topic 1 - atomic structure and the periodic table

Question 1

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

Answer 1

1.8x10^24

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

Question 2

isotope definition

Answer 2

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

Question 3

what is the relative mass of an electron

Answer 3

1/1840

Question 4

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

Answer 4

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

Question 5

definition of relative atomic mass

Answer 5

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

Question 6

definition of relative isotopic mass

Answer 6

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

Question 7

why do isotopes have similar chemical properties but different physical properties

Answer 7

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

Question 8

what is mass spectroscopy used to determine

Answer 8

-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

Question 9

what are the 4 stages of mass spectroscopy

Answer 9

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

Question 10

describe the acceleration stage of mass spectroscopy

Answer 10

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

Question 11

describe the ionisation stage of mass spectroscopy

Answer 11

-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

Question 12

describe the deflection stage of mass spectroscopy

Answer 12

-the positive ions are deflected by a strong magnetic field
-lighter ions are deflected more than heavier ions, causing ions to be separated

Question 13

describe the detection stage of mass spectroscopy

Answer 13

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

Question 14

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

Answer 14

-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

Question 15

what affects how much an ion is deflected in mass spectroscopy

Answer 15

-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

Question 16

what is the ideal gas equation

Answer 16

PV=nRT

Question 17

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

Answer 17

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)

Question 18

how do you convert Celcius into Kelvin

Answer 18

add 273

0 degrees celcius = 273 kelvin

Question 19

how do you convert form cm3 to m3

Answer 19

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

Question 20

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

Answer 20

-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

Question 21

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

Answer 21

-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

Question 22

definition of ionisation energy

Answer 22

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

Question 23

what is the first ionisation energy of an element

Answer 23

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

Question 24

what is the second ionisation energy of an element

Answer 24

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

Question 25

what 3 factors influence ionisation energies

Answer 25

-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

Question 26

what is the order of all the quantum shells possible

Answer 26

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

Question 27

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

Answer 27

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

Question 28

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

Answer 28

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

Question 29

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

Answer 29

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

Question 30

what happens when you give gaseous atoms energy

Answer 30

-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

Question 31

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

Answer 31

-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

Question 32

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

Answer 32

-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

Question 33

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

Answer 33

-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

Question 34

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

Answer 34

-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)

Question 35

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

Answer 35

-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

Question 36

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

Answer 36

-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.

Question 37

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

Answer 37

-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

Question 38

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

Answer 38

-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

Question 39

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

Answer 39

-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

Question 40

state the trend in ionisation energies down a group

Answer 40

there’s a general decrease in ionisation energy

Question 41

state the trend in ionisation energies across a period

Answer 41

there’s a general increase in ionisation energy

Question 42

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

Answer 42

-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

Question 43

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

Answer 43

-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

Question 44

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

Answer 44

period 2 beryllium and boron
period 3 magnesium and aluminium

Question 45

state the trend in atomic radii down a group

Answer 45

atomic radii increase

Question 46

state the trend in atomic radii across a period

Answer 46

atomic radii decrease

Question 47

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

Answer 47

-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

Question 48

what is the definition of periodicity

Answer 48

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

Question 49

how do you work out the atomic radius of an atom

Answer 49

-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

Question 50

what are the 3 types of atomic radii

Answer 50

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

Question 51

explain the trend in atomic radii across a period

Answer 51

-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

Question 52

explain the trend in atomic radii down a group

Answer 52

-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

Question 53

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

Answer 53

-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

Question 54

which has a greater radius Na+ or Na why

Answer 54

-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

Question 55

which has a greater radius Na or Na- why

Answer 55

-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

Question 56

effective nuclear charge definition

Answer 56

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