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Atomic structure Flashcards

1
Q

atomic orbitals

A

regions around nucleus that can hold up to two electrons with opposite spins

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2
Q

fact about the two electrons in the same orbital

A

they have opposite spins (up or down)
due to repulsion forces
between the two negatively charged e-

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3
Q

what is an electron

A

a cloud of negative charge

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4
Q

what does an orbital show

A

predicts (95% accurate) the location of an electron

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5
Q

s orbital explained

A

spherical shaped
every electron shell has 1 s subshell
there is one orbital for s
the s orbital can hold a maximum of 2 electrons
s block starts at

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6
Q

p orbitals explained

A

dumbbell shaped
there are three p subshells
each orbital can hold 2 e- (total 6 electrons)
p-block starts at boron

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7
Q

d orbital explained

A

there are 5 d subshells
each orbital hold 2e-
holds 10 electrons in total
d block starts at scandium

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8
Q

f orbitals explained

A

there are 7 f subshells
two electrons in each orbital
so 14 electrons in total

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9
Q

what is a subshell

A

all of the same type of orbitals

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10
Q

rules for filling atomic orbitals

A

orbitals with lowest energy levels are filled first
each orbital can hold two opposite spin e-
when subshells are getting filled, an electron goes to each orbital before filling an orbital

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11
Q

why do electron get shared across same energy level orbitals before filling up an orbital

A

to reduce the repulsion forces experienced between two same (-) charged electrons in the same orbital

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12
Q

why is the 4s energy subshell filled before 3d

A

4s has a lower energy level than 3d

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13
Q

chromium and copper electron configuration difference

and why

A

the 4s orbital contains one electron
even though there are electrons in the 3d subshell
has only 4s electron so that the 3d subshell can be half full or completely full so it is more stable

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14
Q

where is the s block elements located

A

first two groups including helium
starting from hydrogen (1s^1)

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15
Q

where is the p block elements located

A

right side of the periodic table
starting from boron (2p^1)

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16
Q

where is the d block elements located

A

middle (transition metals)
starting from scandium (3d^1)

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17
Q

how to write the shorthand electron configuration

A

only the number of electrons in the outer shell is needed
put the name of the group 0 (noble gas) in square brackets, that comes before the element
then write out the remainder of the config

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18
Q

what happens to the 4s subshell when it has electrons in it

A

becomes a higher energy level than 3d

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19
Q

so when electrons are removed (ion formed) from a d block element where do they lose electrons from

A

because when a 4s subshell has an electron in it (even though it initially had a lower energy level without any) it has a higher energy level than 3d
so electrons are lost from 4s first

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20
Q

rules for subshells when forming ions

A

electrons are always lost from the highest energy subshell

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21
Q

first ionisation energy defintion

A

energy needed to remove one mole of electrons from one mole of atoms
in their gaseous state
to form a mole of 1+ ions also in their gaseous state

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22
Q

first ionisation energy equation

A

X(g) -> X+(g) + e-

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23
Q

second ionisation energy definition

A

energy needed to remove one mole of electrons
from a mole of a gaseous 1+ ion
to form a 2+ ion also in their gaseous state

