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1 Chemistry 047: Solution Chemistry

Chemistry 047 –Solution Chemistry

A. Solutions and solubility

B. Conductivity of aqueous solutions

C. Molecular polarity and H bonding discussed in last Unit

D. Polar and nonpolar solvents and the nature of solutions of ions

E. Summary of bond types and how to distinguish between them

F. Calculating the concentration of ions in solution

G. Acids and bases and conjugate acid and base pairs

H. Acids, bases, and the pH scale

I. A bit about titration

A. Solutions and Solubility

1. Definitions: solution

Solution = the term solution describes a system in which 1 or more substances are

homogenously mixed or dissolved in another substance.

Solute and solvent: Simple solutions have 2 components

Solute solvent

Dissolved component dissolving agent or least abundant or most abundant component component

example: saltwater

salt water

However, a solution is not just a solid dissolved in water…

3 states of matter can give us 9 different types of solutions:

Solid in solid

Solid in gas

Solid in liquid

Liquid in liquid

Liquid in gas

Liquid in liquid

Gas in solid

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Gas in gas

Gas in liquid

2. Properties of a true solution:

1. Mixture is homogenous, ratio of solute to solvent is variable

2. Dissolved solute, molecular or ionic in size (10-8 to 10-7 cm /0.1-1 nm)

3. Usually transparent

4. Solute won’t settle out

5. Solute can be separated from solvent by physical means (i.e., evaporation,

distillation etc) – but can’t be separated by filtering

Solutions - permanent if system doesn’t change (if no evaporation etc) because both solute

and solvent particles are in constant random motion – this is energetic enough to prevent

solute from settling out.

3. Solubility

a. Definitions

Solubility = Amount of solute dissolved in a specified amount of solvent under

standard conditions (STP = Standard Temp and Pressure)

When referring to exact amount = quantitative amount

- Expressed as grams/gram i.e., 36.0g NaCl in 100.0 g H2O at 20oc

But often referred to in a qualitative way

- i.e, substance is very soluble, slightly soluble, or insoluble

Other terms for solubility of liquid in liquid:

Miscible = liquids that mix and form solution – alcohol and water

Immiscible = do not mix and form solutions –or are insoluble in each other – i.e., oil

and water

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4. Solubility of common ionic compounds

Solubility table – lists the solubilities of various compounds in each other

The following table represents some ‘General Solubility Rules’

Rule not a law, so lots of exceptions!

Soluble in cold water Insoluble Na+, K+, NH4+

Nitrates (NO3-) Acetates (C2H3O2-)

Chlorides (Cl-), Bromides (Br-), Iodides (I-)

Except with Ag+, Hg+, or Pb2+ then they are insoluble

Sulfates (SO42-) (Ag+, Ca2+ are slightly soluble with sulfates)

Except with Ba2+, Sr2+, Pb2+

Are insoluble except with NH4+ and alkali metal cations

Carbonates, phosphates, hydroxides and sulfides are insoluble except with

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4. Concentration (quantitative expression)

Concentration of solution= amount of dissolved solute in a specified amount of solvent

Molarity = M = mole/litre =moles solute/L of solution

5. Factors related to solubility

Predicting solubility is difficult because:

- Variation in : size of ions, charge of ions, interaction between ions, interaction

between solvent and solute, and temperature

- Therefore there are exceptions

a. Nature of solute and solvent

‘like’ dissolves ‘like’

- Polar or ionic substances react with other polar substances

- Nonpolar dissolve in nonpolar substances

i.e., ionic compounds are soluble in water because water is polar

Nonpolar substances form solutions easily – molecules don’t attract or repel each

other ∴ intermingle easily

Polar solutions are more complex i.e, NaCl in water

Water molecules are attracted to each ion on the crystal surface and weaken

attraction between Na+ and Cl- ions.

Positive end of water dipole, attracted to Cl- and negative dipole end of water to Na+

ion, therefore, surface ions become surrounded by water molecules and become

hydrated ions, Na+ and Cl-, and slowly diffuse away from crystal and dissolve in the

solution.

b. Effect of temperature on solubility

Temperature affects solubility (ability to dissolve) of most substances

Most have limited solubility in a specific solvent at a fixed temp

For solids in liquids – as temp rises, solubility usually rises (as usual no single rule

for solubility of solids)

However, solubility of gas in water, usually decreases as temp increases – bubbles

that form when water is heated are due to decreased solubility of air at higher

temps (=Kinetic Molecular Theory = gas molecules form bonds with liquid –as temp

rises, gas energy increases therefore bubbles out)

i.e, cold water has greater concentration of dissolved O2 than warm water – water

heating, gas bubbles form

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c. Effect of pressure on solubility

Small pressure changes have little effect on solids in liquids, or liquids in liquids

However, there is a large effect of gas in liquid – because gases are compressible

Solubility of the gas is directly related to the pressure of that gas above the

solution – i.e, amount of gas dissolved will double if gas over solution doubles.

