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Chemical reactions happen around us, and in us, everyday. To maintain our body temperature at 98.6 °F (37 °C) when the ambient temperature is cooler, requires heat-generating biochemical reactions that are sensitively regulated. In addition to heat formation, there are four other signs that signal the occurance of a chemical reaction: color change, gas (odor) formation, precipitation of a solid, and light. To help chemists (and chemistry students) manage the vast number of chemical reactions, four categories have been created into which most chemical reactions fall:

1. Precipitation
2. Acid-Base
3. Reduction-Oxidation (RedOx)
4. Decomposition

Precipitation reactions occur when the solubility of a product is low and it "falls out of solution" as a precipitate. Observe the precipitation of silver iodide by clicking the Play button in the animation to the left. Since solid reaction products are denser than aqueous solutions, the solid settles to the bottom of the test tube. The saturated solution above the solid is called the supernatant (from Latin - above + swim . . . . float). In the animation, the AgI forms a yellow precipitate while the NaNO₃ dissolves in water and does not form a precipitate. Click Play again to repeat the animation.

Is there a way to predict this behavior without "looking up" each products' solubility data?

Yes . . . . the solubility of hundreds of compounds can be predicted with the six solubility rules below.

Start at Rule 1 and proceed until a condition is met for the compound's cation or anion . . . . then, . . . . a rule with a lower number takes precedence over a rule with a higher number.

If you know the  Soluble_(aq)  salts and the exceptions for chloride, bromide, iodide and sulfate, you will know the solubility rules because the rest of the salts are  Insoluble_(s)↓ . Here's another way to view the Solubility Rules ( when a condition is met) . . . .

Acid-Base Reactions transfer a hydrogen ion, H⁺, from one reactant (the acid) to another reactant (the base). In Section 2.7 we learned that acids are compounds with an H⁺ cation. Aqueous solutions of acids are formed when an acid (i.e. HCl) reacts with water to form the hydronium ion (H₃O⁺) . . . .

HCl_(g) + H₂O_(l) → H₃O⁺_(aq) + Cl^–_(aq)

The reaction above is 100% - there are no H–Cl bonds present in the aqueous solution. All of the H's from HCl are bound to water and have formed H₃O⁺. Acids that react 100% with water are called strong acids . . . . the 7 strong acids are HCl, HBr, HI, HNO₃, H₂SO₄, HClO₃, and HClO₄. Weak acids are acids that are not strong acids . . . HF, HC₂H₃O₂, H₃PO₄, H₂CO₃, etc. To reinforce the fact that weak acids only partially react with water, chemists use a double-arrow (⇌) in the equation:

HC₂H₃O_{2 (l)} + H₂O_(l) ⇌ H₃O⁺_(aq) + C₂H₃O₂^–_(aq)

This reaction is NOT 100% - there are lots of HC₂H₃O₂ molecules are still present in the aqueous solution. However, some of the H's from HC₂H₃O₂ have combined with water to form H–OH₂⁺.

A base is a compound that produces hydroxide ions (OH^–) when it dissolves in water. The strong bases are the hydroxides of the alkali and alkaline earth metals (LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)₂, Sr(OH)₂, and Ba(OH)₂. These strong bases do not "react" with water, instead, they dissolve in water and . . . . break apart into ions . . . cation(s) and anion(s) 100% . . . .

Ca(OH)_{2 (s)} → Ca²⁺_(aq) + 2 OH^–_(aq)

Weak bases, however, do react with water to increase the [OH^–] in the aqueous solution. Sodium bicarbonate, the weak base in baking soda, reacts with water to produce some hydroxide ions . . . .

NaHCO_{3 (s)} + H₂O _(l) ⇌ Na⁺_(aq) + OH^–_(aq) + H₂CO_{3 (aq)}

Bases (weak and strong) ❤ acids (weak and strong) . . . . the attraction is 100% (━❤). Acid-base reactions are called neutralization reactions . . . . the products are a salt + water.

The colors chosen above to represent acids and bases are based on "the  Lit mus  test". The litmus dye turns  red  in acid and  blue  in base. The litmus dye is conveniently available in the form of . . . . strips of paper soaked in litmus and dried.. The litmus test is a quick qualitative measure of acidity - more precise (quantitative) measurements are made with a . The pH of a solution is simply the negative log of the . . . . concentration is expressed as Molarity (moles solute ÷ liter solution). [H⁺]. The pH of acids is 0 to 7 and the pH of bases is 7 to 14.

The litmus dye is isolated from several species of . Chemists use  red litmus paper  to test if a solution is basic and  blue litmus paper  to test if a solution is acidic. A slightly acidic litmus solution is used to make red litmus paper . . . . blue litmus paper is produced from a neutral litmus solution.

Universal pH paper is "more quantitative" than litmus paper. Universal pH paper has 14 colors and 14 pHs while litmus paper has two colors and the pH is characterized as acidic (0-7) or basic (7-14).

A common acid-base reaction is used by millions of people to combat heartburn. Stomach acid contains approximately 0.03 M H⁺ ions pH = -log[H⁺]
pH = - log(0.03)
pH = 1.5 and causes considerable discomfort when it contacts the esophagus. To alleviate this painful situation, we take an antacid (a base) like calcium carbonate the active ingredient in TUMS. Initially the reaction forms carbonic acid, H₂CO₃, when the H⁺ and Ca²⁺ "exchange partners"

CaCO_{3 (s)} + HCl_(aq) ━❤ CaCl_{2 (aq)} + H₂CO_{3 (aq)}

but H₂CO₃ is very unstable and decomposes quickly to water and carbon dioxide gas.

