Chapter 6 Study Guide
6.1
· Know The Wave Nature of Light
· Speed of light=3.00x10^8 m/s, represented as c.
· Electromagnetic radiation is described in terms of wavelength, λ, and frequency, v. λv=c.
· The smaller the wavelength, the higher the frequency.
· Frequency is expressed in cycles per second, hertz (Hz).
· Know where things are placed on the electromagnetic spectrum (The relative wavelengths + frequencies of Gamma rays, X-rays, infrared, etc.)
6.2
· Max Planck- assumed energy is “quantized,” or only comes in fixed “chunks”. It is restricted to certain quantities.
· To describe the energy of EM radiation, Planck proposed this formula: E=hv.
· E= energy, and h = Planck’s constant = 6.626x10^-34 J-s.
· We don’t notice these quantized “steps” in everyday life because they are so small, but they make an important difference on matter at the atomic level.
· The photoelectric effect: Light shining on metal can cause the surface of that metal to emit electrons. Some metals needed a minimum frequency (or energy) of light for this to happen. Einstein assumed that EM radiation acted as a stream of tiny particles, or packets of energy called Photons.
· Light has a Dual Nature – it has the characteristics of both a wave and particles.
6.3
· Monochromatic: radiation composed of a single wavelength
· When radiation from a source is separated into its different wavelengths, it creates a spectrum
· A spectrum with all wavelengths is a continuous spectrum
· When different gases are excited by electrons, they can emit a line spectrum, or a spectrum containing radiation of only specific wavelengths.
· Bohr made a model of an atom based on the line spectrum of hydrogen.
· Scientists assumed that the atom was like a mini solar system, with electrons orbiting around the nucleus.
· According to classical physics, an electrically charged particle that moves in a circular should continuously lose energy by emitting EM radiation.
· If electrons did this, they would gradually spiral into the nucleus of the atom.
· This didn’t happen. Why not?
· Bohr explained this problem with three hypotheses:
o Electrons are only allowed certain orbits, each with a certain corresponding energy.
o An electron in a “permitted orbit” has a specific energy and is in an “allowed” energy state. – (the energy is quantized, and can only have certain values, so the electron cannot lose the tiny amounts that would make it spiral down into the nucleus.)
o Energy is absorbed or emitted by the electron only as it changes from one energy state to another. This energy is emitted or absorbed as a photon. E=hv.
· A good diagram of this idea is on pg. 226.
· Bohr’s model did not explain the spectra of atoms other than hydrogen, and it did not explain the wavelike properties of electrons.
6.4
· Louis De Broglie believed that matter, such as electrons, that was previously supposed to be particles, could be described as waves.
· Heisenberg created the Heisenberg uncertainty principle. This stated that it is impossible to know an electrons momentum (the speed and direction it is traveling in) and location at the same time.
· Heisenberg created the formula ∆x ∙ ∆(mv) ≥ h/4π
· Where ∆x represents position, ∆(mv) represents momentum, and h is planck’s constant.
· ∆(mv) = m∆v, so m (mass) is an important part of uncertainty.
· This equation describes how the smaller something’s mass is, the harder it is to tell its location. The more mass something has, the less uncertain its location is.
6.5
· A physicist, Schrödinger, created an equation that incorporated both the wavelike and particle-like behavior of electrons.
· It led to describing particles with wave functions.
· In the quantum model based on Schrödinger’s equations, we cannot know where an atom is, so its location is described with probability. Orbitals are described with graphs showing the probability that an electron will be found at a certain location. This probability is called probability density or electron density.
· Quantum numbers describe each orbital.
1. The principle quantum number, n, can have positive integral values: 1, 2, 3… etc. The higher n is, the larger the orbital is, and the higher its energy is.
2. The azimuthal quantum number: l. This defines the shape of the orbital. It can have values from 0 to n-1.
3. The magnetic quantum number ml describes the orientation of the orbital in space. It has values between l and –l.
4. Spin: ms can have a value of -1/2 or +1/2, and allows 2 electrons per orbital.
· Electron shell: the collection of orbitals with the same value of n.
· Subshell: the set of orbitals that have the same n and l values.
· 3 important observations:
1. The shell with principal quantum number n will have n subshells.
2. Each subshell has a specific number of orbitals. For a certain value of l, there are 2l+1 allowed orbitals (or values of ml).
3. The total number of orbitals in a shell is n2.
6.6
· This section shows the shapes of the orbitals. Read it over, and know the shapes of the orbitals.
· know the graphs of the radial probability for the s orbitals, and how they show nodes closer to the nucleus than the outside radius.
6.7
· In a many-electron atom, for every value of n, the energy of an orbital increases as the l value increases.
· Degenerate orbitals have the same energy.
· Electron spin allows electrons to occupy the same orbitals.
· The spin magnetic quantum number is represented as ms and can equal either +1/2 or -1/2.
· Pauli Exclusion Principle: no two electrons in an atom can have the same set of four quantum numbers, n, l, ml, and ms.
· Because of the Pauli Exclusion Principle, each orbital can hold up to 2 electrons, each with opposite spins.
6.8
· Electron configuration: the way electrons are distributed among the different orbitals of an atom.
· Orbitals are filled in order of increasing energy, with up to 2 atoms per orbital.
· Know how to write out the electron configuration.
· Paired electrons – electrons in the same orbital.
· Unpaired electrons- electrons that are not accompanied by a partner.
· Hund’s rule: for degenerate orbitals, the lowest energy happens when the number of electrons with the same spin is maximized.
· This means that electrons fill orbitals of the same energy first singly, and then, when all degenerate orbitals have 1 electron, those orbitals are filled doubly.
· Know how to use the noble gases to write condensed electron configurations.
· Core electrons – electrons represented by the noble gas (in condensed electron configurations.
· Outer-shell electrons – the rest of the electrons, written after the noble gas.
· Valence electrons: the electrons involved in bonding.
· Transition Metals: in the middle of the periodic table, electrons in which the d orbitals are being filled.
· Lanthanides and Actinides:
· Lanthanides (rare earth elements) are metals whose 4f orbitals are being filled.
· Actinides are metals whose 5f orbitals are being filled.
6.9
· The periodic table is a good guide as to how orbitals are filled.
· Know the diagram the book gives you (telling which orbitals are being filled in which sections of the periodic table).
· Generally, completely full d or f orbitals are not considered to be among the valence electrons.
· There are a few elements that do not go along with the previous rules – chromium + copper. Know how these differ.