1. Acids & Bases
Learning Outcome C2

2. Student Achievement Indicators
Students should be able to demonstrate the following:
Identify acids and bases using indicators.
Explain the significance of the pH scale, with reference to common substances.
Differentiate between acids, bases and salts by examining chemical properties and formulae.
Recognize the names and formulae of common acids.
Use the periodic table to:
Explain the classification of elements as a metal or non-metal.
Identify the relative reactivity of elements in the alkali metal, alkaline earth metal, halogen and noble gas groups.
Distinguish between metal oxide solutions (basic) and non-metal oxide solutions (acidic).
Use the periodic table and a list of ions to name and write chemical formulae for common ionic compounds.
Convert names to formulae and formulae to names for covalent compounds.

3. Key Terms – Section 5.1 acids bases bromothymol blue concentration
Indigo carmine
litmus paper
methyl orange
pH indicators
phenothphthaleina

4. pH Indicators
pH indicators are chemicals that change color based on the pH of the solution they are placed in.
A pH scale is a numbered scale (0  14) that is used to measure how acidic or basic a solution is.
Acids are chemical compounds that produce a solution with a pH of less than 7 when they are dissolved in water.
Bases are compounds that produce a solution with a pH of greater than 7 when they are dissolved in water.
If a solution has a pH of 7, it is neutral.

5. Types off pH Indicators
Litmus paper
blue litmus  acid (red)/base (no color change)
red litmus  base (blue)/acid (no color change)
Phenophthalien
Is a colorless chemical compound that turns pink in a slightly basic to a highly basic solution.
Bromothymol Blue 
Is a yellow solution that turns blue between a pH of 6.0 – A highly basic solution will be blue in color.

6. Types of pH Indicators Methyl Orange
Is a red solution that changes from red to yellow.
Acids with a pH of 0 to 3.2 will appear red; an acid with a pH of 4.4 will appear orange.
As the solution becomes more basic it will appear yellow.
Indigo Carmine
Is a blue solution that turns yellow at a very high pH (approx. 11)
Methyl Red
Is a red solution that turns yellow as pH increases.
This indicator will start to change an orange color between a pH of 4.8 and 6.0.
As the pH increases the indicator will become yellow.

7. Acids Can be identified by chemical formula
Examples – HCl (hydrochloric acid)/HNO3 (nitric acid)
Many acids only take on acidic properties when mixed with water.
When an acid contains carbon, the H (hydrogen) may be written on the right side.
Example – CH3COOH (Acetic Acid: aka vinegar)

8. Naming Acids Acids may be named in several ways:
Acids in an aqueous state end in “ic” acid
Example – HCl (hydrogen chloride) vs. HCl(aq) (hydrochloric acid) Example – H2SO4 (hydrogen sulphate) vs. H2SO4(aq) (sulphuric acid)
Oxygen-containing acids start with H (hydrogen), and if they end with the suffix “ate”, drop hydrogen and end with suffix “ic” acid.
Example – H2CO3 (hydrogen carbonate) vs. H2CO3(aq) (carbonic acid)
If the name begins with hydrogen and ends with the suffix “ite” change the end to “ous” acid.
Example – H2SO3 (hydrogen sulphite) vs. H2SO3(aq) (sulphurous acid)

9. Naming Bases OH (hydroxide) is on the right side of the formula.
Some can be safe to ingest.
Example – Mg(OH)2 is in antacids and is used to neutralize stomach acid.
Some bases are extremely dangerous and reactive with human skin and tissue.
Highly reactive bases are known as caustic.
Example – drain cleaner and oven cleaner

10. Production of Ions
Acidic and basic solutions have free moving ions that allow both to conduct electricity.
When dissolved in water acids release H+ ions.
When dissolved in water bases release OH- ions.
pH concentrations refers to the concentration of hydrogen ions.
A high concentration of H+ ions = acidic solution (low pH).

11. Production of Ions
H+ and OH- ions readily react with one another to form water, so a solution cannot have a high H+ concentration and a high OH- concentration.
An acidic solution and basic solution can neutralize one another  neutral solution.

12. Properties of Acids & Base
See chart in text
Taste
Touch
Indicator Tests
Reaction with some metals, such as magnesium and zinc
Electrical conductivity pH
Production of ions

13. Key Terms – Section 5.2 metal oxide neutralization non-metal oxide
salts

14. Salts
Neutralization is the name given to a reaction where an acid and base neutralize one another.
This is a type of double replacement reaction.
Acids and bases will react to form water and salt.
Example – HCl + NaOH  NaCl + H2O
hydrochloric acid + sodium hydroxide (base)  table salt + water
Example - 2CH3COOH + Mg(OH)2  Mg(CH3COO) H2O
acetic acid + magnesium hydroxide  magnesium acetate (salt) + water

15. Metal Oxides & Non-Metal Oxides
Metals react with oxygen to form oxides.
An oxide is a chemical compound that includes at least one oxygen atom/ion along with other elements.
A metal oxide is a chemical compound that contains a metal chemically combined with oxygen.
When a metal oxide dissolves in water the solution become basic

16. Example – Na2O(s) + H2O(l)  2NaOH(aq)
sodium oxide + water  sodium hydroxide
Example - CaO(s) + H2O  Ca(OH)2(aq)
calcium oxide + water  calcium hydroxide
Non-metals may also react with oxygen to form acids such as CO2 and SO2.
Fuels (coals/gas) when burned are combined with oxygen. The products are non-metal oxides which are released into the atmosphere.
Example – SO2(g) + H2O  H2SO3

17. Acids and Metals
 When metals react with acids they usually release hydrogen gas.
Example – 2HCl(aq) + Mg(s)  MgCl2(aq) + H2(g)
acid + metal  salt + hydrogen gas

18. Acids and Carbonates
Much of the carbon dioxide on Earth is trapped in rocks, in the form of minerals such as calcite and limestone.
These minerals contain carbonate ions.
When carbonate ions react with acids the carbonates help neutralize the acids.
This is useful to neutralize acid rain in the environment.
Water systems that do not contain carbonates may have calcium carbonate added to the system to protect the water system from acid precipitation.
Example – H2SO4 + CaCO3  CaSO4 + H2O + CO2
sulphuric acid + calcium carbonate  calcium sulphate + water + carbon dioxide