Glossary of chem test 4 ch 8
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- are an oscillation that moves outward from a disturbance.
- Wavelength (wl)
- is the distance between identical points on successive waves.
- is the vertical distance from the midline of a wave to the peak or trough.
- Frequency (v)
- is the number of waves that pass through a particular point in 1 second (Hz = 1 cycle/s).
- All electromagnetic radiation
- wl x v = c
- one cycle =
- hertz; s^-1; 1/s
- the longer the wavelength
- the shorter its frequency. a small wave has huge frequency.
- visible frequency
- freq Hz is at about 10^-13
wl NM is at about 10^4
- PLANCK’S QUANTUM THEORY
- In 1900 Planck proposed that atoms and molecules can emit (or absorb) energy only in discrete (separate) quantities. The smallest quantity of energy that can be emitted is called a quantum (photon) whose energy is given by: E = hv
- Below the threshold frequency
- nothing occurs. All metals experience this effect, but each has a unique threshold frequency.
- Above the threshold frequency
- the kinetic energy of the ejected electrons is proportional to the frequency of the light. As intensity of the light increases, so does the number of ejected electrons. All metals experience this effect, but each has a unique threshold frequency.
- Albert Einstein 1905 the Photoelectric Effect II
- extended Planck’s idea and proposed that light itself is quantized. Photons are the quanta of electromagnetic energy.
Photons have an energy equal to:
E = hv
h = Plank’s Constant
- When an element in the gas state is heated or an electric current is passed through it...
- ...a line spectrum is produced. Each element produces a unique line spectrum. Solids produce continuous spectra.
- Maxwell (1873), proposed that
- visible light consists of electromagnetic waves.
- Electromagnetic radiation
- is the emission (absorption) and transmission of energy in the form of electromagnetic waves.
All electromagnetic radiation
wl x v = c
- Bohr's Model of the hydrogen Atom (1913.)e- can only have specific (quantized) energy values
light is emitted as e- moves from one energy level to a lower energy level. n (principal quantum number) = 1,2,3,…
RH (Rydberg constant)
E = RH( 1/n^2i - 1/n^2f)
i= initial energy level
you can never have negative energy. sign just tells you if energy is eing emitted(-) or absorbed (+)
- de Broglie
- 1924. hydrogen electrons behave like waves as well as particles.A vibrating guitar string is quantized. Only certain motions (standing waves) are allowed and only certain frequencies of sound are produced. an electron bound to a nucleus behaves like a standing wave.An allowed energy state is one in which an integral number of waves will fit around the circumference of the orbit. Thus de Broglie suggested that the wave length of a matter wave is given by:
m= mass of particle
- The Heisenberg Uncertainty Principle
- It is impossible to know simultaneously both the momentum and position of a particle with certainty.
- the region in space where the electron is most likely to be found.
- QUANTUM MECHANICS
- In the 1920’s Heisenberg, Schrödinger, and Dirac developed the modern theory of the atom. QM tells us nothing about the path of the electron. Only probability information is given.
- tells the principle energy level of the electron and its average distance from the nucleus; principal quantum #; shell
- (n-1)the sub-level of the electron and the shape of the orbital; azimuthal quantum #; subshell
- It tells the number of orbitals and their orientation in space; magnetic quantum #; orbital (actual loc of e-)
- magnetic spin; +1/2; -1/2; which way e- is moving.
- the total number of orbitals for a given value of n
- is n^2
- e- density/ nucleus dist relationship
- e- density (1s orbital)falls off rapidly as distance from nucleus increases
- l = 0
- (s orbital)
- l = 1
- 3p orbitals; degenerate orbitals- are all equal in energy
- l = 2
- 5 d orbitals; diff energy levels; are not degenerate orbitals
- degenerate orbitals
- are all equal in energy; like p subshell; unlike d subshell
- The Pauli Exclusion Principle
- No two electrons in an atom can have the same
four quantum numbers.
This means that no more than two electrons can exist in
the same orbital and these electrons must have opposite
Note: the first three quantum numbers designate a
particular orbital in the atom.Electrons with the same spin keep apart in space whereas electrons of opposite spin may occupy the same region of space.
- at least one unpaired electron for an atom (arrows)
- no unpaired e- for an atom (arrows)
- Energy of orbitals in a single electron atom depends on (up to 4)
- principal quantum number n
- Energy of orbitals in a multi-electron atom (above 4)
- Energy depends on n and l
- The most stable arrangement of electrons in subshells
- is the one with the greatest number of parallel spins (Hund’s rule).
- Cr: [Ar] 4s1 3d5
Cu: [Ar] 4s1 3d10
Nb: [Kr] 5s1 4d4
Mo: [Kr] 5s1 4d5
Ru: [Kr] 5s1 4d7
Rh: [Kr] 5s1 4d8
Ag: [Kr] 5s1 4d10
Pt: [Kr] 6s1 4f14 5d9
Au: [Xe] 6s1 4f14 5d10
- All metals lose electrons to form
- Transition metal ions
- rarely attain a noble gas configuration.
They always lose the outer s electrons first. Usually they
lose one or more d electrons as well.
- at least 2 species have the same # of e-; same arrangement of e-; same electronic configuration
- atomic size
- measurable quantity
- atomic radii
- decreases to the right;
- is always smaller than atom from which it is formed; positive
- is always larger than atom from which it is formed.
- Ionization energy
- is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state; left to right inc. up to down dec
- Electron affinity
- is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
- inc to the right; dec down (just like ionization) BUT its about the density acquired. the most electroneg is F and the least is Fr (excluding noble gases)
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