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168 Cards in this Set
- Front
- Back
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How do you work out the number of particles?
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mol*Avogadro constant (6.02*10^23)
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At RTP what volume does 1 mol of gas occupy? |
1 mol of any gas occupies a volume of 24dm^3 at RTP.
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Explain properties of a gas. |
gases are all in continuous motion and have no intermolecular forces. They all exert pressure when they collide. To maintain kinetic energy all collisions act like elastic. They are considered identical as they are all so small in comparison to their container. |
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What are the conditions for RTP?
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1 atm 101325 pa 298K 1 moldm^-3 |
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what is concentration?
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The amount of solute within a solution. worked out by dividing moles by volume.
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what is an anhydrous salt?
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A salt formed once dehydrated. |
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Describe an exothermic reaction.
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In an exothermic reaction energy is released into the surroundings so temperature change is positive. Although the temperature change is positive the enthalpy change is negative as energy is lost to the surroundings.
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What is the limiting reagent? |
the substance with the smaller amount of moles. worked out by dividing energy per mole by energy transferred.
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Define the enthalpy change of formation
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The enthalpy change that takes place when 1 mole of a compound is formed from its constituent elements in their standard states under standard conditions. |
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Define the enthalpy change of neutralisation
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The enthalpy change which accompanies the neutralisation of an aqueous acid by an aqueous base to form 1 mole of water under standard conditions.
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Define the enthalpy change of reaction
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The enthalpy change which accompanies a reaction in the molar quantities expressed in a chemical equation under standard conditions with all reactants and products in their standard states.
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What is Hess' law?
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The idea that a reaction should be able to take place by multiple routes if the initial and final conditions are the same. This means that the total enthalpy change should be the same for each route.
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Define bond enthalpy
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The enthalpy change that takes place when breaking by homolytic fission 1 mol of a given bond in the molecules of a species in the gaseous state.
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define average bond enthalpy
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The average enthalpy change that takes place when breaking by homolytic fission 1 mol of a given bond in the molecules of a species in the gaseous state.
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Why are bond enthalpies positive? |
Energy is absorbed when the bonds are broken.
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Define rate of reaction?
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The change in concentration of a reactant or product over time. Rate = Change in conc/ time
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What conditions are required for a successful collision? |
Particles must collide with correct orientation and have sufficient energy to overcome the activation energy.
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List 5 factors which effect reaction rate |
pressure temperature surface area catalyst concentration |
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How do you work out energy released? |
-MC(delta)T
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how do you work out yield?
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(actual yield/ theoretical yield) * 100
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What is the ideal gas equation?
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pV = nRT T = pV/nR |
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How do you work out the mol of a gas?
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mol = volume(dm^3)/ 24 |
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how do you work out atom economy?
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(RFM of desired product/ RFM total product) * 100
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what was Daltons atomic theory? |
atoms are tiny particles that make up elements. atoms cannot be divided. all atoms of a given element are the same. atoms of one element are different from those of every other element. |
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what did J.J Thomson's do? |
Discovered that a cathode ray acted as a stream of particles that had a negative charge, could be deflected by both a magnet and an electric field and had a very smal mass. Thomson concluded that these were electrons and that the atom could be split and these properties came from the atoms. Thomson proposed that atoms are made up of e- moving around in a 'sea' of positive charge with a positive charge. He named this the plum-pudding atom. |
