OpenLLM-France/wikimedia huggingface dataset:
# A-level Chemistry/AQA/Module 2/Kinetics
## Collision theory

For a reaction to take place between two particles they must collide and do so with enough energy to break bonds. The collision must take place between the parts of the particles that are going to react together, the orientation of the particles also affects the subsequent reaction. To get a lot of collisions you need a lot of particles in a small volume and for the particles to have enough energy to break bonds they need to be moving fast. So for a reaction rate which is fast you need plenty of particles moving rapidly in a small volume.
Most collisions between molecules or other particles do not lead to a reacion as they do not have enough energy or are orientated wrongly.
Factors that affect rate of chemical reactions are as followed.
* Increasing the temperature this increases the speed of the molecules which in turn increases both their energy and the number of collions.
* Increasing the concentration of a solution if there are more particles present in a given volume collisions are more likely therefore the reaction rate would be faster. However as the reaction proceeds the reactants are used up and their concentration falls resulting in most reaction rates to fall as the reaction continues.
* Increasing the pressure of a gas reaction this has the same effect as increasing the concentration of a sloution there are more molecules or atoms in a given volume so collisions are more likely
* Increasing the surface area of solid reactants the greater the total surface area of a solid the more of its particles are available to collide with molecules in the gas/liquid. This can be achieved by breaking up the solid or grinding down into a powder making more reaction sites available thus speeding up the rate of reaction
* Using a catalyst a catalyst is a substance that can change the rate of reactions without being used up throughout the reaction. It also reduces the activation energy needed to kick-start the reaction.
# A-level Chemistry/AQA/Module 2/Redox Reactions
## Oxidation and Reduction

Redox reactions
The word redox is short of reduction-oxidation reaction. An example of an oxidation reaction is that of copper becoming copper oxide with the addition of oxygen.
Cu + ½O$_2$ → CuO
The reverse process is known as the reduction reaction which in this case would involve the removal of oxygen from copper in the following reaction.
CuO + H$_2$ → Cu + H$_2$O
Within these reactions there are what's known as reducing and oxidising agents. The case of the top reaction the oxidising agent is the oxygen and the bottom equation the reducing agent would be the hydrogen.
Should we describe what has happened to the above reaction in term of electrons, we can say that when something loses electrons it is oxidised and when it gains electrons it is reduced. This premise can be easily remembered by the definition that Oxidation is loss and reduction is gain condensed into the acronym OILRIG.

## Oxidation states

We use the idea of oxidation states to allow us to see what is being oxidised and what is being reduced. An alternative name is oxidation number meaning exactly the same thing.
To understand oxidation states the following rules must be taken into consideration.
* Every element in its uncombines state has an oxidation state of 0
* A positive number shows us that element has lost an electron thus has been oxidised e.g. Mg$^{2+}$ has an oxidation state of +2.
* A negative number shows us the element has gained an electron thus has been reduced e.g. Cl$^-$ has an oxidation number of -1
* The more positive or negative number the more it has been oxidized or reduced
* Atoms of element in compounds can lose or gain more than one electron
* The sum of all the oxidation states in a compound equal 0 since all compounds are neutral
* The sum of the oxidation states of an ion e.g. NH$_4$$^+$ add up to the charge on the ion in this case +1
* In a compound the most electron negative element always has a negative value
# A-level Chemistry/AQA/Module 4/Amino Acids

