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118 Cards in this Set
- Front
- Back
State electron repulsion theory. |
Each pair of electrons around an atom will repel all other electron pairs The pairs of electrons will therefore take up positions as far apart as possible to minimise repulsion |
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Electron pairs may be... |
...a shared pair or a lone pair |
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State the priniciple of electron repulsion theory. |
The shape adopted by a simple molecule or ion is that which keeps repulsive forces to a minimum |
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Describe the trigonal bypyramid |
If there are five pairs of electrons, the shape usually adopted is that of a trigonal bypyramid |
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If there are five pairs of electrons, the shape usually adopted is that of a trigonal bypyramid, it will be an... |
octahedral structure |
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State bonding pair-lone pair repulsion. |
Angles of a regular tetrahedron are all 109.5 degrees but lone pairs affect these angles. |
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What holds protons and neutrons together in the nucleus? |
Strong nuclear force which is much stronger than electrostatic forces of attraction holding electrons and protons together in an atom. |
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Define the term atomic/proton number. |
The average mass of an atom of an element relative to one twelfth of the mass of an atom of carbon-12 |
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Define the term of element |
Material made up of one type of an atom |
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Define the term compound |
Made up of two or more elements are chemically bonded together |
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Define the term mixture |
Made up of two or more elements aren't chemically bonded |
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Define the term isotope |
An element that has the same number of electrons and protons but a different number of neutrons |
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Define the term relative isotopic mass |
Mass of a single isotope of an element relative to one twelfth the mass of an atom of carbon-12 |
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Define the term relative formula mass |
Mass of an ionic compound or ions relative to one twelfth the mass of an atom of carbon-12 |
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State a fact about isotopes. |
Different isotopes of the same element react chemically in exactly the same way as they have the same electron configuration |
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State another fact about isotopes |
Atoms of different isotopes of the same element varies in mass number because of different number of neutrons in their nuclei |
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What is the mass spectrometer used for? |
It's the most useful instrument for the accurate determination of relative atomic masses. |
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What are relative atomic masses measured on? |
A scale on which the mass of an atom of carbon-12 is defined as exactly 12 |
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What does the mass spectrometer determines? |
The mass of separate atoms (or molecules). |
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What happens in a TOF mass spectrometer? (Outline the steps) |
The substance(s) in the sample are converted to positive ions accelerated to high speeds (depending on their m/z) and arrive at a detector.
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State step 1 of TOF |
Vacuum - whole apparatus is kept under a high vacuum to prevent ions that are produced colliding with molecules from the air |
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Step 2 of TOF |
Ionisation - Sample to be investigated is dissolved in a volatile solvent and forced through a fine hollow needle connected to positive terminal of high voltage supply. |
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Continue step 2... |
Produces tiny positively charged droplets having lost electrons to the positive charge of the supply. Solvent evaporates from droplets into vacuum and droplets get smaller until become single positively charged ion. |
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State Step 3 of TOF |
Acceleration - Positive ions attracted towards negatively charged plate and accelerate towards it. Lighter ions and more highly charged ions achieve higher speed |
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Step 4 of TOF |
Ion drift - Ions pass through a hole in negatively charged plate, forming beam and travel along a tube, celled flight tube to a detector. |
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Step 5 of TOF |
Detection - ions with same charge arrive at the detector, lighter ones are first and have higher velocities. Flight times are recorded, positive ions pick up electron from the detector causing a flowing current. |
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What can mass spectrometer used to identify as well? |
Different isotopes that makes up an element. Detection of individual ions so different isotopes are detected separately have different masses. |
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How many mass spectrometers can measure relative atomic masses up to how many decimal places? |
