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From: OpenStax College Chemistry book

Earth’s atmosphere contains about 20% molecular oxygen, O₂, a chemically reactive gas that plays an essential role in the metabolism of aerobic organisms and in many environmental processes that shape the world. The term oxidation was originally used to describe chemical reactions involving O₂, but its meaning has evolved to refer to a broad and important reaction class known as oxidation-reduction (redox) reactions. A few examples of such reactions will be used to develop a clear picture of this classification.

Some redox reactions involve the transfer of electrons between reactant species to yield ionic products, such as the reaction between sodium and chlorine to yield sodium chloride:

2Na(𝑠)+Cl2(𝑔)⟶2NaCl(𝑠)2Na(s)+Cl2(g)⟶2NaCl(s)

It is helpful to view the process with regard to each individual reactant, that is, to represent the fate of each reactant in the form of an equation called a half-reaction:

2Na(𝑠)⟶2Na+(𝑠)+2e−Cl2(𝑔)+2e−⟶2Cl−(𝑠)2Na(s)⟶2Na+(s)+2e−Cl2(g)+2e−⟶2Cl−(s)

These equations show that Na atoms lose electrons while Cl atoms (in the Cl₂ molecule) gain electrons, the “s” subscripts for the resulting ions signifying they are present in the form of a solid ionic compound. For redox reactions of this sort, the loss and gain of electrons define the complementary processes that occur:𝐨𝐱𝐢𝐝𝐚𝐭𝐢𝐨𝐧𝐫𝐞𝐝𝐮𝐜𝐭𝐢𝐨𝐧==loss of electronsgain of electronsoxidation=loss of electronsreduction=gain of electrons

In this reaction, then, sodium is oxidized and chlorine undergoes reduction. Viewed from a more active perspective, sodium functions as a reducing agent (reductant), since it provides electrons to (or reduces) chlorine. Likewise, chlorine functions as an oxidizing agent (oxidant), as it effectively removes electrons from (oxidizes) sodium.𝐫𝐞𝐝𝐮𝐜𝐢𝐧𝐠 𝐚𝐠𝐞𝐧𝐭𝐨𝐱𝐢𝐝𝐢𝐳𝐢𝐧𝐠 𝐚𝐠𝐞𝐧𝐭==species that is oxidizedspecies that is reducedreducing agent=species that is oxidizedoxidizing agent=species that is reduced

Some redox processes, however, do not involve the transfer of electrons. Consider, for example, a reaction similar to the one yielding NaCl:

H2(𝑔)+Cl2(𝑔)⟶2HCl(𝑔)H2(g)+Cl2(g)⟶2HCl(g)

The product of this reaction is a covalent compound, so transfer of electrons in the explicit sense is not involved. To clarify the similarity of this reaction to the previous one and permit an unambiguous definition of redox reactions, a property called oxidation number has been defined. The oxidation number (or oxidation state) of an element in a compound is the charge its atoms would possess if the compound was ionic. The following guidelines are used to assign oxidation numbers to each element in a molecule or ion.

1. The oxidation number of an atom in an elemental substance is zero.
2. The oxidation number of a monatomic ion is equal to the ion’s charge.
3. Oxidation numbers for common nonmetals are usually assigned as follows:
   • Hydrogen: +1 when combined with nonmetals, −1 when combined with metals
   • Oxygen: −2 in most compounds, sometimes −1 (so-called peroxides, O2²⁻),O2²⁻), very rarely −1/2 (so-called superoxides, O²⁻),O²⁻), positive values when combined with F (values vary)
   • Halogens: −1 for F always, −1 for other halogens except when combined with oxygen or other halogens (positive oxidation numbers in these cases, varying values)
4. The sum of oxidation numbers for all atoms in a molecule or polyatomic ion equals the charge on the molecule or ion.

