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chemistry > module 3: periodic table and energy > Flashcards

module 3: periodic table and energy Flashcards

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1
Q

what is periodicity?

A

a repeating trend in properties of the elements across each period of the periodic table

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2
Q

what is the first ionisation energy?

A

the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions

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3
Q

what are the factors affection ionisation energy?

A
  • atomic radius
  • nuclear charge
  • electron shielding
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4
Q

explain how atomic radius affects ionisation energy

A
  • the greater the distance between the nucleus and the outer electrons, the less the nuclear attraction
  • the larger the radius, the smaller the amount of energy required to remove the electron from the outermost orbital, so first ionisation energy decreases
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5
Q

explain how nuclear charge affects ionisation energy

A
  • the more protons there are in a nucleus of an atom, the greater the attraction between the nucleus and the outer electrons
  • therefore, as the nuclear charge increases, so does the ionisation energy
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6
Q

explain how electron shielding affects ionisation energy

A
  • the greater the shielding effect, the weaker the attraction between the nucleus and the outer electrons
  • the higher the shielding effect, the lower the ionisation energy
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7
Q

what is the second ionisation energy?

A

the energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

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8
Q

explain the trend in first ionisation energy down a group

A

first ionisation energies decrease down a group, because:
- atomic radius increases
- more inner shells so shielding increases
- nuclear attraction on outer electrons decreases
- first ionisation energy decreases

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9
Q

explain the trend in first ionisation energy across a period

A

first ionisation energies increase across a period, because:
- nuclear charge increases
- same shell: similar shielding
- nuclear attraction increases
- atomic radius decreases
- first ionisation energy increases

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10
Q

explain the fall in ionisation energy between beryllium and boron

A
  • the fall in the first ionisation energy between beryllium and boron marks the start of filling the 2p sub-shell
  • the 2p sub-shell in boron has a higher energy than the 2s sub-shell in beryllium
  • therefore, in boron, the 2p electron is easier to remove than one of the 2s electrons in beryllium
  • the first ionisation energy of boron is less than the first ionisation energy of beryllium
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11
Q

explain the fall in ionisation energy between nitrogen and oxygen

A
  • the fall in the first ionisation energy between nitrogen and oxygen marks the start of electron pairing in the p-orbitals of the 2p sub-shell
  • in nitrogen and oxygen, the highest energy electrons are in a 2p sub-shell
  • in oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom
  • therefore, the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen
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12
Q

what is metallic bonding?

A

the strong electrostatic attraction between cations and delocalised electrons

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13
Q

what are the properties of metals?

A
  • strong metallic bonds
  • high electrical conductivity
  • high melting and boiling points
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14
Q

the reaction of barium with bromine is more vigorous than the reaction of calcium with bromine, explain why

A
  • barium has a larger atomic radius than calcium
  • therefore, outer electrons experience less shielding
  • so the nuclear attraction in barium is weaker than in calcium
  • so it is easier to remove outer electrons in barium than calcium
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15
Q

explain which block in the periodic table sodium and magnesium belong to

A

s-block, outer electron is in a s-orbital

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16
Q

the reactivity of the Group 2 elements Mg–Ba increases down the group, explain why

A
  • atomic radius increases
  • shielding increases
  • nuclear attraction decreases
  • easier to remove an electron
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17
Q

explain the differences in the melting points of sodium and magnesium, using the model of metallic bonding

A
  • magnesium has a higher melting point than sodium
  • magnesium has more outer electrons
  • magnesium ions have a greater positive charge
  • magnesium has a greater attraction between ions and delocalised electrons
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18
Q

what is the name given to this repeating pattern of properties?

A

periodicity

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19
Q

explain the differences in the melting points of phosphorus (P4) and chlorine (Cl2)

A
  • phosphorus has more electrons
  • therefore stronger London forces
  • more energy required to overcome the intermolecular forces
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20
Q

describe the bonding in magnesium and silicon

A
  • magnesium has metallic bonds
  • cations with a +2 charge and delocalised electrons
  • silicon has covalent bonds
  • between atoms
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21
Q

explain why the first ionisation energy of strontium is less than the first ionisation energy of calcium

A
  • strontium has a larger atomic radius
  • increased nuclear charge is outweighed by increased shielding
  • less nuclear attraction
  • so it is easier to remove an electron
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22
Q

state the structure and bonding in silicon dioxide

A

giant covalent lattice

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23
Q

which groups are classified as the s-block in the periodic table?

