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Chemistry A-level > Dr Mannock Yr 12 > Flashcards

Dr Mannock Yr 12 Flashcards

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1
Q

What is the molar mass equation?

A

molar mass (g/mol)=mass (g)/moles (mol)

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2
Q

What is the relative atomic mass of an element?

A

The mean mass of an atom compared to the mass of carbon-12 which is 12

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3
Q

How do you calculate empirical formula?

A

Find the molar ratio between the two reactants, which will be the amount of that substance in the empirical formula (if percentages do 100g of the products)

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4
Q

What is the concentration equation?

A

concentration (mol/dm^3 or g/dm^3) = mass (g) or moles (mol) / volume (dm^3)

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5
Q

How do you calculate the mass of a product/reactant with all reactants total weight?

A
  1. You find the Mr ratio
  2. You add them together to find the total Mr of the reactants
  3. You multiply the total reactants/products weight by the wanted reactant Mr
  4. You then divide this number by the total Mr of reactants/products and find out the weight of the wanted reactant or product
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6
Q

What is the equation for Dilution factor?

A

Dilution factor=Stock concentration or Concentration factor/Diluted concentration

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7
Q

What is the equation for the Original stock volume?

A

Original stock volume=Total volume needed/ Dilution factor

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8
Q

What is the ideal gas equation?

A

pV=nRT
p = pressure in pascals
V = Volume in m3
n = number of moles
R = 8.31 the gas constant
T = temperature measured in kelvin

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9
Q

How do you convert from *C to Kelvin?

A

Add 273

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10
Q

What is one atm in Pa?

A

101,325 Pascals of pressure

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11
Q

What is the Percentage yield equation?

A

% yield = actual yield/theoretical yield x 100

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12
Q

What is the Atom economy equation?

A

atom economy = sum of molar masses of desired products / sum of molar masses of all products x 100

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13
Q

What is atom economy?

A

A measure of the amount of starting materials that end up as useful products

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14
Q

Why is your theoretical yield not always the same as your actual yield?

A

Because some reactants may not react, some product may be left in the reaction container and some product may be lost during any purification processes used.

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15
Q

What are bases?

A

Bases are proton (H+ ions) acceptors as they donate 2e- to form a dative bond with the protons. They can have either a lone pair or negative charge, e.g. all metal oxides.

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16
Q

What are acids?

A

Proton donors, when mixed with water all acids release H+ ions (in aqueous solutions).

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17
Q

What is a salt?

A

Compound formed when a hydrogen ion in an acid is replaced by a metal ion.

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18
Q

What are the general formulae of reactions with acids?

A

Metal + Hydrochloric acid -. Metal chloride+H
Metal + Sulphuric acid -. Metal sulphate + H Metal + Nitric acid - Metal nitrate + H
Metal +Phosphoric acid->Metal phosphate+H Metal oxide + Acid -> Metal salt + H2O Metal hydroxide + Acid -> Metal salt + H2O Metal carbonate+Acid->Metal salt+CO2+H2O NH3 + Acid -> NH4 (ammonium salt)
Metal hydroxide + organic acid, e.g. ethanoic acid -> Organic salt + H2O

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19
Q

What is an ionic equation?

A

It shows the ions that change in a reaction, ignoring the spectator ions. The charge must be equal on both sides of the equation.

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20
Q

What is titration?

A

The process of adding a known amount of solution of known concentration to determine the concentration of another solution.

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21
Q

What is the process of doing titration calculations?

A

1.Identify the known amounts/substances
2.Calculate the number of moles of the known solution-remember to convert between cm^3 to dm^3
3.Using the chemical equation to work out the molar ratio and therefore the moles of the other reactant
4.Using c=n/v to work out the unknown concentration

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22
Q

What are examples of polyatomic ions and their charges?

A

Ammonium/NH4 = charge of +, Oxide/O = charge of 2-, Sulphide/S = charge of 2-, Nitride/N = charge of 3-, Hydroxide/OH = charge of -, Nitrate/NO3 = charge of -, Nitrite/NO2 = charge of -, Hydrogencarbonate/HCO3 = charge of -, Permanganate or Manganate (VII)/MnO4 = charge of -, Carbonate/CO3 = charge of 2-, Sulphate/SO4 = charge of 2-, Sulphite/SO3 = charge of 2-, Dichromate (VI)/Cr2O7 = charge of 2-, Phosphate/PO4 or PO3 = charge of 3-

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23
Q

What is accuracy?

A

A measure of how close a measured value is to the true value.

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24
Q

What is precision?

Study These Flashcards
A

A measure of how close a series of measurements are to one another.