Note: The proper convention for reporting charge is to write the number first, followed by the sign (e.g., 2+), while oxidation number is written with the reversed sequence, sign followed by number (e.g., +2). This convention aims to emphasize the distinction between these two related properties.

EXAMPLE 4.5

Assigning Oxidation NumbersFollow the guidelines in this section of the text to assign oxidation numbers to all the elements in the following species:

(a) H₂S

(b) SO₃²⁻

(c) Na₂SO₄

Solution(a) According to guideline 1, the oxidation number for H is +1.

Using this oxidation number and the compound’s formula, guideline 4 may then be used to calculate the oxidation number for sulfur:

charge on H₂S=0=(2×+1)+(1×x)

x=0−(2×+1)=−2

(b) Guideline 3 suggests the oxidation number for oxygen is −2.

Using this oxidation number and the ion’s formula, guideline 4 may then be used to calculate the oxidation number for sulfur:

charge on SO₃²⁻ =−2=(3×−2)+(1×𝑥)

𝑥=−2−(3×−2)=+4

(c) For ionic compounds, it’s convenient to assign oxidation numbers for the cation and anion separately.

According to guideline 2, the oxidation number for sodium is +1.

Assuming the usual oxidation number for oxygen (−2 per guideline 3), the oxidation number for sulfur is calculated as directed by guideline 4:

charge on SO₄²⁻−=−2=(4×−2)+(1×𝑥)

𝑥=−2−(4×−2)=+6

From: OpenStax College Chemistry book

An acid-base reaction is one in which a hydrogen ion, H⁺, is transferred from one chemical species to another. Such reactions are of central importance to numerous natural and technological processes, ranging from the chemical transformations that take place within cells and the lakes and oceans, to the industrial-scale production of fertilizers, pharmaceuticals, and other substances essential to society. The subject of acid-base chemistry, therefore, is worthy of thorough discussion, and a full chapter is devoted to this topic later in the text.

For purposes of this brief introduction, we will consider only the more common types of acid-base reactions that take place in aqueous solutions. In this context, an acid is a substance that will dissolve in water to yield hydronium ions, H₃O⁺. As an example, consider the equation shown here:

HCl(𝑎𝑞)+H₂O(𝑎𝑞)⟶Cl−(𝑎𝑞)+H3O⁺(𝑎𝑞)HCl(aq)+H₂O(aq)⟶Cl−(aq)+H3O⁺(aq)

The process represented by this equation confirms that hydrogen chloride is an acid. When dissolved in water, H₃O⁺ ions are produced by a chemical reaction in which H⁺ ions are transferred from HCl molecules to H₂O molecules.

The nature of HCl is such that its reaction with water as just described is essentially 100% efficient: Virtually every HCl molecule that dissolves in water will undergo this reaction. Acids that completely react in this fashion are called strong acids, and HCl is one among just a handful of common acid compounds that are classified as strong (Table 4.2). A far greater number of compounds behave as weak acids and only partially react with water, leaving a large majority of dissolved molecules in their original form and generating a relatively small amount of hydronium ions. Weak acids are commonly encountered in nature, being the substances partly responsible for the tangy taste of citrus fruits, the stinging sensation of insect bites, and the unpleasant smells associated with body odor. A familiar example of a weak acid is acetic acid, the main ingredient in food vinegars:

CH₃CO₂H(𝑎𝑞)+H₂O(𝑙)⇌CH₃CO₂⁻(𝑎𝑞)+H₃O⁺(𝑎𝑞)CH₃CO₂H(aq)+H₂O(l)⇌CH₃CO₂^{−(aq)+H₂O+(aq)}

When dissolved in water under typical conditions, only about 1% of acetic acid molecules are present in the ionized form, CH_{3CO2(Figure 4.6). (The use of a double-arrow in the equation above denotes the partial reaction aspect of this process, a concept addressed fully in the chapters on chemical equilibrium.)}

A base is a substance that will dissolve in water to yield hydroxide ions, OH⁻. The most common bases are ionic compounds composed of alkali or alkaline earth metal cations (groups 1 and 2) combined with the hydroxide ion—for example, NaOH and Ca(OH)₂. Unlike the acid compounds discussed previously, these compounds do not react chemically with water; instead they dissolve and dissociate, releasing hydroxide ions directly into the solution. For example, KOH and Ba(OH)₂ dissolve in water and dissociate completely to produce cations (K⁺ and Ba²⁺, respectively) and hydroxide ions, OH⁻. These bases, along with other hydroxides that completely dissociate in water, are considered strong bases.

