AP Chemistry: Chapter 7 – Atomic Structure & Periodicity
7.1 Electromagnetic Radiation______
-Light travels through space as a wave. Waves are made up of crests (which rise above the midline) and troughs which sink below it. Waves have three primary characteristics: wavelength, frequency, and amplitude. Two of these characteristics are of special focus in regards to atomic structure:
§ Wavelength (greek letter lamda, λ) – the distance from two consecutive crests or troughs, most often measured in meters or nanometers.
§ Frequency (greek letter nu, ν) – the number of wave cycles that pass a given point in unit time, measured in Hertz (1 Hz = 1cycle/s).
-Wavelength and frequency are related by the following formula:
c = the speed of light (in a vacuum) = 2.998 x 108 m/s
Reference Figure 7.1 (pg 286)
-Visible light is a tiny portion of the electromagnetic spectrum, covering a wavelength region of 400 nm to 700nm.
Reference Figure 7.2 (pg 286)
7.2 The Nature of Matter______
-Hundreds of years ago it was believed that all properties of light could be explained in terms of its wave nature. Max Planck & Albert Einstein conducted a series of experiments that discredited this idea.
-Max Planck started studying radiation profiles emitted by solid bodies heated (for example, black body radiation). He accounted for these observations by simply postulating that energy can be gained or lost in whole number multiples, in other words energy can be quantized.
-Albert Einstein proposed that electromagnetic radiation itself is quantized and that it can be viewed as a stream of particles called photons. As a result of this theory, today it is considered that light is generated as a stream of particles called photons, whose energy E is represented by the following equation:
Ephoton = hν = hc
λ
E = energy in Joules (J) per photon
h = Planck’s constant = 6.626 x 10-34 J s
ν = frequency = s-1 (reciprocal seconds or the SI unit Hz)
-Einstein arrived at his conclusion through analysis of the photoelectric effect. The photoelectric effect refers to the phenomenon in which electrons are emitted from the surface of a metal when light strikes it. Basically, when electrons absorb energy in the ground state they will jump to an excited state. When they return to there ground state energy is released in the form of light.
-The kinetic energy of emitted electrons increases linearly with the frequency of the light. And in a related development Einstein derived the famous:
E = mc2
-Electromagnetic radiation is not only wave-like in nature but particle like as well with consideration to photons, and this is referred to as wave particle duality.
7.3 The Atomic Spectra of Hydrogen & 7.4 The Bohr Model______
-When visible light is passed through a prism a continues spectrum appears as a result (consisting of the spectrum colors, in order of increasing energy: red, orange, yellow, green, blue, indigo, and violet).
-Elements emission spectrum, when passed through a prism, will not give off all 7 colors but a specific line spectrum (or atomic spectrum) of colors indicating various regions of energy.
Reference Figure 7.7 (pg 294)
-Niels Bohr assumed that the hydrogen atom consists of a central proton around which an electron moves in a circular path. He related that the electrostatic force of the proton for the electron (effective nuclear charge) to the centrifugal force due to the circular motion of the electron. Bohr was able to express the energy of that tome in terms of the radius of the electron’s orbit.
-There are three points to be made with the Bohr Model:
- Bohr designated zero energy as the point at which the proton and electron are completely separated.
- Ordinarily the hydrogen electron is in its lowest energy state, referred to as the ground state (n=1). When an electron absorbs enough energy, it moves to a higher, excited state. For hydrogen, the first excited state is n=2, then n=3.
- When an excited electron drops back to its lower energy state it gives off energy as a photon of light.
-Using these three points it is possible to relate the frequency of light emitted to the quantum numbers, nhi and nlo, of the two states:
ν = RH ( _1 - 1 ) OR 1 = RH ( _1 - 1 )
h nlo2 nhi2 λ nhi2 nlo2
RH (Rydberg’s Constant) = 2.180 x 10-18 J
h (Planck’s Constant) = 6.626 x 10-34 J*s
-All of the lines in the Balmer series come from transitions to n = 2 from higher levels (n = 3,4,5,….)
Figure 7.9 (pg 295)
7.5 The Quantum Mechanical Model_& 7.9 Polyelectronic Atoms______
-The Bohr modeled had a flaw in that it could only be applied to hydrogen. The idea that the electron moved around the nucleus in a well defined orbit at a fixed distance from the nucleus was not sufficient.
-The Bohr model was not applicable to polyelectronic atoms (atoms with more that one electron), it only applied to hydrogen. So another model was necessary.
-Louis de Broglie reasoned that if light could behave has particles (photons) as well as waves, then why can’t electrons behave like a wave. This idea led to the development of wave mechanics, now called quantum mechanics. This deferred from the Bohr model in a number of ways:
§ The kinetic energy of an electron is inversely related to the volume of the region to which it is confined.
§ It is impossible to specify the precise position of an electron in an atom at a given instant.
-Erwin Schrodinger mathematically (this mathematical development is to complex to be discussed at this level of chemistry) justified this model using wave function.
-Werner Heisenberg also contributed to this model by introducing the Heisenberg uncertainty principle, in which he basically stated that it is not possible to know both the position and the momentum of a particle (the electron) at a given time. So there is always some degree of uncertainty as to the specific location of an electron.
-The modern atomic model suggests that electrons are found in regions of energy and not specific orbitals like Bohr suggested. And electrons can move from one energy level to another in an excited state.
7.6 Quantum Numbers & 7.8 Electron Spin & The Pauli Exclusion Principle______
-Erwin Schrodinger proposed a rather complex differential equation to express wave properties of an electron in an atom. The Shrodinger equation can be solved, approximately, for atoms with two or more electrons. There are many solutions for the wave function, each associated with a set of quantum numbers.