Therefore, as pressure decreases, gas will bubble out

d. Saturated vs unsaturated vs supersaturated

1) Saturated

There is a limit to how much solute will dissolve in a given amount of

solvent (at a standard temp).

When a limit is reached – the solution is saturated (If any more solute

was added it would not dissolve)

2 processes happening simultaneously when a solution is saturated –

solid is dissolving in solution and the dissolved solute is crystallizing out

Expressed as:

Solute (undissolved) ⇋ solute (dissolved)

When these are occurring at the same rate, the amount of solute in

solution is constant and equilibrium is established between dissolved

and undissolved solute

Therefore, a saturated solution = dissolved solute in equilibrium with

undissolved.

i.e, if we put 40.0g KCl in 100 g H2O at 20oc, and 34.0g of KCl dissolved

and 6.0 g is undissolved – this solution would be saturated.

Temperature is important because it may be saturated at one temp but

not at another.

2) Unsaturated

Unsaturated contains less solute per unit volume than saturated solution

Meaning: there is room for more solute to dissolve in the solution

3) Supersaturated

Unstable

Contains more dissolved solute than needed to make a saturated solution

at a certain temp

Disturbances (jarring, stirring etc ) will return it to saturation level –

excess will crystalize out– releasing heat in the process

Difficult to make

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Example Will a solution made by adding 2.5g of to10g of H2O be saturated or

unsaturated at 20oc given that the solubility of CuSO4 is 21g/100g H2O? (= to

saying saturation is 2.1 g in 10 g)

So 2.5> 2.1 so solution is saturated and 0.4g will be undissolved

e. Rate (per unit time) of dissolving solids The rate at which substances dissolves

depends on 4 factors:

1) Particle size - Solids dissolve only where its surface is in contact with solvent

As size increases, surface area to volume ratio decreases, therefore, small crystals

dissolve quicker than large ones i.e granular sugar vs sugar cube

2) Temperature- In most cases – rate increases as temp increases. This is due to

kinetic effects – solvent molecules move faster so they hit the surface of the solid

more often – causes rate to increase.

3) Concentration - As concentration increases, rate of dissolving decreases

When solute and solvent are first mixed, the rate of dissolving is at its max – as it

comes closer to saturation, rate of dissolving decreases. Dissolving becomes very

slow when it gets close to saturation.

4) Agitation or stirring – there is a kinetic effect. Stirring distributes the solute through

the solvent – If it isn’t stirred, as it dissolves the more and more dissolved solute

surrounds the undissolved so rate of dissolving goes down greatly – if its stirred it

increases exposure of undissolved solute to the solvent.

B. Conductivity of aqueous solutions

Remember what an ion is? – a charged atom…

When we are looking at a solution, we can find out if it is conductive or non-conductive. It

is related to the amount of ions in the solution and the phase of the solution.

Phase: A solution must be melted or liquid – but ions must also be present, not all melted

substances break into ions. Solids do not conduct ionically (This is just talking about Ionic

conduction, NOT metallic conduction!)

Concentration: the more ions in solution, the greater the conductivity of the solution.

Note examples in Hebden page 195-196

C. Molecular polarity and Hydrogen bonds

Explanations in the last unit…

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D. Polar and nonpolar solvents and the Nature of solutions of ions

Like dissolves Like – which means polar dissolves polar and nonpolar dissolves

nonpolar.

Solubility discussed above is talking about this.

Simple version - In the case of polar dissolves polar, this makes sense, as both the

molecules of the solvent and solute are polar and they will all be attracted to each other

causing the atoms of the solute to be pulled apart.

A nonpolar solvent would not have the ‘pull’ (only has weak London forces) to be able

to attract the atoms of a polar solute. Nonpolar molecules do not have positive and

negative ends they only have weak London forces. However the pull of the solvents

London forces are strong enough to attract the weak London forces of nonpolar

solutes.

Short story…Like dissolves like.

Some definitions:

Disassociation = reaction that involves ions being separated from an ionic solid – in

this case, ions already exist in the ionic solid and the reaction just pulls them apart

NaCl Na+ + Cl-

Ionization = is a reaction that involves a neutral molecule into ions that did not exist

before putting the molecule in the solvent.

