0.45 Eq of H2SO4 how many grams of acid are in the solution? x = # of grams à x ∙
! !"# !! !!!
!" ! !! !!!
= 0.45 Eq x = 44 g H2SO4 3. What is the gram equivalent of H3PO4? Molar mass 98 g H! PO!
→
= 33 g − Eq # of ions
3 H ! ions 4. A 225 mL solution contains 7.3 g of H3PO4. a) What is the molarity of this solution? 1 mol
7.3 g H! PO! = 0.074 mol H! PO! 98 g H! PO!
0.074 𝑚𝑜𝑙 H! PO!
= 0.33 𝑀 0.225 𝐿
6. A 25 mL solution of 0.20 M HI is titrated with 0.30 M NaOH. How many milliliters of NaOH are used? 25 mL ∙ 0.2 M = 0.3 M ∙ V! V! = 17 mL 7. A solution containing 45 mL of NH3 is titrated with 50 mL of 0.60 M HBr. What was the original concentration of NH3? 45 mL · M1 = 50 mL · 0.60 M M1 = 0.67 M 8. Equal volumes of acid and base are used to titrate an acidic solution of 0.50 M. What is the concentration of the base used? V1 = V2 à M1 = M2 à Base = 0.50 M 9. A 0.35 M sample of H2S is titrated with 0.065 L of 0.15 M NaOH. a) What is the initial volume of the acid? 0.35 M · V1 = 0.65 L · 0.15 M V1 = .0028 L b) What is the normality? N = M · # of ions à 0.33 M · 3 ions = 0.99 N c) What is the concentration of this solution in milliequivalents per liter? b) What is the normality of the acid? 0.35 M · 2 H+ = 0.70 N c) What are the microequivalents per liter? !"
Normality = ! à ! !"#$
0.074 mol ∙ ! !"# = 0.223 Eq !.!" !"
0.70 N = ! !
· !·!"! !!"
! !"
= 70000 µμEq/L 0.223 Eq ∙
!""" !"#
! !"
= 223 𝑚𝐸𝑞 223 mEq
mEq
= 991 0.225 L
L
5. A 400 mL solution has a concentration of 0.56 M H2SO4, how many equivalents are in this solution? !"#
M = ! à !.!" !"#
400 𝑚𝐿 ∙ !""" !" = 0.224 𝑚𝑜𝑙 H! SO! ! !"
0.224 mol H! SO! ! !"# = 0.448 𝐸𝑞 10. 35 mL of titrant (KOH) is used neutralize a 75 mL solution containing 4.5 grams of H2SO4. How many grams of KOH were used? ! !"# ! !!
4.5 g H2SO4 ∙ !" ! ! !!! ! = 0.046 mol H! SO! !
!.!"# !"# !! !!!
M
.!"# !
!
= 0.612 M H! SO! M1V1 = M2V2 à 0.612 M · 75 mL = 35 mL · M2 M2 = 1.3 M KOH !.! !"# !"#
.035 L · ! !
