Reading Assignment 1: Read/KTU Ch. 10.1-10.2, pg 304-323; answer questions 1-32

A.     The Mole Concept

1. Representative Particles-  a way to describe the common representation of a substance. Atoms, compounds, molecules.

 Most elemental forms are found in the atomic state.  Exception: Diatomic Molecules

(H₂, N₂, O₂, F₂, Cl₂, Br₂, & I₂)

2. Mole- The number of particles found in a substance which is equivalent to the number of atoms in 12 grams of Carbon-12

a. Avogadros Number (N)- 6.022 x 10²³ particles/mole

A chemical formula also indicates the number of moles of atoms in 1 mole of a compound

b. Avogardros Law.  Equal volumes of gases at the same temperature and pressure contain the same number of particles

Question:  How big is a mole, really?. TEDEd video.  

3. Molar mass. Based on Molecular Weight (contains practice problems for finding molecular weight)

If the masses of samples of two different elements have the same ratio as the ratio of their atomic masses, the samples contain identical numbers of atoms.

The mass of 1 mole of atoms is numerically equal to the atomic mass of the element in terms of grams- (gram atomic mass)

Molecular (compound) molar mass- the sum of the atomic molar masses of the individual elements that comprise the compound

Additional Practice: Molar Mass

Handout: Mole Map -- --Poster Version

4.     Gas Density

the mass per unit volume (density) of specific gases at STP (standard temperature & pressure- 0^oC and 1 atmosphere)

a. Molar volume- One mole of any gas occupies 22.4 L @ STP

5.     Percent Composition

The percent by mass of an element within a compound or sample

Assignment 2. Calculating Percent Composition.

Reading Assignment 2 : Read/KTU Ch. 10.3 pg 325-333; answer questions 33-45 

B. Types of Formulas

1. Empirical Formulas:  The lowest whole-number ratio of elements in a compound

Solving for the empirical formula

a. find the number of moles of each element

b. divide the number of moles of each by the lowest number of moles.

Practice.  Finding the empirical formula from Percent Composition.

2. Molecular Formulas

- the actual ratio of elements in a compound

Finding the molecular formula

a. Divide the formula mass by the empirical mass

b. Multiple the empirical mass ratio by the value from (a)

Practice. (How to determine a chemical formula)

3. Structural Formulas

-defines the bonds associations and the shape of the molecule (see VSEPR & Hybridization model)

Assignment 3. Writing Empirical & Molecular Formulas

Reading Assignment 3: Read/KTU Ch. 11.1 pg. 346-354; answer questions 1-11

C.     Chemical Equations

1. Chemical Reaction- A process where chemical bonds are broken and the atoms are rearranged forming new chemical bonds. This defines a Chemical Change.

What are the 10 signs of chemical change?

a.       Law of Conservation of Matter-  During an ordinary chemical change no detectable changes in mass occurs

b.   Law of Conservation of Energy- During an ordinary chemical change no detectable changes in energy occurs.

** This defines the Universe as a closed system.  All matter & energy is fixed.

2. Chemical equation- shorthand representation that describes the relationship between reactants and products during a chemical changes.

a.       skeletal equation- shows the simplest relationship between reactants and products

1. Reactants- materials that undergo the chemical change, therefore are not conserved in the reaction. These are written on the left side of the arrow and each separated by a "+" sign

2. Products- materials that result from the chemical change.  These are written on the right side of the arrow and each separated by a "+" sign.

3. Arrow - used to separate the reactants from products. Use the words "yields" or "produces"

4. Catalyst- atoms/compounds that aid in the reaction but do not undergo the chemical change.

5. Other symbols

a. subscripts (g), (l), (s), & (aq) correspond to gas, liquid, solid & aqueous states respectively

b. "up arrow"- gas is generated, on product side only

c. "down arrow"- precipitate is formed, on product side only.

b. Balanced chemical equation. 

Coefficients are used in an equations to show that all atoms in the reaction are conserved, establishing the ratio of reactants to products. 

**Rules for Balancing chemical equations.  Be sure :

1.      All formulas are written correctly

2.      Reactants/products are separated by an “à” and individual reactants/products with “+”s

3.      Perform an atom inventory- count the number & type of each element in the reactants and products.     

      **Unchanged polyatomic ions can be counted as individual units.

4.      Use coefficients to “balance” the number of atoms in the reactants to products

a.       don’t change the formulas (leave the subscripts alone!!!!)

b.      start with an element which is present in 1 reactant & 1 product where available

c.   leave elements that are non-bonded (single atoms or diatomic elements) until the end

5.      Perform a second atom inventory to verify if the number and types of atoms are conserved.

6.      Reduce the coefficient to their lowest ratio

In-CLass Practice:  Chembalancer.  Practice Balancing Equations.  Balancing Chemical Equations: PhET simulation

on-Your-Own Practice.  Balancing Chemical Equations

Video: Tutorial on Balancing Equations. Khan Academy

Assignment 4. Balancing Chemical Equations Worksheet.  Print off the worksheet and answer each question.

c.       Classification of Chemical Compounds

1.      Salts- compounds composed of cations and anions.  Soluble salts can be dissolved in water (polar molecule) by dissociation

Handout:  Solubility of Ionic Compounds

a. Rules for Solubility

1. Most nitrates and acetates are water soluble; silver acetate, chromium(II)acetate, and mercury(I)acetate are slightly soluble.