5 and high resolution mass spectrometry |
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Electrons are no longer considered to be a p_______ but a c______ of n__________ c______ |
particles cloud negative charge |
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What is an atomic orbital. |
An electron fills a volume in space |
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State some facts concerning electron spin |
Two electrons in the same orbital must have opposite signs Electrons are usually represented by arrows pointing up or down showing different directions of spin |
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What are the three rules for allocating electrons to atomic orbitals. |
Atomic orbitals of lower energy are filled first - lower main level's filled first and in this level, sub-shells of lower energy's filled first. Atomic orbitals of same energy fill singly before pairing starts. Electrons repel each other. No atomic orbital can hold more than two electrons |
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Define the term of ionisation energy |
Energy required to remove a mole of electrons from a mole of atoms in the gaseous state and is measured in kJ mol^-1 |
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State Successive Ionisation Energies |
First electron needs least energy to remove it due to it being removed from a neutral atom. Second electron needs more energy than the first being removed from +1 ion etc. |
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Describe trends in ionisation energies across a period in the Periodic Table |
Ionisation energies usually increases across a period because the nuclear charge increases and makes it more difficult to remove an electron |
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Describe the trends in ionisation energies down a group in the Periodic Table |
Shows a general decrease in the first ionisation energy because outer electrons in a main level gets further from nucleus in each case |
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Continue about trends in ionisation energies down a group of Periodic Table |
Nuclear charge increases and it would be more difficult to remove an electron. The actual positive charge felt by an electron in outer shell less than full nuclear charge because effect of inner electrons shielding nuclear charge |
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State the equation for the relative atomic mass |
average mass of an atom of an element __________________________________________ 1/12 mass of one atom of carbon-12 or average mass of an atom of an elementx12/mass of one atom of carbon-12 |
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Define the term relative molecular mass |
Mass of a molecule compared to 1/12 the relative atomic mass of an atom of carbon-12 |
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What is the equation of relative molecular mass? |
average mass of one moleculex12/mass of one atom of carbon-12 |
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What Avogadro's constant? |
Number of atoms in 12g of carbon-12 |
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What is the number of Avogadro's Constant? |
6.022x10^23 |
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State what a mole is. |
Amount of substance contains 6.022*10^23 particles |
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The relative atomic mass of any element in... |
...grams contains one mole of atoms |
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The relative molecular mass/relative formula mass of a substance... |
...in grams containing one mole of entities |
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You can have moles of... |
...ions or electrons |
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State the equation of the number of moles |
number of moles (n) = mass/m (g) ÷ mass of 1 mole/M (g) |
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State the equation of concentration |
no. of moles/n ÷ volume/V (dm^3) |
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State the equation of number of moles in a given volume of solution |
concentration/c (mol dm^-3) * volume/V (dm^3) ÷ 1000 |
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Why do chemical bonds form? |
When atoms bond together they share or transfer electrons achieving more stable electron arrangement - often a full outer main level of electrons. |
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What are the three types of strong chemical bonds |
Ionic Covalent Metallic |
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How many electrons does metals have in their outer main levels? |
One, two or three |
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What is the easiest way for metals to attain the electron structure of a noble gas? |
Lose their outer electrons |
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State some facts concerning ionic bonding |
Occurs between metals and non-metals Electrons transferred from metal atoms to non-metal aoms Positive and negative ions formed |
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Describe ionic bonding using the example of sodium chloride |
Sodium ion's positively charged (lost a -ve electron) Chloride ion's negatively charged (gained a -ve electron) Two ions attracted to each other and to other oppositely charged ions by electrostatic forces. |
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Ionic bonding is a result of... |
Electrostatic attraction between oppositely charged ions.