= 0.455 mol KOH !" ! !"#
0.455 mol ∙ ! !"# !"# = 2.5 g KOH 17 1.3 – Worksheet
pH – pOH – [H + ] – [ – OH]
1. What is the concentration of [H+] in a solution whose pH = 4.3? 2. What is the pH of a solution that has a Hydronium concentration of 3.4 x 10-­‐3? 3. What is the pOH of a solution that has a pH of 6.8? 4. What is the concentration of Hydroxide Ions in a solution that has a pOH of 2.9? 5. What is the pOH of a solution that has a [–OH] of 5.7 x 10-­‐5? Please complete the following table pH [H+] 3.5 -­‐4 5.8 x 10
4.2 x 10-­‐5 5.1 6. A solution has a pH of 6.1. What is the concentration of Hydroxide Ions? 7. A solution has a concentration of Hydrogen Ions of 2.8 x 10-­‐6. What is the pOH of this solution? 8. A solution has a [–OH] of 5.8 x 10-­‐7. What is the pH of this solution? 9. A 450 mL beaker is 0.00045 M HCl. What is the pH of this solution? 10. A 320 mL beaker contains 2.30 mg of NaOH. What is the pH of this solution? [–OH] -­‐2 4.2 x 10
7.2 x 10-­‐3 pOH 8.2 2.4 1.3 – Worksheet – KEY
pH – pOH – [H + ] – [ – OH]
1. What is the concentration of [H+] in a solution whose pH = 4.3? [H+] = 10-­‐4.3 10-­‐pH = [H+] +
[H ] = 5.0 x 10-­‐5 2. What is the pH of a solution that has a Hydronium concentration of 3.4 x 10-­‐3? pH = -­‐log(3.4 x 10-­‐3) pH = -­‐log[H3O+] pH = 2.5 Note: H3O+ = H+ 3. What is the pOH of a solution that has a pH of 6.8? 6.8 + pOH = 14 pH + pOH = 14 pOH = 7.2 4. What is the concentration of Hydroxide Ions in a solution a pOH of 2.9? 10-­‐2.9 = [–OH] .0013 = 1.3 x 10-­‐2 = [–OH] 5. What is the pOH of a solution that has a [–OH] of 5.7 x 10-­‐5? pOH = -­‐log(5.7 x 10-­‐5) pOH = -­‐log[–OH] pOH = 4.24 6. A solution has a pH of 6.1. What is the concentration of Hydroxide Ions? 14 – 6.1 = pOH 14 = pH + pOH 7.9 = pOH -­‐7.9= [–OH] 10
10-­‐pOH = [–OH] 1.2 x 10-­‐8 = [–OH] 10-­‐pOH = [–OH] Please complete the following table pH [H+] 3.5 3.2 x 10-­‐4 3.2 5.8 x 10-­‐4 12.6 2.4 x 10-­‐13 5.8 1.6 x 10-­‐6 4.4 4.2 x 10-­‐5 11.6 2.5 x 10-­‐12 10.1 7.9 x 10-­‐11 11.9 1.3 x 10-­‐12 7. A solution has a concentration of Hydrogen Ions of 2.8 x 10-­‐6. What is the pOH of this solution? -­‐log(2.8 x 10-­‐6) = pH +
-­‐log[H ] = pH 5.6 = pH 14 = pH + pOH 14 – 5.6 = pOH 8.4 = pOH 8. A solution has a [–OH] of 5.8 x 10-­‐7. What is the pH of this solution? -­‐log(5.8 x 10-­‐7) = pOH -­‐log[–OH] = pOH 6.2 = pOH 14 = pH + pOH 14 – 6.2 = pH 7.8 = pH 9. A 450 mL beaker is 0.00045 M HCl. What is the pH of this solution? Because this is a strong acid the -­‐log(.00045) = pH 3.3 = pH molarity = the [H+] Also, volume is not important here 10. A 320 mL solution contains 2.30 mg of NaOH. What is the pH of this solution? 2.30 mg NaOH ∙
! !
!""" !"
= 0.0023 g NaOH 1 mol NaOH
= .000058 mol NaOH 40 g NaOH
mol
Concentration (M) = L
.!!!!"# !"# !"#$
M NaOH = = .00018 M NaOH = [–OH] .!"# !
-­‐log[–OH] = pOH à -­‐log(.00018) = 3.7 14 – pOH = pH à 14 – 3.7 = 10.3 à pH = 10.3 . 0023 g NaOH ∙
[–OH] 3.2 x 10-­‐11 1.7 x 10-­‐11 4.2 x 10-­‐2 6.3 x 10-­‐9 2.5 x 10-­‐10 4.0 x 10-­‐3 1.3 x 10-­‐4 7.2 x 10-­‐3 pOH 10.5 10.5 1.4 8.2 9.6 2.4 3.9 2.1 19 1.4 – Worksheet
Buffers
1. A buffer solution contains 0.12 M Acetic Acid (Ka = 1.8 x 10-­‐5) and 0.28 M Sodium Acetate. What is the pH of this solution? 2. A buffer solution contains 0.48 M Benzoic Acid (Ka = 6.3 x 10-­‐5) and 0.32 M Potassium Benzoate. What is the pH of this solution? 3a. Describe how the pH level in blood is regulated through the use of buffers -­‐ specifically with Carbon Dioxide and Hydrogen Carbonate. 3b. Where (which organs) in your body does this regulation occur? 3c. Why are buffers important in the body? 3d. Describe what Alkalosis is and how your body uses hyperventilation to return the body to homeostasis. 4. A buffer solution contains 0.45 M formic acid (Ka = 1.8 x 10-­‐4) and an unknown concentration of sodium formate. The concentration of Hydrogen Ions in this solution is found to be 3.6 x 10-­‐4. What is the concentration of sodium formate? 5. A buffer solution contains an unknown concentration of nitrous acid (Ka = 4.5 x 10-­‐2) and 0.75 M potassium nitrite. After equilibrium has established, the pOH measured was 11.3. What is the initial concentration of the nitrous acid? 1.4
– Worksheet – KEY
Buffers
1. A buffer solution contains 0.12 M Acetic Acid (Ka = 1.8 x 10-­‐5) and 0.28 M Sodium Acetate. What is the pH of this solution? [!! ]
!.!"
