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chemistry > topic 1 - atomic structure and the periodic table > Flashcards

topic 1 - atomic structure and the periodic table Flashcards

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1
Q

The Avogadro constant is 6.023 x 10^23mol-1. What is the number of atoms in 1 mole of carbon dioxide?

A

1.8x10^24

6.023 x 10^23 x 3 =1.8x10^24

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2
Q

isotope definition

A

different atoms of the same element which have the same number of protons but different number of neutrons
so same atomic number but different mass number

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3
Q

what is the relative mass of an electron

A

1/1840

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4
Q

how do you calculate the relative atomic mass of an element, when given its isotopes and abundance

A

(percentage abundance x relative isotopic mass) + (percentage abundance x isotopic mass)/100

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5
Q

definition of relative atomic mass

A

the average mass of of an atom of an element compared to 1/12th the mass of one carbon-12 atom

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6
Q

definition of relative isotopic mass

A

the mass of an atom of an isotope of the element compared to 1/12th the mass of one atom of carbon-12

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7
Q

why do isotopes have similar chemical properties but different physical properties

A

-isotopes have similar chemical properties because they have the same electronic configuration
-they have different physical properties because isotopes have different mass numbers

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8
Q

what is mass spectroscopy used to determine

A

-the molecular mass of a compound and the mass of its fragments, which helps to determine the structure of the compound
-the mass of each isotope present in an element, and the relative abundance of each isotope

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9
Q

what are the 4 stages of mass spectroscopy

A

I - ionisation
A - acceleration
D - deflection
D - detection

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10
Q

describe the acceleration stage of mass spectroscopy

A

-the positive ions pass through positively charged electrodes, which create an electric field and cause the positive ions to accelerate

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11
Q

describe the ionisation stage of mass spectroscopy

A

-first, the sample is vaporised
-the vaporised sample is bombarded by high energy electrons from an electron gun
-the sample is ionised, electrons are removed leaving positive ions

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12
Q

describe the deflection stage of mass spectroscopy

A

-the positive ions are deflected by a strong magnetic field
-all ions start at the same speed and direction, the amount they get deflected is determined by their mass
-lighter ions are deflected more than heavier ions, causing ions to be separated

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13
Q

describe the detection stage of mass spectroscopy

A

-the ions hit a detector connected to a computer
-a mass spectrum is formed

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14
Q

why is it important that the ions in mass spectroscopy pass through a vacuum rather than in air

A

-air particles would collide with the ions and deflect their normal paths
-or air particles may be detected by the detector and create extra peaks on the mass spectrum

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15
Q

what affects how much an ion is deflected in mass spectroscopy

A

-lighter ions are deflected more than heavier ions
so 35Cl+ will be deflected more than 37Cl+
-ions with a greater charge are deflected more
so 16O2+ is deflected more than 16O+, because the m/e value for 16O+ is 8 (mass 16 divided by charge 2) whereas the m/e value for 16O+ is 16

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16
Q

what is the ideal gas equation

A

PV=nRT

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17
Q

what do the letters in the ideal gas equations stand for and what are their units

A

PV=nRT

P=pressure, Pa (Pascals)
V=volume , m^3 (cubic metres)
n=number of moles
R=the gas constant, 8.31JK-1mol-1 (given in data book)
T=temperature, K (Kelvin)

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18
Q

how do you convert Celcius into Kelvin

A

add 273

0 degrees celcius = 273 kelvin

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19
Q

how do you convert form cm3 to m3

A

1m3=1,000dm3
1dm3=1,000cm3
so 1m3=1,000,000cm3

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20
Q

how do you work out the relative molecular mass of a compound using a mass spectrum

A

-find the M peak (usually the largest m/z value unless a M+1 peak is present)
-the m/z value at the M peak equals the relative molecular mass of the compound

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21
Q

on a mass spectrum, what type of peak may be present next to the M peak

A

-in molecules with large masses there may be a M+1 peak next to the M peak
-because there is always a percentage of carbon-13 present in compounds

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22
Q

definition of ionisation energy

A

energy needed to remove one mole of electrons from one mole of atoms of an element in the gaseous state, to form one mole of singly charged positive ions

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23
Q

what is the first ionisation energy of an element

A

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

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24
Q

what is the second ionisation energy of an element

A

the energy required to remove 1 mole of electrons from 1 mole of 1+ positively charged ions of an element in gaseous state to form 1 mole of 2+ ions

25
Q

what 3 factors influence ionisation energies

A

-number of shells, the greater the distance between the electron and the nucleus the weaker the force of attraction, therefore a smaller ionisation energy is required

-nuclear charge, the greater the number of protons in the nuceus, the greater the positive charge of the nucleus so the greater the force of attraction, therefore a higher ionisation energy is required

-shielding, the more electrons there are in the same orbital or quantum shell the more electron-electron repulsion, which causes the outermost electron to increase in energy, so ionisation energy decreases