Attraction extends throughout the compound and every +ve ion attracts every -ve ion, vice versa. |
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Ionic compounds always exist in a structure called a l________ |
lattice |
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State properties of ionically bonded compounds |
Always solids at room temp. Have giant structures High melting temp. because to melt an ionic compound must be supplied breaking up lattice of ions |
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Ionic compounds conduct electricity is molten or solid and explain why |
Molten/dissolved because ions carry current are free to move in the liquid state aren't free in the solid state |
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Why are ionic compounds are brittle and shatter easily when a sharp blow's given? |
(Ionic Compounds) Forms a lattice alternating positive and negative ions. A blow in the direction shown may move the ions and produces contact of ions with like charges |
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State some facts about covalent bonding |
Bond forms between pair of non-metal atoms Atoms share some of their outer electrons so each atom has a stable noble gas arrangement Covalent bond's a shared pair of electrons |
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Atoms with covalent bonds are held together by...? |
...electrostatic attraction between the nuclei and the shared electrons taking place within the molecule |
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Describe double covalent bonds. |
Four electrons shared. E.g. two atoms in an oxygen molecule sharing two pairs of electrons so oxygen atoms have a double bond between them. |
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Describe why molecules have low melting points |
Due to strong covalent bonds between the atoms within molecules. Weak attractions between molecules as they do not need much energy to move apart from each other. |
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Why is the poor conductors of electricity? |
The molecules are neutral overall and no charged particles carrying the current. |
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State what occurs in co-ordinate bonding/dative covalent bonding. |
Atom accepts electron pair is an atom doesn't have a filled outer main level of electrons - atom's electron-deficient Atom that donates electrons has a pair of electrons not being used in a bond (lone pair) |
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Define the term electronegativity. |
The power of an atom attracts the electron density in a covalent bond towards itself |
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Describe the term electron densirt |
Often used to describe the way the -ve charge is distributed in a molecule |
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What is the Pauling scale used for? |
Measure of electronegativity. 0 to 4, greater the number, more electronegative the atom |
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State some factors of electronegativity |
Nuclear charge Distance between nucleus and outer shell electrons Shielding of nuclear charge by electrons in inner shells |
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Explain some factors of electronegativity in order |
Smaller the atom, closer the nucleus is to the shared outer main level electrons and greaer its electronegativity The larger the nuclear charge, the greater the electronegativity |
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Going up a group in the Periodic table, what happened with the electronegativity? |
Increases (atoms get smaller) and less shielding by electrons in inner shells |
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What is the trend of electronegativity going across a period. |
The electronegativity increases and the nuclear charge increases, number of inner main levels remain the same and atoms become smaller |
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Describe the polarity of covalent bonds |
Polarity's about the unequal sharing of electrons between atoms that are bonded covalently. Property of the bond. |
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What else do you need to form a hydrogen bond? |
Very electronegative atom with lone pair of electrons. These will be attracted to the partially charged H atom in another molecule and forms the bond. |
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What are the only atoms that are electronegative enough to form hydrogen bonds? |
Oxygen, Nitrogen, Fluorine |
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Discuss the trend for boiling points of the hydrides (in noble gases) |
Shows a gradual increase in boiling point because only forces acting between atoms are van der Waals forces Increases with number of electrons present |
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Discuss the boiling points of H20, NH3. |
Higher than the hydrides of other elements in their groups because hydrogen's present between molecules in each of these compounds Stronger intermolecular forces of attraction making molecules more difficult to seperate. |
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Discuss the importance of hydrogen bonding |
They are weaker than covalent bonds and can break or make under conditions where covalent bonds are unaffected. |
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Discuss molecules with lone pairs of electrons |
No part of a covalent bonds and lone pairs affect shape of the molecule. |
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Discuss the shape of ammonia |
Has four pairs of electrons and one groups of a lone pair With four pairs of electrons around N atom, ammonia molecule has a shape tetrahedron. Has a triangular pyramid. |
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Bonding pair-lone pair repulsion description. |
Angles of a regular tetrahedron all 109.5degrees but lone pairs affect the angles. The angles compress due to repulsion |
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Provide a summary for repulsion between electron pairs |
repulsion increases ↓ bonding pair - bonding pair ↓ lone pair-bonding pair ↓ lone pair-lone pair |
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Solid What is the arrangement of particles and evidence? |
Arrangement: Regular Evidence: Crystal shapes have straight edges. Solids have definite shapes |
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Solid What is the spacing of particles and state the evidence for it? |
