Quiz 2 – Chem1410 – Flashcards
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| Li+ |
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| Always soluble |
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| Na+ |
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| Always soluble |
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| K+ |
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| Always soluble |
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| NH4^+ |
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| Always soluble |
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| NO3^- |
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| Always soluble |
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| C2H3O2^- |
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| Always soluble |
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| Cl- |
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| Insoluble when paired with Ag+, Hg2^2+, or Pb^2+. |
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| Br- |
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| Insoluble when paired with Ag+, Hg2^2+, or Pb^2+. |
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| I- |
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| Insoluble when paired with Ag+, Hg2^2+, or Pb^2+. |
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| SO4^2- |
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| Insoluble when paired with Sr^2+, Ba^2+, Pb^2+, Ag+, or Ca^2+. |
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| OH- |
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| Soluble when paired with Li+, Na+, K+, NH4^+. Slightly soluble with Ca^2+, Sr^2+, and Ba^2+. |
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| s^2- |
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| Soluble when paired with Li+, Na+, K+, NH4^+, Ca^2+, Sr^2+, and Ba^2+. |
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| CO3^2- |
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| Soluble with Li+, Na+, K+, or NH4^+. |
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| PO4^3- |
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| Soluble with Li+, Na+, K+, or NH4^+. |
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| Oxidation |
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| loss of electrons |
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| Reduction |
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| Gain of electrons |
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| The oxidation state of an atom in a free element |
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| 0 |
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| Oxidation state of monoatomic ion |
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| equal to its charge. E.g. Ca^2+ has an oxidation state of +2. |
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| Sum of oxidation state of all atoms in neutral molecule or formula unit. |
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| 0 |
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| Sum of oxidation states of all atoms in ion. |
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| Equal to its charge. E.g. NO3^- has an overall oxidation state of -1. |
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| Oxidation state of group 1A metals |
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| +1 |
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| Oxidation state of group 2A metals |
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| +2 |
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| Oxidation state of F |
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| -1 (priority - 1) |
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| Oxidation state of H |
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| +1 (priority - 2) |
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| Oxidation state of O |
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| -2 (priority - 3) |
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| Oxidation state of group 7A |
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| -1 (priority - 4) |
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| Oxidation state of group 6A |
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| -2 (priority - 5) |
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| Oxidation state of group 5A |
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| -3 (priority - 6) |
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| Frequency in relation to wavelength |
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| V = c/l |
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| Electromagnetic spectrum from highest energy to lowest |
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| Gamma, X, Ultraviolet, Visible light, Infrared, Microwave, Radio |
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| Constructive interference |
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| Two in phase electromagnetic waves combining to form wave with 2X the amplitude. |
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| Energy of photon in terms of wavelength |
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| E = hc/l |
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| Threshold frequency condition |
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| Energy of photon = Binding energy of emitted electron: hv=phi. |
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| Emission spectrum |
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| Graphic representation of the wavelengths of light which when applied to specific elements results in the emission of electrons. |
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| Bohr model |
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| The idea that electrons exist in "stationary states". Changes in energy only occur when they jump from one stationary state to another. |
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| deBroglie relation |
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| l=h/mv (l=wavelength, m=mass, v=velocity of electron) |
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| Complementary properties |
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| properties which inherently exclude each other. I.e. the position, and velocity of a particle. Also, position and energy of a particle. Position has to do with particle nature of electron, whereas energy and velocity relate to wave nature of electrons. |
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| Heisenberg's uncertainty principle |
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| deltax * m*deltav >= h/4pi. Here we simply need to know that the more accurately you know the position of an electron, the less accurately you know its velocity. |
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| Orbital |
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| probability distribution map showing where an electron is likely to be found |
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| Principal quantum number |
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| n. Determines overall size and energy of orbital. Ranges from 1 to infinity. |
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| Angular momentum quantum number |
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| l. Determines shape of orbital. l=1, s orbital: l=2, p orbital. l=3, d orbital. l=4, f orbital. Ranges from 0 to n-1. |
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| Magnetic quantum number |
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| m(l). Specifies orientation of orbital. Ranges from -l, to +l. |
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| Orbital with equal values n are said to be in the same: |
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| Principal level (or principal shell) |
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| Orbitals with equal values of n and l are said to be in the same: |
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| Sublevel (or subshell) |
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| Number of sublevels in any level: |
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| equals n. |
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| Number of orbitals in any sublevel: |
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| 2l + 1 |
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| Number of orbitals in any level: |
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| n^2 |
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| Wavelength of photon in relation to its energy |
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| l=hc/E |
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| Node |
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| A point where the wave function = 0. In other words a point where the probability of an electron being present = 0. |
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| Spin quantum number |
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| m(s). A quantum number which determines whether an electron has an up spin (+1/2) or a down spin (-1/2). electrons with up spin will be equal to electrons with down spin within an atom. |
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| Ground state |
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| The lowest energy state of an atom |
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| Pauli exclusion principle |
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| No two electrons in an atom can have the same four quantum numbers. |
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| Effective nuclear charge |
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| The charge of the nucleus, minus the charge of all electrons shielding it from whatever we are attempting to calculate the Effective nuclear charge of. |
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| Aufbau principle |
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| The combined ideas that electrons will always fill lower energy orbitals first, and that only two electrons are allowed in each orbital. |
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| Hund's rule |
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| when filling degenerate orbitals, electrons fill them singly first, with parallel spins. |
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| No. of orbitals (and electrons) in f sublevel |
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| 7 - sublevels, 14 - electrons |
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| No. of orbitals (and electrons) in d sublevel |
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| 5 - sublevels, 10 - electrons |
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| No. of orbitals (and electrons) in p sublevel |
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| 3 - sublevels, 6 - electrons |
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| No. of orbitals (and electrons) in s sublevel |
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| 1 - sublevel, 2 - electrons |
