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Atomis Structure
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- 1Structure of an AtomNucleus: positive protons (1+ mass) and neutral neutrons (1+ mass).
Electrons in shells that have a negative charge and negligible mass (1/2000). - 2Mass Number (big number)The number of protons and neutrons in the nucleus of an atom (it's total mass).
- 3Atomic NumberThe number of protons in a nucleus, and for an atom with no overall charge, it's also the number of electrons.
- 4IonsAtoms formed by the loss or gain of electrons.
- 5IsotopesThe same number of protons but a different number of neutrons. (giving it a different mass but the same relative charge)
- 6Bohr's model (just background). Electrons only exist in fixed orbitals
. Each orbital has a fixed energy
. EM radiation is emitted/absorbed when electrons change shells
. The given frequency of the radiation is fixed due to these shells - 7Relative Atomic MassThe average mass of an element on a scale where carbon-12 is exactly 12.
- 8Relative Isotopic MassThe mass of an atom of an isotope when on a scale where carbon-12 is exactly 12.
- 9Calculating Relative Atomic Mass(isotopic mass x percentages) / 100
- 10Relative Molecular MassThe average mass of a molecule on a scale where carbon-12 is exactly 12.
- 11Relative Formula MassThe average mass of a formula unit on a scale where carbon-12 is exactly 12.
- 12Mass SpectroscopyGives information in the relative molecular and atomic masses and the relative abundance of each isotope.
- 13Electrospray IonisationSample is dissolved in a volatile solvent and pushed through a hypodermic needle with a high voltage applied to it. This causes the sample to gain a proton and become a 1+ ion.
- 14Electron GunElectrons are fired at a sample at high speeds that cause an electron to be knocked of the atoms in the sample and cause them to form 1+ ions. (Can lead to fragmentation)
- 15AccelerationPositive ions are accelerated by an electric field. Need to be positive in order to be accelerated. Gives them all the same kinetic energy but due to the differences in masses, the lighter isotopes will have a greater acceleration.
- 16Ion DriftLeave the electrical field with a constant speed and kinetic energy. They drift through at the same speed they left the electric field, so isotopes with a lower mass will be drifting at a higher speed.
- 17DetectionDetector detects current created by the ions when they hit the plate. It records how long it took for them to pass through the spectrometer.
- 18Interpreting Mass SpecPeaks/lines where new isotope is. Y axis shows abundance, X shows relative mass.
For molecules, the highest peak that is furthest to the right is the relative molecular mass. - 19Electron shellsElectrons can only exist in these shells. The further away from the nucleus the higher the energy it holds.
- 20Sub ShellsShells are divided into sub shells: s (holds 2, 1 orbital), p (holds 6, 3 orbitals), d (holds 10, 5 orbitals) and f (holds 14, 7 orbitals).
- 21Sub-shell NotationEnergy level followed by sub shell letter and the number in that level and then continued.
- 22Arrows in boxesOne spin up, one spin down in each box. Each box represents an orbital.
- 23Energy Level DiagramsEnergy along y axis. Boxes on lines.
- 24Rule 1: Electrons fill up the lowest energy levels first4s has a lower energy level than 3d, so 4s gets filled up before 3d.
- 25Rule 2: They fill up singly before they start sharingThey fill in spin up first and then they fill in the spin down (electron pair repulsion).
- 26Rule 3: Ions you just add or take away the outer electronRemove from the highest energy level occupied (if positive) or add to the highest occupied (if negative).
- 27Chromium and CopperCr: 3d5, 4s1.... they are more stable with a half full d sub shell than with only 4 in it.
Cu: 3d10, 4s1.... they are more stable with a full d sub shell than a nearly full sub shell.
(lose the 4s first when ionised) - 28S block elementsThey easily lose electrons to form inert gas configurations.
- 29P block elementsThey easily gain electrons to form negative ions with inert gas configurations.
- 30Group 0 elementsTheir P block is completely full so they don't need to lose or gain any electrons, this is what makes them inert.
- 31First Ionisation EnergyThe energy needed to remove 1 electron from each atom in mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
(g symbol must be present, lower it is, the easier to form a positive ion). - 32Nuclear chargeThe more protons there are, the more positively charged so the stronger the attraction to the electrons so the higher the ionisation energies.
- 33Distance from NucleusAttraction falls rapidly with distance. So the further away the easier for an electron to be lost so thus reducing the ionisation energy.
- 34ShieldingAs the number of electrons between the nucleus and the outer electron increases, the outer electrons feel less attraction to the nucleus. This lessens the pull of the nucleus causing the reduction in the ionisation energy.
- 35Second Ionisation EnergyEnergy required to remove 1 electron from each ion in 1 mole of gaseous plus ions to form 1 mole of gaseous 2+ ions.
