Chem121test4 Flashcards
(96 cards)
1
Q
electronic structure
A
the arrangement and energy of electrons
2
Q
Electromagnetic radiation
A
moves as waves through space at the speed of light
3
Q
wavelength (λ)
A
The distance between corresponding points on adjacent waves
4
Q
frequency (ν)
A
The number of waves passing a given point per unit of time
5
Q
The speed of light (c)
A
3.00 × 10^8 m/s
6
Q
Three observed properties associated with how atoms interact with electromagnetic radiation can NOT be explained by waves
A
the emission of light from hot objects (blackbody radiation)
the emission of electrons from metal surfaces on which light is shone (the photoelectric effect)
emission of light from electronically excited gas atoms (emission spectra)
7
Q
blackbody radiation
A
the emission of light from hot objects
8
Q
the photoelectric effect
A
the emission of electrons from metal surfaces on which light is shone
9
Q
emission spectra
A
emission of light from electronically excited gas atoms
10
Q
frequency equation
A
v = c/λ
11
Q
The Nature of Energy—Quanta
A
Max Planck explained it by assuming that energy comes in packets called quanta
12
Q
energy Einstein equation
A
E = hν
13
Q
Planck’s constant
A
6.626 × 10^−34 J∙s
14
Q
The Photoelectric Effect
A
when photons hit a surface electrons are released
15
Q
Atomic Emissions
A
energy emitted by atoms and molecules
16
Q
continuous spectrum
A
the rainbow
17
Q
line spectrun
A
discrete wavelenghts (colors) are observed depending on which element it is
18
Q
ground state
A
electrons in the lowest energy state
19
Q
excited state
A
any higher state than the ground state
20
Q
The Bohr model only works for
A
hydrogen
21
Q
circular motion is not
A
wave like in nature
22
Q
positive ΔE
A
A photon is absorbed in this instance. This happens if nf > ni.
23
Q
negative ΔE
A
A photon is emitted in this instance. This happens if nf < ni.
24
Q
Important Ideas from the
Bohr Model
A
Electrons exist only in certain discrete energy levels, which are described by quantum numbers.
Energy is involved in the transition of an electron from one level to another.
25
quantum mechanics
a mathematical treatment into which both the wave and particle nature of matter could be incorporated
26
The solution of Schrödinger’s wave equation for hydrogen yields
wave functions for the electron
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The square of the wave function gives the
electron density
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orbitals
Solving the wave equation gives a set of wave functions
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An orbital is described
by a set of three quantum numbers
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Angular Momentum Quantum Number (l)
This quantum number defines the shape of the orbital.
Allowed values of l are integers ranging from 0 to n − 1.
Letter designate the different values of l. This defines the shape of the orbitals.
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l = 0
s
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l = 1
p
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l = 2
d
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l = 3
f
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Magnetic Quantum Number (ml)
The magnetic quantum number describes the three-dimensional orientation of the orbital.
Allowed values of ml are integers ranging from −l to l including 0:
−l ≤ ml ≤ l
Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, and so forth.
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electron shell
Orbitals with the same value of n form an
electron shell
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subshells
Different orbital types within a shell are subshells
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s Orbitals
The value of l for s orbitals is 0.
They are spherical in shape.
The radius of the sphere increases with the value of n
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p Orbitals
The value of l for p orbitals is 1.
They have two lobes with a node between them
barbell shaped
40
d Orbitals
The value of l for a d orbital is 2.
Four of the five d orbitals have four lobes; the other resembles a p orbital with a doughnut around the center
41
f Orbitals
l = 3
Seven equivalent orbitals in a sublevel
Very complicated shapes (not shown
in text)
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Energies of Orbitals—Hydrogen
For a one-electron hydrogen atom, orbitals on the same energy level have the same energy.
Chemists call them degenerate orbitals
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degenerate orbitals
orbitals on the same energy level have the same energy
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As the number of electrons increases
so does the repulsion between them
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Therefore, in atoms with more than one electron
not all orbitals on the same energy level are degenerate
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Orbital sets in the same sublevel
are still degenerate
47
Energy levels start to overlap in energy
Energy levels start to overlap in energy (e.g., 4s is lower
in energy than 3d.)