A

groups 1 and 2

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24
Q

which groups are classified as the p-block in the periodic table?

A

groups 3 to 0

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25
Q

which group is classified as the d-block in the periodic table?

A

transition metals

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26
Q

what is the correct order to carry out the tests for anions?

A
  1. carbonate
  2. sulfate
  3. halides
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27
Q

why does the carbonate test need to be carried out before the sulfate test?

A
  • in the sulfate test, you add a solution containing barium ions and are looking for a white precipitate of barium sulfate
  • barium carbonate is white and insoluble in water, so if you carry out a sulphate test on a carbonate, you will get a white precipitate too
  • therefore, it is important to carry out the carbonate test first and only proceed to the sulfate test when you know that no carbonate ions are present
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28
Q

state and explain the trend in atomic radius from Li to F

A
  • atomic radius decreases
  • because the nuclear charge increases
  • outer electrons are in the same shell
  • so there is a greater nuclear attraction on the outer electrons
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29
Q

what is the most common type of reaction for group 2 elements?

A

redox reactions

30
Q

are group 2 elements oxidising or reducing agents and why?

A

reducing agents, because it loses two electrons and reduces another species

31
Q

explain the trend in reactivity down group 2

A
  • reactivity increases going down the group
  • because the ionisation energies decrease down the group
  • due to the decreased nuclear attraction on the outer electrons
  • as a result of increasing atomic radius and increasing shielding
32
Q

what is the trend in solubility of the metal hydroxides going down group 2?

A
  • solubility of the hydroxides in water increases down the group
  • so the resulting solution has more hydroxide ions and are more alkaline
33
Q

explain the trend in boiling points down group 7

A
  • going down the group, the halogen molecules have more electrons
  • stronger london forces
  • more energy required to overcome these intermolecular forces
  • so boiling point increases
34
Q

why are metals good electrical conductors?

A

the delocalised electrons can move throughout the structure and carry charge, in both solid and liquid states

35
Q

are giant covalent structures soluble or insoluble, and why?

A

insoluble, because the covalent bonds holding together the atoms in the lattice are too strong to be broken down by interaction with solvents

36
Q

what is the trend in melting points across period 2 and 3

A
  • the melting point increases from group 1 to group 4
  • there is a sharp decrease in melting point between group 4 and group 5
  • the melting points are comparatively low from group 5 to group 0
37
Q

why are silicon, carbon, oxygen and chlorine all classified as p-block elements?

A

their highest energy electron is in a p-orbital

38
Q

explain how to carry out the test for carbonates

A
  • in a test tube, add dilute nitric acid to the solid or solution to be tested
  • if bubbles are produced, the unknown compound could be a carbonate
  • to prove the gas is carbon dioxide, bubble the gas through lime water
  • if the limewater turns cloudy, carbon dioxide is present and the compound is a carbonate
39
Q

explain how to carry out the sulfate test

A
  • first, add dilute hydrochloric acid to the solid or solution to be tested
  • then add aqueous barium chloride
  • if a sulfate is present, a white precipitate of barium sulfate will form
40
Q

explain how to carry out the halide tests

A
  • add aqueous silver nitrate to an aqueous solution of a halide
  • the silver halide precipitates are different colours: silver chloride is white, silver bromide is cream, and silver iodide is yellow
  • add aqueous ammonia to test the solubility of the precipitate: silver chloride is soluble in dilute aqueous ammonia, silver bromide is soluble in concentrated aqueous ammonia, and silver iodide is insoluble in concentrated aqueous ammonia
41
Q

first ionisation energies decrease down groups in the periodic table, explain this trend and the effect on the reactivity of groups containing metals

A
  • going down groups, the atomic radius increases
  • there are more inner shells so increased shielding
  • nuclear attraction on outer electrons decreases
  • so the first ionisation energies decrease
  • for groups containing metals, this means that the reactivity increases going down the group
  • as it is easier to remove an outer electron due to the decreased nuclear attraction
42
Q

what is the name given to this repeating pattern of properties?