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Outline Ernest Rutherford's gold-leaf experiment |
He directed alpha particles towards a sheet of very thin gold foil and measured any deflection of the particles. Rutherford calculated that a plum-pudding atom would hardly deflect alpha particles. in 1911 he proposed a new model showing that negative e- orbit this nucleus similar to planets orbiting the sun and the positive charge and most of its mass are concentrated in a nucleus. Most of the atoms volume would be the space between the nucleus and orbiting e- and charges must balance. This was known as the nuclear atom. |
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outline Niel Bohr's planetary model and Henry Moseley's work on atomic numbers |
Bohr altered Rutherford's model to allow e- to follow certain paths. Otherwise e- would spiral into the nucleus. Electrons orbit a central nuclear 'sun' in shells. Bohr's model explained spectral lines seen in emission spectra and the energy of e- at different distances from the nucleus. Mosely added to this idea by discovering a link between X ray frequencies and an elements atomic number. |
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How did Rutherford aid Moseley's work on atomic numbers? |
He discovered the proton and was able to explain Moseleys finding that an atom's atomic number was linked to x-ray frequencies. |
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Use research theories to decribe different beliefs on particle behaviour |
Louise de Broglie suggested that electrons could have the nature of both waves and particles. However Erwin Schrodinger suggested that an electron had wave like properties in an atom. He also introduced atomic orbitals. |
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Outline James Chadwick's dicovery |
Discovered the neutron when observing radiation emitted from some elements. He found that it had no charge and weighed about the same as a proton. |
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Describe the proton |
Sub atomic particle with a relative mass of 1 and relative charge of 1+ and is found in the nucleus. |
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Describe the neutron |
Sub atomic particle with a relative mass of 1 and relative charge of 0 and is found in the nucleus. |
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Describe the electron |
Sub atomic particle with a relative mass of 1/2000 and relative charge of 10 and is found in the nucleus. |
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How do reactions of isotopes differ? |
They dont as theyre the same element and the differing amounts of neutrons make no difference to chemical reactivity. |
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What is mass spectrometry? |
A piece of apparatus that can be used to find out: the identity of an unknown compund, find the relative abundance of each isotope of an element and determine structural information about molecules. |
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An element has 3 isotopes, use their abundance to determine the relative atomic mass of the element. x-24 79% x-25 10% x-26 11% |
(24*79)+(25*10)+(26*11)/100 |
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What is the difference between G1-3 and G5-7 atoms? |
G1-3 lose electrons to form positive ions whilst G5-7 gain electrons to form negative ions. |
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Why don't Be, B, C and Si usually form ions? |
Too much energy is required to transfer the outer shell electrons to form ions. |
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What is the molecular charge of ammonium, hydroxide, nitrate, carbonate and sulfate? |
NH4^+ OH^- NO3^- CO3^2- SO4^2- |
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What is stoichiometry? |
The amount of substances that are invlolved in a chemical reaction. To work out the stoichiometry a you need a balanced equation. |
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State reasons why a reaction may not have 100% yield |
Reaction may be at equilibrium and may not go to completion. Side reactions may occur leading to by-products. The reactants may be impure. Some of the reactants or products may be left behind in the apparatus. Seperation and purification may lead to loss of product. |
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what type of mechanism has 100% atom economy? |
addition reactions! |
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what happens to an acid when it is added to water? |
the acid releases H+ ions ie HCl(g) -> H+(aq) + Cl-(aq) |
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what is an alkali? |
A form of base that is able to dissolve in water to form OH- ions. |
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What is an amphoteric substance? |
a substance that can act as an acid and a base ie glycine. |
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what is the ionic equation for neuralisation? |
H+(aq) + OH-(aq) -> H2O(l) |
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what is a salt? |
A compound that contains a cation and an anion. |
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What's the difference between a cation and an anion? |
A cation is a positive ion and an anion is a negative ion. |
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How do acid salts react? |
Similarly to a salt the H+ ion can be replaced to form a conventional salt. Ie H2SO4 -> Na2SO4 |
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How are salts formed? |
Salts are formed by neutralising acids with bases. |
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How are ammonium salts formed? |
Neautralising an acid with aqueous ammonia(NH4). |
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How do you carry out a titration? |
Using a pipette add a measured volume of one solution to a conical flask. Add a suitable indicator. Place another solution into a burette. Add the solution in the burette to the conical flask until the reaction reaches its end-point. |
