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Bell work. Grab a whiteboard, marker, and eraser. You’ll need it soon. In your notebook, write the ten prefixes for organic compounds. . Agenda. Bell work Organic nomenclature overview Organic practice Pass out lab books Begin review of chapter 4 Homework:

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Bell work

  • Grab a whiteboard, marker, and eraser. You’ll need it soon.

  • In your notebook, write the ten prefixes for organic compounds.


  • Bell work

  • Organic nomenclature overview

  • Organic practice

  • Pass out lab books

  • Begin review of chapter 4

  • Homework:

    • Pre lab questions (due Tuesday), read through lab so you’re prepared for Tuesday

    • Homework problems (on sheet) due Thursday

Organic Chemistry

  • The chemistry of carbon compounds.

  • Carbon has the ability to form long chains.

  • Without this property, large biomolecules such as proteins, lipids, carbohydrates, and nucleic acids could not form.

Structure of Carbon Compounds

  • There are three hybridization states and geometries found in organic compounds:

    • sp3Tetrahedral

    • sp2 Trigonal planar

    • sp Linear


  • Four basic types:

    • Alkanes

    • Alkenes

    • Alkynes

    • Aromatic hydrocarbons


  • Only single bonds.

  • Saturated hydrocarbons.

    • “Saturated” with hydrogens.


  • Lewis structures of alkanes look like this.

  • Also called structural formulas.

  • Often not convenient, though…


…so more often condensed formulas are used.

Properties of Alkanes

  • Only van der Waals force: London force.

  • Boiling point increases with length of chain.

Structure of Alkanes

  • Carbons in alkanes sp3 hybrids.

  • Tetrahedral geometry.

  • 109.5° bond angles.


  • Carbon can also form ringed structures.

  • Five- and six-membered rings are most stable.

    • Can take on conformation in which angles are very close to tetrahedral angle.

    • Smaller rings are quite strained.


  • Contain at least one carbon–carbon double bond.

  • Unsaturated.

    • Have fewer than maximum number of hydrogens.


  • Contain at least one carbon–carbon triple bond.

  • Carbons in triple bond sp-hybridized and have linear geometry.

  • Also unsaturated.

Nomenclature of Alkynes


  • Analogous to naming of alkenes.

  • Suffix is -yne rather than –ene.

Aromatic Hydrocarbons

  • Cyclic hydrocarbons.

  • p-Orbital on each atom.

    • Molecule is planar.

  • Odd number of electron pairs in -system.

Aromatic Nomenclature

Many aromatic hydrocarbons are known by their common names.

Functional Groups

Term used to refer to parts of organic molecules where reactions tend to occur.


  • Contain one or more hydroxyl groups, —OH

  • Named from parent hydrocarbon; suffix changed to -ol and number designates carbon to which hydroxyl is attached.


  • Tend to be quite unreactive.

  • Therefore, they are good polar solvents.

Carbonyl Compounds

  • Contain C—O double bond.

  • Include many classes of compounds.


At least one hydrogen attached to carbonyl carbon.


Two carbons bonded to carbonyl carbon.

Carboxylic Acids

  • Have hydroxyl group bonded to carbonyl group.

  • Tart tasting.

  • Carboxylic acids are weak acids.


Carboxylic Acids


  • Products of reaction between carboxylic acids and alcohols.

  • Found in many fruits and perfumes.


Formed by reaction of carboxylic acids with amines.


  • Organic bases.

  • Generally have strong, unpleasant odors.

Organic Nomenclature

  • Three parts to a compound name:

    • Base: Tells how many carbons are in the longest continuous chain.

Organic Nomenclature

  • Three parts to a compound name:

    • Base: Tells how many carbons are in the longest continuous chain.

    • Suffix: Tells what type of compound it is.

Organic Nomenclature

  • Three parts to a compound name:

    • Base: Tells how many carbons are in the longest continuous chain.

    • Suffix: Tells what type of compound it is.

    • Prefix: Tells what groups are attached to chain.

Organic Nomenclature

  • 2, 3-dimethylpentane

  • How many carbons are in the main chain?

  • What do the numbers mean?

Organic Nomenclature

  • How many bonds can carbon form?

  • How can you tell how many bonds an atom can form?

