Topic 1 - Atomic Structure & The Periodic Table Flashcards

(40 cards)

1
Q

Meaning Relative Atomic Mass (Ar)

A

The weighted mean mass of an atom of an element, compared to 1/12th of the mass of an atom of carbon-12

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2
Q

Meaning Relative Isotopic Mass

A

The mass of an atom of an isotope, compared to 1/12th of the mass of carbon-12

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3
Q

Meaning Relative Molecular Mass (Mr)

A

The mean mass of a molecule, compared to 1/12th the mass of an atom of carbon-12

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4
Q

Name all of the orbitals

A

s, p, d, f

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5
Q

Describe the s orbital

A

Has 1 orbital, so can hold 2 electrons

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6
Q

Describe the p orbital

A

Has 3 orbitals, so can hold 6 electrons (2x3)

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7
Q

Describe the d orbital

A

Has 5 orbitals, so can hold 10 electrons (5x2)

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8
Q

Describe the f orbital

A

Has 7 orbitals, so can hold 14 electrons (7x2)

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9
Q

Which shape does an S orbital have

A

Spherical shape

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10
Q

Which shape does a p orbital have

A

Dumb-bell shaped

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11
Q

Give the electron configuration of iron (26)

A

1s2 2s2 2p6 3s2 3p6 4s2 3d6

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12
Q

Give the electron configuration of calcium +2

A
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13
Q

Give the electron configuration of iron +3

A
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14
Q

Give the electron configuration of copper (29)

A

1s2 2s2 2p6 3s2 3p6 4s1 3d10

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15
Q

State the different radiations in order of the electromagnetic spectrum

A

Radio waves, micro waves, infrared, visible light, ultraviolet, X-rays, gamma rays

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16
Q

When do electrons become excited

A

When electrons absorb energy (endothermic) as it moves up an energy level

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17
Q

In an atomic emission spectra, why do the lines get closer

A

Because energy and frequency increases

18
Q

When do electrons show ultraviolet radiation

A

When an electron falls to the ground state (n=1)

19
Q

When do electrons show visible radiation

A

When an electron falls to the second energy level (n=2)

20
Q

When do electron show infrared radiation

A

When electrons fall to the third energy level (n=3)

21
Q

Meaning ionisation

A

The minimum amount of energy required to remove 1 mole of electrons from 1 mole of atoms in a gaseous state

22
Q

What are 3 factors of ionisation

A

Shielding, nuclear charge, atomic size

23
Q

How does shielding affect ionisation

A

The more electron shells between the positive nuclei and the negative electron thats is being removed, the less energy required to ionise because theres a weaker attraction

24
Q

How does nuclear charge affect ionisation

A

The more protons in the nucleus, the larger the attraction between the nucleus and the outer electrons. This means more energy is needed for ionisation

25
How does atomic size affect ionisation
The bigger the atom, the further away the outer electrons are from the nucleus. The attractive force between the nucleus and the outer electrons reduces, so less energy is needed to ionise an electron
26
Why does the atomic radii decrease across a period
The further across a period, there higher number of protons in the nucleus. This pulls the outer electrons further towards the nucleus. Shielding has no effect
27
Why does the melting points of metals increase across a period
As the metals have more protons in their nucleus, there’s an increase in positive charge. So, there’s an increased number of delocalised electrons and smaller atomic radii. This means stronger metallic bonds, so more energy is needed to bream the binds.
28
Explain why silicon (S8) has the highest melting point in period 3
Silicon is a giant covalent structure. This means it has many strong covalent bonds that hold the silicon atoms together. Thus, a large amount of energy is needed to overcome these string covalent bonds
29
Why does sulfur have a lower melting point than phosphorus, even though it’s further across period 3
Sulfur has a larger simple molecular structure and has larger london forces than phosphorus
30
Why does argon have the lowest melting point in group 3
Argon only exists as individual atoms as its a noble gas, so it has small london forces
31
describe the trend in reactivity in group 1 and 2 elements
the elements increase in reactivity as you go down the group
32
describe the trend in reactivity in group 6 and 7 elements
the elements decrease in reactivity as you go down the group
33
why do group 1 and 2 elements increase in reactivity as you go down the group
going down the group, increases the distance of outer shell electrons from the nucleus, plus the shielding effect of the inner shell electrons overcomes the increasing nuclear charge. Thus, the outer shell electrons are held less strongly and are easier to remove in a reaction.
34
why do group 6 and 7 elements increase in reactivity as you go up the group
going up the group, reduces the distance of the outer shell from the nucleus and decreases shielding effects on the decreasing positive charge, meaning that the outer shell electrons are more strongly attracted to the nucleus. Thus, the outer shell is more readily accepts another electron in a reaction
35
why does the 1st ionisation energy for the elements increase across a period
as the outer electron is removed from elements of increasing atomic number within a period, it is being removed from the same quantum shell, but the nuclear charge is steadily becoming more positive so holding the electron with a greater force, thus more energy is needed to remove it
36
why does the 1st ionisation energy for the elements decrease down a group
as the electron is removed from an outer shell which is increasing in distance from the nucleus. This increasing distance and the increased shielding effect of the increasing number of inner shells present, overcomes the fact that the nuclear charge is increasing. Thus, the outer electron is held less strongly and less energy is needed to remove it
37
why do the successive ionisation energies increase for electrons taken from the same quantum shell
As you take successive electrons from the same quantum shell, the electrons are a similar distance from the nucleus, with the same inner shell shielding, but the ionic charge increases each time, so the attractive force on the electron being removed increases so it requires more energy to take the away
38
why are there large jumps in successive energies for an elements such as sodium between Em1 to Em2 and Em9 to EM10
As you successfully take electrons away from Na, the 1st electron is from the 3rd Quantum shell, which is further from the nucleus than the 2nd electron, which is from the 2nd Quantum shell, which is nearer to the nucleus so attracts the electron more strongly. After 9 electrons have been removed, the 10th electron is now taken from the first Quantum shell, which is again much closer to the nucleus, so attracts the electron more strongly
39
why is there a drop in the 1st I.E. between the elements Mg and Al in the 3rd period
the outer electron is removed from a 3P orbital in Al, but from the 3s orbital in Mg. The energy of the electrons in the 3p subshell is greater than the energy of the 3s subshell. Thus, it takes less energy to remove a 3p electron in Al, than a 3s electron in Mg
40