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What colour changes will you observe with methyl orange? |
red- acid, yellow- base, orange- end point |
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What colour changes will you observe with bromothymol blue? |
yellow- acid, blue- base, green- end point |
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what could changes will you observe with phenolphthalein? |
colourless- acid, pink- base, pale pink- end point |
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What do oxidation numbers monitor? |
How electrons are being used in bonding. |
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What is the oxidation number of an uncombined element? |
0 |
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What is the oxidation number of a combined Oxygen? |
-2 |
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What is the oxidation number of a combined oxygen in a peroxide(H2O2)? |
-1 |
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What is the oxidation number of a combined hydrogen? |
+1 |
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What is the oxidation number of a combined hydrogen in metal hydride(LiH)? |
-1 |
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What is the oxidation number of a combined fluorine? |
-1 |
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What is the oxidation number of a simple ion? |
The charge given on the ion. |
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What is an Oxyanion? |
A negative ion that contains an element along with Oxygen ie SO4^2-. |
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What is the different between oxidation and reduction? |
Oxidation is loss of e- and Reduction is gain of e-. |
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What is a redox reaction? |
A reaction that has both an oxidation and reduction reaction taking place. Within redox reactions oxidation has an increase in oxidation number whilst reduction decreases in oxidation number. Ie Mg + Cl2 -> MgCl2 Mg -> Mg^2+ + 2Cl- Cl2 + 2e- -> 2Cl- |
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How do redox reactions occur between reactive metals and acids? |
The metal is oxidised forming a positive metal ion. The hydrogen in the acid is reduced forming H2(g) and a salt. |
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For the first 4 shells how many e- can each shell hold? |
shell 1- 2 shell 2- 8 shell 3- 18 shell 4- 32 |
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How are quantum numbers used? |
To describe the electrons in shells. Different shells have different quantum numbers and the larger the quantum number the further it is from the nucleus and the higher the energy level. |
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describe an s-orbital |
Spherical in shape and has 2 e- per shell. (1s2) |
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describe a p-orbital |
3D dumbell shape and contains three p-orbitals, Px, Py,Pz each at right angles to one another. Each p-orbital holds 6 e- (two per orbital). (1s2,2s2,2p6) |
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How many electrons can a d-orbital hold? |
10 as each shell contains 5 d-orbitals. (1s2,2s2,2p6,3s2,3p6,3d10) |
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How many electrons can a f-orbital hold? |
14 as each shell contains 7 f-orbitals. (1s2,2s2,2p6,3s2,3p6,3d10,4s2,4p6,4d10,4f14) |
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Whats the difference between electron configuration in an atom and orbitals occupied? use an example |
Electron shows overall orbitals whereas orbitals goes into further detail. ie element B orbitals occupied(1S^2,2S^2,2Px^1,2Py^1) electron configuration(1S^2, 2S2, 2P2) |
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What would be the electronic configuration of Li+ ion if it was originally 1S^2, 2S^1? Why? |
1S^2 because e- in the highest energy levels are lost first when a positive ion is formed. |
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What would the electronic configuration of F- ion be if it was orginally 1S^2,2S^2,2P^5? Why? |
1S^2,2S^2,2P^6 because in the formation of a negative ion an e- is added to the highest energy level. |
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What is ionic bonding? What diagram is used? |
The bonding between a metal and non-metal. Electrons are transferred from the metal atom to the non-metal atom to form oppositely charged ions that attract each other meaning they have an electrostatic attraction between positive and negative ions. They use dot and cross diagrams with square brackets and their charge at the top corner. |
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What is covalent bonding? What diagram is used? |
The bonding between two non-metals to share electrons. Has a stong electrostatic bond between a shared pair of electrons and the nuclei of both atoms. Shared dot and cross diagram is used. |
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What is metallic bonding? |
The bonding between metals that shares electrons between all metal atoms. |
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How do giant ionic lattices work? |
Two ions with a balancing charge(ie Na and Cl) that fully surrounds each other to form a 3D lattice. Each atom is surrounded by opposite charges and ions are attracted for each direction. |
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Why can't solid giant ionic lattices conduct electricity? |
Ions are free to move about in liquids and conduct electricity whereas in a sold ions are in fixed positions so no ions can move. |
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Why is a lot of energy required to break bonds between atoms in a solid giant ionic lattice? |
Solid giant ionic lattices have strong electrostatic bonds that hold oppositely charged ions together. |
