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49 Cards in this Set
- Front
- Back
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Describe the structure of an atom.
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PROTONS - mass: 1 charge: +1
NEUTRONS - mass: 1 charge: 0 ELECTRONS - mss: 1/2000 charge: -1 MASS NUMBER - (top number) total number of protons and neutrons ATOMIC NUMBER - (bottom number) number of protons (+ electrons) |
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What is an Isotope?
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ISOTOPES are atoms with the same number of protons but different numbers of neutrons.
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What is the definition of Relative Atomic Mass?
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Is the average mass of an atom of an element on a scale where an atom of carbon-12 is 12.
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What is the definition of Relative Isotopic Mass?
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Is the mass of an atom of an isotope of an element on a scale where an atom of carbon-12 is 12.
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What is the definition of the Relative Molecular Mass, Mr?
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Is the average mass of a molecule or formula unit on a scale where an atom of carbon-12 is 12.
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How do you calculate the Isotopic Abundance?
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Multiply relative isotopic mass by it % of abundance
Add up the results Divide by 100 |
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What is the Avogadro Constant, Ma and how do you calculate the number of moles from the number of atoms or molecules?
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6.02 x 10^23
No. of moles = No of particles you have/ No of particles on 1 mole |
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How do you calculate the no. of moles?
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General:
No of moles = mass of substance / molar mass (molar mass = Mr) Moles of solution: No of moles = Conc' x Volume (in cm3) / 1000 No of moles = Conc' x volume (in dm) Moles of gases: No of moles = volume (dm3) / 24 No of moles = volume (cm3) / 24 000 |
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How do you calculate an Empirical and Molecular Formula?
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Empirical - smallest whole number ratio of atoms in a compound
Molecular - actual whole number of atoms in a molecule find number of moles of each element divide all by smallest number of moles Round off Divide molecular mass by empirical mass Multiply each ratio by the answer to provide molecular formula |
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What are Acids, Bases and Salts?
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Acids - produce H+ (aq) ions - i.e PROTON DONORS
Bases - remove H+ (aq) ions - i.e PROTON ACCEPTORS Salts - a LATTICE of posative and negative ions |
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Acid (aq) + Base (s) = ....
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Salt (aq) + Water (l)
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Carbonate (s) + Acid (aq) = ....
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Salt (aq) + Carbon Dioxide (g) + Water (l)
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How to calulate the mass of the salt when it is HYDRATED?
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work or mass of water lost
calculate no. of moles of water lost Calculate moles of anhydrous salt work out ratio of salt : water Divide my smallest mole and round |
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What is Oxidation and Reduction?
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Oxidation - loss of electrons, reducing agent
Reduction - gain of electrons, oxidising agent |
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Rules of Oxidation Numbers....
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uncombined elements = 0 (e.g Na, Fe)
elements bonded to identical atoms = 0 (e.g. H2, O2) sum of oxidation numbers for a neutral compound = 0 combined oxygen = -2 combined hydrogen = +1 increase in oxidation numbers = more electrons lost decrease in oxidation number = more electrons gained metals usually donate electrons and so normally posative non-metals usually gain electrons and so normally negative |
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What are the four different shub-shells and what are their maximum electrons and also what is are the numbers given to each shell know as?
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S = 2
P = 6 D = 10 F = 14 PRINCIPAL QUANTUM NUMBERS |
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Give some characteristics of Orbitals....
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An orbital is the bit of SPACE that the electrons move around in.
Electrons in the orbitals spin in the opposite directions - know as SPIN-PAIR. S orbitals are spherical shaped and P orbitals are dumbbell shaped. ( there arew 3 P orbitals and they are at right angles to each other). |
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State the Electron Configuration....
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1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p.....
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What is the definition for First Ionisation Energy?
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Is th energy needed to remove 1 electron from EACH ATOM in 1 MOLE of GASEOUS atoms to form 1 mole of gaseous 1+ ions.
e.g O(g) ----> O+(g) + e- |
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What facts affect ionisation energy and how?
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NUCLEAR CHARGE: the more protons there are - the more positively charged the nucleus is - the stronger the attraction for the electrons
DISTANCE FROM NUCLEUS: closer to the nucleus- the more strongly the electrons are attacted then the ones further away. SHEILDING - the number of electrons between the nucleus and the outer shell - the less attraction there is towards the nuclear charge. |
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Whatis the definition for Second Ionisation Energy?
