year 11 chemistry – unit 2 area of study 1: water

Year 11 Chemistry – Unit 2Area of Study 1: Water

Properties of

• Taste sour (don't taste them!). The word 'acid' comes from the Latin word ‘acidus’, which means 'sour’.

• Have relatively low pH. pH < 7

(Low pH means closer to 1 – this gets confusing!)

•Corrosive.

•Change litmus (a blue vegetable dye) from blue to red.

• Can conduct electricity when in solution (Aqueous).

The International Safety Symbol for Corrosive materials such as

Acids.

Examples of Common Acids

Solutions of Acids can be found around our homes:

• Ethanoic or Acetic acid (In Vinegar).

• Citric acid (In Citrus fruits).

• Ascorbic acid (We know simply as Vitamin C).

• Carbonic acid (In Soft drinks).

• Tannic acid (In Tea).

The three most commonly used Acids in the laboratory are:

•Hydrochloric Acid (HCl)•Sulfuric Acid (H2SO4)•Nitric Acid (HNO3)

Properties of• Taste bitter (Don't taste them!).

• Feel slippery or soapy (Don't touch them!).

• Have a relatively high pH (closer to 14)

•Can react to counter the effect of acids.

• Are Caustic (Corrosive)

• Bases don't change the color of litmus; but they can turn red (acidified) litmus back to blue.

• Can conduct electricity when in solution (Aqueous).

Examples of Common Bases

Washing Powders.

Can also react with fats and oils to form soaps.

Oven Cleaners.

Detergents.

Many cleaning agents contain bases:

Common Acids and Some Everyday Uses

Table 13.1

Common Bases and Some Everyday Uses

Table 13.2

• Hydrochloric Acid is a really strong acid so it has a very low pH value.

• Vinegar is a relatively weak acid with a pH of about 3 (mid-way between strong acids and neutral substances).

• Water has a pH of 7. A substance with a pH of 7 is Neutral.

• Drain Cleaner is an extremely strong base so it has a high pH value.

• Soap is a weaker base with a ph value of about 10.

LOW pH = Very Acidic

HIGH pH = Very

Basic

Safety

Acids and Bases need to be treated with caution:

•Avoid contact with skin and eyes,

•Wear safety glasses and lab coat,

•Label all bottles and containers,

•When diluting acids, add the acid to the water (not the water to the acid), (Some acids can react violently with water and if there is a lot of acid with water being added to it, it can produce a large energy releasing reaction that can spit acid out of the container. )

•Notify the teacher if a spill occurs – or if any substance comes into contact with your skin.

Indicators • Indicators are often extracted from plant dyes and their colour changes with the pH of a solution. Indicators are often acids or bases themselves.

•Common indicators include methyl orange, phenolphthalein, and litmus.

•Litmus is an indicator with its dye obtained from lichen. In the presence of acids, the litmus turns red.

• A Universal indicator is a mixture of many indicators and changes through a range of colours from red to green to violet.

•Indicators are used to indicate when an acid or base is in excess because many reactions involve colourless solutions.

•A pH meter can be used to determine the pH of solutions that do not change colour.

Common indicators and

their pH ranges.

The effects of Acids and Bases on Litmus paper

Neutralisation

Some examples of Neutralisation:

• Farmers use lime (Calcium Oxide - Base) to neutralise acid soils.

• Antacid tablets which contain bases can be used to neutralise excess hydrochloric acid in our stomachs that causes indigestion.

• Bee stings are acidic. They can be neutralised using baking powder, which contains sodium hydrogen carbonate.

• When an acid and a base are mixed they are said to neutralise each other eliminating each others properties.

General reaction types involving Acids Reaction 1

13.2p.236


From the Ionic Equation, we can see that the SO4

2- and Na+ both remain in the aqueous state from reactant to product so they are spectator ions and are not written in the equation


Why are the Nitrate ions Spectator Ions?

They stay in the aqueous state on both sides of the equation!!

• Vinegar was possibly the only acid known in ancient times.

• It was not until after the 12th century that stronger acids were discovered. These include: sulfuric, nitric and hydrochloric acids.

BrØnsted-Lowry Acids and Bases

Johannes BrØnsted and Thomas Lowry described the reactions

of acids as involving the donation of protons (H+)

• A substance is an ACID if it DONATES a proton (H+).

• A substance is a BASE if it ACCEPTS a proton (H+).

• The H+ is a Hydrogen Ion that has lost its only electron – This a Proton. (Not a proton from the atoms nucleus).

BrØnsted-Lowry Acids and Bases

+ Cl -- (aq)

+ Cl -- (aq)

Acid

H3O+

H3O+

(aq)

(aq)

Base

H+

H+

•It is a reciprocal relationship as the acid donates a proton to the base who accepts it.

This explains why Acids and Bases react together!

HCl (g) + H2O (g)

HCl (g) + H2O (g)

Acid-Base Conjugate Pairs

• If substances can be formed from each other by donating or accepting a proton they are said to be a CONJUGATE ACID-BASE PAIR (pg 243).

Acid Base Conjugate

Acid

Conjugate

Base

+ Cl -- (aq)

+ Cl -- (aq)

HCl (g) + H2O (g)

HCl (g) + H2O (g)

H3O+

H3O+

(aq)

(aq)

Hydrogen Ion

• The Hydrogen Ion H+ (aq) (Proton) in a solution is also represented as H3O+ (aq).

•This is called a HYDRONIUM ION

•They both work in the same way - Work like Acids.

They are just different ways of saying the same thing.H+ = H3O+

Amphiprotic Substances• Some substances can act as acids or bases

depending on what they are reacting with.

• This means they can Accept AND Donate Protons.