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Why does MgO have a higher melting point than NaCl? |
Mg has 2+ ions and O has 2- ions meaning they have a greater charge and stronger electrostatic forces between the ions. Therefore more energy is required to break these bonds than it would to break bonds in MgO. |
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How does water break down an ionic lattice? |
The polar water molecules break down an ionic lattice by surrounding each ion to form a solution. the slight charges within the polar substances are able to attract the charged ions in the giant ionic lattice so the lattice is disrupted and ions are pulled out. |
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What are lone pairs? Use an example |
A pair of electrons that are not used in bonding and gives a concentrated region a negative charge around that atom. H2O has 2 lone pairs as only 2 pairs of electrons are used in covalent bonding. |
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What is a dative covalent bond? |
A covalent bond has an atom that supplies both the shared electrons. NH4 has 3 covalent bonds and 1 dative covalent bond. Once formed dative bonds are the same as covalent bonds. |
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When can the octet rule not be used? |
when there isnt enough e-. when more than four e- pair up in bonding. |
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what is the octet rule? |
The octet rule reflects the observation that atoms of main-group elements tend to combine in a way so that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. |
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how can the octet rule be improved? |
state that unpaired e- pair up. The maximum number of e- that can pair up is equal to the number of e- in the outer shell. |
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What is a simple molecular structure? |
A simple solid molecular lattice which is solely bonded using covalent bonds. The different molecules are held together by weak intermolecular forces and weak london forces. Eg Iodine. |
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List properties of simple molecular structures |
Low melting/ boiling points as weak intermolecular forces are easily broken. Doesnt conduct electricity as they contain no charged free particles to move around. Mostly soluble in non-polar solvents such as hexane as weak london forces form between the solvent and covalent molecule helping the lattice break down. |
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List properties of giant covalent structures |
High melting/ boiling points as strong intermolecular forces are tough to break. Doesnt conduct electricity as they contain no charged free particles to move around except in graphite. Insoluble in both polar and non-polar solvents as the covalent bonds in the lattice are too strong to be broken. |
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What is the electron pair repulsion theory? |
As electrons all have negative charge each e- pair repels other e- pairs. The shape used will allow all the pairs of e- to be as far apart as possible. |
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How many bonded electron pairs does a trigonal planar have around its central atom and what is the angle between these atoms? |
3 bonded electron pairs around its central atom at 120 degrees to one another. |
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How many bonded electron pairs does a tetrahedral have around its central atom and what is the angle between these atoms? |
4 bonded electron pairs around its central atom at 109.5 degrees to one another. |
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Whats the difference between a bonded pair than a lone pair? |
lone pairs are more electron dense and repel more. |
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Decribe a methane molecule |
Has a tetrahedral shape with bond angles at 109.5 degrees as its a molecule with lone pairs. |
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Decribe an ammonia molecule |
Has a pyramidal shape with bond angles at 107 degrees as its a molecule with lone pairs. |
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Describe a water molecule |
Has a non-linear shape with bond angles at 104.5 degrees as its a molecule with lone pairs. |
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What is electronegativity? |
The attraction of a bonded atom for the pair of electrons in a covalet bond. Top right of periodic table is the most electronative(fluorine). |
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Why is Hydrogen(H2) non-polar? |
It is an identical molecule meaning each atom is identical resulting in a 100% covalent bond. As the nucleus is equally attracted to each pair the e- in the bonds are evenly distributed. Due to this H is the least electronegative element. |
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Why is HCl polar? |
Cl is more electronegative and has a greater attraction for the bonding pair of e- compared to hydrogen resulting in the bonding electrons to be held closer to the Cl atom. Cl has a slight negative charge and H has a slight positive charge. |
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Why do dipole bonds in a symmetrical molecule cancel out? |
The dipoles act in different directions. |
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Why do intermolecular forces occur? |
They occur due to constant random movements of e- in the shells of the atoms within the molecules. They do not involve sharing or transferring e-. |
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What is the relative strength of each bond type? |
Ionic and covalent- 1000 hydrogen bonding- 50 permanent dipole-dipole - 10 london dispersion forces-1 |
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Why do some molecules have permanent dipoles? |