25
How do you improve the accuracy and precision of an experiment?
Accuracy: -Repeat the experiment -Calculate a mean from the results Precision: -Use finer measures -Use class B rather than class A
26
How do you show the uncertainty of a measurement?
Use l to equal the uncertainty and display it as l=measured value ± uncertainty, e.g., if the value was 8.25 and the containers uncertainty ± 0.05 it could be written as l=8.25 ± 0.05. To add uncertainty's add both so that 0.05+0.05= an uncertainty of ± 0.1
27
What is the percentage error equation?
% error=Maximum error/quantity measured x100. When finding the overall % error of an experiment find it for each value before adding them together.
28
What is the titre value?
The final reading -initial reading to give you the overall amount measured in the acid titration.
29
What are oxidation numbers?
Oxidation numbers are a measure of the amount of valence e^- that an atom uses to form a bond with an atom of a different element. They can determine molecular formulae and balance redox equations.
30
What are the 10 rules about oxidation numbers?
1. The oxidation number of a neutral compound or 1 atom is 0 2. The oxidation number of a monoatomic ion is the same as its charge 3. The sum of all oxidation numbers in a neutral compound is 0, eg, HCl, H=1+, Cl=1- 4. The sum of all oxidation numbers in a polyatomic must equal its charge, eg, SO4=2- 5. Oxidation numbers usually correspond with periodic group Group 1 = +1, Group 2 = +2, Group 3 = +3, etc 6. Oxidation number of H usually +1, unless in a metal hydride, eg, NaH/Sodium hydride 7. F always has an oxidation number of +1, due to its electronegativity 8. Cl, Br + I usually have oxidation numbers of -1, except when bonded to O, when it changes to +1 due to O's electronegativity 9. O will have oxidation number of -2, unless it forms a peroxide (H2 O2 ) when it is -1, or bonded to F when it is +2, due to F's electronegativity 10. In transition metals it can vary so work them out last
31
What are redox reactions?
Reactions that involve the transfer of e^-, resulting in both the oxidation and reduction of atoms.
32
What happens when an element is oxidised or reduced?
Oxidation: -The element loses electrons -This causes an increase in the elements oxidation number Reduction: -The element gains electrons -This causes a decrease in the elements oxidation number (Note: You may be asked to identify what element has been reduced/oxidised. To do this work out the oxidation numbers of elements before and after the reaction and see which has increased and/or decreased)
33
What does isoelectronic mean?
Isoelectronic refers to an ionic compound with the same charge or electronic structure.
34
What are disproportionation reactions?
Reactions in which an element or compound is both reduced and oxidised.
35
How do water molecules dissolve substances?
The water molecules are attracted to ionic compound due to their partial charges and cluster around the individual ion, with the O attracted to + charge, and the H to -. They then break apart the lattice/crystalline structure.
36
How does the melting and boiling point of ionic compounds change?
As the charges increase, eg, 2+ and 2- not - and +, the electrostatic attraction is stronger so the melting and boiling points increase.
37
How does atomic radii change across a period and down a group?
Across period: -Atomic radii decreases across a period due to the increase in the number of protons in the nucleus -This increases the nuclear charge, and so the electrostatic forces of attraction between the valence e^- and the nucleus, causing the radius of the atom to decrease Down a group: -Atomic radii increases down a group as the number of shells increases This means He will be the smallest radius, and Fr the largest radius
38
What is the ionisation energy of an element?
It is the energy required to remove one electron from the ground/normal state of each atom, in a mole of gaseous atoms of an element, to form a positive ion.
39
How does ionisation energy/enthalpy change across a period and down a group?
Across a period: -It increases across a period due to an increase in nuclear charge -This means that the electrostatic force of attraction is stronger and so more energy is required to release an electron Down a group: -It decreases down a group due to an increase in atomic radius, and an increase in electron shielding -This means less energy is required to release one electron
40
What does a jump in ionisation enthalpy indicate?
That the atom had reached a noble gas configuration, so reaches a more stable configuration, so more energy is required to break this electron configuration.
41
What does periodicity refer to?
The trends in the physical and chemical properties across a period and down a group, relating to atomic structure.
42
How does melting point change along period 3?
-It originally increases in metal elements due to an increase in nuclear charge, so metallic bonding will be stronger -Then it changes depending on the size of the simple covalent molecules, so that a larger molecule, and the number of e^- that the molecule contains will have stronger London forces between molecules than smaller covalent structures -This mean silicon has the highest melting point as it is a giant covalent structure, and sulphurs melting point is higher than phosphorus's, as it forms a larger simple molecule and has more e^- than P -Chlorine only forms a diatomic molecule, so therefore has the lowest boiling point of all elements in period 3 (of those that are not noble gases)
43
