Download Acids and Bases: Definitions, Theories, and Calculations and more Study notes Chemistry in PDF only on Docsity!
Acids and Bases
Unit 11
Let’s start our discussion of acids and bases by defining some terms that are essential to the topics that follow. Arrhenius acids and bases are: acid—a substance that increases the concentration of protons (H+^ ) in water base—a substance that increases the concentration of hydroxide ions in water (OH-) These definitions are limited to aqueous solutions.
Brønsted and Lowry acids and bases as: acid—a substance that donates a proton to another substance base—a substance that accepts a proton These definitions can also apply to reactions that are not aqueous, so they are more accurate.
Lewis acids and bases are: acid—a substance that accepts an electron pair base—a substance that donates an electron pair
Here are some other terms that you’ll need to be familiar with: hydronium (H 3 O+)—H+^ riding “piggyback” on a water molecule; water is polar, and the positive charge of the naked proton is greatly attracted to one of the negative electron pairs on adjacent oxygen monoprotic — describes acids that can donate one H +^ (HCl, HBr, HF, HI or HNO 3 ) diprotic — describes acids that can donate two H +^ ions (Sulfuric acid - H 2 SO 4 , carbonic acid - H 2 CO 3 , hydro sulfuric acid - H 2 S, chromic acid - H 2 CrO 4 , and oxalic acid - H 2 C 2 O 4 ). polyprotic — describes acids that can donate more than one H+^ ion (phosphoric acid - H 3 PO 4 and citric acid - C 6 H 8 O 7 ) amphiprotic— describes a substance that can act as either an acid or a base. This means it can either lose a proton or gain one. Water (H 2 O) is amphiprotic: it can form either a hydroxide ion or a hydronium ion. Other examples of amphiprotic
substances are HCO 3 - , HSO 4 - , or HPO 4 -2^ (conjugate Acid-Bases). It is also
called autoionization:
2H 2 O H 3 O+^ + OH-
Conjugate Acid-Base Pairs
Look at the generic acid-base reaction below: HX(aq) + H 2 O(l) ↔X - (aq) + H 3 O(aq)
All acids have a conjugate base, which is formed when their proton has been donated; likewise, all bases have a conjugate acid, formed after they have accepted a proton.
Example Apply the appropriate acid-base theory to first identify the acid and base reacting and then identify the conjugate acid-base pairs in the examples below:
- HNO 3 + H 2 O → H 3 O +^ + NO 3 -^ 2. NH 4 +^ + H 2 O → H 3 O +^ + NH (^3)
Name _________________________ Block ______
Mr. B’s
Chemistry
Explanation
- In this first reaction, we see that HNO 3 gives a proton to water, which then forms a hydronium ion. This makes HNO 3 the acid in the forward reaction, and water acts as the base. HNO 3 ’s conjugate base is NO 3 - , and water’s conjugate acid is the hydronium ion, or
- Here NH 4 +donates the proton to water, so in the forward reaction it acts as the acid, and water is still the base. NH 4 +’s conjugate base is NH 3 , and water’s conjugate acid is again the hydronium ion, H 3 O +.
Relative Strengths of Acids and Bases
Certain acids are stronger than other acids, and some bases are stronger than others. What this means is that some acids are better at donating a proton, and some bases are better proton acceptors. A strong acid or base dissociates or ionizes completely in aqueous solution. A weak acid or base does not completely ionize.
Strong Acids Strong Acids and Strong Bases completely dissociate in solutions There are six strong acids that you’ll need to memorize for the SOL Chemistry test:
Hydrochloric. Hydrobromic and Hydroiodic acids: HCl, HBr, HI Nitric acid: HNO (^3) Sulfuric acid: H 2 SO 4 Perchloric acid: HClO 4
Strong Bases
There are fewer strong bases to memorize for the exam. These are hydroxides (—OH), oxides of 1A and 2A metals (except Mg and Be), H ¯, and CH 3 ¯.
Example:
LiOH - lithium hydroxide *Ca(OH) 2 - calcium hydroxide NaOH - sodium hydroxide *Ba(OH) 2 - barium hydroxide KOH - potassium hydroxide *Sr(OH) 2 - strontium hydroxide RbOH - rubidium hydroxide CsOH - cesium hydroxide
Acid–Base Reactions: Neutralization Reactions
When a strong acid and a strong base solution are mixed, a neutralization reaction occurs, and the products do not have characteristics of either acids or bases. Instead, a neutral salt and water are formed. Look at the reaction below:
HCl (aq) + NaOH (aq) → H 2 O (l) + NaCl (aq)
Example
Classify each of the salts listed below as acidic, basic, or neutral.
1. Fe(NO 3 ) 3 _ Salt - Neutral _ 4. HNO 3 _ Acidic _
2. RbOH _ Basic _ 5. MgSO 4 _ Salt - Neutral _
3. Ni(ClO 4 ) 2 _ Salt - Neutral _ 6. H 2 ClO 4 _ Acidic _
Redox and Electrochemistry
Oxidation-reduction (redox) reactions are another important type of reaction that you will see questions about on the SOL Chemistry test. The test writers will expect you to be able to identify elements that are oxidized and reduced, know their oxidation numbers. The following is a brief overview of the basics.
