Left-click: rotate, Mouse-wheel/middle-click: zoom, Right-click: pan, Escape: close
Generic Electron Configuration:
- This is the basic electron configuration that you are taught in Yr 8
- Essentially, the first shell of an atom can hold 2 electrons.
- The second shell can hold 8 electrons.
- The third shell usually holds 8 electrons, however it has space for 10 more.
- The formula $2n^2$ describes the total possible electrons in an electron shell.
- Here, n represents the principal quantum number, or the energy level notation.
Subshells - Electronic Configuration:
-
Electrons are arranged in energy levels or shells
-
Energy of shells increases as distance from nucleus increases.
-
n is the principle quantum number which shows energy
-
The maximum number of each shell is $2n^2$
-
Shells are divided into sub-shells.
-
n = 1 has a single s subshell
-
n = 2 has s and p subshells
-
n = 3 has s, p and d subshells
-
The subshells are:
-
The 's' subshells holds a maximum of 2 electrons.
-
The 'p' subshells holds a maximum of 6 electrons.
-
The 'd' subshells holds a maximum of 10 electrons.
-
The 'f' subshells holds a maximum of 14 electrons
Valency:
- The combining power of an atom, is equal to the number of hydrogen atoms it could combine/displace.
- Valency vs Valence electrons
- Valence def stated above, whereas Valence electrons refer to the number of electrons in an atom's Valence (outer) shell.
- e.g. Nitrogen: Valency of 3(-), but has a number of Valence electrons of 5.
- Principal quantum shells – the shells of an atom.
- e.g. Sodium's principal quantum shells are 2, 8, 1.
- Nuclear attraction: Effective nuclear attraction is derived by substracting the number of inner (core) electrons (which 'shield' the valence electrons) from the nuclear charge of an atom (how many protons in an atoms nucleus)
Electronegativity: a higher density of atoms? ==research this==
'-'block elements:
- To be an 'x'-block element, the outer subshell must be an 'x' shell (x being s, p, d or f)
- E.g. Neon's electronic configuration (using spdf notation) is:
- Here, the last subshell is a p subshell. Thus, Argon is a p-block element.
Recall: Aufbau's principle!!!! There is not a linear direction to the spdf blocks.
This is because, certain subshells dont have (sorta) consistent levels of energy.
E.g. Titanium: $1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{2}$
Here, the 4s subshell has a lower energy than the 3d subshell, thus we write it first.
Isoelectronic species - chemical species that have the same number of electrons/ same electronic structure/ same electronic configuration
- In general, this applies when atoms have the same electron configuration, but varying atomic mass (e.g. S-2 and Ar)
Absorption and Emissions Spectrum:
- Bohr's model describes the existence of absorption/emissions spectra.
- *From the idea that electrons occupy energy levels, and can be excited to higher energy levels
- An absorption spectrum represents different wavelengths of photons being absorbed.
- In comparison, an emissions spectrum represents wavelengths of photons being emitted
- The lines we see in absorption/emission spectrums are called spectra.
- This occurs as electrons absorb energy to be excited to higher energy levels.
- *Note: A spectrum of light is not the only way for electrons to be excited. They can be excited through heat or electrical discharge
- Here, electrons only choose photons with a specific frequency to excite them, and when they drop back to the ground state, they release photons.
- Note that all energy levels of an atom have specific, discrete energy quantities.
- Here, the energy of the photons is equal to: the difference between the two energy levels that an electron falls from.
- This means multiple photons can be produces from a single electron being excited.
- e.g. Lets say that an electron is excited from the ground state to n = 3. There are 2 ways this electron can fall down, and 3 possible photons that can be produced.
- From n = 3, straight to n = 1 (ground state). This will produce a single photon, with the energy of the difference in energy between n = 3 and n = 1.
- From n = 3, to n = 2, then n = 1. This produces 2 photons, each varying in energy (diff. between n = 3 and 2, and another with diff. between n = 2 and 1)
- What do the 2 spectrums look like?
- Absorption spectrum: Black lines on coloured background.
- Emissions spectrum: Coloured lines on black background.
- Absorption and emissions spectrum (of the same species) will always superimpose exactly.
- Every elements spectrum is unique!
- In a spectrum, the intensity of a spectra indicates the abundance of an element.
- We use this to identify concentration in a mixture.
- Why is light absorbed by electrons?
- Einstein proposed that light behaved as a wave.
- In his theory, the energy of a photon would be related to the frequency of the photon's (electromagnetic) wave.
- Presently, we known that this relation is derived from the Planck constant
Very minor details about absorption/emissions spectrum:
- Any photons produced by an excited electron dropping back from n = 2 to n = 1 are not visible.
- In general, the visible light spectrum for a hydrogen atom is called the Balmer series.
- The Balmer series does not describe any other atoms, but the principle is the same. ==asking pranav==
Good Resources to Review:
- any worksheets given by Ms Pilling
- ==research similarities / differences of flame test to AAS==
Related Class Notes:
