CHAPTER 6 Electronic Structure and Periodic Properties of Elements

electromagnetic radiation

energy transmitted by

waves that have an electric-field component and a

magnetic-field component

electromagnetic spectrum

range of energies that

electromagnetic radiation can comprise, including radio, microwaves, infrared, visible, ultraviolet, X-rays, and gamma rays

wave

oscillation of a property over time or space;

can transport energy from one point to another

constant speed of light (c)

2.998 × 108 m/s

What is wavelength (λ)?

The distance between two consecutive peaks or troughs in a wave.

What determines light's color?

The wavelength and its frequency of the light.

In what units is wavelength measured?

nanometers (nm) or millimeters (mm), m,

frequency (v and nu)

number of wave cycles (peaks or troughs) that pass a specified point in space per unit time

amplitude

extent of the displacement caused by a wave

Hert (Hz)

the unit of frequency, which is the number of cycles per second, s^-1

Brightness

Amplitude is related to the intensity of the wave for light

Loudness

Amplitude related to the intensity of the wave for sound

speed of the wave

= speed of light = wavelength x frequency

Increasing Energy (E)

Increasing frequency and decreasing wavelength

increasing frequency (v)

increasing energy and decreasing wavelength

increasing wavelength

decreasing frequency, energy, mass

interference patterns

A pattern typically consisting of alternating bright and dark fringes; it results from constructive and destructive interference of waves

Standing waves (stationary waves)

remain constrained within some region of space

quantization

A limitation of some property to specific discrete values, not continuous

nodes (n-1)

any point of a standing wave with zero amplitude

blackbody

idealized perfect absorber of all incident electromagnetic radiation; such bodies emit electromagnetic radiation in characteristic continuous spectra called blackbody radiation

Planck's Quantum Theory

assume that the vibrational energies of atoms are not continuous but are quantized

Planck's Equation

E =nhv

E

energy of the vibrating atom, = hc/wavelength

n

integer (1, 2, 3, ...) representing the discrete energy levels

h

Planck's constant (approximately 6.626×10^−34 Js)

photon

smallest possible packet of

electromagnetic radiation, a particle of light

line spectra

the light emission only at specific wavelengths

Atomic Theory

Niels Bohr developed a theory that posits that all matter is composed of tiny particles called atoms.

force of attraction

A force that pulls objects together equals centrifugal force

ground electronic state

state in which the electrons in an atom, ion, or molecule have the lowest energy possible

excited electronic state

state having an energy greater than the ground-state energy

ionization energy

The amount of energy required to remove an electron from an atom (change in E)

ionization energy equation

|change E| = |Ef-Ei| = hv = hc/wavelength

Light

electromagnetic radiation, a form of energy that is made up of oscillating electric and magnetic fields

7 rays of light (low to high energy)

Radio, microwave, infrared, visible light, ultraviolet, x-ray, gamma ray

Red wavelength

620-750 nm

orange wavelength

590-620 nm

yellow wavelength

570-590 nm

green wavelength

500-570 nm

blue wavelength

450-500 nm

violet wavelength

380-450 nm

number of photons equation

#P = Etotal / Ephotons

Avogadro's number

6.022x10^23 ( 1 moles, particles, molecules)

Heisenberg uncertainty principle

It is fundamentally impossible to determine simultaneously and exactly both the momentum and the position of a particle.

wavefunctions

mathematical description of an

atomic orbital that describes the shape of the orbital; it can be used to calculate the probability

of finding the electron at any given location in the orbital, as well as dynamical variables such as the energy and the angular momentum

principle quantum number

number having only specific

allowed values and used to characterize the

arrangement of electrons in an atom

shell

atomic orbitals with the same principal quantum number, n (higher shell, higher energy)

s orbital (sublevel)

a spherical region of space with high electron density, describes orbitals with l= 0

p orbital (p sublevel)

A dumbbell-shaped region of space with high electron density, describes orbitals with l = 1

f orbital (f sublevel)

multilobed region of space with high electron density, describes orbitals with l = 3

d orbital (d sublevel)

region of space with high electron density that is either four lobed or contains a dumbbell

and torus shape; describes orbitals with l = 2.

principal quantum number

n, shell, the general region for the value of energy for an electron on the orbital

angular momentum or azimuthal quantum number

l, subshell, the shape of the orbital

l equation

0 ≤ l ≤ n - 1

magnetic quantum number

ml, orientation of the orbital

ml equation

- l ≤ ml ≤ l

spin quantum number

ms, direction of the intrinsic quantum "spinning" of the electron, Up (1/2) or down (-1/2)

Pauli Exclusion Principle

No two electrons in the same atom can have exactly the same set of all the four quantum numbers

Electron Configurations

listing that identifies the electron occupancy of an atom's shells and

subshells

What is the order of orbital energy levels from lowest to highest?

s < p < d < f

What does it mean when orbitals are less penetrating?

Less electron density is found near the nucleus.

electron density

A measure of the probability of locating an electron in a particular region of space.

How is electron density calculated

It is equal to the squared absolute value of the wave function ψ.

the Aufbau Principle

A procedure in which the electron configuration of the elements is determined by 'building' them in order of atomic numbers.

How does the Aufbau Principle work?

By adding one proton to the nucleus and one electron to the proper subshell at a time.

valence electrons

The electrons occupying the outermost shell orbital(s) (highest value of n)

core electrons

electrons occupying the inner shell orbitals

covalent/atomic radius

one-half the distance between the nuclei of two identical atoms when they are joined by a covalent bond

atomic radius increase

down and to the left

Ionic Radii

Ion size based on the amount of electrons.

ionic radii increasing

down a column.

What is the order of ionic radii size from largest to smallest?

Anion, neutral, cation.

effective nuclear charge (Zeff)

amount of charge felt by the most recently added electron (Zeff = Z- S (core e-))

isoelectronic

A group of ions or atoms that have

identical electron configurations

first ionization energy (IE1)

energy required to remove the first electron, always endothermic and positive, kj/mol

first ionization energy (IE1) equation

X(g) -> X⁺(g) + e⁻

electron affinity (EA)

energy change associated with

addition of an electron to a gaseous atom or ion, is always negative

electron affinity (EA) equation

X⁺(g) + e⁻ -> X-

electron affinity increasing

up and to the right

ionization energy exceptions

Group 2A (higher, more unstable) to Group 3A and Group 5A (higher) to Group 6A, Ionization energy decreases

Why does removing a full or half-full orbital require more ionization energy?

Because it is more stable and requires more energy to remove an electron.

electron affinity exceptions

1A to 2A (lower) and 4A (lower) to 5A and 8A (lower)

Why does adding to a full or half-full orbital require less electron affinity?

because it is causing more repulsion, making the element bigger, further from the nucleus