Higher than the first because the electron is being removed from a positive ion, so the attraction will be larger. - 36Ionisation trends down Group 2Decreases: One more electron shell than the one above, so more shielding from the attraction of the nucleus. Also increases the atomic radius, reducing the attraction further meaning electrons can be more easily removed, resulting in the reduction in ionisation energies.
- 37Ionisation trends across a PeriodGeneral increase due to the increasing number of protons.
Falls at Al: breaks into another sub shell (p) so slightly more shielding and distance.
Falls at S: electron pair repulsion. - 38Ionisation energies and Shell StructureBig jumps when a new shell is started.
- 39The Mole6.02x10^23 particles. Number of particles = the number of mole x Avagadro's constant
- 40Calculations with MolesMass / Mr
- 41Moles and ConcentrationNumber of moles = (concentration x volume (in cm3)) / 1000
Number of moles = concentration x volume (in dm3) - 42The ideal gas equationpV = nRT
Volume in m3
Pressure in Pa
Temperature in K - 43Conversions1m3 = 1x10^6 cm3 = 1x10^3 dm3
Kelvin = *C 273
kPa = 1000Pa - 44Calculating Masses1. balanced equation
2. how many moles of reactant
3. use the molar ratio to work out the number of mole of product
4. calculate the mass of that many moles - 45State Symbolss = solid
aq = aqueous (dissolved in water)
l = liquid
g = gas - 46Neutralisation ReactionsAcid and alkali react giving a salt and water. This is used in titration.
- 47Making a Standard Solution1. Work our how many moles you need using concentration x volume.
2. Work out grams needed (mass/mr).
3. Weigh it out.
4. Add distilled water to the beaker and stir until all the solid is dissolved.
5. Funnel into a volumetric flask.
6. Rinse funnel and stirrer into the flask and fill it up with more water until the meniscus meets the line.
7. Stopper it and mix. - 48Making a Standard Solution (making it more precise)Weigh the weighing boat before and after to measure exactly how much solid was added.
Funnel the solution to make sure none is spilt. Then rinse the funnel to remove any remaining solution from the edges. Rinse the stirrer as well.
Calculate real concentration. - 49Titrations1. Standard solution in burette. Unknown concentration in the flask (measured into it using a pipette).
2. Allow standard to drip through until other side of the tap is filled.
3. Add phenolphthalein. Add standard until the unknown changes colour. Record the titre (volume added).
4 Repeat it until two concordant results are found. - 50Titration CalculationsFind the moles of the standard solution. Use molar ratio to find the moles of the unknown solution. Use the volume and moles of the unknown to find the concentration.
- 51Molecular FormulaThe actual numbers of atoms in a molecule.
- 52Empirical FormulaGives the smallest whole number ratio of the atoms in a compound.
- 53Finding the Molecular Formula from the Empirical FormulaFind the empirical mass and divide the molecular mass by the empirical mass. This tells you how many multiples of the empirical formula are in the molecular formula. Multiply the empirical formula by that number to find the molecular formula.
- 54Calculating Theoretical Yield1. Work out how many moles of reactant you have
2. Use the equation to work out how many moles of product you would expect (ratios)
3. Calculate the mass of that many moles of product - 55Calculating Percentage YieldActual Yield / Theoretical Yield
x100 - 56Atom EconomyThe measure of the proportion of reactant atoms that became part of the desired product in the balanced chemical equation.
- 57Advantages of High Atom EconomyMore efficient use of raw materials. Less money needs to be spent on disposing of waste. More sustainable. Processes would cost less money so the end products could have a reduced price.
- 58Atom economy (%)Molecular Mass of Desired Product / Sum of Molecular Masses of all Reactants
x 100 - 59IonsFormed when electrons are transferred from one atom to another, in order to gain a full electron shell.
- 60Ionic CompoundsElectrostatic attraction holds positive and negative ions together- it's very strong. Held together in a lattice.
- 61Giant Ionic LatticesIonic crystals are giant lattices of ions. Regular structure, giant because it's made of he same repeated basic unit over and over again.
- 62Electrical ConductivityConduct when molten or dissolved (when their electrons are free to move). In a solid they're fixed in position by the strong ionic bonds.
- 63Melting pointThey have high melting points because of the strong electrostatic forces that hold the ions together. It takes a lot of energy to overcome these forces and therefore they have a high melting and boiling point.
- 64SolubilityThey tend to dissolve in water because of the polar nature of the compounds. The water molecules pull the ions away from the lattice and cause it to dissolve.
- 65MoleculesThey form when two or more atoms bond together. The atoms can be the same or different.
- 66Single BondsTwo atoms share electrons so they've both got full outer shells of electrons. Both the positive nuclei are attracted electrostatically to the shared electrons.
- 67Double and Triple bondsMultiple pairs of shared electrons.
- 68Simple Covalent CompoundsMade up of lots of individual molecules. The atoms in the molecules are held together by covalent bonds, but the molecules are held together by much weaker intermolecular forces.