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24
Q

second ionisation energy equation

A

X+(g) -> X2+ (g) + 2e-

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25
factors affecting ionisation energy
- atomic radius - number of protons (atomic charge) - sheilding
26
affect of atomic radius on first ionisation energy
as atomic radius increases, first ionisation energy decreases due to weaker electrostatic forces of attraction between positive nucleus and negative electrons, as the distance between them is larger
27
atomic charge definition
the total positive charge of the nucleus of an atom
28
affect of number of protons (atomic charge) on first ionisation energy
As proton number increases, first ionisation energy increases as the more positive protons in the nucleus the greater the electrostatic forces of attraction to the negative outer shell electrons
29
what is shielding
repulsion experienced between electrons in inner and outer shells, pushing outer energy levels further away from nucleus
30
affect of shielding on first ionisation energy
the more shielding experienced, the lower the first ionisation energy electrons experience repulsion pushing them further away from the nucleus so distance between nucleus and electrons increases weaker electrostatic forces of attraction between protons and electrons
31
why is their a gradual increase in ionisation energy when removing electrons from the same shell
remaining outer shell electrons are pulled closer in to nucleus so greater electrostatic forces of attraction between + nucleus and - electrons
32
what explains the massive energy increase on a first ionisation energy graph
when you remove the last electron from the outer shells the graph spikes to the next shell indicating high increase in f.i.e as new shell is closer to nucleus
33
first ionisation energy trend down a group
first ionisation energy decreases as number of shells increases so larger atomic radius and more shielding experienced so more distance means weaker electrostatic forces of attraction between (-) electrons and (+) nucleus
34
first ionisation energy across a period
usually increases across a period moving from left to right the nuclear charge (protons) increases, which decreases atomic radius increasing attraction of (-) electrons towards (+) nucleus
35
why does shielding not affect first ionisation energy when going across a period
same period = same number of shells so shielding effect the electrons experience is similar for elements in the same period not making a difference in the trend going across the period
36
exception in the trend in first ionisation energy across the period (group 2 to group 3)
energy drop from group 2 to group 3 (Be to B) as Beryllium has outer electron in subshell 2s a lower energy subshell and more stable (higher energy needed to remove e-) than Boron 2p, a higher energy subshell so easier (less energy needed) to remove electron as less stable
37
exception in trend in first ionisation energy (group 5 to 6)
energy drop electrons in nitrogen are more stable as it has 3 electrons in a 2p orbital each (harder to remove electron) oxygen has one filled subshell increasing the repulsion forces between electron pair (easier to remove an electron: less energy needed)
38
1 atom history: 400 BC
Democritus theorised 'atomos' proposed indivisible particles
39
2 atom history: Dalton 1803
theorised billiard ball model proposed indivisible atom proposed law of conservation of mass
40
3 atom history: Thompson 1897
theorised plum pudding model discovered electrons using cathode ray tube experiment proposed atom is a positive ball with negative electrons embedded into it
41
4 atom history: Rutherford 1911
theorised nuclear model using gold foil experiment discovered atoms are made up of mostly empty space as alpha particles when straight through thin sheet discovered small, dense positive nucleus as some alpha particles deflected
42
5 atom history: Bohr 1913
theorised fixed energy levels in which electrons orbit explaining atomic spectra
43
final atom history: Schrodinger 1926
theorised orbital (electron exist in probability clouds) theorised electrons behave as both waves and particles
44
mass number (A)
sum of the protons and neutrons in an atoms nucleus
45
atomic number (z)
the number of protons in the nucleus of an atom
46
relative atomic mass (Ar)
the average mass of an atom of an element relative to 1/12th the mass of an atom of Carbon-12
47
isotopes definition
atoms of the same element (same number of protons) but different number of neutrons (different mass number)
48
isotopes chemical properties
as they have the same proton number, neutral isotopes will have the same electron configuration so they have the same chemical properties (react the same way)
49
isotopes physical properties
isotopes have differing mass numbers so they have differing physical properties
50
ions
formed when an atom gains or loses an electron to have a net charge
51
what is mass spetrometry
analytical technique used to identify different isotopes used to find Ar of an element
52
TOF mass spectrometry: two types of ionisation
electron impact electrospray
53
what is electron impact ionisation used for
used for smaller molecules (low Mr) as it can cause fragmentation
54
what is electrospray ionisation used for
used for larger molecules like polypeptides as this does not fragment molecule
55
how does electrospray ionisation happen
sample dissolved in volatile solvent (with H+) passed through a fine needle with high voltage forming positively charged droplets that evaporate to single positively charged ions as they gain a H+ from the solvent (X+)
56
equation for electrospray ionisation charge on ions
X(g) + H+(aq) -> XH+(g)
57
rule for calculating mass of ion that has gone through electrospray ionisation
subtract one from the calculated mass because their is an extra proton from the addition of a H+ in the beginning
58
how does electron impact ionisation work
high energy electron fired from electron gun knocks off electron from vapourised atom making positive ion (X+)
59
equation for electron impact ionisation to form ion from atom
X(g) -> X+(g) + e-
60
how are ions accelerated in TOF
electric field applied to accelerate all ions to the same kinetic energy
61
equation for kinetic energy
Ek = 1/2 mv^2
62
ion drift explained
ions travel through a flight tube that has no electric field, after being accelerated considering all ions have the same kinetic energy, their speed through the flight tube depends on mass
63

chemistry A-LEVEL (12 decks)

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