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Is the energy needed to removed 1 electron from EACH ATOM in 1 MOLE of GASEOUS 1+ ions to form 1 mole gaseous 2+ ions.
e.g O+(g) ----> O2+(g) + e- |
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What is Ionic Bonding?
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Ionic Bonding - It is the ELECTROSTATIC attraction between two OPPOSITELY charged ions.
Group 1 - 1+ charge Group 2 - 2+ charge Group 6 - 2- charge Group 7 - 1- charge |
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Explain the behaviour of Ionic Compounds......
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CONDUCT ELECTRICITY when dissolved or molten but not when solid - as in liquid are free to move carrying the charge.
HIGH MELTING POINT - giant lattices are held together by strong ELECTROSTATIC forces, alot of energy is required to break the forces. DISSOLVE in water - water is polar and parts of the lattice are negative and posative ( very slightly) and so the water molecules pull ions away causing it to dissolve. |
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What is a Covalent Bond?
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Is a shared pair of electrons.
Does it to become more stable. |
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What is Dative Covalent Bonding?
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It is where BOTH ELECTRONS come from ONE ATOM.
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Diamond.....
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Is a giant covalent lattice.
Very high melting point Extremly hard Cant conduct electicity - as outer bonds held in localised bonds Good thermal conductor - vibrations easily carried through Doesnt dissolve in any solvent. Each carbon bond bonded to 4 other carbons. Have a tetrahedrial shape. |
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Graphite.....
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Is a giant covalent lattice.
Carbon atoms are aranged in sheets of flat hexagons covalently bonded - have 3 bonds each. The 4th outer electron sheel is delocalised between teh sheets. the sheets are bonded together via weak VAN DER WAALS forces. Delocalised electrons are free to move and so transports electricity. Strong and light weight as layers are quiet far apart and so is dense. Strong covalent bonds means high melting point Insoluble as covalent bonds are hard to break. |
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Metals.....
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Giant metallic structures.
Electrons in the out shell are delocalised - electrons move freely around the posativly charged ions. Posative metal is attracted to the negative electrons and so stay packed together in a sea of delocalised electrons. More the delo calised electrons - stronger the bonding is - higher the melting point. As no bonds holding the metal in place it is DUCTILE and MALLEABLE. Pass kinetic energy through - good thermal conductor good at electical conduct - due to delocalised elecetrons. Insoluble - (except in liquid metals) because of the strength on the metalic bonds. |
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How do Lone pair and Bond pair angles differ?
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lone pair/lone pair = biggest bond angle
lone pair/bond pair = second biggest bond angle bond pair/bond pair = smallest bond angle |
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Linear Molecules .....
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..... 180 degrees (2 electrons on central atoms)
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Trigonal Planar.....
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.....120 degrees (3 electrons on central atoms)
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Trigonal Pyramidal....
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....107 degrees (4 electrons on central atom - 1 of which is a lone pair)
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Tetrahedral.....
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.....109 degrees (4 electrons on central atom)
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What is Electronegativity?
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Ability to attact bonding electrons in a covalent bond.
Bondng electrons are pulled more towards the more electronegative atom - making it polar. Polar bonds form DIPOLES between the atoms. Greater the difference in electronegativity the more polar the bond. |
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What are Intermolecular Forces?
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The charges on the polar molecules cause there to be a weak electrostatic force of attraction between the molecules.
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What are Van Der Waals Forces?
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They cause ALL atoms to be attracted to each other.
Electrons in a charge cloud move ver quickly - at a particular moment there is likely to be more electrons on one side than on the other of the atom - forming a TEMPORARY DIPOLE. This dipole causes there to be another dipole in the OPPOSITE direction to be formed on the neighbouring atom - therefore the two dipoles attract to each other. Dipoles are constantly being formed and destroyed but the over all effect is that all the atoms are attracted to one another. The stronger the van der waals forces - hight boiling points - greater the surface area the greater the van der waals forces as they are expose a greater electron cloud. As you go down the group - intermolecular forces increases - as atomic/molecular size increases and also the number of shells of electrons increases. |
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What happens during Melting and Boiling?
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covalent bonds dont need to be broken - only the van der waals forres and the hydrogen bonds need to be overcome.