To Name a Compound…

  • Find the longest chain in the molecule. That’s your base

  • Number the chain from the end nearest the first substituent encountered.

  • List the substituents as a prefix along with the number(s) of the carbon(s) to which they are attached.

To Name a Compound…

If there is more than one type of substituent in the molecule, list them alphabetically.

Sample problems

  • Going from condensed formula to structural formula

  • When there are parenthesis, it means that it comes off that specific carbon. It is not a part of the main chain.

  • Follow nomenclature guidelines

Pass out lab books

  • In pen, write your name in the front of the book

  • Do not write in these books otherwise. You will be able to copy tables into your actual lab notebooks. You’ll get these on Tuesday.

  • Read through lab 1 well enough so you have a solid idea of what you’re going to be doing on Tuesday. Tuesday should not be the first time you read through the lab.

  • Complete the pre-lab questions. They’re due on Tuesday.

Bell work 9/7

  • Turn your pre-lab questions into the tray.

  • What are some questions you have about the lab?

  • When you are finished, put your science notebook back. Grab and read a “Techniques for Handling Crucible” hand out.


  • Bell work

  • Turn in pre-lab questions

  • Pass out lab notebooks

  • Overview use of lab notebooks

  • Lab edits

  • Homework:

    • Complete first 5 organic questions (due Thursday)

    • Work on post-lab questions and analysis (due Tuesday)

Lab notebook

  • Write your first and last name clearly on the front

  • BEWARE – these are carbon copy sheets, so use the periodic table pull out to keep from marking on several sheets at once

  • On the first page, create a table of contents

  • Skip to page 6 and fill in spaces for name, experiment, date, etc

  • Refer to my college lab notebooks

Lab edits

  • You are not going to repeat the experiment. Instead, we are going to pool data.

  • You would do 3 trials—instead do the lab once and then get data from two other groups


  • Copy the data table into your fancy lab notebook

  • Follow the directions outlined in your pink book. Record your actions and observations in your lab notebook.


  • Your lab analysis and post-lab questions will be due next week Tuesday.

  • Make sure you get the data from two other groups, so you can do the full analysis of Trial 1, Trial 2, and Trial 3

Bell work

  • A compound has 1.534 moles of Fe and 2.045 moles of O. Use the ratio between the number of moles of iron and number of moles of oxygen to calculate the empirical formula of iron oxide.


  • Bell work

  • Questions over the lab

  • Chapter 4

    • Sample problems throughout

  • Solubility rules song

  • Revisit organic

  • Homework:

    • Finish organic nomenclature problems

    • Textbook problems from Chapter 4 (written on top of organic sheet)

Chapter 4

  • Aqueous reactions and solution stoichiometry


  • Homogeneous mixtures of two or more pure substances.

  • The solvent is present in greatest abundance.

  • All other substances are solutes.

  • Solutions in which water is the dissolving medium are called aqueous solutions


  • When an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them.

  • This process is called dissociation.

Ionic Compounds in Water

  • Ionic solid dissociates into its component ions as it dissolves

  • Water is an effective solvent

    • Oxygen is electron-rich and carries partial negative charge

    • Hydrogen has a partial positive charge

    • Polar molecule: uneven distribution of electron density

  • Imagine preparing two aqueous solutions—one by dissolving a teaspoon of table salt (sodium chloride) in a cup of water and the other by dissolving a teaspoon of table sugar (sucrose) in a cup of water.

    • How can you tell them apart?

Electrolytic Properties

  • Water itself is a poor conductor of electricity

  • Dissociated ions carry electrical charge

  • Electrolyte: a substance whose aqueous solution contains ions

    • Ex: NaCl

  • Nonelectrolyte: a substance that does not form ions in solution

    • Ex: C12H22O11 (sucrose)

  • As ionic compounds dissociate, their ions become solvated (surrounded) by water molecules

  • This process helps stabilize ions in solution

Molecular compounds

Ionic compounds

Molecular Compounds in Water

  • Solution usually consists of intact molecules dispersed throughout the solution

  • No ionsnonelectrolytes

  • There are a few important exceptions, and most are acids and bases

    • HCl(g) ionizes to form H+(aq) and Cl-(aq)

    • NH3(g) reacts with water to form NH4+(aq) and OH-(aq)

Strong and Weak Electrolytes

  • A strong electrolyte dissociates completely when dissolved in water.