Polar molecules are present so it has a slightly positive and slightly negative charge. |
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How do permanent dipoles induce non polar molecules? |
Shifts the electrons slightly by repelling with the negative end or attracting them with the positive end. This makes the molecules become slightly polar and an attraction occurs. |
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How does a permanent dipole to permanent dipole interact? |
Opposite ends attract like magnets. |
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Why do non polar molecules have weak intermolecular forces? |
They have no dipole! |
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What causes London dispersion forces? |
Constant random movement of electrons in an atoms shell. This movement unbalances the distribution of charge within the electron shells. An instantaneous dipole is formed and induces a dipole in neighbouring molecules. The small induced dipoles attract one another causing weak intermolecular forces/ LDF. |
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How do you increase the size of london dispersion forces? |
Increase the amount of electrons so the induced dipole is larger and has greater attractive forces between molecules. |
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Why do non-polar molecules have low melting/boiling points? |
They only have london forces between molecules which are weak and can easily be broken. |
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What is a hydrogen bond? |
A strong bond between hydrogen and a highly electronegative element which have permanent dipoles. |
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Why is ice less dense than water? |
When ice forms water molecules arrange themselves into a pattern and a hydrogen bond forms between each molecule. Ice has an open lattice with hydrogen bonds holding the water molecules apart so when the ice melts the rigid hydrogen bonds collapse allowing the H2O molecules to move closer together. |
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Why does H2O have an unusually high boiling point? |
The hydrogen forces are stronger than the other forces so these forces must be overcome before the water can be boiled. |
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Define valency
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Available electrons for bonding ie carbon has a valancy of 4
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Why does bromine react with unsaturated molecules?
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Double/ triple bonds are highly reactive as they have a higher density making them more reactive.
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What can't polar and non polar molecules do?
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mix!
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How do volcanoes emit greenhouse gases?
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When the volcano erupts it releases greenhouse gases naturally. These gases vibrate due to absorbing infrared radiation. The vibrating molecule can re-emit this radiation for other molecules to absorb. Alternatively it can be lost into space or absorbed by the Earth's surface.
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What conditions are required to react an alkane and a halogen?
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UV light and temperature of 300 degrees C.
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What are the radical substitution steps?
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Initiation- UV breaks down Cl-Cl bonds homolytically to generate two radicals. Propagation- Cl radical is then free to react further with another Cl2 molecule and firm the monosubstituted product. Cl radical is regenerated in this process. Termination- two radicals randomly collide forming a stable product. |
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Use ethane and bromine to show radical substitution using equations.
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Br-Br -> 2Br* Br* + C2H6 -> HBr + *C2H5 2Br* + *C2H5 -> BrC2H5 + Br* Br* + *C2H5 -> BrC2H5 |
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What are the rules for cis/trans isomerism?
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C=C double bond is required and each C atom must be bonded to the same group.
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What are the rules for E/Z isomerism using Cahn-Ingold-Prelog rules?
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C=C double bond and each C atom must be bonded to a different group. Group with the highest RFM takes priority.
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How do you classify a secondary alcohol?
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Check whether the Carbon atom is attached to 2 bonds/groups ie butan-2,3,di-ol
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What are the trends you would expect to find actross a period? |
Decreasing atomic radius. Structure will change from a giant lattice to a simple molecule. Will decrease in electrical conductivity. Melting/ boiling point will show a bell curve. First ionisation energy will increase. |
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What are the trends you would expect to find down a group? |
Atomic radius will increase as will ionic radius. Melting/ boiling point will decrease as will first ionisation energy. |
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Explain why the atomic radius will decrease as you move across a period. |
The charge of the nucleus increases whilst the electron shielding remains the same. Therefore the effect the nuclear charge has on the outer electrons increases and the atomic radii decreases. |
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Explain why the atomic radius will increase as you move down a group. |
