CHEMISTRY: THE CENTRAL SCIENCE chapter 6

electronic structure

arrangement of electron in atom

(section 6.0)

quantum theory

theory that explains behavior of electrons in an atoms

(section 6.0)

quantum mechanics

"new" physics developed to describe atoms correctly

(section 6.0)

electromagnetic radiation

form of energy that has wave characteristics and that propagates through a vacuum at the characteristic speed of 2.998×10⁸ m/s (speed of light)

examples (decreasing wavelength, increasing energy)

∙ radio

∙ microwave

∙ infrared (heat)

∙ visible light

∙ ultraviolet

∙ X-rays

∙ gamma rays

(section 6.1)

periodicity

pattern of peaks and troughs that is repeated at regular intervals in a wave-like fashion

(section 6.1)

wavelength

distance between identical points on successive waves

this cycle of a wave is usually measured peak-to-peak or trough-to-trough

inverse relationship exists between wavelength of light and its energy

(section 6.1)

frequency

number of complete cycles of a wave that pass a given point per unit of time (unit: s⁻¹, "per second")

positive correlation exists between frequency of light and its energy

(section 6.1)

relationship between wavelength (λ) and frequency (v)

relationship is inverse because

c = λv

∙ c (speed of light; 3.00×10⁸ m/s)

∙ λ (lambda) (wavelength, m)

∙ v (nu) (frequency, s⁻¹, "per second")

as waves get shorter they get more frequent

(section 6.1)

electromagnetic spectrum

various types of electromagnetic radiation arranged in order of increasing wavelength

(section 6.1)

gamma radiation

energetic electromagnetic radiation emanating from the nucleus of a radioactive atom

∙ ↑v, ↑E

∙ ↓λ

∙ λ < 10 pm, 10⁻¹² m

(sect 21.1)

X-rays

type of electromagnetic radiation used in bone and tissue imaging

∙ ↑v, ↑E

∙ ↓λ

∙ λ = 5 pm to 10 nm = 5×10⁻¹² m to 10⁻⁸ m

∙ unit: Angstrom Å = 10⁻¹⁰ m = 0.1 nm

ultraviolet light

(UV)

electromagnetic radiation with a wavelength from 400 nm to 10 nm, shorter than that of visible light but longer than X-rays.

∙ λ = 10-400 nm = 1×10⁻⁸ m 4×10⁻⁷ m

∙ unit: nanometer, nm = 10⁻⁹ m

visible light

type of electromagnetic radiation that is visible to human beings (unit: nm)

(long→short) ROY G. B(I)V

Red

Orange

Yellow

Green

Blue

(Indigo)

Violet

∙ λ = 400-750 nm = 4×10⁻⁷m 7×10⁻⁷m

∙ unit: nanometer, nm = 10⁻⁹ m

infrared radiation

(IR)

electromagnetic radiation with longer wavelengths than those of visible light.

most of the thermal radiation emitted by objects near room temperature is infrared.

∙ λ = 700 to 1000 nm = 7×10⁻⁷ tp 7×10⁻⁶ m

∙ v = 430 THz to 300 GHz

black-body radiation

emission of light from hot objects

(so called because the objects studied appear black before heating)

(section 6.2)

emission spectra

emission of light from electronically excited gas atoms

(section 6.2)

quantized

restricted to certain quantity

discrete or step-wise, as opposed to continuous

(section 6.2)

quantum

smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation

E = hv = energy of photon

∙ h (Planck's constant, 6.626×10⁻³⁴ J-s, "Joule-second")

∙ v (nu) (frequency, s⁻¹, "per second")

literal translation: "fixed amount"

(section 6.2)

Planck's constant

(h)

6.626×10⁻³⁴ J-s

(section 6.2)

photoelectric effect

emission of electrons from metal surfaces on which light shines

explained by Albert Einstein (1879-1955) in 1905

(section 6.2)

photon

smallest increment (quantum) of radiant energy (unit: J)

energy of ______ = E = hv

∙ h (Planck's constant, 6.626×10⁻³⁴ J-s, "Joule-second")