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Spin Quantum Number, ms
`
49
anthanide elements
atomic numbers
57 to 70) have electrons entering the 4f sublevel
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actinide elements
including Uranium, at no. 92, and Plutonium, at no. 94) have electrons entering the 5f sublevel
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Pauli Exclusion Principle
No two electrons in the same atom can have the same set of four quantum numbers
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electron configuration
The way electrons are distributed in an atom
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ground state
The most stable organization is the lowest possible energy
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Each component consists of
a number denoting the energy level
a letter denoting the type of orbital
a superscript denoting the number of electrons in those orbitals
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Hund’s Rule
When filling degenerate orbitals the lowest energy is attained when the number of electrons having the same spin is maximized
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Heisenburg uncertainty principle
it is impossible to know both the position and the momentum of an electron
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Periodicity
the repetitive pattern of a property for elements based on atomic number
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Effective Nuclear Charge
The effective nuclear charge, Zeff, is determined using: Zeff = Z − S
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nonbonding atomic radius
van der Waals radius
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van der Waals radius
half of the shortest distance separating two nuclei during a collision of atoms
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bonding atomic radius
half the distance between nuclei in a bond
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Sizes of Atoms
left to right > top to bottom ↑
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Sizes of Ions
Ionic size depends on
the nuclear charge.
the number of electrons.
the orbitals in which electrons reside
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Cations
are smaller than their parent atoms
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Anions
Electrons are added and repulsions between electrons are increased
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isoelectronic series
ions have the same number of electrons
size decreases with an increasing nuclear charge
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Ionization Energy (I)
the minimum energy required to remove an electron from the ground state of a gaseous atom or ion
68
The higher the ionization energy
the more difficult it is to remove an electron
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Ionization direction
increases to the right > increases up ↑
70
Electron affinity
the energy change accompanying the addition of an electron to a gaseous atom
71
Electron affinity is typically exothermic
exothermic
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Three notable exceptions toElectron affinity
group 2a, 5a, 8a
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General Trend in Electron Affinity
Not much change in a group, it generally increases across a period
>>>>>
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Irregularities in the General Trend
The trend is not followed when the added valence electron in the next element
enters a new sublevel (higher energy sublevel);
is the first electron to pair in one orbital of the sublevel (electron repulsions lower energy).
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Metals
Metals tend to form cations
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Properties of metals
Shiny luster
Conduct heat and electricity
Malleable and ductile
Solids at room temperature (except mercury)
Low ionization energies/form cations easily
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Metal Chemistry
Compounds formed between metals and nonmetals tend to be ionic. Remember that Ionic Salts tend to have high melting point, typically have good solubility in water, etc
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Metal oxides
tend to be basic and react with acids
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Nonmetals
Nonmetals are found on the right hand side of the periodic table
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Properties of nonmetals
Solid, liquid, or gas (depends on element)
Solids are dull, brittle, poor conductors
Large negative electron affinity, so they form anions readily
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Nonmetal Chemistry
Substances containing only nonmetals are molecular compounds.
Most nonmetal oxides are acidic
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Sc2O3(s) + 6 HNO3(aq)
2 Sc(NO3)3(aq) + 3 H2O(l) (makes salt and water)
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Metalloids
Metalloids have some characteristics of metals and some of nonmetals
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Group Trends
Elements in a group have similar properties.
Trends also exist within groups
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Alkali metals
They have low densities and melting points.
They also have low ionization energies.
soft, metallic solids.
86
Alkali Metal Chemistry
Their reactions with water are famously exothermic
87
Differences in Alkali Metal Chemistry
Lithium reacts with oxygen to make an oxide:
4 Li + O2 2 Li2O
Sodium reacts with oxygen to form a peroxide:
2 Na + O2 Na2O2
K, Rb, and Cs also form superoxides:
M + O2 MO2
88
Flame Tests
Qualitative tests for alkali metals include their characteristic colors in flames
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Alkaline Earth Metals
Alkaline earth metals have higher densities and melting points than alkali metals.
Their ionization energies are low, but not as low as those of alkali metals.
They readily form +2 cations, losing the 2 valence electrons
90
Beryllium
does not react with water
91
magnesium
reacts only with steam
92
polonium
is most likely to have a positive charge
93
Oxygen can exist as
Oxygen gas, O2 (technically called dioxygen)
Ozone gas, O3
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Group 7A—Halogens
typical nonmetals.
95
Group 8A—Noble Gases
The noble gases have very large ionization energies.
Their electron affinities are positive (can’t form stable anions).
Therefore, they are relatively unreactive.
They are found as monatomic gases
96
Hydrogen
Is 1s1 a metallic electron configuration like the other ns1 elements?
We do think of acid compounds, like HCl, as having H+, however they are really covalent in nature.
When reacting with metals, hydride anions (H–) form.