Spacing: close Evidence: Solids are not easilt compressed |
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Solid What is the movement of particles of solids and the evidence for it? |
Movement: vibrating about a point Evidence: diffusion is slow. Solids expand when heated |
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Liquid What is the arrangement of particles and what is the evidence for it? |
Arrangement: Random Evidence: Liquid changes shape to fill bottom of its container |
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Liquid What is the spacing of particles and what is the evidence for it? |
Spacing: Close Evidence: Liquids aren't easily compressed |
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Liquid What is the movement of particles and what is the evidence for it? |
Movement: Rapid 'jostling' Evidence: Diffusion is slow. Liquids evaporate |
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Gas What is the arrangement of particles and what is the evidence? |
Arrangement: Random Evidence: None direct but a gas will fill its container |
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What does a Maxwell-Boltzmann distribution show? |
It shows distribution of energies at given temperatures |
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When a factor increases a rate of reaction what happens? |
A higher proportion of molecules will reach the activation energy |
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Define the term standard enthalpy of combustion |
The enthalpy change when one mole of substance is completely burnt in oxygen under standard conditions, all reactants and products being in their standard states. |
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What is the structure of a haloalkane? |
Contains the functional group C-X where X is a halogen (F, Cl, Br or I) |
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What are the types of haloalkanes? |
Haloarenes Haloalkanes |
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Describe haloalkanes |
A halogen is attached to an aliphatic skeleton - alkyl group |
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Describe haloarenes |
Halogen's attached directly to a benzene ring |
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Describe the physical properties of haloalkenes |
Boiling points increase with mass. The greater the branching the lower the boiling point. Haloalkanes are soluble in organic solvents but insoluble in water - they are not polar enough and do not exhibit hydrogen bonding. |
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What is the standard molar enthalpy of formation? |
Enthalpy change when one mole of substance is formed its constituent elements under standard conditions, all reactants and products being in their standard states |
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What is the standard molar enthalpy of combustion? |
Enthalpy change when 1 mole of substance's completely burned in oxygen with all reactants and products |
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What happens when an enthalpy change occurs? |
Energy is transferred between system and surroundings |
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Define enthalpy change |
The amount of heat energy taken in or given out during any change in a system provided the pressure is constant |
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Define standard enthalpy change of formation |
Energy transferred when 1 mole of the compound formed from its elements under standard conditions (298K and 100kPa), reactants and products bringing their standard states |
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State some facts exothermic reactions |
Energy is given out to surroundings Enthalpy change is negative Products have less than reactants |
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State some facts about endothermic reactions |
Energy taken in from the surroundings Enthalpy change is positive Products have more energy than reactants |
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What are isomers? |
Molecules have the same molecular formula but atoms are arranged differently |
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What are alkanes? |
Saturated hydrocarbons - contain only carbon-carbon and carbon-hydrogen single bonds |
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What are physical properties of alkanes? |
Polarity - alkanes are almost non-polar as electronegativities are similar Boiling points - increasing intermolecular force is why boiling points increase as chain length increases Solubility - insoluble in water. Water molecules are held together by H bonds which are stronger than van der Waals' forces |
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Describe alkanes reactivity |
Unreactive. They have strong carbon-carbon and carbon-hydrogen bonds. They don't react with acids, bases and oxidising agents |
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Why are shorter chain hydrocarbons more in demand? |
They are more reactive and economically valuable |
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What two useful results does cracking lead to? |
Shorter, more useful chains are produced, especially petrol Some products are alkenes, they are more reactive than alkanes |
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(thermal cracking) What are the conditions required for cracking? |
High temperature, 700-1200 K, under high pressure, up to 7000kPa |
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What does thermal cracking produce? |
High proportion of alkenes. |
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(Catalytic cracking) What are the conditions required for cracking? |
Lower temperature (approx 720 K), lower pressure (but more than atmospheric) and in a presence of a catalyst. |
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State the catalyst used in catalytic cracking? |
Zeolite catalyst (silicon dioxide and aluminium oxides) - honeycomb structure with enormous SA |
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What products are a result of catalytic cracking? |
Produce motor fuels. Mostly branched alkanes, cycloalkanes (rings) and aromatic compounds |
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When does incomplete combustion occur? |
When there is a limited supply of oxygen. |
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What is flue gas? |
Gases given put by power stations |