- 69PropertiesMelting: low due to the weak intermolecular forces that hold the molecules together.
Conductivity: don't conduct (no free electrons). - 70Graphite (three carbons bonded together).Weak bonds between layers so they can slide (lubricant)
.Delocalised electrons so electric current can flow
.Low density due to the layers being spread out
.High melting point due to strong covalent bonds between the carbons
.Insoluble, covalent bonds are too difficult to break - 71Diamond (carbon bonded 4 times).High melting point
.Hard
.Good thermal conductor, but can't conduct electricity
.Won't dissolve - 72Co-ordinate bondsOne atom donates both of the electrons in a covalent bond.
- 73Charge CloudsBonding pairs and lone pairs exist as charge clouds. This is an area where there's a large chance of finding an electron. Lone pairs are unshared electrons. Bonding pairs are shared electrons.
- 74Electron pair repulsionElectrons are all negatively charged so charge clouds repel each other until they are as far apart as possible. Lone pair charge clouds repel more than bonding pair charge clouds so bonds angles are reduced because bonding pairs are pushed together by lone pair repulsion. Lone/lone are the biggest, lone/bonding are second biggest and bonding/bonding are the smallest.
- 75Central with two electron pairsLinear shape with a bond angle of 180
- 76Central atom with three electron pairsTrigonal planar with a bond angle of 120
- 77Central atom with four bonding pairsTetrahedral with a bond angle of 109.5
- 78Central atom with 3 bonding pairs and 1 lone pairTrigonal pyramidal with a bond angle of 107
- 79Central with two bonding pairs and wo lone pairsBent with a bonding angle of 104.5
- 80Central atom with 5 bonding pairsTrigonal bipyramidal with bonding angles of 120 and 90
- 81Central atom with 4 bonding pairs and 1 lone pairSeesaw shape with bond angles of 102 and 86.5
- 82Central atom with 3 bonding pairs and 2 lone pairsT-shaped with bond angles of 87.5
- 83Central atom with 6 bonding pairsOctahedral with bond angle of 90
- 84Central atom with 5 bonding pairs and 1 lone pairSquare pyramidal with bonding angles of 90
- 85Central atom with 4 bonding pairs and 2 lone pairsSquare planar with bonding angles of 90
- 86Awkward moleculesTreat multiples bonds like they are a normal single bond when working out the shape.
- 87ElectronegativityThe ability to attract the bonding electrons in a covalent bond. (Fluorine is the most electronegative).
- 88Polar bondsWhen the atoms in a covalent bond have different electronegativity's so one atoms pulls the bonding pair closer to the atom. The difference in electronegativity's causes a dipole.
- 89Van der WaalsCauses all atoms to become attracted to each other. Electrons move around in the charge clouds and at any moment electrons could all be at one side. This causes a temporary dipole. This repels electrons in the atom next to it causing a dipole in that molecule too. This causes all the positive ends and negative ends to attract.
- 90Van der Waal propertiesLarger mass means larger Van der Waals. Closer the two molecules are the stronger the forces (i.e. unbranched). Molecules with stronger v.d.w's have a higher boiling point.
- 91Permanent Dipole DipoleMolecules with permanent dipoles will have electrostatic attractions between the delta + and delta - charges.
- 92Permanent Dipole PropertiesThe more polar the higher the electrostatic attraction. Higher boiling and melting points. Stronger than V.d.W.s.
- 93Hydrogen BondingStrongest intermolecular force. Only happens when hydrogen and NOF are bonded. They are very electronegative so they draw the bonding pair away from the hydrogen atom.
- 94Simple covalent properties (strong covalent bonds within molecules atoms but weak intermolecular forces between molecules)Electrical conductivity: don't conduct because there are no free ions
Melting point: low melting because intermolecular forces are easily overcome
Solubility: some can depending on polarity - 95Metallic BondingGiant metallic lattices. The outermost shell of electrons are delocalised. The positive metal ions are attracted to the delocalised negative electrons. They form extermeley closely packed lattices.
- 96Melting PointHigh melting point because of strong electrostatic attraction. The more delocalised electrons per atom, the higher the melting point.
- 97Ability to be shapedCan easily slide due to there being no specific bonds holding atoms together. so metal are malleable (shaped) and ductile (drawn into wire easily).
- 98ConductivityDelocalised electrons that can carry energy (thermal conductors) and charge around the metal and therefore carry current to conduct electricity.
- 99SolubilityMetals are insoluble, expect as liquid metals, because of the strength of the metallic bonds.
- 100Changing statesPacked closely together as solids. When energy is provided the atoms move around and bonds break causing them to melt and become liquid. Then more energy is supplied and all bonds break turning it into a gas state.
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- AS - Level AS Chemistry Fichas sobre Atomis Structure, creado por beckycurtis el 16/05/2016.
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