Consider as an example the dissolution of lye (sodium hydroxide) in water:

NaOH(𝑠)⟶Na⁺(𝑎𝑞)+OH⁻(𝑎𝑞)

This equation confirms that sodium hydroxide is a base. When dissolved in water, NaOH dissociates to yield Na⁺ and OH⁻ ions. This is also true for any other ionic compound containing hydroxide ions. Since the dissociation process is essentially complete when ionic compounds dissolve in water under typical conditions, NaOH and other ionic hydroxides are all classified as strong bases.

Unlike ionic hydroxides, some compounds produce hydroxide ions when dissolved by chemically reacting with water molecules. In all cases, these compounds react only partially and so are classified as weak bases. These types of compounds are also abundant in nature and important commodities in various technologies. For example, global production of the weak base ammonia is typically well over 100 metric tons annually, being widely used as an agricultural fertilizer, a raw material for chemical synthesis of other compounds, and an active ingredient in household cleaners (Figure 4.7). When dissolved in water, ammonia reacts partially to yield hydroxide ions, as shown here:

NH₃₂O(𝑙)⇌NH₄⁺(𝑎𝑞)+OH⁻(𝑎𝑞)NH₃(aq)+H₂O(l)⇌NH₄⁺(aq)+OH⁻(aq)

This is, by definition, an acid-base reaction, in this case involving the transfer of H⁺ ions from water molecules to ammonia molecules. Under typical conditions, only about 1% of the dissolved ammonia is present as NH4+ ions.

EXAMPLE 4.4

Writing Equations for Acid-Base Reactions. Write balanced chemical equations for the acid-base reactions described here:

• (a) the weak acid hydrogen hypochlorite reacts with water
• (b) a solution of barium hydroxide is neutralized with a solution of nitric acid

Solution(a) The two reactants are provided, HOCl and H₂O. Since the substance is reported to be an acid, its reaction with water will involve the transfer of H⁺ from HOCl to H₂O to generate hydronium ions, H₃O⁺ and hypochlorite ions, OCl⁻.

HOCl(𝑎𝑞)+H₂O(𝑙)⇌OCl−(𝑎𝑞)+H₃O⁺

A double-arrow is appropriate in this equation because it indicates the HOCl is a weak acid that has not reacted completely.

(b) The two reactants are provided, Ba(OH)₂ and HNO₃. Since this is a neutralization reaction, the two products will be water and a salt composed of the cation of the ionic hydroxide (Ba²⁺) and the anion generated when the acid transfers its hydrogen ion (NO₃⁻).

Ba(OH)₂(𝑎𝑞)+2HNO₃(𝑎𝑞)⟶Ba(NO₃)₂(𝑎𝑞)+2H₂O(𝑙)

From OpenStax College Chemistry book.

By the end of this section, you will be able to:

• Define three common types of chemical reactions (precipitation, acid-base, and oxidation-reduction)
• Classify chemical reactions as one of these three types given appropriate descriptions or chemical equations
• Identify common acids and bases
• Predict the solubility of common inorganic compounds by using solubility rules
• Compute the oxidation states for elements in compounds

Humans interact with one another in various and complex ways, and we classify these interactions according to common patterns of behavior. When two humans exchange information, we say they are communicating. When they exchange blows with their fists or feet, we say they are fighting. Faced with a wide range of varied interactions between chemical substances, scientists have likewise found it convenient (or even necessary) to classify chemical interactions by identifying common patterns of reactivity. This module will provide an introduction to three of the most prevalent types of chemical reactions: precipitation, acid-base, and oxidation-reduction.