Amino Acids, important class of organic compounds that contain both the amino (8NH2) and carboxyl (8COOH) groups. Of these acids, 20 serve as the building blocks of proteins. Known as the standard, or alpha, amino acids, they comprise alanine, arginine, asparagine, aspartic acid, cysteine, glutamic acid, glutamine, glycine, histidine, isoleucine, leucine, lysine, methionine, phenylalanine, proline, serine, threonine, tryptophan, tyrosine, and valine. All 20 are constructed according to a general formula:
As the formula shows, the amino and carboxyl groups are both attached to a single carbon atom, which is called the alpha carbon atom. Attached to the carbon atom is a variable group (R); it is in their R groups that the molecules of the 20 standard amino acids differ from one another. In the simplest of the acids, glycine, the R consists of a single hydrogen atom. Other amino acids have more complex R groups that contain carbon as well as hydrogen and may include oxygen, nitrogen, or sulphur, as well.
When a living cell makes protein, the carboxyl group of one amino acid reacts with the amino group of another to form a peptide bond. The carboxyl group of the second amino acid similarly reacts with the amino group of a third, and so on, until a long chain is produced. This chainlike molecule, which may contain from 50 to several hundred amino acid subunits, is called a polypeptide. A protein may be formed of a single polypeptide chain, or it may consist of several such chains held together by weak molecular bonds. Each protein is formed according to a precise set of instructions contained within the nucleic acid—the genetic material of the cell. These instructions determine which of the 20 standard amino acids are to be incorporated into the protein, and in what sequence. The R groups of the amino acid subunits determine the final shape of the protein and its chemical properties; an extraordinary variety of proteins can be produced from the same 20 subunits.
The standard amino acids serve as raw materials for the manufacture of many other cellular products, including hormones and pigments. In addition, several of these amino acids are key intermediates in cellular metabolism.
Most plants and micro-organisms are able to use inorganic compounds to make all the amino acids they require for normal growth. Animals, however, must obtain some of the standard amino acids from their diet in order to survive; these particular amino acids are called essential. Essential amino acids for humans include lysine, tryptophan, valine, histidine, leucine, isoleucine, phenylalanine, threonine, methionine, and arginine. They are found in adequate amounts in protein-rich foods from animal sources or in carefully chosen combinations of plant proteins.
In addition to the amino acids that form proteins, more than 150 other amino acids have been found in nature, including some that have the carboxyl and amino groups attached to separate carbon atoms. These unusually structured amino acids are most often found in fungi and higher plants.
# A-level Chemistry/AQA/Module 4/Aromatic Chemistry

Aromatic chemistry is the study of reactions involving the benzene (phenyl-) group C$_6$H$_6$. An aromatic compound containing this group - such as benzene, methylbenzene or TNT (trinitrotoluene) - is called an arene, and a functional group containing a benzene ring is called an aryl group.

## Structure

From around 1865, the accepted model of benzene was the Kekulé structure - effectively cyclohexa-1,3,5-triene. It was thought that there were three double bonds in each benzene ring, with each of the 6 carbon atoms being attached to a single hydrogen nucleus. This model was later revised to suggest that the double bonds oscillated rapidly between the two possible positions (i.e. 1,3,5 and 2,4,6) in the molecule before the 1930s...
...when it was discovered through experiment that bonding in benzene was actually hybrid - neither single nor double - for all C-C links. It was found that, rather than localised π bonds between carbon atoms, a delocalised ring is formed - with each carbon atom donating one p-orbital electron. From this, the skeletal symbol for benzene (a hexagon with internal circle) is derived - and although rarely used, benzene can also be drawn out in full with each carbon atom shown and an internal ring representing the bond.
As there are three areas of electron density around each carbon atom, repulsion theory dictates that the molecule is planar with a 120° bond angle.
This structure is evidenced by:
* Bond length data - experiment has found that C-C bonds are 0.154 nm long, whereas C=C bonds are 0.134 nm. Observation of benzene showed that only one bond length of 0.139 nm was present - uniform to all carbon-carbon links in the molecule.
* thermochemical stability - hydrogenation of cyclohexene (i.e. cyclohexane with one double bond) produces an enthalpy change of -120kJ per mole. Hydrogenation of benzene would therefore be expected to produce an enthalpy change of -360kJ per mole if the structure contained three double bonds - but in reality only -208kJ pre mole is observed. The ring structure makes benzene more stable.