  • A weak electrolyte only dissociates partially when dissolved in water.

  • Differ in the extent to which they conduct electricity

  • Soluble ionic compounds are strong electrolytes

    • Ionic compounds are composed of metals and non-metals or compounds containing ammonium (NH4)

Strong electrolyte

HCl(aq) H+(aq) + Cl-(aq)

Weak electrolyte

HC2H3O2(aq) H+(aq) + C2H3O2-(aq)

Strong Electrolytes Are…

  • Strong acids

  • Strong bases

  • Soluble ionic salts

Strong electrolytes are…

  • Strong acids

  • Strong bases

  • Soluble ionic salts

Solubility Songs!

3 major types of chemical processes occur in aqueous solutions

  • Precipitation reactions

    • Exchange (Metathesis) Reactions

  • Acid-base reactions

    • Neutralization reactions and salts

    • Acid-base reactions with gas formation

  • Oxidation-reduction reactions

Precipitation Reactions

  • Reactions that result in an insoluble product that is called a precipitate

  • Solubility guidelines allow you to predict the formation of a precipitate

    • Refer to table 4.1 in text

Steps to predict whether a precipitate forms

  • Note ions present in reactants

  • Consider possible combinations of cations and anions

  • Use Table 4.1 to determine if any of those combinations is insoluble

Mg(NO3)2(aq)+ 2 NaOH(aq)

Mg(OH)2(s)+ 2 NaNO3(aq)

Exchange (Metathesis) Reaction

  • Two cations exchange anions

  • Chemical formulas of the products are based on the charges of the ions

  • Conform to the general equation:



Mg(NO3)2(aq)+ 2 NaOH(aq)

Mg(OH)2(s)+ 2 NaNO3(aq)

Balancing a metathesis equation

  • Use chemical formula of reactant to determine the ions that are present

  • Write chemical formulas of products by combining the cations of one reactant with anion of the other (use charges of ions to determine the subscripts)

  • Finally, balance the equation

Molecular equation

  • The molecular equation lists the reactants and products in their molecular form

AgNO3 (aq) + KCl(aq) AgCl(s) + KNO3 (aq)

Ionic equation

  • All strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions

  • This more accurately reflects the species that are found in the reaction mixture

Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq) 

AgCl(s) + K+ (aq) + NO3- (aq)

Net Ionic Equation

  • To form the net ionic equation, cross out anything that does not change from left side of the equation to the right

Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)

AgCl(s) + K+(aq) + NO3-(aq)

  • Those things that didn’t change (and were deleted from the net ionic equation) are called spectator ions.

Net Ionic Equation

  • The only things left in the equation are those things that change (i.e., react) during the course of the reaction.

Ag+(aq) + Cl-(aq) AgCl(s)

Writing a net ionic equation

  • 1. Write chemical formulas of reactants and products

  • 2. Predict insoluble product

  • 3. Balance this equation

  • 4. Write soluble strong electrolytes as ions (complete ionic equation)

  • 5. Cross out spectator ions (net ionic equation)

Sample problems

  • Flip to the back of the first page of your sample problems sheet.


  • Substances that increase the concentration of H+ when dissolved in water (Arrhenius).

  • Proton donors (Brønsted–Lowry).

There are only seven strong acids:

  • Hydrochloric (HCl)

  • Hydrobromic (HBr)

  • Hydroiodic (HI)

  • Nitric (HNO3)

  • Sulfuric (H2SO4)

  • Chloric (HClO3)

  • Perchloric (HClO4)


  • Substances that increase the concentration of OH− when dissolved in water (Arrhenius).

  • Proton acceptors (Brønsted–Lowry).

The strong bases are the soluble salts of hydroxide ion:

  • Alkali metals

  • Calcium

  • Strontium

  • Barium

Acid-Base Reactions

In an acid-base reaction, the acid donates a proton (H+) to the base.

Neutralization Reactions

Generally, when solutions of an acid and a base are combined, the products are a salt and water.