Precipitation Reactions and Solubility Rules

A precipitation reaction is one in which dissolved substances react to form one (or more) solid products. Many reactions of this type involve the exchange of ions between ionic compounds in aqueous solution and are sometimes referred to as double displacement, double replacement, or metathesis reactions. These reactions are common in nature and are responsible for the formation of coral reefs in ocean waters and kidney stones in animals. They are used widely in industry for production of a number of commodity and specialty chemicals. Precipitation reactions also play a central role in many chemical analysis techniques, including spot tests used to identify metal ions and gravimetric methods for determining the composition of matter (see the last module of this chapter).

The extent to which a substance may be dissolved in water, or any solvent, is quantitatively expressed as its solubility, defined as the maximum concentration of a substance that can be achieved under specified conditions. Substances with relatively large solubilities are said to be soluble. A substance will precipitate when solution conditions are such that its concentration exceeds its solubility. Substances with relatively low solubilities are said to be insoluble, and these are the substances that readily precipitate from solution. More information on these important concepts is provided in a later chapter on solutions. For purposes of predicting the identities of solids formed by precipitation reactions, one may simply refer to patterns of solubility that have been observed for many ionic compounds.

A vivid example of precipitation is observed when solutions of potassium iodide and lead nitrate are mixed, resulting in the formation of solid lead iodide:

2KI(𝑎𝑞)+Pb(NO3)2(𝑎𝑞)⟶PbI2(𝑠)+2KNO3(𝑎𝑞)2KI(aq)+Pb(NO3)2(aq)⟶PbI2(s)+2KNO3(aq)

This observation is consistent with the solubility guidelines: The only insoluble compound among all those involved is lead iodide, one of the exceptions to the general solubility of iodide salts.

The net ionic equation representing this reaction is:

Pb2+(𝑎𝑞)+2I−(𝑎𝑞)⟶PbI2(𝑠)Pb2+(aq)+2I−(aq)⟶PbI2(s)

Lead iodide is a bright yellow solid that was formerly used as an artist’s pigment known as iodine yellow (Figure 4.4). The properties of pure PbI₂ crystals make them useful for fabrication of X-ray and gamma ray detectors.

The solubility guidelines in Table 4.1 may be used to predict whether a precipitation reaction will occur when solutions of soluble ionic compounds are mixed together. One merely needs to identify all the ions present in the solution and then consider if possible cation/anion pairing could result in an insoluble compound. For example, mixing solutions of silver nitrate and sodium fluoride will yield a solution containing Ag⁺, NO3−,NO3−, Na⁺, and F⁻ ions. Aside from the two ionic compounds originally present in the solutions, AgNO₃ and NaF, two additional ionic compounds may be derived from this collection of ions: NaNO₃ and AgF. The solubility guidelines indicate all nitrate salts are soluble but that AgF is one of the exceptions to the general solubility of fluoride salts. A precipitation reaction, therefore, is predicted to occur, as described by the following equations:

NaF(𝑎𝑞)+AgNO3(𝑎𝑞)⟶AgF(𝑠)+NaNO3(𝑎𝑞) (molecular)

Ag+(𝑎𝑞)+F−(𝑎𝑞)⟶AgF(𝑠) (net ionic)

EXAMPLE 4.3

Predicting Precipitation ReactionsPredict the result of mixing reasonably concentrated solutions of the following ionic compounds. If precipitation is expected, write a balanced net ionic equation for the reaction.