The charge of the nucleus increases whilst the electron shielding increases. Therefore the effect the nuclear charge has on the outer electrons decreases and the atomic radii increases. |
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How does bonding down a group differ include a n example. |
Bonding changes from covalent to metallic. ie Group 14 Carbon has covalent bonds whilst Lead has metallic bonds. |
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How does bonding differ across a period include an example. |
Bonding changes from metallic to covalent. ie Period 3 Sodium has metallic bonds whilst Chlorine has covalent bonds. |
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Why is Sodium a good electrical conductor? |
Sodium has metallic bonds and therefore has mobile outer shell electrons. These electrons are delocalised which allows metals to conduct heat and electricity. |
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Why is Graphite a good electrical conductor? |
It has mobile outer shell electrons which are delocalised. The electrons are able to freely move around and conduct electricity. |
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Why do simple covalent molecules have lower boiling points than giant lattices? |
They have weak intermolecular forces such as induced dipole-dipole interaction forces so small amounts of energy is required to overcome them making melting point low. Whereas giant lattices have strong bonds throughout the lattice all of which must be broken. This requires a lot of energy making the boiling point high. |
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Why does melting point increade from group 1 to group 13? |
They all have giant metallic lattices with strong bonds between all atoms however metallic bonds in G13 elements are stronger then G1 and G2 elements. |
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Explain the melting point trend in group 17. |
The only forces that need to be broken are induced dipole-dipole forces which depend on the number of electrons. As halogens decend the number of electrons increases so the melting point also increases. |
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Describe trends in first ionisation energies. |
When removing an electron the attraction between the electron and the nucleus must be overcome. The further the electron is from the nucleus the easier it is to remove it- this decreases ionisation energy. As the number of protons in the nucleus increases its attraction for outer most electrons increases- this increases ionisation energy. Electrons in inner shells exert a repelling effect on the electrons in an outer shell reducing the pull of the nucleus on the outer electrons. This reduced the effect of the nucleus and decreases ionisation energy. |
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Why does ionisation energy decrease down a group? |
Atomic radius increases moving the outer electron away from the nucleus. Nuclear charge and shielding effect increases down the group. Although nuclear charge increasing down a group increases ionisation energy atomic radius and shielding effect outweighs this causing the ionisation energy to decrease. |
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Why does ionisation energy increase across a period? |
Atomic radius decreases pulling electrons closer to the nucleus. Nuclear charge increases down the group and shielding effect remains the same. |
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Which ion if formed by ionisation energy? |
The positive ion/ cation. |
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Describe physical properties of group 2 elements. |
As they're all metals they can all conduct electricity and have high melting points. Other than beryllium they form colourless ionic compounds with that also have high melting points. These compounds are good conductors of electricy when molten or in aqueous solution but poor when solid. |
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Describe chemical properties of group 2 elements. |
As you move down the group reactivity increases. They react by losing 2 electrons forming 2+ ions which are isoelectronic with a nobel gas. When reacting with oxygen they undergo redox reactions forming metal oxides. Similarly when reacting with water they undergo a redox reaction forming metal hydroxide and water. All G2 elements form white precipitates when reacting with chlorine. When reacted with a dilute acid metal salt and hydrogen gas forms. |
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State what colour a flame will turn when a group 2 element is reacted with oxygen. |
Magnesium - white Calcium - brick red Strontium - crimson Barium - green |
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What type of agent is a group 2 element and why? |
Reducing agent as it's oxidised shown by the oxidation state increasing from 0 to +2. |
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Name some purposes of a group 2 metal oxide. |
MgO is used as a refractory ceramic to line furnaces as it has a high melting point. CaO is known as quicklime and produces Ca(OH)2 when reacted with water, this is used to reduce acidity in soil. |
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Why does the state of water matter when reacting a group 2 element? |
If the water is in liquid form the metal will react slower and will form a white metal hydroxide precipitate. Whereas if the water is a gaseous state it will react readily to form a metal oxide. |
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How can magnesium hydroxide be used commercially? |
Used as an antacid to treat indigestion by reacting with water. |
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Why are group 2 oxide reactions not redox reactions? |
Oxidation numbers of all the elements remain the same. |
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What properties does group 2 caronates have? |
