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Chemistry > Atomic Structure > Flashcards

Atomic Structure Flashcards

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1
Q

What is the top number when representing an element?

A

Mass number/nucleon number

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2
Q

What is nucleon number/mass number?

A

Total number of protons and neutrons in the nucleus of an atom

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3
Q

What is the bottom number when representing an element?

A

Atomic number/proton number?

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4
Q

What is atomic number/proton number?

A

number of protons in the nucleus of an atom

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5
Q

What is the characteristics of an atom?

A

It is electrically neutral, it has the same number of electrons and protons

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6
Q

What are isotopes?

A

Atoms of the same element with the same number of protons but different number of neutrons

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7
Q

What are isotopes of the same element similar in?

A
  • Number of electrons
  • Number of protons/ Atomic number
  • Chemical Properties
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8
Q

What do isotopes of the same element differ in?

A
  • Number of neutrons
  • Mass number
  • Relative atomic mass
  • Physical properties
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9
Q

What is isotopic?

A

Atoms or ions that contain the same number of protons

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10
Q

What is isoelectronic?

A

Atoms or ions that contain the same number of electrons

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11
Q

What is isotonic?

A

Atoms or ions that contain the same number of neutrons

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12
Q

What is the extent of deflection directly proportional to?

A

Charge/mass where the mass is the mass number

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13
Q

What happens to the extent of deflection when the charge is greater?

A

When the charge is greater, the extent of deflection is greater because of the stronger electrostatic forces of attraction that pulls it downwards/upwards towards the charge plate

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14
Q

What happens to the extent of deflection when the mass is greater?

A

The extent of deflection is smaller because the heavier ion have larger inertia and is harder to pull it downwards to the negatively charged plate

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15
Q

What is an orbital?

A

A region of space where there is a high probability density or electron density

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16
Q

What is the chance of finding an electron in an orbital?

A

95%

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17
Q

What are the characteristics of s orbital?

A
  • Each s subshell contains only one orbital
  • Spherical shape
  • Non-directional nature
  • Equal probability of finding an electron at any given distance
  • 0 probability at the nucleus
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18
Q

What are the characteristics of p orbital?

A
  • each p subshell has 3 degenerate orbitals (px, py, pz)
  • dumbbell shape
  • symmetrical along axis
  • directional nature
  • 0 probability at the nucleus
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19
Q

What are the characteristics of d orbital?

A
  • each d subshell has 5 degenerate orbitals
  • gives transitional metals their special properties
  • 0 probability of finding an electron along dxy, dxz, dyz axes
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20
Q

What is the maximum number of electrons each orbital can accommodate?

A

2

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21
Q

What happens to the size of the orbital when it is of a higher energy?

A

Increases

22
Q

What does a larger n value (principal quantum number) indicate about the energy?

A

A larger n value indicates a larger orbital which is further away from the nucleus and hence have a higher energy

23
Q

Within a subshell, what energy do orbitals have?

A

Degenerate (same energy)

24
Q

What is the relative energies of the orbitals?

A

1s (lowest energy), 2s, 2p, 3s, 3p, 4s, 3d (exception), 4p

25
Q

What is the aim of orbital filling?

A

Achieve an electronic configuration of the lowest possible energy at the ground state

26
Q

What are the 3 rules in orbital filling?

A
  • Aufbau principle
  • Pauli Exclusion Principle
  • Hund’s Rule
27
Q

What is Aufbau Principle?

A

Electrons reside in the orbital of lowest possible energy first

(draw the arrow diagram)

28
Q

What is the Pauli Exclusion Principle?

A

an orbital can accomodate a maximum of 2 electrons and they must be in opposite spins

29
Q

What is the Hund’s rule?

A

Electrons occupy degenerate orbitals singly and with the same spin before pairing occurs

30
Q

What is the ground state?

A

The most stable and lowest energy state

31
Q

What are the exceptions to orbital filling?

A

Cr and Cu because half-filled or fully-filled d subshell tend to be more stable

32
Q

What indicates the number of principal quantum shell?