• (a) potassium sulfate and barium nitrate
• (b) lithium chloride and silver acetate
• (c) lead nitrate and ammonium carbonate

Solution(a) The two possible products for this combination are KNO₃ and BaSO₄. The solubility guidelines indicate BaSO₄ is insoluble, and so a precipitation reaction is expected. The net ionic equation for this reaction, derived in the manner detailed in the previous module, is

Ba2+(𝑎𝑞)+SO2−4(𝑎𝑞)⟶BaSO4(𝑠)Ba2+(aq)+SO42−(aq)⟶BaSO4(s)

(b) The two possible products for this combination are LiC₂H₃O₂ and AgCl. The solubility guidelines indicate AgCl is insoluble, and so a precipitation reaction is expected. The net ionic equation for this reaction, derived in the manner detailed in the previous module, is

Ag+(𝑎𝑞)+Cl−(𝑎𝑞)⟶AgCl(𝑠)Ag+(aq)+Cl−(aq)⟶AgCl(s)

(c) The two possible products for this combination are PbCO₃ and NH₄NO₃. The solubility guidelines indicate PbCO₃ is insoluble, and so a precipitation reaction is expected. The net ionic equation for this reaction, derived in the manner detailed in the previous module, is

Pb2+(𝑎𝑞)+CO2−3(𝑎𝑞)⟶PbCO3(𝑠)

Remember to ask questions! That’s why I am here. Feel free to let me know if you run into new issues. You can create a free account to add comments to the bottom or you can email me any time: mbriscoe@myedme.com.

App Development

Next Steps – Reading/Testing!

1. Persist data with SQLite: Persist data with SQLite (needed for requirement #1)
2. Read and write files (needed for requirement #1)
3. Store key-value data on disk (needed for requirement #1)
4. all subsections of Forms (helps with screen #2, Requirement B)
5. all subsections of Lists (helps with screen #3)
6. Mixed-list example on flutter.dev
7. Find information on Github or a blog about adding a timer widget
8. Find cool components/options in the Flutter Gallery

App Requirements

• A. Display a weekly and daily schedule of classes, meetings, and assignments.
• B. Allow users to add events to the app.
• C. Allow users to access a timer they can use to help them stay on task while executing assignments.

Outline of App

1. Homescreen: Static information displaying high priority information

2. Adding Event screen
3. Displaying Event screen (uses colors to sort visually the classes/events/meetings)
4. Timer
   1. NOTE: Data Model
      1. Unique ID (UID; required): A uniquely generated name for each event.
      2. EventName (required): User-entered name for each event
      3. DateDue (required): Calendar generated
      4. TimeDue (required): Selected
      5. GoalDate (optional): Allows user to enter date assignment should be done by
      6. EventType (required): select: Assignment, Meeting,
      7. EventDetails (optional):
      8. EventLink (optional):

Flutter related readings and sources

• “Cookbook” (list of pieces for building apps): https://flutter.dev/docs/cookbook
• Google Maps directions: https://codelabs.developers.google.com/codelabs/google-maps-in-flutter/#5
• Database design textbook: https://opentextbc.ca/dbdesign01/

Here is a 30-minute video that covers everything you need to understand using the terminal (for Macs) or command line interfaces (Windows).

Math

The Unit Circle interactive helps you see the trends in the sin and cos values at your pace. Remember that radians are just when we measure by pi (and the circumference of the unit circle is 2ãr.

Proving trig identities usual takes about 3 steps and requires one or more of these formulas and identities. The most important identity is from the unit circle and Pythagoras; the square of sine plus the square of cosine equals 1.

From: OpenStax College Chemistry

Formula Mass

An earlier chapter of this text described the development of the atomic mass unit, the concept of average atomic masses, and the use of chemical formulas to represent the elemental makeup of substances. These ideas can be extended to calculate the formula mass of a substance by summing the average atomic masses of all the atoms represented in the substance’s formula.

