Anglia Ruskin University Acid Base and Neutralization Reactions Discussion Response

You have been exposed to the first six chapters of the material for this course. What concepts have you struggled with in these chapters, and what tools or ways have you used to overcome the obstacle to learn the material? Are you still struggling with a concept with which you need help? Share these answers and any other questions you have

UNIT II STUDY GUIDE
Reactions
Course Learning Outcomes for Unit II
Upon completion of this unit, students should be able to:
4. Relate chemical behavior to atomic structure.
4.1 Calculate formula weights.
4.2 Distinguish between empirical and molecular formulas of a compound from percentage
composition and molecular weight.
4.3 Identify limiting reactants, amounts produced, and percent yields of a reaction.
6. Distinguish between reactions involving different types of compounds.
6.1 Identify simple combination, decomposition, combustion, and oxidation-reduction reactions.
6.2 Identify compounds as acids, bases, and as strong, weak, or nonelectrolytes.
8. Describe molecular interactions in the condensed phase.
8.1 Detail how to carry out a dilution to achieve a desired solution concentration.
8.2 Detail how to perform and interpret the results of a titration.
8.3 Express molarity as a conversion between moles of a solute and volume of the solution.
Course/Unit
Learning Outcomes
4.1
4.2
4.3
6.1
6.2
8.1
8.2
8.3
Learning Activity
Unit Lesson
Chapter 3, pp. 83–110
Unit II Quiz
Unit Lesson
Chapter 3, pp. 83–110
Unit II Quiz
Unit Lesson
Chapter 3, pp. 83–110
Unit II Quiz
Unit Lesson
Chapter 3, pp. 83–110
Unit II Quiz
Unit Lesson
Chapter 4, pp. 121–154
Unit II Lab Assignment
Unit II Quiz
Unit Lesson
Chapter 4, pp. 121–154
Unit II Quiz
Unit Lesson
Chapter 4, pp. 121–154
Unit II Quiz
Unit Lesson
Chapter 4, pp. 121–154
Unit II Quiz
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Required Unit Resources
Chapter 3: Chemical Reactions and Reaction Stoichiometry, pp. 83–110
Chapter 4: Reactions in Aqueous Solution, pp. 121–154
Unit Lesson
Chapter 3
Chemical Equations (3.1)
Stoichiometry is the study of the quantitative or measurable relationships that exist in chemical formulas and
chemical reactions. A great deal of information relies on stoichiometric calculations; for example, such actions
as measuring the concentration of chlorine in water and determining the amount of oxygen in air depend on
stoichiometric calculations.
One of the most important concepts in stoichiometry is the law of conservation of matter. Matter, like energy,
is neither created nor destroyed in any process. In a balanced chemical equation, the ratio of the coefficients
can be interpreted both as the relative number of particles involved in the reaction and as the relative number
of moles. The ratio of moles, or molar ratio, in a balanced equation is essential for solving any stoichiometric
problem. Using the molar ratio, you can determine the number of moles of any substance in the reaction if
you know the number of moles of at least one other substance in the reaction.
The above example shows how to balance an equation.
(Brown et al., 2018, p. 85)
Patterns of Chemical Reactivity (3.2)
Three general types of chemical reactions are discussed in the textbook.
1. Combination reactions are one or more reactants that come together to form a single product. The
general form for a combination reaction is
A + B  AB
where A and B stand for either elements or compounds, and AB stands for a compound consisting of A
and B.
2. Decomposition is a single compound that is broken down into two smaller compounds or elements.
The general form for a combination reaction is
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AB  A + B
where AB represents a compound. A and B represent elements or compounds.
3. Combustion reactions are rapid reactions that produce a flame; this type of reaction usually involves
oxygen from the air as a reactant. For example, a hydrocarbon, or a related compound, will react
with oxygen in the air to form carbon dioxide and water. This is the general form of a singlereplacement reaction.
Hydrocarbon + O2  CO2 + H2O
Magnesium combustion in air produces a flame when the magnesium reacts
with the oxygen in the air. The final product is magnesium oxide (MgO).
(Brown et al., 2018, p. 88)
Formula Weights (3.3)
The formula for a compound indicates the number and kind of each atom in a representative particle of the
compound. The formula weight is one of the most important ideas in chemistry.
The formula weight of a compound is the sum of all the atomic weights of the elements present in the formula
of the compound. For example, the formula weight of sodium chloride (table salt) is found as follows:
•
•
•
Atomic mass of Na
Atomic mass of Cl
Formula mass of NaCl
22.990 amu
35.453 amu
58.443 amu
For more examples on how to calculate the formula weights of compounds, be sure to review the examples
on page 85 and also the sample exercise 3.5 on page 91 of the textbook.
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Avogadro’s Number and the Mole (3.4)
The amount of a compound expressed in
grams equal to its formula mass is called a
mole (abbreviated mol). For example, 18
grams of water equals 1 mole of water,
because the quantity 18 is equal to the
formula mass of water (18 amu). The
number of moles of a compound can be
found by using the following equation:
Mass of compound (g) / Formula mass
(g/mol) = Number of moles (mol)
The number of particles in 1 mole of any
substance is always the same—Avogadro’s
number. If you know how many moles are in
a given amount of a substance, then the
23
number of particles will be 6.02 X 10 times
that number of moles.
For example, to find how many molecules
are in 2 moles of water, you simply multiply
the number of molecules in 1 mole of
23
water (6.02 X 10 ) by 2. The answer is
23
12.04 X 10 .
Empirical Formulas (3.5)
Mass comparison using different units: 1 molecule of
water has a mass of 2.99 x 10-23 g (18.0 amu), while
1 mole of water has a mass of 18.0 grams.
(Brown et al., 2018, p. 94)
We use empirical formulas to denote the ratio of ions in a compound. An empirical formula uses element
symbols to indicate the atoms or ions in a compound, with subscripts added to indicate their ratios.
For example, calcium fluoride has the formula CaF2, which indicates that the compound contains 1 calcium
ion for every 2 fluoride ions. If we know the percentage composition of a compound, we can compute the
empirical formula of the compound. In covalent compounds, the molecular formula is either the same as the
empirical formula or is some whole-number multiple of it.
Procedure for calculating an empirical formula from percentage composition
(Brown et al., 2018, p. 99)
Chemical equations tell us the number of moles of each substance that is involved in a given reaction. For
example, the equation for the formation of ammonia is as follows.
N2 (g) + 3H2 (g)  2NH3 (g)
This formula tells us that hydrogen gas combines with nitrogen gas in a ratio of 1 mole of nitrogen to 3 moles
of hydrogen, and that 2 moles of ammonia are produced for every 1 mole of nitrogen entering the reaction.
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To calculate the number of grams of a product from the number of grams of a reactant, first convert grams of
