Chem121test4 Flashcards
electronic structure
the arrangement and energy of electrons
Electromagnetic radiation
moves as waves through space at the speed of light
wavelength (λ)
The distance between corresponding points on adjacent waves
frequency (ν)
The number of waves passing a given point per unit of time
The speed of light (c)
3.00 × 10^8 m/s
Three observed properties associated with how atoms interact with electromagnetic radiation can NOT be explained by waves
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)
blackbody radiation
the emission of light from hot objects
the photoelectric effect
the emission of electrons from metal surfaces on which light is shone
emission spectra
emission of light from electronically excited gas atoms
frequency equation
v = c/λ
The Nature of Energy—Quanta
Max Planck explained it by assuming that energy comes in packets called quanta
energy Einstein equation
E = hν
Planck’s constant
6.626 × 10^−34 J∙s
The Photoelectric Effect
when photons hit a surface electrons are released
Atomic Emissions
energy emitted by atoms and molecules
continuous spectrum
the rainbow
line spectrun
discrete wavelenghts (colors) are observed depending on which element it is
ground state
electrons in the lowest energy state
excited state
any higher state than the ground state
The Bohr model only works for
hydrogen
circular motion is not
wave like in nature
positive ΔE
A photon is absorbed in this instance. This happens if nf > ni.
negative ΔE
A photon is emitted in this instance. This happens if nf < ni.
Important Ideas from the
Bohr Model
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.
quantum mechanics
a mathematical treatment into which both the wave and particle nature of matter could be incorporated
The solution of Schrödinger’s wave equation for hydrogen yields
wave functions for the electron
The square of the wave function gives the
electron density
orbitals
Solving the wave equation gives a set of wave functions
An orbital is described
by a set of three quantum numbers
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.
l = 0
s
l = 1
p
l = 2
d
l = 3
f
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.
electron shell
Orbitals with the same value of n form an
electron shell
subshells
Different orbital types within a shell are subshells
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
p Orbitals
The value of l for p orbitals is 1.
They have two lobes with a node between them
barbell shaped
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
f Orbitals
l = 3
Seven equivalent orbitals in a sublevel
Very complicated shapes (not shown
in text)
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
degenerate orbitals
orbitals on the same energy level have the same energy
As the number of electrons increases
so does the repulsion between them
Therefore, in atoms with more than one electron
not all orbitals on the same energy level are degenerate
Orbital sets in the same sublevel
are still degenerate
Energy levels start to overlap in energy
Energy levels start to overlap in energy (e.g., 4s is lower
in energy than 3d.)
Spin Quantum Number, ms
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anthanide elements
atomic numbers
57 to 70) have electrons entering the 4f sublevel
actinide elements
including Uranium, at no. 92, and Plutonium, at no. 94) have electrons entering the 5f sublevel
Pauli Exclusion Principle
No two electrons in the same atom can have the same set of four quantum numbers
electron configuration
The way electrons are distributed in an atom
ground state
The most stable organization is the lowest possible energy
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
Hund’s Rule
When filling degenerate orbitals the lowest energy is attained when the number of electrons having the same spin is maximized
Heisenburg uncertainty principle
it is impossible to know both the position and the momentum of an electron
Periodicity
the repetitive pattern of a property for elements based on atomic number
Effective Nuclear Charge
The effective nuclear charge, Zeff, is determined using: Zeff = Z − S
nonbonding atomic radius
van der Waals radius
van der Waals radius
half of the shortest distance separating two nuclei during a collision of atoms
bonding atomic radius
half the distance between nuclei in a bond
Sizes of Atoms
left to right > top to bottom ↑
Sizes of Ions
Ionic size depends on
the nuclear charge.
the number of electrons.
the orbitals in which electrons reside
Cations
are smaller than their parent atoms
Anions
Electrons are added and repulsions between electrons are increased
isoelectronic series
ions have the same number of electrons
size decreases with an increasing nuclear charge
Ionization Energy (I)
the minimum energy required to remove an electron from the ground state of a gaseous atom or ion
The higher the ionization energy
the more difficult it is to remove an electron
Ionization direction
increases to the right > increases up ↑
Electron affinity
the energy change accompanying the addition of an electron to a gaseous atom
Electron affinity is typically exothermic
exothermic
Three notable exceptions toElectron affinity
group 2a, 5a, 8a
General Trend in Electron Affinity
Not much change in a group, it generally increases across a period
»»>
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).
Metals
Metals tend to form cations
Properties of metals
Shiny luster
Conduct heat and electricity
Malleable and ductile
Solids at room temperature (except mercury)
Low ionization energies/form cations easily
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
Metal oxides
tend to be basic and react with acids
Nonmetals
Nonmetals are found on the right hand side of the periodic table
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
Nonmetal Chemistry
Substances containing only nonmetals are molecular compounds.
Most nonmetal oxides are acidic
Sc2O3(s) + 6 HNO3(aq)
2 Sc(NO3)3(aq) + 3 H2O(l) (makes salt and water)
Metalloids
Metalloids have some characteristics of metals and some of nonmetals
Group Trends
Elements in a group have similar properties.
Trends also exist within groups
Alkali metals
They have low densities and melting points.
They also have low ionization energies.
soft, metallic solids.
Alkali Metal Chemistry
Their reactions with water are famously exothermic
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
Flame Tests
Qualitative tests for alkali metals include their characteristic colors in flames
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
Beryllium
does not react with water
magnesium
reacts only with steam
polonium
is most likely to have a positive charge
Oxygen can exist as
Oxygen gas, O2 (technically called dioxygen)
Ozone gas, O3
Group 7A—Halogens
typical nonmetals.
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
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.