-The principal quantum number:
§ Symbolized by n, basically the energy level the electron is in.
§ Also called an energy level and energy shell
§ The energy of the electron depends mainly on this number. As n increases, so does the value of the electron.
§ n = 1,2,3,4,5,…….
-The angular momentum quantum number (also referred to as the orbital quantum number):
§ Symbolized by l, basically represents the sublevel the electron is in s,p,d, or f.
§ Also commonly called a subshell
§ s: l = 0, p: l = 1, d: l = 2, f: l = 3
§ l = 0,1,2,…… (n-1)
§ Sublevels increase in energy as well. Sublevel 2p is higher in energy than sublevel 2s.
*The letters s,p,d, and f come from the adjectives used to describe spectral lines: sharp, principal, diffuse, fundamental.
-The magnetic quantum number:
§ Symbolized by ml, this determines the direction in space of the electron cloud surrounding the nucleus.
§ Sublevel s has only 1 possible orbital
§ Sublevel p has 3 possible orbitals
§ Sublevel d has 5 possible orbitals
§ Sublevel f has 7 possible orbitals
§ All of the orbitals in a sublevel have the same energy
s __ p ______d ______f ______
Reference Figures 7.14, 7.15, 7.17, 7.18 (pgs 305 – 307)
-The spin quantum number:
§ Symbolized by ms, this represents the electron spin.
§ An electron has magnetic properties that correspond to those of a charged particle spinning on its axis, either clockwise or counterclockwise.
§ ms can equal +1/2 or -1/2
Reference Figure 7.20 (pg 308)
-The Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers – it requires that only 2 electrons can fit into an orbital.
7.11 The Aufbau Peinciple & Periodic Table (Also Electron Configurations in Atoms)____
-The Aufbau Principle: As protons are added one by one to the nucleus to build up elements, electrons are added similarly starting with hydrogen like orbitals (basically all configurations start at hydrogen and build up).
-Electron configuration describes the arrangement of electrons.
For example: Ne 1s22s22p63s23p6
-Abbreviated electron configurations may be used by preceding with the noble gas for the element. For example: [Ne] 3s23p4
-Electron Configuration & The Periodic Table
§ The atoms of elements in a group of the periodic table have the same distribution of electrons in the outermost principal energy level (also known as valence electrons).
§ The elements in groups 13-18 fill p sublevels (starts at 2p).
§ The transition metals fill d sublevels (starts at 3d).
§ The two sets of elements listed separately at the bottom of the table are filling f sublevels (starts at 4f).
Reference Figure 7.28 (pg 315)
Reference Figure 7.29 (pg 316)
-Orbital diagrams show how electrons are distributed using arrows (to represent electrons) and parentheses or lines (to represent the orbitals).
-Hund’s Rule: The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of orbitals (basically up arrows fill first).
§ When several orbitals of equal energy are available, as in a given sublevel, electrons enter singly with parallel spins.
§ In all filled orbitals, the two electrons have opposed spins.
§ In accordance with this rule, within a given sublevel, there are as many half filled orbitals as possible.
-Paramagnetic: If there are unpaired electrons present the solid will be attracted into the field
-Diamagnetic: If the atoms in the solid contain only paired electrons it is slightly repelled by the field.
-Electron configurations of ions show the arrangement of electrons when an electron has been lost or gained.
-In general when a monatomic ion is formed electrons are added to or removed from sublevels in the highest principal energy level. Many monatomic ions (especially the main group elements) have noble gas configurations.
-Transition metal cations to the right of the scandium group do not form ions with noble-gas configurations, they would have to lose four or more electrons to do so. In transition metals the outer s electrons are usually lost first to form positive ions.
For example: Mn:
Mn+2:
-In ions like Fe, electrons will be lost from 4s first then the 3d. This is usually referred to as the “first in, first out” rule.
7.10 History of the Periodic Table______
-The periodic table was originally constructed to represent the patterns observed in the chemical properties of the elements.
-Two scientists are credited with independently originating the periodic table: Julius Meyer and Dimitri Medeleev (although Mendeleev is given most of the credit).
-Mendeleev originally put elements known at the time in order of atomic masses and saw a reoccurring pattern.
Reference Figure 7.25 (pg 311)
-Another scientis by the name of Henry Moseley put the elements in order of increasing atomic number and say that patterns were more apparent and gave rise to the periodic law.
-The Periodic Law: The chemical and physical properties of elements are a periodic function of atomic number.
7.12 Periodic Trends in Atomic Properties______
-Ionization Energy (IE): the energy required to remove an electron:
X à X+ + e-
§ Increases across the periodic table from left to right
§ Decreases down the periodic table
§ Once the first electron is removed (first ionization energy), more energy is required to remove remaining electrons.
Reference Figures 7.5, 7.7, & 7.7 (pgs 319-320)
-Electron Affinity (EA) is the energy change associated with the addition of an electron:
X + e- à X—
-Electronegativity has the same trend as EA and is defined as an atoms ability to attract an electron.
§ Increases across the periodic table from left to right
§ Decreases down the periodic table
-Atomic Radius: one half the distance of closest approach between atoms in an elemental substance.
§ Decreases across a period from left to right on the periodic table
§ Increases down a group on the periodic table
Reference Figures 7.34 & 7.35 (pgs 322-323)
-Ionic Radius:
§ Positive ions are smaller than the metal atoms from which they are formed
§ Negative ions are larger than the nonmetal atoms from which they are formed
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