Insoluble in water and decompose to form carbon dioxide and metal oxide. Decomposition rate decreases as you go down the group therefore Barium carbonate requires strong heating to start decomposition. BeCO3 is so unstable it does not exist at RTP. No group 2 carbonate reacts in a redox reaction as the oxidation number remains the same. |
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What properties does group 2 sulphates have? |
Solubility decreases down the group and they are all white solids. |
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What uses can group 2 sulphates have? |
Hydrated magnesium sulphate is used as an epsom salt laxative. Barium sulphate is insoluble in water and can be used to test for sulphate ions. Barium sulphate also absorbs Xrays strongly and is used to diagnose stomach/ intestine disorders. Soluble Barium sulphate compounds are toxic but as barium sulphate is insoluble it is not absorbed. Xrays cannot pass through the 'barium meal' creating a shadow on the Xray film. |
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How does reactivity differ in halogens? |
Reactivity increases up the group making Fluorine and Chlorine the most hazardous. |
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What can halogens be used for? |
Used in plastics, pharmaceuticals, anaesthetics, dyestuff and chemicals for water treatment. |
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How do halogens bond together? |
Bonded together covalently to form non polar diatomic molecules. |
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What trends are there in halogens? |
Atomic radius increases down the group, ionisation energy decreases down the group and reactivity increases up the group. When a halogen gains an electron a 1+ ion is formed. |
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What characteristic physical properties does a halogen have? |
Poor conductors as they are non-metals. At room temperature fluorine is a yellow gas, chlorine is a green gas, bromine is brown liquid and iodine is a black solid which sublimes to form a violet gas. As you move down the group there is an increase in induced dipole-dipole interactions as the amount of electrons in the molecule is increasing, this leads to a reduced volatility and an increase in boiling point. Fluroine is the most electronegative element and is a powerful oxidising agent. |
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List a use for each halogen. |
Fluorine - teflon non stick polymer Chlorine - PVC or as a bleach Bromine - medicines Iodine - used within our diet so the thyroid gland can make thyroxine ie milk |
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What are the chemical properties of halogens? |
They react by gaining an electron to form halide anions. As atomic radius and shielding increases (down the group) the nuclear charge decreases allowing electrons to gain less readily. Fluorine is the most electronegative element and a powerful oxidising agent. |
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What halogens displace each other? |
Fluorine displaces Cl-, Br- and I- from solution. Chlorine displaces Br- and I- from solution. Bromine displaces I- from solution. Iodine does not displace any halide. |
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What can be determined about halogen oxidation using displacement? |
That as you move down a group oxidation power decreases. |
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What solvents can halogens dissolve in and why? |
Non-polar solvents such as hexane as it is immiscible with water and will form two layers. The solutions remain the same colour as the halogen vapour. |
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What can chlorine be used in? name some advantages and disadvantages of this |
Water treatment by reacting reversibly with water and forming a mixture that kills bacteria. Cl2 + H20 -> HCl + HClO The chlorine undergos both oxidation and reduction. Adv: reduces risk of disease by passing HClO through the molecules inside bacteria cells that then break up and kill the organism. Dis: by products that form can be harmful and lead to cancers. |
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How is bleach formed? |
Cl2 + 2NaOH -> NaCl + NaClO + H20 a disproportionate reaction occurs between the chlorine so many different chlorates form as well as sodium chloride and water. The chlorine is reduced and oxidised. |
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How can you test for halide ions? |
Acidifying a solution with dilute nitric acid and silver nitrate. As AgCl, AgBr and AgI are all insoluble in water these acids will detect the halide ions by forming a precipitate. The identity of the precipitate can be confirmed by its solubility in ammonia. AgCl will form a white precipitate; AgBr will form a cream precipitate; AgI will form a yellow precipitate. As these colour changes aren't distingtive the addition of different concentrated ammonias can detect the ion. Dilute ammonia will dissolve AgCl, Concentrated ammonia will dissolve AgBr and AgI doesn't disolve in ammonia at all. |
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Define enthalpy change |
the difference between the enthalpy of the reactants and the enthalpy of the products. |
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What occurs in an exothermic reaction? |
Energy transfer from the reaction mixture to the surroundings. Chemical energy is released by the reactants and the temperature of the surroundings increases. Enthalpy change is negative as reactants have lost energy to the surroundings. |
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what occurs in an endothermic reaction? |
Energy transfer from the surroundings to the reaction mixture. Enthalpy change is positive as reactants have taken heat and energy from the surroundings. |
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Define activation energy |
minimum energy required to initiate a reaction. |
