Atomic structure

Define isoelectronic

Isoelectronic is defined as atoms / ions with the same number of electrons.

Define isotonic

Isotonic is defined as atoms / ions with the same number of neutrons.

Define isotopes

Isotopes are atoms of the same element with a different number of neutrons. They have the same atomic number but different mass numbers

Properties of isotopes

1. Isotopes have the same chemical properties but different physical properties

2. Isotopes have different nucleon numbers, relative atomic masses and number of neutrons

3. Isotope with more neutrons = Heavier isotope

4. Isotopes of some elements are radioactive as the nuclei of these atoms are unstable and breakdown spontaneously, emitting radiation which may be alpha particles, beta particles or gamma rays

Uses of isotopes

• Biochemical tracers in nuclear medicine

• Used as chemical clocks in geological and archeological dating

• PET (positions emission tomography) scanners give 3d images of tracer concentration in the body and is used to detect cancers

• Determination of relative isotopic masses, relative molecular masses and structural features of organic compound as well as identification of unknown compounds, e.g. in forensic science

What is mass spectrometry

• A_r of an element is the weighted average atomic mass of its naturally occurring isotopes

• Modern mass spectrometry determines A_r from its isotopic composition

• Each peak in mass spectrum indicates an isotope of the element and m/z values can be used to identify isotopes

• Height of peak in mass spectrum indicates relative abundance of the isotopes

Define electromagnetic spectrum

Electromagnetic spectrum is termed as the arrangement of all the electromagnetic radiations in increasing order of wavelengths or decreasing order of frequencies

Frequency of radiation equation

c = λf

c = speed of light
λ = wavelength
f = frequency

What is E = hf

Energy of a photon and its wavelength
E = energy
h = Planck’s constant (in data booklet)

f = frequency. can be obtained from c = λf
c = speed of light
λ = wavelength

What is the Bohr model

The Bohr model depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus. Energy of the electron in an atom is quantized.

Define quantization

Quantization means that a quantity cannot vary continuously to have any arbitrary value but can change only discontinuously to have specific or discrete values

Define ground state

Ground state is the lowest energy state of an electron and electrons in this energy level are closest to the nucleus

Define hydrogen emission spectrum

Hydrogen emission spectrum is a line spectrum and the lines in the spectrum correspond to the photons of a particular wavelength (Frequency)

How does hydrogen emission spectrum work

• When a sample of H₂ (g) at low temperature is subjected to high amount of energy (Electricity / Heat), the atoms will be able to emit electromagnetic radiation

• Electron in the hydrogen atom can absorb energy and undergo an electronic transition (jump) to higher energy level

  • Electrons at higher energy level are said to be “excited”

• When electrons return to a lower energy level. energy is emitted as light / photon

  • E.g. Electron goes from n=5 to n=4, then n=2 then n=1 or n=3 to n=1

• Electron cannot change its energy in a continuously way, only in discrete amounts

• An emitted photon from the atom on passing through a very thin slit (diffraction grating) and then through a prism

  • Emission lines correspond to photons of energies that are emitted when the excited electrons transit from higher energy levels (n= 4,5,6) to a lower energy level (n = 1,2)

  • Consists of four colored lines separated by dark bands

    • Four lines are 656.3 nm (red), 486.1 nm (blue-green), 434.0 nm (Indigo) and 410.1 nm (violet)

  • This set of lines are found in the visible region and are associated with electron transitions to n = 2 energy level

• There are other lines in the ultraviolet and infrared red region, associated with electronic transitions to n=1 and n=3 energy level respectively

• Each series in the hydrogen emission spectrum consists of lines of fixed frequencies

• Shows quantized energy of electron, having a fixed amount of energy as the electron is in a discrete energy level

• These energy levels are characterized by a whole number called the principle quantum number, n.

Applications of emission spectrum

• For each element, the emission spectrum is characteristic and unique to the element and it can be used to identify the element

• Emission spectrum provides crucial evidence for the existence of electrons in discrete energy levels which converge at higher energies (frequencies)

Define principal quantum number

Principal quantum number (n) determines the energy and average distance of an electron from the nucleus. The main energy level is given an integer number and can hold up to 2n² electrons

Shape, directionality, number of orbitals in s sublevel

Spherical
Non directional
1

Shape, directionality, number of orbitals in p sublevel

Dumbbell shape
Directional
3 per quantum shell excluding n = 1

Number of orbitals in d sublevel

5

Number of orbitals in f sublevel

7

Define degenerate orbitals

Orbital with the same energy are called degenerate

When are electrons stable

Electrons are stable when the total attractive interactions are greater than the total repulsive interactions

What is aufbau principle

Aufbau Principle: Electrons are added progressively to the orbitals starting with the lowest energy

What is pauli exclusion principle

Pauli Exclusion Principle: Each orbital can hold a maximum of 2 electrons. Paired electrons can only be stable when they have opposite spin so magnetic attraction from opposite spin counterbalances the electrical repulsion

What is hunds rule

Hund’s Rule: When filling a sub-level, each orbital must occupied singly before they are occupied in pairs

Define electronic configuration

Electronic configuration of an atom or ion refers to the arrangement of electrons in its main energy levels, sub-levels and orbitals.