pH = pKa + log![!"]! pH = -­‐log(1.8 x 10-­‐5) + log! ! [A-­‐] = Sodium Acetate !.!"
pH = 4.74 + 0.368 [HA] = Acetic Acid 2. A buffer solution contains 0.48 M Benzoic Acid (Ka = 6.3 x 10-­‐5) and 0.32 M Potassium Benzoate. What is the pH of this solution? !.!"
[!! ]
pH = -­‐log(6.3 x 10-­‐5) + log!!.!"! pH = pKa + log![!"]! pH = 4.2 + (-­‐0.176) [A-­‐] = Potassium Benzoate pH = 4.02 [HA] = Benzoic Acid 3a. Describe how the pH level in blood is regulated through the use of buffers -­‐ specifically with Carbon Dioxide and Hydrogen Carbonate. Carbonic Acid is unstable and exists in equilibrium as –HCO3 and CO2. When there is excess CO2 the equilibrium shifts to the right and the pH drops due to an increased amount of H3O+ in the blood. CO2 (aq) + H2O(l) H2CO3 (aq) –HCO3 + H3O+ 3b. Where (which organs) in your body does this regulation occur? The equilibrium with CO2 is regulated in the lungs, and the regulation of –HCO3 occurs in the kidneys. 3c. Why are buffers important in the body? In order to maintain Homeostasis, our bodies need to stay at a pH of around 7.4 à even variations of even a few tenths of a pH unit can produce severe physiological symptoms! 3d. Describe what Acidosis is and how your body uses hyperventilation to return the body to homeostasis. Acidosis is when your blood’s pH is below 7.35. Hyperventilation causes the body to expel excess CO2 thereby reducing the amount of acid in the blood. 4. A buffer solution contains 0.45 M formic acid (Ka = 1.8 x 10-­‐4) and an unknown concentration of sodium formate. The concentration of Hydrogen Ions in this solution is found to be 3.6 x 10-­‐4. What is the concentration of sodium formate? -­‐log(3.6 x 10-­‐4) = pH 3.44 = pH [!! ]
pH = pKa + log![!"]! [!! ]
3.44 = -­‐log(1.8 x 10-­‐4) + log! !.!" ! [A-­‐] = Sodium Formate [A-­‐] = 0.22 M Sodium Formate [HA] = Formic Acid 5. A buffer solution contains an unknown concentration of nitrous acid (Ka = 4.5 x 10-­‐2) and 0.75 M potassium nitrite. After equilibrium has established, the pOH measured was 11.3. What is the initial concentration of the nitrous acid? 14 – 11.3 = pH [!! ]
2.7 = pH pH = pKa + log![!"]! !.!"