∙ v (nu) (frequency, s⁻¹, "per second")

(section 6.2)

spectrum

electromagnetic radiation separated into its component wavelengths

(section 6.3)

line spectrum

spectrum containing only specific wavelengths

(section 6.3)

continuous spectrum

spectrum containing light of ALL wavelengths

example

∙ rainbow

(section 6.3)

Bohr Model

E = -h×c×R(H)×(1/n²) = (-2.18×10⁻¹⁸ J)×(1/n²)

and

∆E = E₁ - E₀

important ideas

∙ 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

postulates

∙ for orbit of H atom electron, only orbits of certain radii are permitted

∙ an electron in a permitted orbit is in an "allowed" energy state and does not radiate energy (and thus does not spiral into the nucleus)

∙ energy is emitted or absorbed by the electron only as it changes from one allowed energy state to another. This energy is emitted or absorbed as a photon that has energy "E = hv"

(section 6.3)

principal quantum number

(n)

whole-number (integer) value corresponding to electron orbits

orbit radius increases as "n" increases

(section 6.3)

ground state

lowest-energy state (n=1) of an atom

(section 6.3)

excited state

higher-energy (than ground) state (n≥2) of an atom

(section 6.3)

wavelength of particle

λ = h / m×v = h / p

∙ h (Planck's constant, 6.626×10⁻³⁴)

∙ p = m×v (momentum, kg-m/s)

∙ m (mass, kg)

∙ v (velocity, m/s)

developed by Louis de Broglie (1892-1987) along with idea of "matter waves"

(section 6.4)

momentum

product of the mass, m, and

velocity, v, of an object (unit: kg-m/s)

p = mv

∙ p (momentum, kg-m/s)

∙ m (mass, kg)

∙ v (velocity, m/s)

(section 6.4)

matter waves

term used to describe the wave characteristics of a moving particle based on assumption that particles may move in wave-like fashion

example of "wave-particle duality"

(section 6.4)

uncertainty principle

principle stating there is an inherent uncertainty in the precision with which we can simultaneously specify the position and momentum of a particle

∆x ∙ ∆(mv) ≥ h / (4×π×m×∆v)

this uncertainty is significant only for particles of extremely small mass, such as electrons

developed by Werner Heisenberg

(section 6.4)

wave functions

a mathematical description of an allowed energy state (an orbital) for an electron in the quantum mechanical model of the atom

its usually symbolized by the Greek letter Ψ

(section 6.5)

orbital

an allowed energy state of an electron in the quantum mechanical model of the atom

the term _____ is also used to describe the spatial distribution of the electron

a(n) _______ is defined by the values of the three quantum numbers: n, l, and m₁

(section 6.5)

degenerate

name for orbitals with same energy

(section 6.7)

electron spin

property of the electron that makes it behave as though it were a tiny magnet

the electron behaves as if it were a tiny sphere spinning on its axis

________ ____ is quantized

(section 6.7)

spin magnetic quantum number

m(s)

a quantum number associated with the electron spin

it may have values of +1/2 or -1/2

(section 6.7)

Pauli exclusion principle

statement that no two electrons in an atom can have the same set of four quantum numbers n, 1, m₁ and m(s)

for a given orbital, the values of n, l, and m₁ are fixed

(section 6.7)

electron configuration

the arrangement of electrons in the orbitals of an atom or molecule

(section 6.8)

Hund's rule

rule stating that electrons occupy degenerate orbitals in such a way as to maximize the number of electrons with the same spin

in other words, each orbital has one electron placed in it before pairing of electrons in orbitals occurs

for degenerate orbitals, the lowest energy is attained when the number of electrons having the same spin is maximized

(section 6.8)

transition elements

a.k.a. transition metals

(section 6.8)

representative element

an element from within the s and p blocks of the periodic table

a.k.a. main-group elements

(section 6.9)

f-block metals

lanthanide and actinide elements in which the 4f and 5f orbitals are partially occupied

(section 6.9)