Acid

Acid

Base

Base

Therefore HCO3- is Amphiprotic

Therefore H2O is Amphiprotic

Acid and Base Strength

•Different acid solutions of the same concentrations do not have the same strength

•The tendency of an acid or base to accept or donate protons from water can be used as a measure of their strength

Eg. 1M HCl, 1M H2SO4 and 1M CH3COOH

Strong Acids•A Strong acid very readily donates a proton

•Acids that completely ionise in solution are called Strong Acids (pg246)

•If an acid readily donates all its protons (H+s) this means it completely ionises

Eg. HCl, H2SO4 and HNO3

Weak Acids

•Acids that partially ionise in solution are called Weak Acids (pg246)

•A Weak Acid only donates some of its protons

•This means it only partially ionises (or Dissociates) - This means that not every molecule breaks apart

Eg. CH3COOH and HF

Strong and Weak Bases

•Strong bases easily accept protons (or an H+)

•Weak bases can accept protons (or an H+) but do not do so readily

Eg. OH- and HCO3-

We can see these examples have negative charges so will gain a positive proton.

Eg. SO42- and H2PO4

-

Polyprotic Acids•Polyprotic Acids are acids that can donate more than one proton

•The number of protons an acid can donate depends on its structure

That is, how many Hydrogen's it has to start with

There can be MONOPROTIC - 1 H DIPROTIC - 2 H TRIPROTIC - 3 H

They donate their Hydrogen's in steps not all at once

Strength vs Concentration

It is important not to confuse the terms Strong and Weak with Concentrated and Dilute

•Strength is a measure of the ionisation of an Acid or Base - Strong and Weak

•Concentration is how much of an Acid or Base actually dissolves in a solution - Concentrated and Dilute

Strength vs Concentration

- Strength is Qualitative - Describing what it happening

- Concentration is Quantitative - Measuring how much is there

Remember:

Acidic, Basic and Neutral Solutions

•The acidity of a solution is a measure of the concentration of Hydrogen Ions (H+) present.

•The higher the concentration of H+ ions the more acidic the solution.

Water acts as both an acid and a base

In this reaction it is both a WEAK ACID and a WEAK BASE

Acidic, Basic and Neutral Solutions

Pure water is a neutral substance because it contains equal concentrations of H3O+ and OH-

10-7

•Acidic Solutions contain greater concentrations of H3O+ than OH-.

•Neutral Solutions contain equal concentrations of H3O+ and OH-.

•Basic Solutions contain lower concentrations of H3O+ than OH-.

Measuring Acidity - Ionic Product

Experiments have shown that all aqueous solutions (aq) contain both H3O+ and OH- ions

Their product is always 10-14

This relationship is called the Ionic Product and is represented by:

[H3O+] x [OH-] = 10-14


Water is neutral so for this to be correct [H3O+] = [OH-]

[H3O+] x [OH-] = 10-14

Indices Laws: 10-7 x 10-7 = 10-14

Therefore, since [H3O+] x [OH-] = 10-14

[H3O+] = 10-7 and[OH-] = 10-7

So if [H3O+] goes up, [OH-] must go down so that the product of the two is still 1 10 –14.


We can therefore define acidic, neutral and basic solutions in terms of the concentrations of [H3O+] and [OH-] present.

[H3O+] x [OH-] = 10-14


Example:

0.1 mol of Hydrogen Chloride (HCl) gas was bubbled into sufficient water to produce 1L of solution.

Calculate the solution concentration of

a) H3O+ ions b) OH-

[H3O+] x [OH-] = 10-14

pH = -log[H+]

The pH scale is used to compare the strength of different acids and bases.

The strength of an acid is based on the [H3O+]

But these values are often very small numbers

To express small numbers conveniently, chemists use the “p scale” which is based on logarithm (base 10)

E.g. 1.0 10-9 mol L-1 = 9.00 on the p scale

pH = -log [H+] = 14 pH = -log [H+] = 1

[H+] = 1 10 -14

= 0.00000000000001 mol L-1

[H+] = 1 10 -1

= 0.1 mol L-1

So you can convert the concentrations of [H+] and [OH-] in solutions in terms of pH and pOH

pH = - log [H+] pOH = - log [OH-]

pH + pOH = 14

So if you know the [H+] then you can calculate the pH of a solution.

Remember [H+] = [H3O+] so pH = - log [H3O+]

•pH is related to hydrogen ion concentration

•As [H+] increases the pH decreases…more acidic

•As [OH-] increases the pH increases…more basic

pH comes from the french pouvoir hydrogene

meaning ‘hydrogen power’

•Since the pH scale is a logarithmic scale, increasing the concentration of H+ ions by a factor of 10 increases the pH only by a factor of 1.

•Eg. Find the pH if [H+] = 0.001 M and if [H+] = 0.01 M

•What is the pOH of the above solution?

pH + pOH = 14

pOH = 14 - pHpOH = 14 – ……..

•What about calculating the concentration of H+ ions when we are

given the pH?

•pH = 5, [H+] = ?

pH = -log [H+]

•If the concentration of OH- ions is given then this needs to be converted to H+ ions to calculate the pH.

Eg. Calculate the pH of a solution that contains 0.0004 M OH-

pH = -log [H+]

Eg. 30 ml of 0.100 M HNO3 is added to 50 ml of Water. What is the pH of the diluted

solution? (Remember 1000ml = 1L)

n = CV

n = 0.100 x 0.03

n = 3 x 10-3 mol

Total Volume = 80 ml

We need to find new concentration…..

C = n V

C = 3 x 10-3

0.08C = 0.0375 M

Now calculate pH…..

year 11 chemistry – unit 2 area of study 1: water

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