Exceptions to aufbau principle

Energy difference between 3d and 4s sublevels are small, leading to exceptions among heavier transition elements

• Fully filled and half filled orbitals have extra stability, leading to p³, p⁶, d⁵, d¹⁰, f⁷, f¹⁴ configuration to have extra stability

• Hence for copper / chromium, one of the 4s electrons enters the 3d orbital to become half or fully filled, leading to greater symmetry and hence greater stability

Why are half filled/fully filled degenerate orbitals more stable

• Half filled / Fully filled degenerate orbitals have a greater number of exchanges and hence a larger exchange energy of stabilization

• The exchange refers to the shifting of electrons from one orbital to another within the same sub-level

• E.g. 3d⁴4s² vs 3d⁵4s¹, in 3d⁴, there are 6 electrons exchanges / six possible arrangments with parallel spin, while in 3d⁵ there are 10 electron exchanges, hence there is higher total amount of electron exchanges, lending it relatively greater stability

Define cation

Formed by loss of electrons

Define anion

Formed by gain of anions

What is principal quantum number

Principal quantum number (n) represents the main energy levels, or shells, the electrons can occupy in an atom, with a whole number value.

What is secondary quantum number

Secondary quantum number (l), represents sub-levels and it has value ranging from 0 to n-1 and the letters s, p, d, f. l describes the shape of the orbital

What is magnetic quantum number

Magnetic quantum number (m_l) describes the amount of energy levels in a sub-level and has value -l to +l. It determines the number of orbitals and their orientation within a sub-level.

What is spin quantum number

Spin quantum number (m_s) represents electron spin and has values +1/2 or -1/2 . The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers (n,l,m_l,m_s)

Define 1st ionisation energy

1st Ionization energy is defined as the minimum energy required in removing 1 mole of valence electrons from 1 mole of gaseous atoms to form 1 mole of singly positively charge gaseous ion

Factors affecting ionisation energy

1. Nuclear Charge

   1. Nuclear Charge is the total charge of protons in the nucleus. (same number as protons) Given atoms or ions with the same number of filled quantum shells, the greater the nuclear charge, the greater the ionization energy

2. Shield effect by inner electrons

   1. Electrons in inner shells repel valence electrons, giving rise to a shielding effect. Shielding effect is usually given by number of electrons in inner quantum shells

   2. Greater shielding effect means lower ionization energy

What is effective nuclear charge

Effective nuclear charge (Z_eff) is difference between nuclear charge and shielding effect. This gives measure of the actual nuclear charge experienced by electrons in the valence shell, hence how tightly the electrons are attracted to the nucleus.

Trend in successive ionisation energies

• Successive ionization energies of an atom increase with the removal of each electron

• This is due to increase in successive ionization energy is due to increasing amount of energy required to remove successive electrons from an increasingly positive ion due to increasing electrostatic force of attraction between nucleus and valence electrons

Trends in 1st ionisation energy

Across the period

• Increase in effective nuclear charge as nuclear charge increases but shielding effect remains relatively constant since inner quantum shells of electrons remain the same

• Valence electrons are drawn closer to the nucleus so electrostatic forces of attraction between valence electron and nuclei increase

• More energy required to remove a valence electron from the gaseous atom

Down the group

• Number of filled quantum shells increase

• Size of atoms increase

• Valence electrons occupy energy levels increasingly further from nucleus

• Increased distance reduces electrostatic forces of attraction between protons in nuclei and valence electrons

• Electrostatic forces of attraction of the nucleus for the valence electrons decreases and less energy is required to remove them

Deviations from trend of 1st IE

1. First Ionization energy of aluminum is lower than that of magnesium

   1. Less energy required to remove 3p electron in Al than 3s electron in Mg as 3p sub-level is of higher energy than 3s and 3p is further from nucleus

   2. 3p also experiences shielding from 3s electrons

   3. Weaker electrostatic forces of attraction from nucleus

1. First Ionization energy of sulfur is lower than phosphorus

   1. S has 2 electrons occupying same 3p orbital

   2. Inter-electronic repulsion cause less energy to be required to remove a paired 3p electron from S compared to unpaired 3p electron from P