HCl(aq) + NaOH(aq)  NaCl(aq) + H2O (l)

When a strong acid reacts with a strong base, the net ionic equation is…

HCl(aq) + NaOH(aq)  NaCl(aq) + H2O (l)

H+ (aq)+ Cl- (aq)+ Na+ (aq) + OH-(aq)

Na+ (aq)+ Cl- (aq)+ H2O (l)

H+ (aq)+ Cl- (aq)+ Na+ (aq) + OH- (aq)

Na+ (aq) + Cl- (aq) + H2O (l)

Gas-Forming Reactions

  • These metathesis reactions do not give the product expected.

  • The expected product decomposes to give a gaseous product (CO2 or SO2).

  • You end up with the expected salt, carbon dioxide or sulfur dioxide, and water

  • Primarily carbonates and sulfites

Gas-Forming Reactions

CaCO3 (s) + HCl(aq) 

CaCl2 (aq) + CO2 (g) + H2O (l)

NaHCO3 (aq) + HBr(aq) 

NaBr(aq)+ CO2 (g) + H2O (l)

SrSO3 (s) + 2 HI(aq)

SrI2 (aq) + SO2 (g) + H2O (l)

Na2S (aq) + H2SO4 (aq)

Na2SO4 (aq) + H2S (g)

More sample problems

  • Turn to the last page of your sample problems


  • Due Tuesday

  • Lab analysis

  • Organic problems and Chapter 4 textbook problems (#s on organic sheet)

Bell work 9/14

  • Turn in your lab analysis and homework (organic problems, textbook problems).

  • Write the ions present when the following compounds are mixed in water:

    Aluminum nitrateNickel (II) sulfate

    Mercury (II) iodideCalcium hydroxide

  • What is oxidation? What is reduction?


  • Bell work

  • Finish Chapter 4: Redox, concentrations, titrations (solution stoichiometry)

  • Revisit stoich

  • HW: Written on top of sample sheet (Chapter 4 #50, 52, 56, 70, 80, 86)

  • Review solubility rules for lab on Thursday

Oxidation-Reduction Reactions

  • Reactions that involve the transfer of electrons

  • AKA redox reactions

  • Oxidation: loss of electrons

  • Reduction: gain of electrons

LEO the lion says GER

Loss of electrons = oxidation

Gain of electrons = reduction


Oxidation involves loss of electrons

Reduction involves gain of electrons

Oxidation-Reduction Reactions

  • One cannot occur without the other.

  • In all redox reactions, one species is being reduced a the same time as another is oxidized.

Reducing agent

Oxidizing agent

Oxidation Numbers/States

  • Help keep track of electrons during chemical reactions

  • Oxidation – oxidation number goes up (lost something negative)

  • Reduction – oxidation number goes down (gained something negative)

Rules for oxidation numbers

  • Any atom in elemental form has oxidation number of zero

  • For any monotomic ion, the oxidation number equals the charge on the ion

  • Nonmentals usually have negative oxidation numbers

    • Oxidation of oxygen is usually -2 (exception is peroxides, which contain O22- giving each oxygen an oxidation number of -1)

    • Oxidation number of hydrogen is +1 when bonded to nonmetals and -1 when bonded to metals

    • Oxidation number of flourine is -1 in all compounds (other halogens are usually -1 in most binary compounds. Exception in oxyanions when they have positive oxidation states)

  • Sum of oxidation numbers of all atoms in neutral compound is zero

  • Sum of oxidation numbers in a polyatomic ion equals the charge of the ion

Sample problems

  • Try to first 3 with the person sitting next to you

  • You have 5 minutes

Displacement Reactions

  • In displacement reactions, ions oxidize an element.

  • The ions, then, are reduced.

General pattern:

A + BXAX + B

Activity Series

  • Decreasing ease of oxidation

Active metals

  • Any metal on the list can be oxidized by the ions of elements below it

  • Lower elements get reduced and oxidize the ones above them

  • Pg. 143

Noble metals


  • In a reaction between solid copper and silver ions, what will be oxidized? What will be reduced?

Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2Ag(s)


  • Which of the following metals will be oxidized by Pb(NO3)2 : Zn, Cu, Fe?

  • Where do you go to find this information?