26
Q

what is the order of all the quantum shells possible

A

1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 4d^10

27
Q

how does adding electrons to an atom affect the energy needed to remove the atom

A

the electrons repel eachother
so the energy levels of electrons increases
the increased repulsion means it takes less energy to make an electron in one of the orbitals a free electron

28
Q

how does a greater number of protons affect the energy level of electrons

A

the more protons the greater the nuclear charge, so the stronger the electrostatic charge between the outer electron and protons, so a greater ionisation energy is required

29
Q

what is the meaning of ground state
(e.g. ground state electronic configuration)

A

ground state - the lowest possible energy the electrons in an atom can have
e.g. helium’s ground state electronic configuration is 1s^2

30
Q

what happens when you give gaseous atoms energy

A

-the electrons move to higher energy levels
-they are said to be in an “excited” state
-eventually they return to their original lower energy levels and emit electromagnetic radiation
-only specific frequencies are emitted which are unique to each element
-a spectroscope analyses the electromagnetic radiation and a line emission spectra is produced

31
Q

what is the evidence that electrons can only have fixed energy values

A

-as electrons drop to their original energy levels after being in an excited state they release electromagnetic radiation
-which is only at specific frequencies and unique to each element
-showing that electrons only have fixed energy levels rather than a continuous range

32
Q

explain why is sodium the first ionisation energy have the lowest value

A

-the 1st ionisation energy removes the outermost electron, so this electron has the highest energy level out of all the electrons
-so it needs to gain a lower amount of energy to reach the required energy needed to be removed from the atom

33
Q

explain why is sodium there is a big jump in ionisation energy between the 1st and 2nd

A

-the 2nd electron removed is in a quantum shell of a lower energy level, so the electron needs to gain more energy to reach the required energy to be removed

34
Q

even though the electrons are in the same quantum shell, why is there a steady increase in successive ionisation energies

A

-as each successive electron is removed from the shell, the shielding (electron-electron repulsion) within the shell decreases
-resulting in a decrease of energies of the remaining electrons
-therefore an increase in ionisation energy (as the electrons need to gain more energy to reach the required amount)
-the electrons are in quantum shells closer to the nucleus so electrostatic force of attraction increases
-nuclear charge remains the same, so electrostatic force of attraction shared between fewer electrons

35
Q

what is shielding/electron-electron repulsion
what increases and decreases it
what effect does increases shielding/electron-electron repulsion have

A

-electron-electron repulsion increases the energy of the electrons involved aboeve the value they would have if there was no repulsion.
-electron-electron repulsion exists between 2 electrons in the same orbital, and between electrons in different orbitals but same quantum shell
-increased electron-electron repulsion is achieved when more electron are added to the atom
-decreased electron-electron repulsion occurs when electrons are removed from the atom

36
Q

what effect does increasing and decreasing shielding/electron-electron repulsion have

A

-if electron-electron repulsion is increased the electrons involved will have higher energies, so their ionisation energies will decrease, because they need to gain less energy to reach the required energy.
-if electron-electron repulsion decreases the electrons involved will have less energy, so their ionisation energies will increases, because they need to gain more energy to reach the required energy.

37
Q

which has a higher first ionisation energy hydrogen or helium
H 1s^1
He 1s^2

A

-although both their outer electrons are in the 1s orbital, they don’t have the same first ionisation energy
-in helium there is a higher electron-electron repulsion because there are 2 electrons is the 1s orbital (each electron shields the other from the nuclear charge) so this increases the energy of the electrons.
-however the nuclear charge of helium is double that of hydrogen, this decreases the energy of the electrons because they are attracted more strongly to the nucleus.
-the effect if the increased nuclear charge is greater than increased shielding, so the 1st ionisation energy of helium is greater than hydrogen’s

38
Q

which has a higher first ionisation energy helium or lithium
He 1s^2
Li 1s^2 2s^1

A

-the nuclear charge of lithium (3 protons) is greater than lithium’s (2 protons). (this should increase lithium’s ionisation energy)
-the outer electron in lithium is in a 2s orbital so have a higher energy level than the outer electron in helium. (should decrease lithium’s ionisation energy)
-the outer most electron in lithium experiences repulsion from the 2 inner electrons so it experiences shielding from the nuclear charge meaning the electron has increased energy. (should decrease lithium’s ionisation energy)
-the higher energy orbital and increased shielding are more significant than increased nuclear charge so the ionisation of lithium is smaller than that of helium

39
Q

what are the 3 factors which affect the energy an electron has

A

-the orbital the electron is in (if an electron is in an orbital further from the nucleus it will have a higher energy)
-nuclear charge of the atom (the greater the nuclear charge the less energy the electrons have)
-the repulsion it experiences from all other electrons present in the atom

40
Q

state the trend in ionisation energies down a group

A

there’s a general decrease in ionisation energy

41
Q

state the trend in ionisation energies across a period

A

there’s a general increase in ionisation energy

42
Q

explain the trend in ionisation energies across a period, by using period 2 lithium to neon as an example

