ABSE_Acids and Bases Lesson

Acids and Bases

In the chemical reactions module we studied the most fundamental definition of acids and bases, the Arrhenius definition. The Arrhenius definition describes an acid as a substance that produces H3O+ in water and a base as a substance that produces OH- in water. There was a problem with this definition however. Some substances that behaved as acids or bases did not fit the Arrhenius definition. Therefore, a new, broader definition was needed.   Two scientists, Johannes Brønsted, a Danish chemist, and Thomas Lowry, a British scientist, determined that a better way to define acids and bases was in terms of proton transfer. Note that a proton is the same as a hydrogen ion, H+.

A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor.

Let's look at ammonia, NH3, which did not fit the Arrhenius definition of a base. NH3 can react with HCl by accepting the H+ from HCl, therefore it is a Brønsted-Lowry base.

HCL + NH₃ → NH₄⁺ + Cl⁻

See that the H+ (from the HCl) was donated to the NH3.

So, the HCl is the Brønsted-Lowry acid and the NH3 is the Brønsted-Lowry base. (From here on, we will just use the terms acid and base.)

Many reactions between acids and bases are actually equilibrium reaction where we can label not only the acid and base of the forward reaction, but we can also label the acid and base of the reverse reaction. To differentiate between the acid and base in the forward and reverse reactions, we use the terms conjugate acid and conjugate base.   Below is the reaction between acetic acid and ammonia. We will go through how to identify the acid, base, conjugate acid, and conjugate base here as an example.

Identify the acid, base, conjugate acid, and conjugate base. HC₂H₃O₂ + NH3 C₂H₃O₂ + NH₄⁺ The easiest way to answer this question is to first identify the conjugate acid-base pairs. This refers to one reactant and one product that differ only by one H+. Above you can see that the conjugate acid-base pairs are: HC₂H₃O₂/C₂H₃O2 NH₃,/NH₄⁺ The only difference between each member of a pair is one H+. In the balanced equation, we identify the pairs by using brackets. These brackets can be written above or below the reaction. Their position has no significance. HC₂H₃O₂ + NH₃ ↔️ C₂H₃O₂ + NH4+ Now, let's label the species. Remember that an acid donates a proton (H+). So the acid will have one more H+ than its partner in the pair. The acid in each pair is circled. HC₂H₃O₂ CeH₃O₂ NH₃/NH₄⁺ The acid that appears on the reactant side as written is labeled acid. The acid that appears on the product side is labeled conjugate acid. HC₂H₃O₂ + NH3 acid ↔️ C2₃3O₂ NH₄⁺ Conjugate acid The base is the reactant in the pair that accepts a proton (H+). It has one less H+ than its partner the pair. HC₂H₃O₂ + NH3 acid base ₂ ↔️ C₂H₃O₂ + NHA₄⁺ Conjugate base Conjugate acid

Many times you will be asked to label the species as we did above, but the reactions will look more complicated.   If you will always start by identifying the conjugate acid-base pairs, you will easily be able to determine each species.

For example:

H₂PO₄⁻ acid+ CO₃²⁻base ↔️ HPO₄²⁻ conjugate base + HCO₃⁻ conjugate acid

Amphoteric

Did you notice anything about how water was labeled in b) and c) above? Go back and look if you didn't!

Water behaved like a base in b) and an acid in c). A substance, like water, that can be either an acid or base depending on the reaction is called amphoteric. Sometimes this is called amphiprotic to emphasize that a proton is involved. Water is not the only substance that is amphoteric. Ions and other compounds can also be amphoteric.

Acid Strength

The strength of an acid is determined by its ability to donate a proton. Also, the more complete the reaction, the stronger the acid. There are some trends in acid strength that are visible using the periodic table.   

Binary Acid Strength

Acid Strength Chart for the periodic elementsRecall from module 3 that binary acids are those that contain hydrogen and a non-metal. Moving left to the right within a period on the periodic table, the strength of binary acids increases. For example, HCl is stronger acid than H2S which is a stronger acid than PH3. This trend is consistent with the trend in electronegativity. Electronegativity also increases from left to right. The greater the electronegativity of X, the more polar the H-X bond. The stronger the partial charge, the easier it is for the H+ to separate from the X-, thus the stronger the acid.

The strength of binary acids increases from top to bottom on the periodic table.  

HF   <   HCl   <   HBr   <   HI

This is opposite the trend in electronegativity.   Read about the reason for this contradiction in your book.

Oxoacid Strength

Oxoacids are those composed of hydrogen, oxygen, and one other element. When the central atom holds the same number of oxygen atoms, the acid strength increases from the bottom to top within a group and from left to right within a period.   For example, look at the acid strength of the following:

 Acid strength: HClO4 > HBrO4 > HIO4

For a given central atom, the acid strength of an oxoacid increases with the number of oxygens held by the central atom. For example sulfuric acid (with four oxygens) is stronger than sulfurous acid (with three oxygens).

Acid strength: H2SO4 > H2SO3

You will need to remember how to name and write formulas for acids and bases. This lesson from module 3 is located at the end of this module.

Base Strength

One easy way to comment on the strength of bases is to look at them from the perspective of their conjugate acid.   Stronger acids and bases tend to react with each other to produce their weaker conjugates. The stronger a Brønsted acid is, the weaker is its conjugate base. The weaker a Brønsted acid is, the stronger is its conjugate base. This is shown in the diagram below.

Acid Base Strength list of very strong acids to very weak acids and very weak bases to very strong bases

Remember to work on the module practice problems as you complete each section of content.  

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