2. All chlorides are soluble except mercury(I), silver, lead(II), and copper(I); lead(II) chloride is soluble in hot water

3. All sulfates except those of Sr, Ba, and Pb(II); Ca & Si sulfates are slightly soluble

4. Carbonates, phosphates, borates, arsenates, and arsenites are insoluble, except those of ammonium and alkali metals

5. The hydroxides of the alkali metals and of barium and strontium are soluble, and other hydroxides are insoluble; calcium hydroxide is slightly soluble

6. Most sulfides are insoluble, except for the sulfides of the alkali metals which react with water to give solutions of the hydroxide and hydrogen sulfide ion, HS^-. 

2.      Acids & Bases-

a.       Bronsted Acid- a compound that donates Hydrogen ions (protons) to solution

-Hydronium- H₃O⁺.

b.      Bronsted Base- a compound that accepts Hydrogen ions in solution

c.       Amphoteric compound- acts as an acid and base.  Water

3.      Electrolytes & Nonelectrolytes

a.       Electrolyte- substance that dissolves in water to form ions

b.      Strong Electrolyte- compound that dissolves to yield 100% ions

c.       Weak Electrolyte- compound that give low percentage of ions

d.      Nonelectrolyte- compound that does not give ions to solution

4.      Polymers- large molecules made of repeating smaller units called monomers.

d. Ionic Equation-  each reactant and product is written in the predominant form in which it occurs in the equation

1. Symbols used in chemical equations to represent the ionic substance as it would be found in the reaction

a. (aq)- aqueous.  dissolved in water

b. (s)- solid. insoluble salt

c. down arrow- precipitate.  formation of an insoluble salt

2. Types of Ionic Equations

a. Complete Ionic Equation- Shows the formulas of all electrolytes as ions  

b. Net ionic Equation- Shows only the ions who have undergone a change in oxidation states

c. Half-Reactions- Shows individual atoms undergoing a change in oxidation states (electron movement)

Tutorial: Writing molecular, ionic & net ionic equations

Tutorial: Ionic Equations: Khan Academy

On-Your-Own Practice: Writing Ionic Equations; Questions 1-6 of Replacement section & question 1 of Ionic section.

Assignment 5:  Writing ionic equations

D.     Chemical Reactions.

1. Combination  (Synthesis)-  A reaction where two or more substances combine to produce another

2. Decomposition- A reaction where a substance is broken down into two or more substances.

3. Single Replacement- Atoms of one element replaces atoms of a second element in a compound

a.       Activity Series of Metals

A more reactive metal will replace a less reactive metal in solution but a less reactive cannot replace a more reactive.

Handout: Activity Series of Common Metals

4. Double Replacement (Metathesis)- Atoms of positive ions exchange position in their respective compounds.

Driven by:

a.       formation of a partially soluble or insoluble product

b.      formation of a gas

c.       formation of a molecular compound such as water

5. Combustion- A reaction where one reaction is oxygen with a release of energy as either heat or light.

-Hydrocarbons usually combust to carbon dioxide and water

6. Others

a.       Oxidation-Reduction- reaction which involves a change in oxidation numbers ( exchange of electrons)

1.      Oxidation- a loss of electrons

2.      Reduction- a gain of electrons

** For every oxidation there exists a reduction.

 b. Acid-Base- reaction where a hydrogen ion is transferred from a Brönsted acid to a Brönsted base.

** Always results in the formation of a salt and water.  

c. Reversible- reaction that can proceed in either direction

Assignment 6. Predicting Products from chemical reactions. Print off the questions, write the formulas for the products & balance the equations. 

Reading Assignment 5:  Read Ch. 12.1-12.3 pg. 382-409; answer questions 1-6, 11-23, 26-33.

E. Interpreting Chemical Reactions

1. Molar Ratios:  The coefficients in chemical equations represent the mole ratio of reactants to products and vice versa

a. particles

b. Compounds (molecules & ions)

c. Mass

d. Volume- not conserved in the equation

Remember:   Mass and Particles are conserved in a chemical reaction- Law of Conservation of Matter

2. Stoichiometric Analysis

a. Stoichiometric Mole Map:  Molar Conversions

1.      Mass à Mass, Volume, & Particles

2.      Particles à Mass, Volume, & Particles

3.      Volume à Mass, Volume, & Particles

Handout. Stoichiometric Mole Map

Explanation. Stoichiometric Applet

Assignment 7. Stoichiometry Problems

3. Limiting Reagent

a. Limiting Reagent- limits the amount of the products formed in a chemical reaction

b. Excess Reagent- extra reactant that does not get completely consumed in a chemical reaction

Assignment 8. Limiting Reactant Problems

4. Percent Yield

a. Theoretical yield- the predicted value of products formed from stoichiometric predictions

b. Actual yield- the measured amount of product formed by experimental processes

c. Percent yield- the ratio of actual yield to theoretical yield

Assignment 9. Ch. 12 Review Questions. p 262: answer questions 56, 58, 60, 65, 68, & 73

Additional Practice.

Unit V. Review Worksheet.  Answers