Quiz 2 – Chem1410

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

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