reactant to moles of reactant. Then use the coefficients in the balanced equation to convert the number of
moles of reactant to moles of product. Finally, convert moles of product to grams of product.
The reactant that limits the amount of product formed in a chemical reaction is called the limiting reactant. The
quantities of products formed in a reaction are always determined by the quantity of the limiting reactant.
The amount of a product that should be produced based on calculations is called the expected yield. The
amount of product that is really obtained from a chemical reaction is called the actual yield. To calculate the
percent yield, you divide the actual yield obtained in the laboratory by the expected yield predicted by
stoichiometry calculations.
Quantitative interpretation of a balanced chemical equation
(Brown et al., 2018, p. 102)
Chapter 4
Properties of Aqueous Solutions (4.1)
A solution is a homogeneous mixture of two or more substances in a single physical state. In a solution, one
substance is usually considered to be dissolved in another. The substance that is dissolved is called the
solute, while the substance that does the dissolving is called the solvent. Solutions with water as the solvent
are given a special name—they are called aqueous solutions.
A substance that dissolves in water to form a solution that conducts an electric current is called an electrolyte.
Sodium chloride solution, pictured below, is an electrolyte, as is any soluble ionic compound. A solution that
contains neutral solute molecules does not conduct an electric current because no charged particles are
available. This type of solution is called a nonelectrolyte. A solution of sugar in water, such as the sucrose
solution pictured below, is a nonelectrolyte.
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Electrical conductivities of water and two aqueous solutions—sucrose solution and sodium
chloride solution
(Brown et al., 2018, p. 122)
Precipitation Reactions (4.2)
A reaction in which two solutions are mixed and a precipitate forms is called a precipitation reaction. This type
of reaction is described by a balanced equation. Chemical equations also show whether dissolved substances
are present in a solution predominantly as ions or molecules.
An equation that shows all soluble ionic substances as ions is called a complete ionic equation. A complete
ionic equation is found by writing each original solution as the sum of its constituent ions. A net ionic equation
includes only those compounds and ions that undergo a chemical change during a reaction in an aqueous
solution. When the complete chemical formulas of all reactants and products are used, the equation is called
a molecular equation.
There are solubility rules that you can use to predict whether an ionic compound will be soluble in water. A list
of these rules is provided in Table 4.1 on page 127 of the textbook.
A precipitation reaction
(Brown et al., 2018, p. 126)
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Acids, Bases, and Neutralization Reactions (4.3)
An acid is any substance that can donate H+ ions (proton donor). A base is a substance that can accept H+
ions (proton acceptor).
Strong acids and strong bases are strong electrolytes, because they react completely to form ions.
Hydrochloric acid (HCl) is a strong acid. It transfers H+ ions to water to form H3O+ ions. OH- is a strong base,
and so we consider compounds such as NaOH or Ca(OH)2, which readily yield OH- ions when they dissolve
in water, to be strong bases as well.
The reaction between an acid and a base is called an acid-base neutralization reaction. During such a
reaction, the acid neutralizes the base, and the base neutralizes the acid. A neutralization reaction results in a
solution that has none of the distinctive properties of either an acid or a base.
Proton transfer
(Brown et al., 2018, p. 131)
Oxidation-Reduction Reactions (4.4)
Oxidation is the process by which a substance loses one or more electrons. Reduction is the process by
which a substance gains one or more electrons. Oxidation and reduction always occur together.
If electrons are lost by one substance, they are gained by another substance. Because oxidation and
reduction occur together, reactions in which electrons are transferred between reactants are called oxidationreduction reactions, or redox-reactions.
Oxidation of calcium metal by molecular oxygen—during the oxidation process,
electrons transfer from the calcium metal to the O2, forming CaO.
(Brown et al., 2018, p. 138)
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Many metals are oxidized by O2, acids, and salts. The redox reactions between metals and acids, as well as
those between metals and salts, are called displacement reactions. The products of these displacement
reactions are always an element (H2 or a metal) and a salt. Metals have different reactions as seen by their
tendency to give up electrons and become oxidized. The relative reactivity of metals is reflected in their
abilities to replace one another from their compounds.
Metals are ranked according to their relative reactivity in a listing that is called the activity series. An activity
series of metals in aqueous solution is provided in Table 4.5 on page 142 of the textbook. Any metal on the
list can be oxidized by the ions of metals below it in the series.
Concentrations of Solutions and Solution Stoichiometry (4.5)
The concentration of a solution is the amount of solute in a given amount of solvent or solution. The most
commonly used expression for solution concentration is molarity (M).
The molarity of a solution is defined as the number of moles of solute dissolved in each liter solution. A
volumetric flask is used for making a solution of a precise molarity. A balance is used to obtain the desired
number of moles of solute, which is then added to the flask. Solvent is added to the flask until the solution’s
desired volume is reached.
The concentration of an acid or a base can be determined with a procedure called acid-base titration. In this
process, we combine a solution of known concentration (a standard solution) with a solution of unknown
concentration to determine the unknown concentration, or the quantity of solute, in the unknown. The point
in the titration at which stoichiometrically equivalent quantities of reactants are brought together is called the
equivalence point. Chemists typically use one additional substance in a titration: an acid-base indicator to
show the endpoint of the titration. The image below shows the preparation of 0.250 L of a 1.00 M solution
of CuSO4.
(Brown et al., 2018, p. 144)
CHM 1301, General Chemistry I
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Reference
Brown, T. L., LeMay, H. E., Jr., Bursten, B. E., Murphy, C. J., Woodward, P. M., & Stoltzfus, M. W. (with
Lufaso, M. W.). (2018). Chemistry: The central science (14th ed.). New York, NY: Pearson.
Suggested Unit Resources
In order to access the following resources, click the links below.
ChemGuy on YouTube further explains some of the more challenging concepts/problems from Chapters 3
and 4.
This video discusses combustion reactions and balancing.
1chemguy. (2010, March 10). Chemguy chemistry A4F1G1 [Video]. https://www.youtube.com/watch?v=rMCz0MnraY
Transcript for Chemguy Chemistry A4F1G1 video
This video discusses empirical and molecular formulas.
1chemguy. (2010, March 15). Chemguy chemistry K4K2F5 [Video].