A

-from Li to Ne across period 2, nuclear charge increases because proton number increases, so there’s a greater attraction between the outermost electron and nucleus, so outermost electron has less energy.
-electrons are added in the same quantum shell (mostly)
-across the period, 1 more electron is added to each element, this increases the electron-electron repulsion within the quantum shell, so the outermost electron has increased energy, so ionisation energy decreases
-the increase in nuclear charge is more significant than the increased shielding, so there’s a greater attraction between nucleus and electrons, making the atomic radius smaller , so a general increase in ionisation energy across period 2
-this is the same for period 3, excluding d-block elements

43
Q

explain the trend in ionisation energies down a group, using group 1 lithium to caesium as an example

A

-down group 1 from Li to Cs, nuclear charge increases because no of protons increase, so increased attraction between nucleus and outermost electron, so energy of electron decreases, which increases ionisation energy
-each time you go down the group a new quantum shell is added, quantum shells increase in energy as there are more of them, so the energy of the outermost electron increases, so ionisation energy decreases
-as each new quantum shell is added the outer electron experiences increased shielding from the inner electrons, so the outermost electron increases in energy, so ionisation energy decreases
-more quantum shell and increased shielding is more significant than increased nuclear charge, so there is a general decreases in ionisation energy down the group

44
Q

what 2 periods do not fit the trend in ionisation energies across a period

A

period 2 beryllium and boron
period 3 magnesium and aluminium

45
Q

state the trend in atomic radii down a group

A

atomic radii increase

46
Q

state the trend in atomic radii across a period

A

atomic radii decrease

47
Q

why is the ionisation energy for sulfur less than the ionisation energy of phosphorous

A

-because when you draw the electrons in box diagram for sulfur it shows it has 1 pair of electrons and 2 single electrons in the the 3p subshell whereas for sulfur the 3p orbital is half-filled
-the presence of 2 electrons in 1 orbital causes the paired electrons to repel each other, so it has greater shielding, so ionisation energy decreases
-as phosphorous the 3p subshell is half-filled, it is stable so requires a higher ionisation energy

48
Q

what is the definition of periodicity

A

a regular repeating trend of atomic, physical and chemical properties with increasing atomic number across a period

49
Q

how do you work out the atomic radius of an atom

A

-the atomic radius of an atom is the distance from the centre of the nucleus to the boundary of the electron cloud
-but the electron cloud doesn’t have a well-defined boundary
-so we measure the distance between 2 nuclei of the same type of element and then divide by 2

50
Q

what are the 3 types of atomic radii

A

covalent radius-measured between 2 atoms covalently bonded together
van der Waals radius-measured between 2 atoms not chemically bonded together
metallic radius-measured between 2 metals metallically bonded together

51
Q

explain the trend in atomic radii across a period

A

-along each period, a proton is added, so the nuclear charge increases
-so the attraction between the nucleus and electrons increases, meaning the electrons are more tightly held so atomic radius decreases
-despite the increase in shielding due to the outer electrons being in the same shell

52
Q

explain the trend in atomic radii down a group

A

-each time we go down the group a new quantum shell is added, so the outer electron is further from the nucleus,
-increased shielding
-this decreases the attraction between the nucleus and electrons
-so the electrons are less tightly held, so atomic radius increases
-despite nuclear charge increasing as more protons are added effective nuclear charge decreases because of increased shielding

53
Q

explain why the first ionisation energy of beryllium is more than that of boron
Be 1s^2 2s^2
B 1s^2 2s^2 2p^1

A

-although the nuclear charge of boron is greater than that of beryllium
-boron’s outer electron has a higher energy level because it’s in the 2p subshell as opposed to the 2s subshell for beryllium
-and boron has greater shielding because it has 2 inner electron subshells as opposed to only 1 is beryllium
-so boron’s outer electron has more energy, so requires a lower ionisation energy compared to beryllium

54
Q

which has a greater radius Na+ or Na why

A

-Na has a greater atomic radius
-Na+ is formed by losing an electron
-the nuclear charge stays the same
-but because there is 1 less electron to share to charge between, the electrostatic force of attraction between the protons and electrons increase
-so the electrons are held closer to the nucleus, making the ionic radius smaller

55
Q

which has a greater radius Na or Na- why

A

-Na- has a greater ionic radius
-Na- is formed by gaining 1 electron
-the nuclear charge needs to be spread between more electrons
-so the electrostatic force of attraction between the protons and electrons decreases
-so the electrons are held less tightly and further away from the nucleus, increasing ionic radius

56
Q

effective nuclear charge definition

A

the net positive charge the outermost electron experiences after taking into account the effects of shielding
its always less than the nuclear charge

57
Q

state how the relative abundance of 2 isotopes can be determined (2 marks)

A

compare the number of particicles for each isotope detected
using a mass spectrometer

58
Q

what is the definition of relative isotopic mass

A

the mass of an atom of that isotope relative to 1/12th the mass of a carbon 12 atom

59
Q

in the mass spectrum of chlorine, what is the ratio of m/z values (abundance for each peak)

A

9:6:1

chemistry (3 decks)

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