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Chemistry > Electronic structure > Flashcards

Electronic structure Flashcards

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1
Q

what is the uncertainty principle

A

you can never know both the exact location and velocity of a subatomic particle at the same time

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2
Q

what are ionization energies

A

ionization energies are measures of the amount of energy needed to remove electrons from atoms (or ions)

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3
Q

what is the definition for first ionization energy

A

the energy required to remove one mole of electrons from one mole of gaseous atoms to for one mole of gaseous psositive ions

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4
Q

why are ionization energies always positive

A

because electrons and protons attract one another so energy must be added to the system to pull the electron away

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5
Q

why does an atom with a high nuclear charge have a high ionization energy

A

the greater the pull of the nucleus the harder it will be to pull an electron away (more energy required) so higher ionization energy

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6
Q

what do ionization trends give evidence for

A

electrons being in shells and subshells

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7
Q

what is electron affinity

A

the amount of energy needed to add electrons to atoms/ions

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8
Q

what is the definition of the first electron affinity

A

the energy required to add one mole of electrons to one mole of a gaseous atom to form one mole of gaseous negative ions

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9
Q

why is first electron affinities negative

A

indicates energy is released on the addition of the electron
the more negative the electron affinity, the more stable the negative ion that is formed

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10
Q

what is the trend of atomic size

A

increases down a group, decreases across a period

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11
Q

what is the trend of first ionization energy

A

decreases down a group and increases across a period

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12
Q

what is the trend of first electron affinity

A

becomes less negative down a group and become more negative across a period

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13
Q

what is an orbital

A

a region in space where there is a given (usually 95%) probability of finding a particular electron
cannot specify the definite or exact position of an electron in an atom

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14
Q

what do quantum numbers do

A

identify the various energy levels available with the atom in which the electrons can reside
they are identification numbers “address” for each electron in an atom
- they specify the position (or location) of an electron in an atom
- predict the direction of spin or rotation of the electron
- determine the energy and angular momentum of an electron

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15
Q

name the four quantum numbers that are required to specify the character of an electron

A
  • principal quantum numbers (n)
  • azimuthal or subsidaiart quantum numbers (l)
  • magnetic quantum numbers (m)
  • spin quantum numbers (s)
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16
Q

in poly-electron atoms what are principal energy levels called and referred to by

A

they are called shells and referred to by letter - K, L, M….

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17
Q

what principle quantum number do principal energy levels correspond to

A

the principal quantum number ‘n’

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18
Q

what does the K equal

A

n = 1

19
Q

what does the L shell equal

A

n = 2

20
Q

which shell doesn’t contain multiple orbitals

A

n = 1

21
Q

what does ‘n’ indicate

A

DISTANCE OF ELECTRONS FROM NUCLEUS
higher the ‘n’ the further away the electron from the nucleus
ENERGY OF ELECTRONS
higher the ‘n’ higher energy electrons - less tightly held so easier to remove

22
Q

what is the rule for the maximum number of electrons a shell can hold
give examples

A 
2(n^2)
K - n = 1 so max = 2 
L - n = 2 so max = 8
M - n = 3 so max = 18
N - n = 4 so max = 32
23
Q

what is a subshell

A

a group of orbitals with the same energy

24
Q

what quantum number describes a subshell

A

azimuthal or subsidiary quantum number ‘l’

25
Q

how does the value of ‘l’ depend on the value of ‘n’

A

‘l’ equals all values from 0 to (n-1)

e.g. if n = 2 l = 0, 1

26
Q

what subshells do the values of l 0, 1, 2, 3 correspond to

A 
0 = s 
1 = p
2 = d
3 = f
27
Q

what is the maximum number of electrons a given subshell can hold

A 
2(2l+1)
e.g 
l = 0 so no. of electrons = 2
l = 1 so no. of electrons = 6
l = 2 so no. of electrons = 10 
l =  3 so no. of electrons = 14
28
Q

what shape does l = 0 s-orbital represent

A

spherical

29
Q

what shape does l = 1 p-orbital represent

A

dumbbell

30
Q

what shape does l = 2 d-orbital represent

A

more complex

31
Q

what shape does 1 = 3 f-orbital represent

A

still more complex

32
Q

magnetic quantum numbers

A
  • represents the orbitals in a given subshell
  • indicates the direction of a particular orbital relative to the magnetic field/axes
  • does not indicate energy
  • tells us the number of orbitals present in the subshell
33
Q

what does the value of m equal

A

m = -l to +l

e.g. l = 2 so m = -2, -1, 0, 1, 2

34
Q

if m = 0 how many orbitals are in each level

A

one

35
Q

if m = 1 how many orbitals are in each level

A

three - Px, Py, Pz

36
Q

if m = 2 how many orbitals are in each level

A

five d-orbitals

37
Q

if m = 3 how many orbitals are in each level

A

7 f-orbitals

38
Q

how to know what ml =

A

ml is where the final electron lands

e.g 3p^5 has 5 electrons and the 5th electron lands on 0 orbital so ml = 0

39
Q

how to know what ms equals

A

ms equals either -1/2 or +1/2 this depends on whether the final electron is pointing up or down
if the arrow is pointing down ms = -1/2 if the arrow is pointing up ms = +1/2

40
Q 
if
n = 3
l = 2 
ml = 1 
ms = -1/2
identify the electron
A 
n = 3 means energy level 3 
l =2 means d-orbital 
d-orbital has 5 orbitals 
ml = 1 so the electron is in 1 orbital and facing down as ms = -1/2 
so 3d9
41
Q

if an ion is negative are electrons added or taken away

A

added

42
Q

if an ion is positive are electrons added or taken away

A

taken away

43
Q

define diamagnetic

A

no unpaired electrons

44
Q

define paramagnetic

A

at least 1 unpaired electron

Chemistry (10 decks)

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