CaCO_{3 (s)} + HCl_(aq) ━❤ CaCl_{2 (aq)} + H₂O_(l) + CO_{2 (g)}

Oxidation-reduction reactions - were originally associated with reactions where oxygen, O₂, was a reactant. It's not surprising that this reactive gas, comprising 21% of the Earth's atmosphere, plays an essential role in biochemical and environmental processes. When you consume "carbs", you are eating carbohydrates (hydrates of carbon) that have the empirical formula of CH₂O. Now, the oxygen that is absorbed through your lungs reacts with these "carbs" and forms CO₂ (which you exhale), water and  E🗲E R G Y . The reaction below shows the oxidation of glucose (C₆H₁₂O₆ or 6CH₂O) by oxygen.

C₆H₁₂O_{6 (aq)} + 6 O_{2 (aq)} → 6 H₂O_(l) + 6 CO_{2 (g)} + Energy

2 Na_(s) + Cl_{2 (g)} → 2 NaCl_(s)

Na_(s) → Na⁺_(s) + e^–        

Cl_{2 (g)} + 2 e^– → 2 Cl^–_(s)      

In the oxidation reaction, electrons are "lost" by the reactant (Na_(s)) while the reduction reaction shows electrons "gained" by Cl_{2 (g)}. Stated another way, sodium functions as a reducing agent since its electrons "reduce" chlorine. Likewise, chlorine functions as an oxidizing agent by removing electrons from sodium. Half-reactions, like other reactions, are balanced by mass (elements) and charge.

Cl₂^o_(g) + 2 e^– → 2 Cl^–_(s)
     (charge = -2)      (charge = -2)  

There are two mnemonics for remembering oxidation and reduction: Leo says Ger and Oil Rig.

In the reactions above, the two half-reactions were written from the balanced molecular equation . . . . it turns out we can do the  reverse ↺ . In fact, seemingly complicated molecular equations are easily balanced from their redox half-reactions. But first, we must be able to determine which element gains electrons and which one loses electrons in the redox reaction. To do this, we must know the "charge" of each reactant element and the "charge" of each product element . . . . only compounds containing elements whose "charge" changes during the reaction are included in the half-reactions. Some "charges" are easily determined - in the example above, both sodium and chlorine are in their elemental (uncharged) state and have a charge of zero. The product of the reaction is NaCl_(s) where Na has a +1 charge and Cl has a -1 charge. The initial half-reactions are . . . .

Na^o_(s) → Na⁺_(s) + 1 e^–   (oxidation)
Cl₂^o_(g) + 2 e^– → 2 Cl^–_(s)   (reduction)

In the oxidation half-reaction, Na^o → Na⁺, a single electron has been lost. In the reduction half-reaction, Cl₂^o → 2 Cl^–, two electrons have been gained – one electron by each Cl atom (in Cl₂^o). While each half-reaction is currently balanced by atom and charge, it's impossible for the Cl_{2 (g)} in the reduction half-reaction to gain 2 electrons if the reducing agent, Na_(s), only loses 1 electron.

Number of electrons lost (oxidation) = Number of electrons gained (reduction)

To account for this inequality, the oxidation half-reaction is multiplied by 2 and added to the reduction half-reaction . . .

2 Na^o_(s) → 2 Na⁺_(s) + 2 e^–   (oxidation)
  Cl₂^o_(g) + 2 e^– → 2 Cl^–_(s)     (reduction)

---

                   2 Na^o_(s) + Cl₂^o_(g) → 2 NaCl^o_(s) (molecular equation)

In the reaction between sodium and chlorine, the charges of the reactant elements (Na and Cl₂) and the product ions (Na⁺ and Cl^–) were assigned based on information learned in Section 2.6. However, when the reactant or product consists of two or more elements, each element is assigned an . . . . the hypothetical charge of an atom in a compound if all of its bonds to other atoms are assumed to be 100% ionic..

Consider the reaction where MnO₄^– is reduced to MnO₂, the oxidation numbers of the elements are . . .

Mn+7O-2_{4 (aq)}– → Mn+4O-2_{2 (s)}

While the oxidation number of O-2 reflects its normal charge in an ionic compound, it is the assignment of +7 and +4 for manganese that needs further explanation . . . . .

Applying these oxidation rules to the previous equation gives +7 and +4 oxidation numbers for manganese.

Mn+7+7O-2-8 = – 1_{4 (aq)}– → Mn+4+4O-2-4 = 0_{2 (s)}

It is only by assigning oxidation numbers to atoms involved in a redox reaction that we can determine which atoms are changing oxidation states. In this reaction, manganese is being reduced (gaining electrons) . . . . Mn+7 + 3 e^– → Mn+4. However, we can't split Mn away from the atoms it is bonded to, so we must write the reduction half-reaction as . . . .

Mn+7O_{4 (aq)}– + 3 e^– → Mn+4O_{2 (s)}

While the added electrons balance the oxidation numbers in the oxidation half-reaction above, the atoms are not balanced. The section below will guide you through the process of balancing both charge and atoms in redox reactions.

Decomposition Reactions are reactions where a single reactant breaks apart into two or more products. Previously, we encountered the weak acid, H₂CO₃, and learned that it readily decomposes to H₂O and CO₂. A similar decomposition occurs when baker's ammonia (ammonium bicarbonate) is heated.

NH₄HCO_{3 (s)} → NH_{3 (g)} + H₂O_(g) + CO_{2 (g)}

If you have applied hydrogen peroxide to a cut, the cleansing bubbles result from the production of oxygen gas . . . .

2 H₂O_{2 (l)} → 2 H₂O_(l) + O_{2 (g)}