Formula Mass for Covalent Substances

For covalent substances, the formula represents the numbers and types of atoms composing a single molecule of the substance; therefore, the formula mass may be correctly referred to as a molecular mass. Consider chloroform (CHCl₃), a covalent compound once used as a surgical anesthetic and now primarily used in the production of tetrafluoroethylene, the building block for the “anti-stick” polymer, Teflon. The molecular formula of chloroform indicates that a single molecule contains one carbon atom, one hydrogen atom, and three chlorine atoms. The average molecular mass of a chloroform molecule is therefore equal to the sum of the average atomic masses of these atoms. 

Moles are the big idea

The identity of a substance is defined not only by the types of atoms or ions it contains, but by the quantity of each type of atom or ion. For example, water, H₂O, and hydrogen peroxide, H₂O₂, are alike in that their respective molecules are composed of hydrogen and oxygen atoms. However, because a hydrogen peroxide molecule contains two oxygen atoms, as opposed to the water molecule, which has only one, the two substances exhibit very different properties. Today, sophisticated instruments allow the direct measurement of these defining microscopic traits; however, the same traits were originally derived from the measurement of macroscopic properties (the masses and volumes of bulk quantities of matter) using relatively simple tools (balances and volumetric glassware). This experimental approach required the introduction of a new unit for amount of substances, the mole, which remains indispensable in modern chemical science.

The mole is an amount unit similar to familiar units like pair, dozen, gross, etc. It provides a specific measure of the number of atoms or molecules in a sample of matter. One Latin connotation for the word “mole” is “large mass” or “bulk,” which is consistent with its use as the name for this unit. The mole provides a link between an easily measured macroscopic property, bulk mass, and an extremely important fundamental property, number of atoms, molecules, and so forth. A mole of substance is that amount in which there are 6.02214076 × 10²³ discrete entities (atoms or molecules). This large number is a fundamental constant known as Avogadro’s number (N_A) or the Avogadro constant in honor of Italian scientist Amedeo Avogadro. This constant is properly reported with an explicit unit of “per mole,” a conveniently rounded version being 6.022 × 10²³/mol.

Consistent with its definition as an amount unit, 1 mole of any element contains the same number of atoms as 1 mole of any other element. The masses of 1 mole of different elements, however, are different, since the masses of the individual atoms are drastically different. The molar mass of an element (or compound) is the mass in grams of 1 mole of that substance, a property expressed in units of grams per mole (g/mol).

Example

Deriving Grams from Moles for an ElementA liter of air contains 9.2 × 10⁻⁴ mol argon. What is the mass of Ar in a liter of air?

SolutionThe molar amount of Ar is provided and must be used to derive the corresponding mass in grams. Since the amount of Ar is less than 1 mole, the mass will be less than the mass of 1 mole of Ar, approximately 40 g. The molar amount in question is approximately one-one thousandth (~10⁻³) of a mole, and so the corresponding mass should be roughly one-one thousandth of the molar mass (~0.04 g):

Moles -> Weights

mol x (grams/mol) = grams

In this case, logic dictates (and the factor-label method supports) multiplying the provided amount (mol) by the molar mass (g/mol): 9.2×10−4mol (39.95g/mol )=0.037g of Ar

The result is in agreement with our expectations, around 0.04 g Ar.

Questions for Moles

• Coming soon!

Summary

The formula mass of a substance is the sum of the average atomic masses of each atom represented in the chemical formula and is expressed in atomic mass units. The formula mass of a covalent compound is also called the molecular mass. A convenient amount unit for expressing very large numbers of atoms or molecules is the mole. Experimental measurements have determined the number of entities composing 1 mole of substance to be 6.022 ×× 10²³, a quantity called Avogadro’s number. The mass in grams of 1 mole of substance is its molar mass. Due to the use of the same reference substance in defining the atomic mass unit and the mole, the formula mass (amu) and molar mass (g/mol) for any substance are numerically equivalent (for example, one H₂O molecule weighs approximately18 amu and 1 mole of H₂O molecules weighs approximately 18 g).