  • Find lead

  • Lead can oxidize the elements above it

  • So Zn and Fe will be oxidized but not Cu

moles of solute

Molarity (M) =

volume of solution in liters


  • Expresses concentration

Concentration of Electrolytes

  • Remember electrolytes dissociate into ions

  • In one mole of the compound, you may have several moles of an ion

    • For a 1.0 M solution of NaCl

      • The solution is 1.0 M in Na+ ions and 1.0 M in Cl- ions

    • For a 1.0 M solution of Na2SO4

      • The solution is 2.0 M in Na+ ions and 1.0 M in SO42- ions

Example (not on your sheet)

  • What are the molar concentrations of each of the ions present in a 0.025 M aqueous solution of calcium nitrate?

moles of solute

Molarity (M) =

volume of solution in liters

Interconvertingmolarity, moles, and volume

  • Remember this formula:

  • If you’re given grams, change to moles

  • If you’re given mL, change to L

  • If you know molarity and volume, you can find # of moles (mol = V x M)


  • You can make less concentrated solutions by mixing a concentrated solution (stock solution) with solvent

  • You can also take solution of known molarity and dilute it with more solvent

  • MconcVconc= MdilVdil

  • MinitialVinitial= MfinalVfinal

  • Note: the M is molarity, NOT moles

Example (not on your sheet)

  • Hector wants to prepare 250.0 mL of 0.100 M CuSO4 by diluting a stock solution containing 1.00 M CuSO4. How should he do it?

  • MconcVconc= MdilVdil

  • (1.00M)Vconc= (0.100 M)(0.250L)

  • Vconc= 0.0250 L

  • To prepare the diluted solution, Hector should remove 0.0250 L of the concentrated solution, place it in a volumetric flask, and dilute it to a final volume of 250.0 mL.


  • The molarity of the more concentrated stock solution is always larger than the molarity of the dilute solution.

  • Because the volume of the solution increases upon dilution, Vdilis always larger than Vconc.

  • Any unit for volume can be used as long as that same unit is used on both sides of the equation

Practice (on your sheet)

  • How many milliliters of 3.0 M H2SO4 are needed to make 450 mL of 0.10 M H2SO4?

  • 15 mL

Practice sheet

  • Problem 4.67

  • Problem at very bottom of first pg

  • 5 minutes

  • If you finish early, try some of the other ones

Solution Stiochiometry and Chemical Analysis

  • Must always convert to moles!!

  • Must be working with a balanced equation!!

It’s the same as regular stoich

  • Steps?

  • Balanced equation

  • Write down what they give you

  • Moles

  • Mole ratio

  • Units

The mole conversion is different

Problem (on back of sheet)

  • How many grams of Ca(OH)2 are needed to neutralize 25.0 mL of 0.100 M HNO3?

  • What kind of reaction is this?

  • Acid-base neutralization reaction

  • What products can I expect?

  • Water and a salt

  • What is the balanced equation?


  • A way to determine the concentration of a solution

  • You have 2 solutions: one with a known concentration and one with an unknown concentration

  • You know what each of them are

  • You let them undergo a reaction (you know the stoichiometry of it)

  • You use the volumes used to determine the number of moles and thus the molarity


  • Let’s say we know the molarity of an NaOH solution and we want to find the molarity of a solution of HCl.

  • What do you need?

    • Known quantity of HCl

    • Acid-base indicator

    • Standard volume of base (NaOH with known molarity)

      • You know volume and molarity, thus of moles



After the titration, you know

  • How much NaOH used

    • Difference between initial and final buret reading

  • Moles of NaOH used

    • Known molarity x known volume

  • Balanced equation for the reaction

    • You chose the reaction and made sure it had known stoichiometry

  • From that balanced equation, you can determine the moles of HCl that reacted

  • Knowing moles and volume of HCl, you can determine its molarity

End point vs equivalence point

  • End point: point where the indicator changes

  • Equivalence point: point where the stoichiometrically correct number of moles of each reactant is present

  • They are not the same, but you would choose an indicator that will change as close to the equivalence point as possible because the color change is what you actually see and how you would know when to stop

Stoichiometry revisit

  • Titrations are stoichiometry, you just have a different way of getting to moles (through molarity and volume)

  • We need to work a little bit on stoichiometry

  • Let’s set up a time for extra practice. When works for you guys?

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