Transcript for Chemguy Chemistry K4K2F5 video
This video discusses solution stoichiometry.
1chemguy. (2010, March 15). Chemguy chemistry Q2A5H8 [Video].

Transcript for Chemguy Chemistry Q2A5H8 video
This video discusses oxidation-reduction reactions.
1chemguy. (2010, March 12). Chemguy chemistry H3E5T9 [Video].

Transcript for Chemguy Chemistry H3E5T9 video
This video discusses oxidation numbers.
1chemguy. (2010, March 17). Chemguy chemistry U7H4F6 [Video].

Transcript for Chemguy Chemistry U7H4F6 video
CHM 1301, General Chemistry I
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Learning Activities (Nongraded)
Nongraded Learning Activities are provided to aid students in their course of study. You do not have to submit
them. If you have questions, contact your instructor for further guidance and information.
To help you better prepare for the unit quiz and final exam, it is highly recommended that you complete the
Unit II Practice Problems Worksheet. The worksheet includes an answer table.
For additional help, the video on Unit II Worked Problems is available.
Transcript for Unit II Worked Problems video
Before attempting the homework assignment, you are encouraged to view the Unit II Handout for additional
help. The knowledge checks in the handout are designed to provide practice on math-heavy portions of your
quizzes and homework assignments.
As you complete the required reading for this unit in the textbook, consider taking the additional time to work
the sample exercises to make sure you understand all the concepts and actions that were covered in that
section.
In Chapter 3, these exercises are found on pp. 86, 87, 89, 90, 91, 92, 93, 94, 96, 97, 98, 99, 100, 101, 104,
105, 107, and 108.
In Chapter 4, these exercises are found on pp. 125, 128, 130, 132, 133, 134, 139, 141, 143, 145, 146, 148,
149, 151, and 152.
Additional practice that will help you master the chapter’s educational content is provided in the exercises on
pp. 111–118 (Chapter 3) and pp. 154–161 (Chapter 4). Answers to the odd-numbered exercises are provided
in the back of the textbook.
If there is anything you do not understand or an exercise you are unable to successfully complete, contact
your instructor for additional explanation or information.
CHM 1301, General Chemistry I
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