CH3COOH CH3COO- + H+

**both these reactions produce electrically conducting ionic solutions

(you don’t need to be able to predict which reaction will happen)

E. Summary of Bond types and how to distinguish between them: (Hebden pg

208)

Ionic bond – compound formed from a metal and non-metal or polyatomic ions.

Covalent bond – compound formed between 2 nonmetals or non-metal and

hydrogen.

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Hydrogen bond – attraction BETWEEN compounds – between H on one molecule

and N,O, or F of another molecule.

Dipole-dipole – attraction between compounds due to asymmetric molecule

London Force – attraction between compounds – all have London forces – it is the

weakest – if there is no Hydrogen bond or Dipole force there is London force.

F. Calculating concentrations of ions in solution

Reminder from mole concept unit: Solutions are a reaction medium - Many chemical

reactions will not occur unless they are put into a solution -they have no way of coming

into contact with each other in such a way that would allow them to react

Molecules or ions must collide to react – if ions are locked inside a crystal structure,

they cannot react with other ions

When crystal broken down – when dissolved – ions can move and react

Solutions can also act as a diluting agent if reaction is too violent

This set of square brackets means concentration [ ]

When calculating concentration of ions in a solution – the simplest way is to look

at the relative amounts of one ion versus the other versus the original solid:

Example: what is the molar concentration of chloride ions in 0.25 M AlCl3

AlCl3 (s) Al3+ (aq) + 3 Cl-(aq)

There is 1 molecule of AlCl3 to 1 molecule of Al3+ to 3 molecules of Cl- so the

concentration of Cl ions is 3 times that of AlCl3

Just have to multiply 0.25M x 3 = 0.75 M Cl-

So concentration of Chloride ions is 0.75M

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Example: What is the concentration of each type of ion in a solution made by mixing

50.0ml of 0.240M AlBr3 and 25.0ml of 0.300M CaBr2?

First calculate new dilution: (remember M means mol/1L) 0.240 mol (0.050L) = 0.012mol/0.075L = 0.160M AlBr3 1L 0.300 mol (0.025L) = 0.0075mol/0.075L = 0.100M CaBr2 1L

Now calculate concentration of ions in reactions:

AlBr3 Al3+ + 3Br-

0.160M 0.160M 3(0.160M) = 0.480M Where concentration of bromine ions is 3 times the amount of aluminum ions

CaBr2 Ca2+ + 2Br- 0.100M 0.100M 2(0.100M)=0.200M Where concentration of bromine ions is 2 times that of calcium ions. G. Acids and bases and conjugate acid base pairs

We will start out this section with an example of a reaction between an acid and water:

HF (aq) + H2O H3O+ + F-

In this reaction, when HF is put into water, it will dissociate into a H+ which will be

accepted by the water to form a hydronium H3O+ (H+ is also referred to as a proton –

remember a H only has 1 proton and no neutrons and H+ also has no electrons –

therefore, is simply a proton) and a F- ion.

This reaction is reversible and will go back and forth. Therefore, these reactants and

products can be paired:

In one direction, there is a substance that loses a proton (H+), then when the

reaction reverses, that substance can gain a proton (H+) and become an acid

again; HF and F- are called a conjugate acid base pair.

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The HF loses a proton (H+) and becomes F-

HF (aq) + H2O H3O+ + F-

And in the reverse direction, F- will gain a proton (H+)

Look at the other pair in the equation, the same thing is happening except reverse:

In this direction, the H2O gains a proton (H+) and becomes a hydronium ion

HF (aq) + H2O H3O+ + F-

And in the reverse direction, the H3O+ will lose a proton (H+) to become H2O

In this case, H2O and H3O+ are also called a conjugate acid-base pair.

These pairs have a special relationship where you can go back and forth between them

just by either gaining or losing one proton (H+)

By the name, conjugate acid-base pair, you can guess that one of the pair will be

considered an acid and the other will be considered a base.

Two guys, Bronsted and Lowry, both independently came up with the definition of which

is considered an acid and which is a base…

The acid is anything that will give away (donate) one H+

The base is anything that will accept one H+

So in our above example:

The HF is the conjugate acid because it will give away a proton (H+)

HF (aq) + H2O H3O+ + F-

Whereas, the F- is the conjugate base because it will accept a proton (H+)

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The same with the other conjugate acid-base pair:

the H2O gains a proton (H+) so it is the conjugate base

HF (aq) + H2O H3O+ + F-

And the H3O+ is the conjugate acid because it will give up a proton (H+)

Note!! If you gain or lose more than one proton (H+), these ARE NOT considered

conjugate acid base pairs – this reaction only applies if there is ONE proton

gained or lost.

