Equivalence Point pH

STEP ONE: Write a balanced equation for what has happened at the equivalence point. This is not an equilibrium reaction, it is neutralisation.

STEP TWO: Use the n, c, V relationship to calculate the concentration of the conjugate being formed in the neutralisation reaction.

STEP THREE: Write an equation for the dissociation of the conjugate with water.

STEP FOUR: Use the weak base (or weak acid) calculations, as appropriate, to find the concentration of hydronium ions.

STEP FIVE: Use this concentration to calculate pH.

Titration Curves

During a titration, the pH does not change uniformly. There are key points that can be calculated, and the remainder of the curve can be sketched from these:

Initial pH

This example starts with a weak acid in the conical flask. If it were a weak base, the calculations are a lot harder, but you can refer to THIS for how to do that calculation.

Equivalence Volume and Half-Equivalence Point

This relies on a good memory of the titration topic from last year. Once you know the equivalence volume, you are ready to mark the pH of the half-equivalence volume (pKa = pH).

Final pH

This is very easy if you remember how to calculate the pH of a strong base (in this example). Put all of these points together to sketch the shape of the curve.

Equivalence Point pH

This is very hard and seldom asked in NCEA. Having said that, maybe this is the year that it will be asked...

We will cover this in Friday's lesson and dedicate an entire blog post to it.

Buffers

The first things we need to master about buffers:

1. Define a buffer
2. Explain how a buffer works
3. Calculate the pH of buffer solutions

On Mondays lesson, we can use our workbooks (Continuing Chemistry), and the internet as our resources, as well as the following videos:

Our next step is to work out how to make a buffer of a desired pH...

pH of Weak Bases

Whenever these questions are asked, they are worth Excellence. This is because there are six key steps to remember, and they need to be remembered in order!

1. Write out the equation for the base acting as a an alkaline solution.
2. Write a K_b expression.
3. Calculate K_b from K_a
4. Calculate [OH^-] from the K_b expression (assuming [HB⁺] = [OH^-]
5. Calculate [H₃O⁺] from [OH^-], using K_W
6. Calculate pH

This video is very long, so use it to work through an example, rather than trying to learn it all in one go:

pH of Weak Acids

Weak acids only partially dissociate, so how do we calculate their pH?

We were encouraged to work through pp180-181 in Continuing Chemistry to check whether or not we have understood this.

pH of Strong Acids and Bases

This is just a recap of last year, but these skills are vital for moving forward in this topic:

Here is the concept being taught to a Year 12 class:

Acids, Bases and Salts

We started with a recap of the Bronsted-Lowry definitions of acids, bases and amphiprotic species:

Then, we looked at the pH of some salts, comparing them to a control of NaCl (known to have a pH of 7.0):

We need to write chemical equations for the ions to justify the observed pH values (alkaline or acidic):

For example:

HCO₃^- + H₂O <=> H₂CO₃ + OH^-

This equation shows an increase in [OH^-], which is expected as the pH > 7.0

Predicting Precipitation

This was an overview of Ionic Product and we use it to predict if a precipitate will form. Then we looked at what happens when a common ion is added to a saturated solution, for example making a limewater solution more alkaline with sodium hydroxide.

Solubility Product

There is a relationship between the Solubility Product and solubility (in mol L^-1). We also need to remember n = m/M_R and c = n/V for calculations regarding solubility.

Aqueous Solution Introduction

We started the lesson with an overview of the aqueous solution topic, linking it to the chemical reactivity topic from last year:

In the worked example above, we use an equilibrium equation to show the sparingly soluble nature of copper (II) carbonate. This means that we can write an equilibrium expression for it, called the solubility product.

One of the ions made can actually act as a base, making the solution slightly alkaline. We can show this with another equilibrium equation (as it is a weak base). Again, there is an equilibrium expression for this, called the base constant.

The base (carbonate) and its conjugate acid (hydrogen carbonate) can be put together in similar concentrations to create a solution called a buffer. Buffer solutions can resist changes in temperature and pH due to the presence of both the base and its conjugate acid (or vice versa).