-­‐2
2.7 = -­‐log(4.5 x 10 ) + log![!"]! [A-­‐] = Potassium Nitrite [HA] = Nitrous Acid [A-­‐] = 0.033 M Nitrous Acid 21 Naming • If there are more than one of equal length • Most substituted wins Parent -­‐ Longest Chain Name all sustituents Assign Numbers • Ex. CH3 -­‐ methyl • Ex. 2,4 -­‐ diethyl • Prekixes cylco-­‐ & iso-­‐ are part of the group and are alphabetized • Ignore di-­‐, tri-­‐, tert-­‐, sec-­‐ Alphabatize • Use commas between #s and dashes between #s and words. Structural Formula Condensed Structure Line Structure # of Carbon Prefixes 1 – Meth-­‐ 2 – Eth-­‐ 3 – Prop-­‐ 4 – But-­‐ 5 – Pent-­‐ 6 – Hex-­‐ 7 – Hept-­‐ 8 – Oct-­‐ 9 – Non-­‐ 10 – Dec-­‐ 11 – Undec-­‐ 12 – Dodec-­‐ 13 – Tridec-­‐ 14 – Tetradec-­‐ 15 – Pentadec-­‐ 16 – Hexadec-­‐ 17 – Heptadec-­‐ 18 – Octadec-­‐ 19 – Nonadec-­‐ 20 – Eicos-­‐ 23 1.5 – Worksheet
Namin g Alkanes
1.5 – Worksheet – KEY
Namin g Alkanes
Pentane 2-­‐methylbutane 2-­‐methylpropane 2,4-­‐dimethylhexane 2,3-­‐dimethylpentane 2,4-­‐dimethylpetane 4-­‐ethyl-­‐3-­‐methylheptane 1-­‐methylcyclopentane 1-­‐ethyl-­‐2-­‐methyl-­‐
cyclohexane 1-­‐isopropyl-­‐3-­‐n-­‐
propylcyclohexane 5-­‐ethyl-­‐3,8,9-­‐
trimethylundecane 1-­‐methyl-­‐2-­‐tert-­‐
butylcyclopentane 3-­‐n-­‐butyl-­‐4-­‐isopropyl-­‐1,2-­‐
dimethylcyclohexane 3-­‐n-­‐butyl-­‐2,2-­‐diethyl-­‐1,5-­‐
dimethylcyclohexane 4-­‐sec-­‐butyl-­‐2-­‐methyl-­‐1-­‐
cyclopropylheptane 3-­‐cyclobutyl-­‐4-­‐cyclohexyl-­‐7-­‐
methylnonane 5-­‐sec-­‐butyl-­‐6-­‐tert-­‐butyl 2-­‐
cyclopentylnonane 6-­‐ethyl-­‐5-­‐isobutyl-­‐4,8-­‐
dimethyl-­‐2-­‐
cylcopentylnonane 25 1.6 – Worksheet
Fun ction al Groups
Identify and circle the functional groups 1.6 – Worksheet – KEY
Fun ction al Groups
Alcohol Ether Carboxylic Acid Keytone Amine Alkyne Amide Aldehyde Anhydride Ester Alkene x 2 Aromatic 27 1.7 – Worksheet
Drawing Structures
Convert to line structure Convert to condensed structure 1.8 – Worksheet – KEY
Drawing Structures
Convert to line structure Convert to condensed structure 29 1.8 – Worksheet
Con formers & Isomers
Identify the isomers and conformers of each compound. Please put a circle around the isomers and a square around the conformers. 1.8 – Worksheet - KEY
Con formers & Isomers
Identify the isomers and conformers of each compound. Please put a circle around the isomers and a square around the conformers. 31 Practice T est – IA
1. Please complete the following table. (12) Name Acid Ammonium Name Base 2-­‐SO
Conjugate Base –HCO
3 Conjugate Acid Name Name Perchloric Acid 4 2. A 345 mL solution contains 4.7 g of H2CO3. (6) a) What is the molarity of this solution? b) What is the normality? c) What is the concentration of this solution in milliequivalents per liter? 3. 67 mL of titrant, NaOH, is used to neutralize a 48 mL solution containing 5.6 grams of HCl. How many moles of NaOH were used? (8) 4. A 540 mL beaker is 7.5 x 10-­‐4 M HBr. What is the pH of this solution? (4) 5. Please complete the following table. (12) pH [H+] 4.5 7.8 x 10-­‐3 [–OH] 2.8 x 10-­‐4 pOH 6.8 6. What is the pH of a buffer solution that contains 0.46 M Hydrocyanic Acid (Ka = 4.9 x 10-­‐10) and 0.37 M Sodium Cyanide? (4) Practice T est – IA