Oxidation-Reduction
Oxidation-reduction reactions involve the transfer of electrons between substances (losing or gain
electrons).
Electrochemistry: The study of the interchange of chemical and electrical energy. Oxidation: The loss of electrons. Since electrons are negative, this will appear as an increase in the charge (e.g., Zn loses two electrons; its charge goes from 0 to +2). Metals are oxidized.
Oxidizing agent (OA): The species that is reduced and thus causes oxidation.
Reduction: The gain of electrons. When an element gains electrons, the charge on the element appears to decrease, so we say it has a reduction of charge (e.g.,
Explanation
- Fe(NO 3 ) 3 —This salt was formed from the reaction of a weak base, iron (III) hydroxide, with a strong acid, nitric acid. This means that the salt will be acidic.
- MgSO 4 —This salt was formed from the reaction of a strong base, magnesium hydroxide, with strong acid, sulfuric acid. This reaction results in a neutral salt.
- Ni(ClO 4 ) 2 —This salt was formed from the reaction of a weak base, nickel (II) hydroxide, with a strong acid, perchloric acid. This is an acidic salt.
Cl gains one electron and goes from an oxidation number of 0 to -1). Nonmetals are reduced.
Reducing agent (RA): The species that is oxidized and thus causes reduction.
Oxidation number: The assigned charge on an atom. You’ve been using these numbers to balance formulas.
Oxidation number: A reaction is considered a redox reaction if the oxidation numbers of the elements in the reaction change in the course of the reaction. We can determine which elements undergo a change in oxidation state by keeping track of the oxidation numbers as the reaction progresses. You can use the following simple rules:
- Hydrogen is assigned an oxidation state of +1. Metal hydrides are an exception: in metal hydrides, H has an oxidation state of -1.
- Oxygen is usually assigned an oxidation state of -2 in its covalent compounds. Exceptions to this rule include peroxides (compounds containing the O 2 -^2 group), where each oxygen is assigned an oxidation state of -1, as in hydrogen peroxide (H 2 O 2 ).
- In compounds, fluorine, chlorine and bromine are assigned an oxidation state of -1.
- The sum of the oxidation states must be zero for an electrically neutral compound.
- For a polyatomic ion, the sum of the oxidation states must equal the charge of the ion.
Now try applying these rules to a problem.
Example Assign oxidation numbers to each element in the following:
- NaCl It is a neutral element Cl is Cl -1^ , so Na is +1. Na +1Cl -1^ = 0 a neutral compound
- MgF 2 It is a neutral element F is F-1^ , so Mg is +2. Mg+2^ F-2= 0 a neutral compound
- PO 4 -^3 It is an anion element O 4 is O -2^ , so P is +5. P +5^ O 4 -8^ = -3 an anion
Examples Oxidation and Reduction: 2 Mg + O 2 → 2 MgO 2 Mg +2^ O - Mg +2^ is oxidized (loses electrons) O -2^ is reduction (gains electrons) Consider: if you gain electrons it makes the element negative (ions) if you loses electrons the element becomes positive (cat ion).
Mg → Mg +2^ + 2 e-
O 2 + 4 e- - → 2 O -
pH S CALE : Use the formulas for pH and pOH to perform the following calculations.
- What is the pH of 6.6 × 10 -4 M HCl? 8. What is the pH of 1.5 × 10 -3 M NaOH? pH = -log [6.6 x 10 -4^ ] pH = 3.18 pOH = -log[1.5 x 10 -3^ ] pOH = 2. pH + pOH = 14 pH = 11.
- A solution of HNO 3 has a pH of 4.5. What is the molarity of HNO 3? We are working backwards.
Molarity [ ] = “2 nd^ log” (- pH) [ ] = 10 -4.5^ = 3.16 x 10 -5^ M
- What is the molarity of KOH in a solution that has a pH of 10.0? pH + pOH = 14 10.0 + pOH = 14 pOH = 4 Molarity [ ] = “2 nd^ log” (- pOH) [ ] = 10 -4^ = 1.0x 10 -4^ M
- Fill in the following table:
[H+^ ] [OH − ] pH pOH acid, base or neutral?
1.0 x 10 -4^ M 1.0 x 10 -10^ M 4 10 Acidic
1.0 x 10 -7^ M 1.0 x 10 -7^ M 7 7 Neutral
1.0 x 10 -12^ M 1.0 x 10 -2^ M 12.0 2 Basic
1 1.0 x 10 -14^ M 0 14.0 Acidic
3.16 x 10 -4^ M 3.16 x 10 -11^ M 3.5 10.5 Acidic
4.6 x 10 -3^ M 2.18 x 10 -12^ M 2.34 11.66 Acidic
1.2 x 10 -3^ M 8.2 x 10 -12^ M 2.91 11.09 Acidic
