Lesson 4


What are the Arrehnius and Bronsted Lowry definitions of acids and bases?


National Standards Addressed: NS912.2


Overarching Goal: The overarching goal of this lesson is to compare the differences the Arrehnius and Bronsted-Lowry definitions of acids and bases.


Learning Objectives:


Hook/ Warm up: How many different ways can one define cold and hot? (Give class time to think and answer)

  We can say, high temperatures mean hot, and low temperatures correspond to cold.  We can also say that if it burns then it is hot, and if it is freezing then it is cold.  The point is that there is more than one way of defining something.  Definitions are modified as more information is being found.  When new evidence is present, we will have to modify our definitions to incorporate this new knowledge.  Similarly, there have been numerous definitions of acids and bases. In addition to the pH scale, we have different ways of identifying if something is an acid or a base.  We will focus on the Arrehnius definition and the Bronsted Lowry definition.  (In a higher level course, say A.P. chemistry, Lewis definitions will be introduced.)



Given the following acids, what do they all have in common?






Given the following bases, what do they all have in common?





From the two questions above, we will see that all the acids listed have hydrogen in them, while all the bases have an OH group in them. 

Basically, Arrehnius’ definition of an acid, is a compound that produces H+ ions, hydrogen ions (hydronium, when combined with water H3O+)

Arrehnius’ definition of a base is any compound that produces hydroxide ions, or OH-.

But what about something like, ammonia.  From our previous lesson, we found that ammonia has a pH that is higher than 7, therefore we concluded that it was a base.  But where is the OH-?  This is where the Bronsted-Lowry definitions come in.  Because the ammonia does not have any OH- in it, the Arrehnius definition does not apply. In 1923 an English scientist, T. M. Lowry and a Danish scientist J. N. Bronsted independently proposed a new definition. They stated that in a chemical reaction, any substance that donates a proton is an acid and any substance that accepts a proton is a base. According to Bronsted-Lowry theory, an acid is a proton donor while a base is a proton acceptor. Free protons are hydrated by water molecule to form H3O+ , the hydronium ion. The two definitions are similar, but the latter one modifies and adds to the first definition.  (We will see this soon, when we analyze acid base reactions in our next lesson.)

Analyze the following equation as a class:


NH3+ H2O à OH-+  NH4+

What happens to the NH3? H2O?

NH3à NH4+

H2O à OH-  


The ammonia (base) becomes an ammonium ion.  The ammonia has gained a proton (H+), and water has lost a proton.  When a species has changed by gaining or losing a proton, the product formed is a conjugate.  Therefore ammonium is the conjugate acid of ammonia.  The hydroxide ion is the conjugate base of water.  The topic of conjugates will be discussed in further detail in the next lesson on acid and base reactions.   This is just an introduction to the concept.


Assignment (In Class):


Given the following reactions:

Tell students to look at what happens to each of the underlined compounds.  Do they gain hydrogen or do they loose a hydrogen ion.

First, determine if the underlined compounds are an acid or a base, then determine if they fit the Arrehnius definition or the Bronsted Lowry Definition, or both.  Then identify conjugate pairs, similar to what we did with the ammonia example.  While students are completing this (they may do so singly or with the person sitting next to them), walk around to make sure students are understanding the concept, and answer any questions they may have.  The concept of conjugates is often difficult and confusing for some students. 

NH4+ + H2OàH3O+ + NH3

ClO4- + NH4+àHClO4 + NH3

HSO4- + H2O à H2SO4 + OH-

HSO4- + H2O àSO4-2 + H3O+


Discuss what is happening in each of the equations.  Help students make the connections between the acid and its conjugate base, and the base with its conjugate acid.

After completing this exercise, ask the following question: What do you notice about the last two equations? 

In the last two equations, the HSO4acts as a base in one equation and an acid in the other.  The same for water.  This is known as amphoterism.  A compound that can act as either an acid or a base is called amphoteric. 




In talking about acids and bases, and defining them, we have mentioned several acids and bases.  We see the same ones in most of our examples.  Ask the class to see if they could name some that, they are familiar with.  Some of these include:

HCl (hydrochloric acid)

H2SO4 (Sulfuric acid)

NaOH (sodium hydroxide)

LiOH (lithium hydroxide)

Mg(OH)2 (magnesium hydroxide)

NH3 (ammonia)

HNO3 (nitric acid)

**(This is also a good time to go over naming compounds)

So, what are these compounds?  What are they used for?  Ask class, what they think.

HCl (hydrochloric acid)-cleaning, sometimes it is used in low concentrations to clean jewelry; it can also be used to clean concrete.

H2SO4 (Sulfuric acid)-battery acid

NaOH (sodium hydroxide)-pipe cleaners, drain cleaners

LiOH (lithium hydroxide)-same as sodium hydroxide

Mg(OH)2 (magnesium hydroxide)-antacids like Mylanta

NH3 (ammonia)-cleaning, many household cleaning detergents have ammonia in it

HNO3 (nitric acid)-used in making other compounds with nitrogen (i.e. TNT)