7. Describe how and where the pH level in blood is regulated through the use of buffers – specifically with Carbon Dioxide and Hydrogen Carbonate. (6) 8. To a 514 mL solution of 0.54 M Hydrofluoric Acid (Ka = 6.8 x 10-­‐4) was added 423 mL of an unknown concentration of Potassium Fluoride. Once equilibrium was established, the concentration of Hydrogen Ions was found to be 6.4 x 10-­‐4. What was the original concentration of Potassium Fluoride before the buffer solution was made? (8) 9. Please name the following compounds. (12) 10. Please identify the functional groups by circling and providing the names of the groups on the following compounds. (10) 33 Practice T est – IA
11. Please complete the following table. (8) Convert to condensed structure 4-­‐ethyl-­‐2,3,6-­‐
trimethylheptane Convert to line structure 3,4-­‐diethyl-­‐2,5-­‐
dimethylhexane 12. Identify the isomers and conformers of the compound in the upper left corner. Please put a circle around its isomers and a square around its conformers. (8) 13. Please label the following Brønsted-­‐Lowry acids and bases and identify their conjugate acid and base pairs. (8) –HCO
3 + H2O +NH + –Br 4
H2CO3 + –HO HBr + NH3 100 35 Practice T est – IA – KEY
1. Please complete the following table. (12) Name Acid +NH
Ammonium 4 Carbonic Acid H2CO3 Name Base Sulfate Ion 2-­‐SO
4 Conjugate Base NH3 –HCO
3 Conjugate Acid Name Ammonia Bicarbonate Ion Name -­‐HSO
4 Hydrogen Sulfate –ClO Perchlorate Ion HClO4 Perchloric Acid 4
2. A 345 mL solution contains 4.7 g of H2CO3. (6) a) What is the molarity of this solution? ! !"#
!.!"# !"# !! !!!
4.7 g H! CO! !" ! ! !! = 0.076 mol H! CO! = 0.22 𝑀 !.!"# !
!
!
b) What is the normality? N = M · # of ions à 0.22 M · 2 ions = 0.44 N c) What is the concentration of this solution in milliequivalents per liter? ! !"#$
!""" !"#
!"# !"#
!"#
0.076 mol ∙ ! !"# = 0.152 Eq 0.152 Eq ∙ ! !" = 152 𝑚𝐸𝑞 !.!"# ! = 440.6 ! 4. A 540 mL beaker is 7.5 x 10-­‐4 M HBr. What is the pH of this solution? (4) -­‐log(7.5 x 10-­‐4) = 3.12 3. 67 mL of titrant, NaOH, is used to neutralize a 48 mL solution containing 5.6 grams of HCl. How many moles of NaOH were used? (4) ! !"# !"#
5.6 𝑔 𝐻𝐶𝐿 ∙ !" ! !"# = 1.6 𝑚𝑜𝑙 𝐻𝐶𝑙 # moles HCl = # moles NaOH 1.6 moles NaOH 5. Please complete the following table. (12) pH [H+] 10.4 3.57 x 10-­‐11 4.5 3.2 x 10-­‐5 2.1 7.8 x 10-­‐3 7.2 6.3 x 10-­‐8 [–OH] 2.8 x 10-­‐4 3.2 x 10-­‐10 1.28 x 10-­‐12 1.6 x 10-­‐7 pOH 3.6 9.5 11.9 6.8 6. What is the pH of a buffer solution that contains 0.46 M Hydrocyanic Acid (Ka = 4.9 x 10-­‐10) and 0.37 M Sodium Cyanide? (4) [!.!"]
[!! ]
pH = -­‐log(4.9 x 10-­‐10) + log![!.!"]! pH = pKa + log!
! [!"]