Revision Questions

With school exams just around the corner, here are some questions to jog your memory:

Electrochemical Cells

We have already met the concept that spontaneous reactions have a positive net electrode (reduction) potential. It would make sense, therefore, that these chemical reactions generate a voltage (electricity).

Electrochemical cells comprise of two half-cells, one where oxidation occurs and one where reduction occurs. There are conventions for writing cell diagrams, which we need to understand.

Electrode Potential Calculations

We are expected to go over the theory behind electrode potentials ourselves, but we were shown how to use these values to justify if a reaction is spontaneous or not.

Identifying Species

Sometimes, it is unclear which species have been produced. We can use electrode potentials to justify which species are made, but some qualitative tests also help:

Half Equations

We started by looking at the reactions of potassium permangate under different conditions. This video shows the results:

We also need to know how to balance the half equations for these reactions:

Half Equations

Another recap of Level 2 work...

If you want some help learning your redox pairs, try these FLASH CARDS

These FLASH CARDS are also excellent.

Oxidation Numbers

This is just a recap of last year's work. The allocation of oxidation numbers:

Oxidation-Reduction Introduction

Today was a recap of the key elements of Level 2 Redox that will be important to use for the topic and assessment in Level 3:

Bonding Types

We have met types of bonding already, but we need to be more specific about "inter-molecular bonding". Once we classify what type of bonding a substance can make, we can use this to explain physical properties such as melting point trends or molar heat of vaporisation trends.

Polarity

This is a very quick overview for how to work out if a molecule is polar or not:

Shapes of Ions/Molecules

We were introduced to the shapes of molecules/ions with 5 or 6 regions of electron density.

Source: http://figures.boundless.com/10397/full/vsepr-geometries.png

Lewis Dot Diagrams

Start with a recap of Level 2 Chemistry:

Then, how about some examples which break the Octet Rule:

Finally, Lewis Dot Diagrams for polyatomic ions:

Transition Metals

Transition metals are a great context to show our understanding of atomic structure and periodic trends. The transition metals are found in the "d block" of the Periodic Table of Elements.

They have some important (but not necessarily unique) properties:

The following website has an excellent summary of these key properties: http://www.chemguide.co.uk/inorganic/transition/features.html

Variable Oxidation States/Numbers

In all metals, the s-orbital electrons are involved in bonding. This is usually an ionic bond (when it is with another element). As d-orbital electrons can also be involved in bonding, there are many permutations for the number of electrons being involved in bonding, so often more than one possible oxidation state/number.

Zinc is an exception, as its ion has a full 3d-orbital. This makes these electrons unavailable for bonding, similar to the p-orbital electrons of Noble Gases being unavailable for bonding.

This is summarised well in this website: http://www.s-cool.co.uk/a-level/chemistry/transition-metals/revise-it/variable-oxidation-states

Coloured Compounds

Substances are coloured when they absorb part of the electromagnetic spectrum (light) and reflect the remainder. Transition metals' ions are usually coloured because they have incomplete d-orbitals. When bound to other ions/molecules, the d-orbital splits into two groups (with slightly different energy levels). The energy levels of these is so close that white light has enough energy to excited electrons from one energy level to the other. The remaining wavelengths/frequencies are the colour we see.

SOURCE: http://www.chemguide.co.uk/inorganic/complexions/absorbyellow.gif

Zinc is an exception, as its ion has a full 3d-orbital. This makes it impossible for the electrons to move into an excited state due to electromagnetic radiation (light) as they would need to "jump" into a completely different orbital or energy level.

This is summarised well in this website: http://www.s-cool.co.uk/a-level/chemistry/transition-metals/revise-it/coloured-compounds

Catalysts

Simplistically, the density of positive charge in transition metal ions, and their ability to make dative bonds (due to incomplete d-orbitals), make these useful catalysts.