A

periodicity

43
Q

state how and explain why the attraction between nuclei and outermost electrons in gaseous atoms varies across period 3

A
  • the attraction increases across the period
  • because, the nuclear charge increases and the atomic radius decreases
44
Q

describe and explain the electrical conductivity of sodium oxide, and sodium in their solid and molten states

A

sodium:
- can conduct in both its solid and molten states
- because the delocalised electrons are free to move

sodium oxide:
- cannot conduct in its solid state as the ions are fixed in place in an ionic lattice
- can conduct in its molten state as the ions are mobile

45
Q

state and explain the trend in atomic radius from Li to F

A
  • atomic radius decreases from Li to F
  • this is because the nuclear charge increases
  • and they have the same number of shells
  • so there is greater nuclear attraction on outer electrons
46
Q

what is the activation energy?

A

the minimum energy required to start a reaction by the breaking of bonds

47
Q

what is enthalpy?

A

a measure of the heat energy in a chemical system

48
Q

what is the enthalpy change?

A

the difference between the enthalpy of the products and the enthalpy of the reactants

49
Q

what is the law of conservation of energy?

A

energy cannot be created or destroyed

50
Q

what does “standard state” mean?

A

the physical state of a substance under the standard conditions of 100kPa and 298K

51
Q

what is the standard enthalpy change of reaction?

A

the enthalpy change that accompanies a reaction in the molar quantities stated in the equation under standard conditions, all reactants and products being in their standard states

52
Q

what is the standard enthalpy change of formation?

A

the enthalpy change that takes place when one mole of a compound is formed from its constituent elements in their standard states under standard conditions

53
Q

what is the standard enthalpy change of combustion?

A

the enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard conditions, all reactants and products being in their standard states

54
Q

what are the standard conditions for enthalpy change?

A

100kPa pressure, 298K temperature and 1.0 mol per cubic decimetre concentration of aqueous solutions

55
Q

what is the average bond enthalpy?

A

the energy required to break one mole of a specified type of bond in a gaseous molecule

56
Q

what are the two conditions that are always true for bond enthalpies?

A
  • bond enthalpies are always endothermic
  • bond enthalpies always have a positive enthalpy value
57
Q

what are the limitations of average bond enthalpies?

A

the actual bond enthalpy can vary depending on the chemical environment of the bond e.g. the composition of the given molecule

58
Q

what is Hess’ Law?

A

the enthalpy change accompanying a chemical reaction is independent of the route by which the change occurs

59
Q

what is the role of a catalyst?

A
  • lowers the activation energy by providing an alternate pathway
  • they increase the rate of reaction, and are not used up during the reaction
60
Q

what is the difference between the activation energy in exothermic and endothermic reactions?

A

the activation energy is endothermic reactions is relatively larger than in exothermic reactions

61
Q

what is a heterogeneous catalyst?

A

a catalyst that is in a different physical state to the reactant

62
Q

what is the Boltzmann distribution?

A

the spread of molecular energies in gases

63
Q

what are the key features of the Boltzmann distribution?

A
  • no molecules have zero energy, it starts at the origin
  • the area under the curve is equal to the total number of molecules
  • there is no maximum energy for a molecule, the curve does not meet the x-axis at high energy
64
Q

what is Le Chatelier’s principle?

A

states that when a system in equilibrium is subjected to an external change the system readjusts itself, the position of equilibrium shifts, to minimise the effect of that change

65
Q

what does dynamic equilibrium mean?

A
  • the rate of the forward reaction is equal to the rate of the reverse reaction
  • the concentration of reactants and products do not change
66
Q

what does the value of the equilibrium constant tell us?

A
  • the exact position of equilibrium
  • indicates the relative proportions of reactants and products in the equilibrium system
67
Q

describe the effect of a catalyst on the position of equilibrium

A
  • a catalyst does not change the position of equilibrium
  • it merely speeds up the rates of the forward and reverse reactions equally
  • but it will increase the rate at which an equilibrium is established
68
Q

describe and explain the effect of decreasing the pressure on the rate of reaction

A
  • concentration is reduced
  • fewer collisions per second
  • and the rate of reaction decreases
69
Q

what are the factors affecting equilibrium position?

A
  • temperature
  • pressure
  • concentration
70
Q

how do you convert cm³ to m³?

A

times ten to the minus 6

71
Q

how do you convert dm³ to m³?

A

times ten to the minus 3

chemistry (9 decks)

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