A

Period number

33
Q

What is the number of valence electrons equal to?

A

Group number

34
Q

How to find group number using spdf notation?

A
  • s-block: sum of valence electrons (sum of number of electrons with the highest principal quantum shell number)
  • p-block: sum of valence electrons + 10
  • d-block: sum of d + s electrons
35
Q

Where are the different blocks located?

A
  • s is left two
  • d is middle
  • p is right
  • f is bottom 2 that is separated
36
Q

What do elements in the same group have?

A

The same valence shell electronic configuration

37
Q

How are electrons removed in cation formation?

A

Electrons are removed from the outermost shell (4s removed before 3d cause 4s is outer)

38
Q

How are electrons added in anion formation?

A

Electrons are added to the empty/partially filled outermost orbitals

39
Q

What is ionisation energy?

A

Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms or ions (for 1st IE it’s atoms, for 2nd IE it’s ions)

40
Q

What is the equation for 1st IE?

A

X(g) → X⁺(g) + e⁻
ΔH = 1st IE of X

41
Q

What are the factors that affect ionisation energy?

A
  • effective charge (nuclear charge and shielding effect)
  • atomic radius
42
Q

What is the formula for effective nuclear charge?

A

Zeff = Z - shielding constant

Z: nuclear charge

43
Q

What affects nuclear charge, Z and what are the implications?

A
  • The larger the number of protons, the greater the nuclear charge
  • as nuclear charge increases, the attraction on electrons increases and more energy is required to remove the valence electron
44
Q

What affects the shielding constant and what are the implications?

A
  • The larger the number of inner principal quantum shell of electrons, the greater the magnitude of shielding constant, the greater the shielding effect
45
Q

What affects atomic radius?

A
  • as number of filled principal quantum shells increases, atomic radius increases
  • as effective nuclear charge increases, atomic radius decreases
46
Q

How does a larger atomic radius related to ionisation energy?

A
  • A larger atomic radius implies that valence electrons are located further away from the nucleus
  • as a result, they experience weaker attractive force from the nucleus and can be removed more easily, hence a lower IE
47
Q

How does a larger effective nuclear charge affect ionisation energy?

A

a larger effective nuclear charge implies that the valence electrons are held more strongly to the nucleus and require more energy to be removed

48
Q

What is the trend of 1st IE across a period?

A
  • 1st IE generally increase across a period
  • across a period, nuclear charge increases while there is a negligible increase in shielding effect as electrons are removed from the same shell, thus effective nuclear charge increases
  • atomic radius also decreases
  • more energy is required to remove the valence electrons that are increasingly more attracted by the positive nucleus
49
Q

What are the anomalies of the trend of 1st IE across a period?

A
  • 2p electrons to be removed from B has a higher energy than 2s electrons to be removed from Be
  • Coulombic repulsion between paired 2p electrons in O makes it easier to remove one of the paired 2p electrons than an unpaired 2p electron from N
50
Q

What is the trend of 1st IE down a group?

A
  • 1st IE decreases down a group
  • down a group, atomic radius increases
  • the increase in shielding effect due to increase in number of principal quantum shells is largely cancelled out by the increase in nuclear charge
  • less energy is required to remove the valence electrons that become increasingly less attracted by the positive nucleus
51
Q

How to describe trend of successive ionisation energy and how to use it to predict group number?

A
  • there is an increase in successive IE as the electron is removed from an ion of increasing positive charge
  • first two ionisation energies increase gradually as they only involve the removal of outermost 3s electrons only
  • sharp increase from 2nd to 3rd IE since it involves the removal of 2p electrons from an inner principal quantum shell
  • only two electrons in outermost shell, group II element
52
Q

Why is there a sharp increase in IE?

A
  • The electron is removed from an inner principal quantum shell which is located nearer to the nucleus
  • The electron also experiences lesser shielding effect and greater effective nuclear charge, hence it is more strongly attracted to the nucleus and will require a larger amount of energy to remove

Chemistry (21 decks)

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