Alternatively, the variable oxidation states of transition metals allow them to be involved in the reaction (alternate pathway provided) then be recovered in their original oxidation state. While being involved, the transition metal is not "used up", so this is still considered to be a catalyst.

There are some excellent examples in these websites: http://www.sky-web.pwp.blueyonder.co.uk/Science/transitionmetalcatalysts.htm and http://www.4college.co.uk/a/ss/catalyst.php

Complex Ions

Explaining the details of this is actually very complicated. However, the reason transition metals can form complex ions is very similar to their ability to be catalysts; a dative bond can be formed with molecules/ions (ligands) due to the charge density of the cation. d-orbitals are involved in this, but the d-orbitals do not need to be incomplete.

This is well explained in this website: http://www.s-cool.co.uk/a-level/chemistry/transition-metals/revise-it/complex-ions

A more in-depth explanation can be found on these pages:

1. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch12/valence.php
2. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch12/crystal.php
3. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch12/ligand.php

Periodic Trends

This week, we need to use our understanding of atomic structure to justify the trends in certain Periodic Trends:

1. atomic radius
2. ionic radius
3. electronegativity
4. 1st ionisation energy

For example:

Electron Configuration and Hund's Rule

We spent more time on using the spdf electron configuration. We also looked at applying Hund's Rule, which helps us understand such things as the electron configuration of Cu:

Using the Periodic Table

In previous years, we focused primarily on using the first 20 elements to understand how the Periodic Table is arranged. From this, we could infer the valency of and ionic charge formed by each element. While this was a good model, it fell down once we got to Sc (#21) and the other Transition Metals.

In this lesson, we were introduced to a more robust model of the atom, and a better way to note electron configurations. The "spdf notation" helps us better explain the properties of the elements.

We also heated some ionic compounds in Bunsen flames, to show that certain elements (and their ions) create very characteristic colours. The electrons move into a much higher Energy Level (excited state), then emit wavelengths of light when they return to their natural Energy Level (ground state).

Entropy

This video sums it up very nicely:

It is worthwhile knowing the sorts of questions we may be asked about entropy as well:

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Hess's Law

We have been looking at Hess's Law over the past week. Here is a video of the key steps being explained:

Spectra Revision

Overview of IR Spectra:

Answering a typical Assessment Task:

Polyamides

The amide bond can sometimes be found in polymers, such as nylon.

Did you know that nylon gets its name from the two cities where it was jointly invented?

Acyl Chlorides

These are wonderful compounds to make as precursors for other organic compounds. We had two challenges today:

1. Explore the preparation and reactions of acyl chlorides. Why are acyl chlorides such useful organic compounds?
   
2. Esters can be made from acyl chlorides. They can also be made from carboxylic acids. We have to compare and contrast these two methods. Which one is better?
   

Carboxylic Acids

Carboxylic Acids are very useful precursor chemicals, as well as being useful chemicals themselves (e.g. ethanoic acid is a good preservative - vinegar).

Haloalkanes

We do not need to know much more about haloalkanes than what we learned last year. We do need to know how to convert haloalkanes into alkenes, however:

Carbonyl Compounds

We are now directing our own learning. We have two lessons to find out as much as we can about aldehydes and ketones:

One idea was to use the aldehyde (or ketone) hexagon and select another hexagon at random. The work out a method to tell the two organic compounds apart.

Optical Isomerism

The challenge was to make an isomer of butanol which had two possible structures, yet both had the same structural formula:

These are called enantiomers or optical isomers. They are very interesting compounds and important in nature. However, the only physical property that can be used to tell them apart is their ability to rotate plane-polarised light - one isomer rotates it clockwise, while the other rotates it anti-clockwise.

Reflux or Distillation?

1-propanol can be oxidised to either propanal (an aldehyde) or propanoic acid, using the same reagents and equipment. The challenge today was to make both propanal and propanoic acid in two experiments, using the same chemicals and equipment. We need to confirm the production of propanal using Benedict's Solution.