Intro to Atomic Structure, Periodicity and VSEPR
25 September 2017 15:24
ATOMIC STRUCTURE
Chemistry is the branch of science If electrons are shared absolutely equally e.g. as in Cl2, then a pure covalent bond exists.
concerned with the substances of which If for a compound XY an electron were completely transferred from Y to X giving Y+ and X- then purely ionic bonding exists.
matter is composed, the investigation of • Pure covalent and pure ionic bonding represent extremes, and something between the two is usually found. The larger
their properties and reactions, and the the difference in electronegativity between the two values in a bond X-Y, the more ionic the bond type will be.
use of such reactions to form new
substances. • In the early 1900s, it was known that atoms contained protons,
neutrons and electrons, but little was known of atomic structure.
Chemical bonding involves the sharing • The assumption was that Newtonian Mechanics held true for atoms
(covalent) or donating/receiving (ionic) of as it did for larger objects.
electrons between two or more atoms. ○ However, this didn't explain the atomic spectrum of hydrogen
as when you expose hydrogen gas to an electric discharge
Rydberg Equation there are missing areas where it has absorbed the light.
• Johannes Rydberg was a Swedish physicist
who found that the energies (or frequencies, v) of When an atom in an excited state n2 emits a photon and falls to a lower energy
the hydrogen lines had a pattern. state n1 the energy lost is given by the Rydberg equation.
• He showed that they could be predicted using the e.g if an electron transitions from excited state n=3 to ground state n=1 you
formula below (the RYDBERG EQUATION): input n1=1 and n2=3 into the Rydberg equation.
To convert cm-1 into m-1, multiply by 100
Emission of radiation takes place when an electron makes a transition from a
13.6 eV or 109,678cm-1. The wavenumber state of energy and the difference which is equal to
is the number of wavelengths in a given
distance e.g. 1cm-1 is one complete is carried away as a photon of energy
wavelength in 1cm.
By equating these two energies we obtain the Rydberg equation.
The Rydberg Equation correctly predicted the position of the spectral lines in the absorption spectrum for
electron transitions within the hydrogen atom from n2 to n1 (n=1 is the innermost electron shell), but the
number of times reasons for its success were not known. There is zero energy when n = infinity as the electron and nucleus
per second that Lyman series, n1=1, ultraviolet
Balmer series, n1=2, visible region are widely separated and stationary.
a wave travels a
complete cycle. Paschen series, n1=3, infrared region
Brackett series, n1=4, infrared region Bohr
1Hz = 1s-1 • Following quantum theory, Danish physicist Niels Bohr proposed a model of the hydrogen atom
(the Bohr atom) in 1913.
Planck
○ Central nucleus with electrons around it
• Max Planck was a German physicist that observed that atoms
○ Electrons are only allowed to orbit at specific radii from the nucleus
and molecules emit energy only in certain discrete quantities
○ Defined energy levels (orbitals) for electrons as distances from the central nucleus.
when heated.
• Energies of orbits given by:
○ "quanta" - packets of energy
○ Electrons in atoms occupy specific or discrete energy levels
• Bohr's orbits explain the success of the Rydberg equation in predicting the atomic spectra of
○ These energy levels are said to be quantized.
certain elements, but could not explain the spectra of many other species.
• Quantum theory suggested that particles at the atomic scale
• The model was also unable to explain why only certain energy levels existed.
behave differently than larger species
• Model doesn't work very well for heavier elements (elements with more than one electron).
De Broglie
• French physicist Louis de Broglie suggested that all matter • According to de Broglie, electrons behave as standing waves.
possesses characteristics of both waves and particles, which led ○ Like waves created when plucking a guitar string, a
to the development of wave mechanics (quantum mechanics). standing wave appears to not be moving when it is. Other
• Evidence exists that electrons have wave-particle duality (wave energies are not allowed, explaining the quantisation of
properties and particle properties). energy levels.
○ They have mass and can be weighed, which is a particle
property The Schrödinger Equation
○ They can be diffracted by a crystal lattice, which is a wave • Used complex mathematics to form the wave equation.
property. • The wave nature of the electron in hydrogen (the location
• Light also possesses both particle and wave properties. and properties of electrons in atoms e.g. how fast it is
• This means that it is impossible to know the linear momentum travelling) can be described by solving Schrodinger's wave
(mass x velocity) and the location of an electron simultaneously. equation to get a wave function psi.
De Broglie equation:
• Duality is difficult to rationalise but it is
only important for very small entities
such as atomic particles or light. • Solution of Schrodinger's wave equation gives
which corresponds to energy levels
• On the left there is a standing wave as the • Each of the wave functions is called an orbital,
line has a continuous pattern with no break. and describe the nature of electrons in orbitals.
• On the right the wave is not a standing • The wave functions themselves have no established physical meaning,
wave as the wave doesn't overlap/doesn't however psi squared is said to represent the probability of finding an electron
superimpose each time in orbit. Each time it at a specific point.
goes round its position moves.
Implications of wave/particle duality
Quantum numbers • Atomic orbitals have to be considered as regions of space where an electron is likely to be
• From the mathematical solution of the wave equation we can define found (the wave function describes atomic orbitals). They are viewed as electron clouds:
three quantum numbers, n, l and ml, where each quantum number ○ Giving the probability (rather than certainty, due to the wave particle duality) of the
specifies a physical property of the electron. electron being at any given point.
○ n = principal quantum number (energy of the electron). Also ○ This defines the shapes of the orbitals.
indicates orbital size - the larger n high-energy orbitals diffuse • s-orbitals are spherically symmetrical.
more than low-n low-energy orbitals. ○ 1s-orbital (n=1, l=0, ml=0)
○ l = orbital angular momentum quantum number (magnitude of ○ 2s-orbital (n=2, l=0, ml=0) etc
the orbital angular momentum - the energy an electron has • The density of the electron clouds at any point are given by psi squared.
when moving in circular orbit) may be understood as the probability of the electron being at a given point.
○
m = magnetic quantum number (spatial orientation of the
Inorganic Chemistry I Page 1
