HL IB Chemistry - Electronic Configurations
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What does a huge increase in ionisation energy between the first and second ionisation of sodium indicate about its electron configuration?

  • It does not have any core electrons.
  • The first electron is in the valence shell. (correct)
  • Its first electron removed is from a full subshell.
  • It has three valence electrons.

Which of the following statements is correct regarding magnesium's ionisation energy?

  • It belongs to group II of the periodic table. (correct)
  • It has three valence electrons.
  • There is a small jump from the second to the third ionisation energy.
  • The second ionisation energy is greater than the first.

What can be inferred about aluminium based on its third to fourth ionisation energy jump?

  • It has one valence electron.
  • The jump reflects the removal of a 3p electron. (correct)
  • The fourth electron removed is significantly easier to take than the others.
  • Aluminium has a stable octet configuration.

Which element has the following electron configuration: 1s2 2s2 2p6 3s2 3p1?

<p>Aluminium (D)</p> Signup and view all the answers

What is the group classification of sodium based on its electron configuration?

<p>Group I (A)</p> Signup and view all the answers

Why is the third ionisation energy of magnesium significantly higher than the second?

<p>It indicates a jump to higher energy subshell. (C)</p> Signup and view all the answers

What characteristic of ionisation energy is reflected in the trends observed for aluminium?

<p>There is a predictable pattern based on subshells. (B)</p> Signup and view all the answers

How does the electron configuration influence the ionisation energies of the elements listed?

<p>Electrons in outer shells are easier to remove than those in filled shells. (A)</p> Signup and view all the answers

What principle states that electrons occupy the lowest energy subshells first?

<p>Aufbau Principle (C)</p> Signup and view all the answers

Which of the following correctly describes the maximum occupancy of an atomic orbital?

<p>Two electrons (C)</p> Signup and view all the answers

The energy levels of orbitals in an atom do not follow a regular pattern beyond which principal quantum number?

<p>n = 3 (A)</p> Signup and view all the answers

Which of the following shapes corresponds to the p orbitals?

<p>Dumbbell (A)</p> Signup and view all the answers

Which statement about the ground state of an atom is true?

<p>It is the most stable configuration at low energy. (C)</p> Signup and view all the answers

What characterizes the d orbitals in relation to the IB Chemistry curriculum?

<p>Their specific shapes are not required for study. (A)</p> Signup and view all the answers

Which of the following correctly describes the role of subshells in ionization energies?

<p>Higher subshells have higher energy and thus lower ionization energy. (C)</p> Signup and view all the answers

Identifying electron configuration helps predict which atomic property?

<p>Ionization energy (B)</p> Signup and view all the answers

What does the principal quantum number (n) indicate about an atom’s electron configuration?

<p>It determines the energy of the electrons in the shell. (A)</p> Signup and view all the answers

Which principal quantum number can hold a maximum of 8 electrons?

<p>n = 2 (A)</p> Signup and view all the answers

What is the maximum number of electrons that can be held in the subshell represented by 'd'?

<p>10 (A)</p> Signup and view all the answers

Which of the following correctly describes the order of energy levels in subshells?

<p>s &lt; p &lt; d (A)</p> Signup and view all the answers

How many total electrons can be held in the principal quantum shell with n = 4?

<p>32 (C)</p> Signup and view all the answers

Which of the following statements about subshells is true?

<p>The energy of the subshells increases in the order s &lt; p &lt; d &lt; f. (B)</p> Signup and view all the answers

What mathematical relationship defines the maximum number of electrons in a principal quantum shell?

<p>2n^2 (A)</p> Signup and view all the answers

Which of the following correctly identifies the pattern of subshell filling across principal quantum shells?

<p>Subshells fill in the order of increasing energy, but this can vary with n. (D)</p> Signup and view all the answers

Flashcards

Atomic Orbitals

Specific energy levels where electrons exist in an atom, with fixed 3D shapes.

Subshells

Groups of atomic orbitals within a specific energy level (shell).

Electron Capacity of Orbitals

Each atomic orbital can hold a maximum of two electrons.

Ground State

The stable lowest energy arrangement of electrons in an atom.

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Aufbau Principle

Electrons fill the lowest energy subshells first when forming an atom.

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s orbitals

These are spherical-shaped atomic orbitals and are the first type of orbital filled.

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p orbitals

These orbitals have specific 3D shapes along the x, y, and z axes.

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Energy Level

The energy level specified for specific orbitals in atoms.

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What is electronic configuration?

The arrangement of electrons in an atom, specifying the energy levels and sublevels they occupy.

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What are principal quantum shells?

Energy levels surrounding the nucleus, with higher numbers indicating greater energy.

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Principal quantum number (n)

A whole number that designates a specific energy level or shell, with lower numbers closer to the nucleus.

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Electron capacity of shells

The maximum number of electrons each shell can hold, determined by the formula 2n^2.

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Subshell energy order

The order of increasing energy for subshells is s < p < d < f.