## Halogenation

When a benzene ring is exposed to HCl in the presence of AlCl$_3$ or FeCl$_3$, halobenzenes are formed.
# A-level Chemistry/AQA/Module 1/Bonding
## Intermolecular Forces

The bonds which act between a molecules are called intermolecular forces. The word intermolecular means, "between molecules". There are three kinds of intermolecular forces.

### Van Der Waals

Temporary Induced Dipole-Dipole Attractions Also known as London forces. These are the weakest of the intermolecular forces.

### Dipole-Dipole

Also known as Dispersion forces.

### Hydrogen Bonding

This is the strongest type of intermolecular force. This only occurs between a lone pair of electrons on a Nitrogen (N), Oxygen (O) or Fluorine (F)atom and a hydrogen atom that has a strong partial charge(δ+). The electronegative atom pulls electrons away from the hydrogen so that, on the opposite side to the bond, the hydrogen appears almost like an unshielded proton.

## Bonds
### Ionic

This occurs between oppositely charged ions (atoms or molecules which have gained or lost electrons). An example of an ionically bonded compound is NaCl.
The sodium exists as a positive ion, $Na^{+}$ (positively charged). It loses the electron in its $3s^{1}$ sub-energy level, giving it a stable electron configuration of $1s^{2}2s^{2}2p^{6}$, or $[Ne]$. This electron is donated into the 2p sub-energy level of the chlorine atom, which then forms a chloride ion, $Cl^{-}$. This also now has a stable electron configuration of $1s^{2}2s^{2}2p^{6}$.
The force of attraction (called the "electrostatic force of attraction") between these two ions, caused by their opposite charges, is the ionic bond.
The ions arrange themselves into "giant ionic lattices", which are regular, repeating structures. Ionic compounds are usually white and crystalline in appearance; they have high melting and boiling points, as a lot of energy is required to overcome the electrostatic forces of attraction. They can, however, be disrupted by polar solvents, such as water; they are, therefore, water soluble.

### Covalent

The term covalent bond is used to describe the bonds in compounds that result from the sharing of one or more pairs of electrons. This is usually indicated as H-F. The electron pair creates a 'bond' between the two atoms because it attracts the nucleus of each atom and therefore resist the separation of the two atoms.
A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.

### Metallic

Positive ions are arranged in a lattice with a sea of delocalised (free) electrons. The force of attraction between the delocalised electrons and positively charged ions holds the structure together.
# A-level Chemistry/AQA/Module 1/Periodicity
## Ionisation Energy

Ionisation Energy is the energy needed to dislodge an electron from the attraction which holds it in an atom/ion/molecule. The first ionisation energy is the energy needed to dislodge the least powerfully held electron. Second ionization energy is the energy needed to dislodge the second least powerfully held electron etc.
The size of the ionisation energy for an electron is related to the size of the attraction between that electron and the nucleus to which it is bound. 4 factors effect the size of this attraction.
The number of protons in the nucleus.
The higher the number of protons in the nucleus the higher its total positive charge and the higher the electrostatic attraction between it and the negatively charged electron.
The distance between the electron and the nucleus it is bound to.
Attraction decreases with distance.
The number of intervening electrons between the electron and the nucleus - 'Screening'
Intervening electrons have a negative charge which repulses other electrons and reduces the attraction between electron and nucleus. This is referred to as screening. Note that intervening electrons do not 'block' attraction as the name suggests rather they reduce attraction by providing an opposing repulsion.
Whether or not the electron under consideration is alone in its orbital or paired with another electron
Paired electrons repulse each other and reduce attraction between the electron and the nucleus.