H. Acids, bases and the pH scale

According to our definitions above, an acid is a substance that donates a hydrogen ion

and a base is a substance that accepts a hydrogen ion.

Therefore, you can see that acids and bases all have to do with hydrogen ions. In

water, a small number of the molecules will split apart and when they do, they become

a OH- ion (hydroxide ion which is considered the base because in the reverse reaction it

would accept an H+) and the H+ ion that split off will join another H2O molecule and

become a H3O+ ion (called a hydronium ion – which would give away a proton and is,

therefore, the acid). For simplicity sake, we simplify this and refer to the H3O+ ion

simply as an H+ ion. In pure water, there is an equal number of hydrogen ions and

hydroxide ions, therefore it is considered neither acid nor base – it is considered

neutral.

How acid (relative amount of H+ions) or base a solution is, is measured on a scale

called a pH scale (pH = potential hydrogen). It is a logarithmic scale because it deals

with huge numbers which are based on the H+ ion concentration relative to pure water.

An example of the size of numbers the pH scale deals with (and hence why it is

calculated as a logarithmic scale): a strong acidic solution can have one hundred

million million, or one hundred trillion (100,000,000,000,000) times more hydrogen

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ions than a strong basic solution. On the flip side, a strong basic solution can have

100,000,000,000,000 times more hydroxide ions than a strong acidic solution1.

pH Scale:

1 This and the pH scale from: http://www.sciencebuddies.org/science-fair-projects/project_ideas/Chem_AcidsBasespHScale.shtml

pH Value H+ Concentration

Relative to Pure Water Example

0 10 000 000 battery acid

1 1 000 000 gastric acid

2 100 000 lemon juice, vinegar

3 10 000 orange juice, soda

4 1 000 tomato juice, acid rain

5 100 black coffee, bananas

6 10 urine, milk

7 1 pure water

8 0.1 sea water, eggs

9 0.01 baking soda

10 0.001 Great Salt Lake, milk of magnesia

11 0.000 1 ammonia solution

12 0.000 01 soapy water

13 0.000 001 bleach, oven cleaner

14 0.000 000 1 liquid drain cleaner

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I. A bit about titration

We have used molarities to convert back and forth between the moles of solute and the volume of their solutions. We can calculate the ion concentrations in solution as we did above. But what happens when we don’t know the concentration (molarity) of a solution? We can use a process called titration to figure this out - this process will allow us to determine the molarity of a solution of an acid or base. In a titration, one solution (solution 1, called the titrant) is added from a buret to another solution (solution 2 - so, we use the titrant to titrate solution 2) until a chemical reaction between the two solutions has run to completion. The end point of this reaction is usually shown by a color change caused by an indicator (an indicator is a substance that does not contribute anything to the reaction other than to change color). A titration usually happens this way: A specific volume of a solution to be titrated (solution 2 described above that has an unknown concentration) is put into an Erlenmeyer flask. Then a solution (the titrant) with known concentration that will react with the solution to be titrated is added to a buret. This is set up like the diagram below. A buret is used because it allows the titrant to be added to the solution in the Erlenmeyer flask in a very controlled manner. An indicator is added to the solution to be titrated in the Erlenmeyer flask (a common indicator is phenolphthalein which changes from pinkish red to clear when a solution becomes acidic). The titrant is slowly added until the indicator changes color – this is called the endpoint.

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The volume of titrant added from the buret is then recorded. We can now use all our information to calculate the molarity of the unknown concentration of the solution that we titrated. The following example will be used to show this: Example: If we wanted to determine the concentration of a nitric acid solution we could put 25.00 ml in an Erlenmeyer flask and titrate it with a 0.115M NaOH solution (in the buret). In this example we will say it took 18.3 ml of 0.115M solution. The following calculation can be used to calculate the molarity of the nitric acid solution – we want to calculate: mol HNO3

1L HNO3 solution

NaOH(aq) + HNO3(aq) → NaNO3(aq) + H2O(l)

First, remember that M is mol/L so convert ml to L (you can do this using dimensional analysis as part of the whole equation, but to simplify the process I will convert ml to L first).

18.3 ml 0.115M NaOH solution = 0.0183L of 0.115 M solution 25.0 ml HNO3 = 0.025 L of HNO3

Use your stoichiometry skills to convert from moles of NaOH to moles of HNO3:

0.0183 L NaOH ( 0.115 mol NaOH

1 L NaOH ) (

1 mol HNO3

1 mol NaOH) = 0.0842 M HNO3

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