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Subshell overlapping

The subshells of higher principal quantum numbers can have lower energy than those in lower shells.

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What are orbitals?

Specific regions within a subshell where electrons are most likely to be found, having fixed 3D shapes.

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Ionisation Energy

The minimum energy required to remove one mole of electrons from one mole of gaseous atoms, forming one mole of gaseous ions.

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First Ionisation Energy

The energy required to remove one mole of electrons from one mole of gaseous atoms, forming one mole of gaseous ions with a +1 charge.

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Second Ionisation Energy

The energy required to remove one mole of electrons from one mole of gaseous ions with a +1 charge, forming one mole of gaseous ions with a +2 charge.

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Trends in Ionisation Energy

Ionisation energy generally increases across a period and decreases down a group. This is due to factors like nuclear charge, shielding and atomic radius.

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Large Jump in Ionisation Energy

A significant increase in ionization energy indicates a change in the shell from which the electron is being removed.

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Group Number and Ionisation Energy

The number of large jumps in ionisation energy corresponds to the group number of the element. This is because it becomes much harder to remove electrons after removing all the valence electrons.

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Sodium's Ionization Energy

Sodium has a large jump between the first and second ionization energy, indicating that it has one valence electron and belongs to Group 1.

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Magnesium's Ionization Energy

Magnesium exhibits a large jump between the second and third ionization energy, indicating that it possesses two valence electrons and resides in Group 2.

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Study Notes

HL IB Chemistry - Electronic Configurations

  • Electromagnetic Spectrum:

    • A range of frequencies covering all electromagnetic radiation, with related wavelengths and energy.
    • Divided into bands, important in analytical chemistry.
    • Shows the relationship between frequency, wavelength, and energy.
    • Frequency is the number of waves passing per second; wavelength is the distance between consecutive peaks.
    • Gamma rays, X-rays, and UV radiation are high-energy, high-frequency, and dangerous.
    • All light waves travel at the same speed, differing only in frequency.

  • Emission Spectra:

    • Electrons orbit the nucleus in energy shells.
    • Increased energy allows electrons to jump to higher energy levels.
    • Electrons return to original energy levels emitting energy.
    • Emitted energy frequency is the same as absorbed energy, just emitted instead of absorbed.
    • Emitted energy in the visible region can be analyzed using a diffraction grating.
    • Resulting pattern is a line emission spectrum.
    • The spectrum of hydrogen shows specific frequencies of light (ie., quantized).

  • Continuous vs. Line Spectra:

    • Continuous spectrum displays all frequencies of light (like a rainbow).
    • Line spectrum displays only specific frequencies of light, indicating quantized energy levels.

  • Speed of Light:

    • Constant (3.00 x 10⁸ ms⁻¹).
    • Inverse relationship between frequency and wavelength (c = fλ).

  • Energy Levels, Sublevels & Orbitals:

    • Electrons arranged in principal energy levels (shells, numbered by principal quantum numbers, n).
    • Lower n value, closer to nucleus and lower energy.
    • Higher n value, farther from nucleus, and higher energy.
    • Each shell has a fixed capacity of electrons.
    • (n = 1, up to 2 electrons; n = 2, up to 8 electrons; n = 3, up to 18 electrons; n = 4, up to 32 electrons).
    • Shells split further into subshells (s, p, d, f).
    • Subshells contain atomic orbitals.
    • s orbitals are spherical; p orbitals are dumbbell-shaped.

  • Electron Configuration:

    • Describes the arrangement of electrons in shells, subshells, and orbitals.
    • Filled in order of increasing energy.
    • Each orbital can hold two electrons with opposite spins (Pauli Exclusion Principle).
    • Hund's Rule states that orbitals in a subshell are singly occupied before any one orbital is doubly occupied.
    • Electron configurations can be represented in full form or shorthand (using noble gas notation).
    • Different blocks in the periodic table correspond to different subshells (s, p, d, f).

  • Ionisation Energy from an Emission Spectrum (HL):

    • Electrons move between shells, absorbing or emitting energy.
    • When an electron falls back to the ground state, it emits a photon corresponding to the energy difference between levels (n = 1 to n = ∞).
    • Measuring the wavelength/frequency of emitted photons reveals the differences in energy levels in the atom.

  • Successive Ionisation Energies (HL):

    • Successive Ionisation Energies (IEs) of elements increase as electrons are removed.
    • IEs increase as fewer electrons remain and the net positive charge of the ion increases.
    • There are notable jumps in IE corresponding to changes in electron shells or subshells.
    • IEs provide information about electronic configurations and shell structure.

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HL IB Chemistry Electronic Configurations PDF

Description

Explore the intricacies of electronic configurations and the electromagnetic spectrum in this HL IB Chemistry quiz. Understand the relationship between frequency, wavelength, energy, and emission spectra. This quiz will enhance your grasp of analytical chemistry concepts and the behavior of electrons within atom energy shells.

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