[A-­‐] = Sodium Cyanide [HA] = Hydrocyanic Acid pH = 9.2 Practice T est – IA – KEY
7. Describe how and where the pH level in blood is regulated through the use of buffers – specifically with Carbon Dioxide and Hydrogen Carbonate. (6) Carbonic Acid is unstable and exists in equilibrium as –HCO3 and CO2. When there is excess CO2 the equilibrium shifts to the right and the pH drops due to an increased amount of H3O+ in the blood. CO2 (aq) + H2O(l) H2CO3 (aq) –HCO3 + H3O+ The equilibrium with CO2 is regulated in the lungs, and the regulation of –HCO3 occurs in the kidneys. 8. A buffer solution contains 0.54 M Hydrofluoric Acid (Ka = 6.8 x 10-­‐4) and an unknown concentration of Potassium Fluoride. The concentration of Hydrogen Ions was found to be 6.4 x 10-­‐4. What is concentration of Potassium Fluoride? (6) pH = -­‐log(6.4 x 10-­‐4) !
[! ]
pH = 3.19 pH = pKa + log![!"]! [!! ]
3.19 = -­‐log(6.8 x 10-­‐4) + log![!.!"]! -­‐
[A ] = Potassium Fluoride [𝐀! ] = 0.57 M Potassium Fluoride [HA] = Hydrofluoric Acid 9. Please name the following compounds. (12) 1-­‐ethyl-­‐2-­‐methyl-­‐3-­‐
propylcyclohexane 3-­‐ethyl-­‐2,5-­‐dimethylheptane 2,5-­‐dimethyl-­‐4-­‐sec-­‐
butyloctane 2-­‐cyclopropyl-­‐3-­‐
methylbutane 6-­‐cyclopentyl-­‐5-­‐isopropyl 2-­‐
methyloctane 4-­‐cyclohexyl-­‐3,5,6-­‐
trimethylnonane 10. Please identify the functional groups by circling and providing the names of the groups on the following compounds. (10) 37 Practice T est – IA – KEY
11. Please complete the following table. (8) Convert to condensed structure 4-­‐ethyl-­‐2,3,6-­‐
trimethylheptane Convert to line structure 3,4-­‐diethyl-­‐2,5-­‐
dimethylhexane 12. Identify the isomers and conformers of the compound in the upper left corner. Please put a circle around its isomers and a square around its conformers. (8) 13. Please label the following Brønsted-­‐Lowry acids and bases and identify their conjugate acid and base pairs. (8) –HCO
3 + H2O Conj. Base
Conj. Acid
+NH + –Br 4
Conj. Acid
Conj. Base
H2CO3 + –HO Acid
Base
HBr + NH3 Acid
Base
39 Practice T est – IB
1. Below is the condensed formula for a compound. (6) a. Please convert to line structure in the box. b. Please provide the name of the compound. 2. Below is the line structure for a compound. (6) a. Please convert to condensed formula in the box. b. Please provide the name of the compound. 3. Please draw the line structure and the condensed formula for 4-­‐sec-­‐butyl-­‐7-­‐ethyl-­‐2,2,8-­‐trimethyldecane. (8) 4. Please provide the names for the following compounds. (9) Practice T est – IB
5. Please complete the following table. (12) Name Hydronium Acid Conjugate Base -­‐H PO
2
4 Name Name Base Conjugate Acid Name Water -­‐CN 6. A 235 mL solution contains 0.00035 M HCl. (8) a. What is the pH of this solution? (2) b. What is the pOH of this solution? (2) c. How many liters would be needed to titrate this solution with 0.0024 M NaOH? (4) 7. Please complete the following table. (9) pH [H+] 5.7 3.7 x 10-­‐3 [–OH] 6.3 x 10-­‐5 pOH 8. A buffer solution contains 0.48 M Acetic Acid (Ka = 1.7 x 10-­‐5) and 0.35 M Potassium Acetate. What is the pH of this solution? (4) 9. Describe how and where the pH level in blood is regulated through the use of buffers – specifically with Carbon Dioxide and Hydrogen Carbonate. (4) 41 Practice T est – IB