## Changes in first Ionisation Energy within groups and periods

The First Ionisation Energy decreases down Group II. This is because each element has a new shell, increasing the amount of shielding and the atomic radius.
The First Ionisation Energy increases across Period III. This is because:
* The effective nuclear charge increases, due to increase in the number of protons.
* The radius of atom decreases with increasing atomic number, hence more energy is required to remove outer electrons.
* The number of shells remains the same; leading to the same amount of shielding.
There is a drop from Mg to Al because there is a new subshell which is higher in energy (more energy is required to remove the outer electrons). There is another drop from P to S because both electrons are now filled in the 3p subshell, leading to some repulsion.

## Changes in Atomic Radius

Across Period 3 (Na to Ar) as the number of protons in the nucleus increases, the effective nuclear charge increases but shielding by inner electrons is kept unchanged as new electrons enter the same shell as they are attracted by the protons. This in turn will decrease the atomic radius.
# A-level Chemistry/AQA/Module 2/Energetics
## Energy Change

Energetics is all about the energy transfer involved in the making and breaking of bonds. The breaking of a bond involves the injection of energy into the system so that the bond can no longer hold the parts together; the creation of bonds involves the removal of energy from the system so that components are captured by the binding forces. In a typical reaction some bonds are broken and some are created so some energy is injected and some is removed. Whether the total(net) of these amounts to an injection or removal determines whether the reaction is endothermic (involving an net injection of energy) or exothermic (involving a net removal of energy.)

## Enthalpy change (ΔH)

The simple idea described above is complicated by the interaction of the system(reaction) being studied and its environment. This interaction amounts to the changes in volume of the system brought about by the reaction. To deal with this complexity we introduce the idea of enthalpy(given the symbol H). Enthalpy is not measured directly: we concern ourselves instead with CHANGES in enthalpy (ΔH) resulting from a reaction.
For a given reaction change in Enthalpy is equal to the heat energy change under conditions of constant pressure.

### Exothermic

* negative ΔH
Such reactions include the combustion of a fuel like methane CH$_4$ (g) + 2O$_2$ (g) → CO$_2$ (g) + 2H$_2$O (l) = -890kJ mol$^{−1}$ Also including the oxicdation of carbohydrates such sa glucose C$_6$H$_{12}$O$_6$

### Endothermic

* positive ΔH
Such reactions include the thermal decomposition of calcium carbonate CaCO$_3$ (s) → CaO (s) + CO$_2$ (g) = +178kJ mol$^{−1}$

#### Enthalpy change

    The heat energy transferred under constant pressure.
    ΔH = ΔH products – ΔH reactants

#### Standard conditions

We often deal with enthalpy change under standard conditions. i.e.
    100kPa, 298K

#### Standard enthalpy of formation

    $\Delta H_{f}^{\ominus }$
    The enthalpy change when 1 mole of compound in its standard state is formed from its elements in their standard states (under standard conditions).

#### Standard enthalpy of combustion

    $\Delta H_{c}^{\ominus }$
    The enthalpy change when 1 mole of substance is completely burned in oxygen under standard conditions.

## Calorimetry
### Calculating energy changes

* $q=mc\Delta T$
** heat energy = mass of substance x specific heat capacity x temperature change
* $\Delta H={\frac {q}{n}}$
** enthalpy change = heat energy / moles

## Hess's Law
### Hess's Law

    The overall enthalpy change is independent of the route the chemical reaction takes place.

## Bond Enthalpies
### Bond Enthalpy

* Energy/enthalpy needed to break/dissociate a bond averaged over different molecules.
* Always positive because breaking bonds requires energy.
* $\Delta H_{r}^{\ominus }=\Sigma \Delta H_{r}^{\ominus }(bonds\ broken)-\Sigma \Delta H_{r}^{\ominus }(bonds\ made)$
# A-level Computing 2009/AQA/Problem Solving, Programming, Data Representation and Practical Exercise/Fundamentals of Programming