What is the general trend of ionisation energy across period 2, and what are the exceptions to this trend?
-The general trend across period 2 is an increase in ionisation energy due to the decrease in atomic radius and increased electrostatic forces of attraction -The exceptions are, which applies to elements in the same position above and below are boron/B and oxygen/O Boron: -Slight decrease between Be and B -This is because the last e^- in Be is in a 2s subshell, whereas in B it is a 2p subshell, which is further from the nucleus Oxygen: -Slight decrease between N and O -This is because the paired e^- in the 2p subshell of O repel, making it easier to remove an e^-
44
What is an enthalpy change?
The energy transferred to and from the environment and system, in standard conditions, is measured as an enthalpy change
45
What are standard conditions?
-1 atmosphere (atm) of pressure -Specified temperature in Kelvin, usually 298K (25*C) -Standard concentration, 1 mol dm^- 3 , for solutions -All elements must be in their RTP states of matter
46
What is the standard enthalpy change of combustion? (ΔH°c)
-The energy transferred during a combustion reaction -To calculate it, change in temperature and volume of H2O must known -This is because it found out by using a spirit burner to heat up a volume of H2O -Draught shields are used to conserve heat exchange with the surroundings, and the moles of fuel must be 1
47
How do you calculate ΔH°c?
q=cmΔT and ΔH°=q/n q = heat exchanged with surroundings c = specific heat capacity of water (given) m = mass/volume of water (cm^3) ΔT = temperature change of the surroundings/H2O (K) n = moles of limiting reactant/H2O for ΔH°c
48
What is standard enthalpy change of reaction, and how do you calculate it? (ΔH°r)
-The enthalpy change when the molar quantities of reactants as stated in the equation react together under standard conditions -It is calculated using the same equation as ΔH°c
49
What is standard enthalpy change of neutralisation, and how do you calculate it? (ΔH°neut)
-Enthalpy change when 1 mole of H^+ ions react with 1 mole of OH^- ions to form 1 mole of H2O under standard conditions and in solutions containing 1mol dm^-3 -Calculated the same way as ΔH°c
50
What is ΔH f, and how is it calculated?
-The enthalpy change of a reaction during which 1 mole of products is formed from its constituent elements -Calculated in the same way as ΔHº
51
How do group 2 elements react with H2O?
If the H2O is cold or RTP: G2 element, eg, 2Mg + 2H2O -> 2Mg(OH)2 + H2 If H2O vapour: Mg + H2O -> MgO + H2
52
How is solubility determined, and how does the group 2 hydroxide and sulphate solubility change down the group?
It is determined in 1 of 2 ways: -Lattice disassociation enthalpy (energy the needed to break a crystal lattice apart) -Hydration enthalpy (energy released when the ions are hydrated, and H2O molecules form a shell) Group 2 hydroxides: -Down the group, lattice disassociation enthalpy decreases -Outweighs the change in hydration enthalpy, so therefore the solubility increases down the group Group 2 sulphates: -The hydration enthalpy decreases much quicker than the lattice disassociation enthalpy down the group, making the sulphates more insoluble down the group
53
What is the medical use of Ba^2+ ions, and why are they safe to use as BaSO4 tablets?
-Ba^2+ ions can be used before an x-ray to provide contrast -It is safe as there are little to no free toxic Ba^2+ ions to cause harm
54
What is the average bond enthalpy?
The average energy required to break 1 mole of a bond to separate all atoms in a gaseous state.
55
What is Hess' law, and when is it used?
Hess' law is if a reaction can take place by more than 1 route, and the initial and final conditions are the same, the total enthalpy change will be same. -Eg, reaction -> products or reactants -> intermediate -> products -Only used if enthalpy change cannot be calculated due to: -High activation energy -Slow ror (rate of reaction) -More than 1 reaction occurs at once
56
What are the equations when working out either ΔH°c or ΔH°r from ΔH°f values?
ΔH°c= ΔH°f reactants - ΔH°f products ΔH°r= ΔH°f products - ΔH°f reactants -When calculating these questions, always write out a ΔH°f Hess cycle to make sure you have it right
57
What is collision theory and what does it state must occur between particles for a successful reaction to occur?
-Collision theory explains conditions for reactants to react to form bonds It states particles must: -Collide -Have sufficient energy -Collide with the correct orientation Most collisions are not successful because they do not meet conditions 2 or 3
58
What factors affect rate of reaction?
Temperature: -Particles have a higher kinetic energy -This increases the proportion of successful collisions, as more collisions occur with sufficient energy, and the frequency of successful collisions, as particles move faster and so collide more frequently Concentration: -More particles occupy the same space, and so the frequency of successful collisions increases, as there are more particles available to collide Surface area (for solids): -Same as conc, as more particles exposed to collide, and so the frequency increases Pressure (for gases): -Same as conc, as more gas particles are available to collide Catalysts: -Reduces the activation energy (Ea), and so increasing the proportion of successful collisions as the minimum energy requirement is lower
59
What is the equation for rate of reaction?
rate of reaction=change in conc (moldm^-3) /time (s)

Chemistry A-level (2 decks)

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