10. Please calculate the gram equivalent for the following: (6) a) Mg(OH)2 b) H3PO4 11. Please calculate the Normality for the following: (6) a) 0.78 M H2CO3 b) 1.46 M H2Cr2O7 12. Identify the isomers and conformers of the compound in the upper left corner. Please put a circle around its isomers and a square around its conformers. (10) 13. Below is a compound called 2-­‐disco-­‐4-­‐everOrganoman. Please identify all of his functional groups. (10) J 100 43 Practice T est – IB – KEY
1. Below is the condensed formula for a compound. (6) a. Please convert to line structure in the box. b. Please provide the name of the compound. 3-­‐ethyl-­‐4-­‐methylhexane 2. Below is the line structure for a compound. (6) a. Please convert to condensed formula in the box. b. Please provide the name of the compound. 4-­‐isopropyl-­‐2,3,7-­‐trimethyloctane 3. Please draw the line structure and the condensed formula for 4-­‐sec-­‐butyl-­‐7-­‐ethyl-­‐2,2,8-­‐trimethyldecane. (8) 4. Please provide the names for the following compounds. (9) 1-­‐ethyl-­‐2,4-­‐dimethylcyclohexane 3-­‐cyclopentyl-­‐4-­‐
ethylheptane 3-­‐tert-­‐butyl-­‐2-­‐
cyclobutylhexane Practice T est – IB – KEY
5. Please complete the following table. (12) Name Hydronium Phosphoric Acid Acid H3O+ H3PO4 Conjugate Base H2O -­‐H PO
2
4 Name Water Dihydrogen Phosphate Name Base Conjugate Acid Name Hydroxide Ion -­‐OH H20 Water Cyanide Ion -­‐CN HCN Hydrocyanic Acid 6. A 235 mL solution contains 0.00035 M HCl. (8) a. What is the pH of this solution? (2) pH = 3.5 -­‐log([H+]) = pH -­‐log(0.00035) = pH 3.5 = pH b. What is the pOH of this solution? (2) pOH = 10.5 14-­‐3.5 = 10.5 c. How many liters would be needed to titrate this solution with 0.0024 M NaOH? (4) M1V1=M2V2 M1 = 0.00035 M, V1 = 0.235 L, M2 = 0.0024 M, Solve for V2 0.00035 Ÿ 0.235 L = 0.0024 Ÿ X X = 0.034 L NaOH 7. Please complete the following table. (9) pH [H+] 5.7 2.0 x 10-­‐6 9.8 1.6 x 10-­‐10 2.4 3.7 x 10-­‐3 [–OH] 5.0 x 10-­‐9 6.3 x 10-­‐5 2.5 x 10-­‐12 pOH 8.3 4.2 11.6 8. A buffer solution contains 0.48 M Acetic Acid (Ka = 1.7 x 10-­‐5) and 0.35 M Potassium Acetate. What is the pH of this solution? (4) !.!"
[!! ]
pH = -­‐log(1.7 x 10-­‐5) + log!!.!"! pH = pKa + log![!"]! [A-­‐] = Potassium Acetate pH = 4.6 [HA] = Acetic Acid 9. Describe how and where the pH level in blood is regulated through the use of buffers – specifically with Carbon Dioxide and Hydrogen Carbonate. (6) Carbonic Acid is unstable and exists in equilibrium as –HCO3 and CO2. When there is excess CO2 the equilibrium shifts to the right and the pH drops due to an increased amount of H3O+ in the blood. CO2 (aq) + H2O(l) H2CO3 (aq) –HCO3 + H3O+ The equilibrium with CO2 is –
regulated in the lungs, and the regulation of HCO3 occurs in the kidneys. 45 Practice T est – IB – KEY
10. Please calculate the gram equivalent for the following: (6) a) Mg(OH)2 !"#$% !"## !".!!!"∙!!!∙!
g-­‐Eq = !"#$%#&'( !"#$%&'#( !"#$ →
= 29.2 g !
b) H3PO4 !".!"!!"∙!!!∙!
g-­‐Eq = = 32.7 g !
11. Please calculate the Normality for the following: (6) a) 0.78 M H2CO3 Normality = M× #ions à 0.78 M × 2 ions = 1.56 N b) 1.46 M H2Cr2O7 1.46 M × 2 ions = 2.92 N 12. Identify the isomers and conformers of the compound in the upper left corner. Please put a circle around its isomers and a square around its conformers. (10) 13. Below is a compound called 2-­‐disco-­‐4-­‐everOrganoman. Please identify all of his functional groups. (10) J 117 119 121 123