* Fundamentals of programming
* Features of Imperative High Level Languages
* Illustrate these features for a particular imperative, third-generation language such as Pascal.
* A program
* Variables
* Comments
* Input and output
* Data Types
* Use the following appropriately.
* Built-in
* Integer, byte, real, boolean, character, string, date/time.
* User-defined
* Enumerated, subrange, sets, records, arrays
* Data types
** Built-in data types
* Constant Definitions
* Built-in functions
** User-defined data types
* Programming statements
* Type Definitions
* Variable Declarations
* Constant Definitions
* Explain the advantages of named variables and constants.
* Assignment
* Iteration
* Selection
* Procedure/Function Declarations
* Procedure and Function calling - Explain the advantages of procedure/functions.
* Procedure and Function Parameters
* Describe the use of parameters to pass data within programs.
* Understand the different mechanisms for parameter passing: by value and by reference
* Arithmetic operators including modulo arithmetic +, –, /, x, DIV, MOD
* Relational operators =, <, >, <>, <=, >=
* Boolean operators NOT, AND, OR
* Logical bitwise operators NOT, AND, OR, XOR
* Set operators Union, difference, intersection, membership
* Built-in functions
* Arithmetic functions: round, truncation.
* String handling functions: length, position, substring, concatenation.
* String conversion functions to/from integer, real, date/time.
* Selection
* Iteration
* Functions and Procedures
* Operators
** Arithmetic operators
** Modulo arithmetic
** Relational operators
** Boolean operators
** Logical bitwise operators
** Set operators
* The Role of Variables
* Recognise the different roles a variable can take:
* fixed value, stepper, most recent holder, most wanted holder, gatherer, transformation, follower, temporary.
* Global and Local Variables
* The Role of Variables
* Fundamentals of Structured Programming
* Understand the structured approach to program design and construction.
* Construct and use structure tables, structure charts and hierarchy charts when designing programs.
* Use meaningful identifier names.
* Use procedures/functions with interfaces.
* Use procedures that execute a single task.
* Explain the advantages of the structured approach.
* Data Structures
* Arrays
* One- and Two-Dimensional Arrays - Use arrays in the design of solutions to simple problems.
* Fields, Records and Files
* Read/write from/to a text file (including csv file)
* Read/write records from/to a file of records
* Validation
* Understand the importance of validation of input data.
* Program simple validation
* Validation
* Debugging

## Variables

questions
# A-level Computing 2009/AQA/Problem Solving, Programming, Data Representation and Practical Exercise/Fundamentals of Data Representation/Bit patterns in a computer

The language that a computer understands is very simple, so simple that it only has 2 different numbers: 1 and 0. This is called Binary. Everything you see on a computer, images, sounds, games, text, videos, spreadsheets, websites etc. Whatever it is, it will be stored as a string of ones and zeroes.
* All stored as binary
* Exercise: Bit patterns in a Computer
* How do computers store data?
* Answer:
* as binary values, using a pattern of 1s and 0s
* What sort of data can be stored in binary?
* Answer:
* Video
* Sound
* Picture
* Text
* Code
* Spreadsheet
* Game
* etc
* What does the following binary string represent: 10011100
* Answer:
* This could be anything:
* sound data
* picture data
* text ASCII for œ
* unsigned integer = 156
* video data
* etc
* How many bits in a byte?
* Answer:
* 8, but it is originally the amount of bits used to represent a character
* How many bits in 7 bytes?
* Answer:
* 7 * 8 = 56
# A-level Computing 2009/AQA/Problem Solving, Programming, Data Representation and Practical Exercise/Fundamentals of Data Representation/Error checking and correction

When you send data across the internet or even from your USB to a computer you are sending millions upon millions of ones and zeros. What would happen if one of them got corrupted? Think of this situation: You are buying a new game from an online retailer and put £40 into the payment box. You click on send and the number 40 is sent to your bank stored in a byte: 00101000. Now imagine if the second most significant bit got corrupted on its way to the bank, and the bank received the following: 01101000. You'd be paying £104 for that game! Error Checking and Correction stops things like this happening. There are many ways to detect and correct corrupted data, we are going to learn two.