Unit 1 - Structure and Properties of Matter Flashcards
Democritus (460 BCE)
Experiment: It was based on intuition, not experimentation.
Theory: Matter is composed of elementary particles called atoms that are indivisible.
Limitation: There was no experiments to back his findings.
John Dalton (1766-1844)
Experiment: He found that gas molecules fit in between the spaces of liquids when they dissolve (solubility of gases).
Theory:
- Atoms are indivisible, indestructible spheres.
- Atoms of the same element are identical.
- Compounds are formed by joining two or more atoms.
- There are large spaces between atoms/molecuies.
Model: Billiard Ball Model of the Atom
Limitation: Scientists discovered that there was something smaller than the atom (i.e. protons, neutrons, electrons).
J.J. Thomson (1856-1940)
Experiment: Cathode Ray Tube
* High voltage was connected to terminals in a vacuum tube.
* The rays that were produced were attracted toward a positive plate.
* The rays were the same no matter what metal composed to cathode (or which gas was in the tube).
Theory:
* Electrons are negatively charged particles present in all elements.
* Different elements have different numbers of electrons embedded in them.
Model: Raisin Bun/Plum Pudding Model - The atom is a positively charged mass with discrete negative particles distributed throughout it.
Limitation: Scientists found particles that could travel through atoms. Therefore, they could not be a solid mass.
Ernest Rutherford (1871-1937)
Experiment: Gold Foil Experiment
* Positively charged α-particles (type of radiation) were shot at a very thin sheet of gold foil.
* A zinc sulfide detection screen showed that most particles went straight through, whereas some were deflected.
Theory:
* Atoms are made up of mostly empty space.
* Atoms contain a small, dense, positively charged “nucleus”.
* Electrons move about in the empty space that makes up the rest of the atom.
Model: Nuclear (Beehive) Model - A positively charged nucleus is encircled with electrons travelling anywhere around it.
Limitation: Moving charges (i.e. electrons) usually emit photons of electromagnetic radiation, therefore losing energy. This should cause their orbit to decrease until they spiral into the nucleus and the atom collapses, but this doesn’t happen.
Niels Bohr (1885-1962)
Experiment: Spectral Line Experiment
- A high voltage was connected to terminals in tubes containing hydrogen gas.
- When the light emitted was viewed through a prism, only four distinct colours of light were seen (red, aqua, blue, purple).
Theory:
- When electricity hits an electron, it gives it some energy; moving it further from the nucleus to an excited state.
- The electron then spontaneously drops down to its normal energy level, called the “ground” state.
- When the electron drops back down, it gives off energy, which can sometimes be seen as a colour of visible light.
Model: Planetary/Solar System Model - A positively charged nucleus with electrons travelling around it, in specific energy levels (like planets around the Sun).
Limitation: While this theory could explain the spectral lines of hydrogen, it could not predict the spectral lines of multi-electron atoms.
spectroscopy
analysis of spectra in order to determine properties of their source
emission spectrum
the spectrum (pattern of bright lines) emitted by an atom
continuous spectrum: an emission spectrum that contains all the wavelengths in a specific region of the electromagnetic spectrum
line spectrum: an emission spectrum that contains only those wavelengths, characteristic of the element being studied; arises when excited electrons emit energy (each coloured bond corresponds to a discrete wavelength)
Heisenburg Uncertainty Principle
- Werner Heisenberg stated that the position and velocity of an electron cannot be measured at the same time because an electron is so small
- by measuring the velocity, the position is disturbed, and vice versa
wave-particle duality
Louis de Broglie stated that sometimes electrons behave like particles, and at other times, they behave like waves
discovery of orbitals
Erwin Schrödinger used the physics of wave mechanics to create complex 3D functions which describe the region in space where electrons are likely found (“orbitals”)
quantum mechanical model
- electrons can be in different orbitals by absorbing or emitting quanta of energy
- the location of electrons is given by a probability distribution
principal quantum number (n)
- describes the size and energy of an atomic orbital
- a higher n value indicates a higher energy level that is larger in size and further from the nucleus, and the electrons are less tightly bound
shell
an atom’s main energy level, where the shell number is given by the principal quantum number (n = 1, 2, 3…)
secondary quantum number (l)
- describes the shape and energy of an atomic orbital
- values between 0 and n-1
s (sharp): l = 0
p (principal): l = 1
d (diffuse): l = 2
f (fundamental): l = 3
subshell
- orbitals of different shapes and energies as given by the secondary principal number; referred to as s, p, d, f
- each orbital can hold a maximum of two electrons
magnetic quantum number (ml)
- describes the orientation of an atomic orbital in space relative to the other orbitals in the atom
- values between -l and +l
spin quantum number (ms)
- relates to the spin of the electron
- values of -1/2 or +1/2
Pauli Exclusion principle
- no two electrons can have the same set of four quantum numbers (n, l, ml, or ms) or be in the same quantum state
- this is due to two electrons in the same orbital needing to have opposite spin
Aufbau principle
an atom is “built up” by the addition of electrons, which fill orbitals starting at the lowest available energy orbital, before filling higher energy orbitals
Hund’s Rule
as you read across each period from left to right, it shows the order in which the sublevels (s, p, d, f) are filled
isoelectronic
- a substance with the same number of electrons as another substance
- e.g. S-2 is isoelectronic with Ar