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What are the Periodic Trends?
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Atomic Radius DECREASES across the period - as the number of proton increases so does the posative charge and so electrons are pulled closer to the nucleus.
Ionisation energy INCREASES across a period and DECREASES down a gorup - ( affect by atomic radius/nuclear charge/electron sheilding) Melting ang boiling points are linked to teh bond strength adn structure - metals have a higher m/b points and the simple structure (02, F2etc) depends on van der waals forces. |
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Group 2 Elements.....
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Reactivity INCREASES down the group - atomic radius increases and so does the sheilding affect.
React with both water adn oxygen - when react with water they produce HYDROXIDES, when react with oxygen they form OXIDES. The hydroxides and oxides are bases and form alkaline solutions in water - oxides are much stronger. Thermal stability INCREASES down the group - (break down by heat) Are used to neutralise acidity. |
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Group 7 - Halogens....
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Highly reactive non-metals
melting and boiling point INCREASE as you go down the group - due to increased strength of the van der waals forces Volatility DECREASES down the group. Reactivity DECREASES down the group - as sheilding increases its harder for larger atoms to attrect to the electrons needed for the reaction to occur as a greater posative charge. They displace less reacive halide ions from solution - Cl displaces with Br and I - Br displaces with I - I displaces with nothing, SILVER NITRATE - test for halides - add nitric acid to get rid of any ions that might interfer with the test - then add siver nitrate solution (AgNO3) - look at precipitate formed - Cl- = white Br- = cream I- = yellow |
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What is Disproportionation?
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Halogens undergo disproportionation with alkalis - where they become reduced and oxidiced
e.g. X2 + 2NaOH ----> NaXO + NaX + H2O |
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How is bleach made?
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mixture of chlorine gas with sodium hydroxide ( at RTP) to form SODIUM CHLORATE (l)
2NaOH + Cl2 -----> NaClO + NaCl + H2O chlorine is used to kill bacteria |
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What are Water Treatments and what are there implications?
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Chlorine is important psrt of water treatmnts - kills dideases - chlorine stays in the ater which means that it will prvent infection further down the supply - prevents groth of algea and eliminates bad taste/smell and removes decolouration caused by the organic compounds.
Risks = chlorine gas is very harmful - liquid chlorine on skin or eyes can cauase severe chemical burns - chlorine reacts with the organic compound in the water which forms cancer causing molecules ( rick is small compared to drinking untreated water). |
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What is Disproportionation?
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Halogens undergo disproportionation with alkalis - where they become reduced and oxidiced
e.g. X2 + 2NaOH ----> NaXO + NaX + H2O |
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How is bleach made?
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mixture of chlorine gas with sodium hydroxide ( at RTP) to form SODIUM CHLORATE (l)
2NaOH + Cl2 -----> NaClO + NaCl + H2O chlorine is used to kill bacteria |
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What are Water Treatments and what are there implications?
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Chlorine is important psrt of water treatmnts - kills dideases - chlorine stays in the ater which means that it will prvent infection further down the supply - prevents groth of algea and eliminates bad taste/smell and removes decolouration caused by the organic compounds.
Risks = chlorine gas is very harmful - liquid chlorine on skin or eyes can cauase severe chemical burns - chlorine reacts with the organic compound in the water which forms cancer causing molecules ( rick is small compared to drinking untreated water). |
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What is Disproportionation?
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Halogens undergo disproportionation with alkalis - where they become reduced and oxidiced
e.g. X2 + 2NaOH ----> NaXO + NaX + H2O |
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How is bleach made?
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mixture of chlorine gas with sodium hydroxide ( at RTP) to form SODIUM CHLORATE (l)
2NaOH + Cl2 -----> NaClO + NaCl + H2O chlorine is used to kill bacteria |
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What are Water Treatments and what are there implications?
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Chlorine is important psrt of water treatmnts - kills dideases - chlorine stays in the ater which means that it will prvent infection further down the supply - prevents groth of algea and eliminates bad taste/smell and removes decolouration caused by the organic compounds.
Risks = chlorine gas is very harmful - liquid chlorine on skin or eyes can cauase severe chemical burns - chlorine reacts with the organic compound in the water which forms cancer causing molecules ( rick is small compared to drinking untreated water). |
