Chapter 5 ARRANGEMENT OF ELECTRONS IN ATOMS THE

Chapter 5 ARRANGEMENT OF ELECTRONS IN ATOMS

THE ELECTROMAGNETIC SPECTRUM � Light is electromagnetic radiation (energy that exhibits wavelike behavior). � ER moves at a constant speed (c) of 3. 0 x 108 m/s (through a vacuum).

THE ELECTROMAGNETIC SPECTRUM

PROPERTIES OF LIGHT � Light consists of waves � Wavelength (λ) is distance b/t corresponding points on a wave. � Frequency (ν) is number of waves that pass a given point in a specific time (usually in one second).

PROPERTIES OF LIGHT (CON’T) � Wavelength and frequency are inversely proportional to one another; as wavelength increases, frequency decreases. � Speed of light= C = 3. 00 x 108 m/s �C= λν �Rearrange the eqn. to solve for frequency and wavelength! � Practice Problem � A certain green light has a frequency of 6. 26 x 1010 Hz. The speed of light is 3. 00 x 108 m/s. What is the wavelength? � λ = c/v = 3. 00 x 108 m/s 6. 26 x 10101/s = 4. 79 x 10 -3 m

THE PHOTOELECTRIC EFFECT � Light has a dual wave-particle nature; it does exhibit wave like properties but it can be thought of as a stream of particles with each photon carrying a quantum of energy. �E = hv �h = Plank’s constant (6. 626 x 10 -34 J/Hz) �E = energy of a quantum �v = frequency in hertz

SAMPLE PROBLEM #2 � Violet light has a wavelength of 4. 10 x 10 -7 m. The speed of light is 3. 00 x 108 m/s. What is the frequency of the light? �λ = 4. 10 x 10 -7 m � c = 3. 00 x 108 m/s �V = ? �C =λν �v = c / λ = 3. 00 x 108 m/s = 7. 32 x 1014 1/s or Hz 4. 10 x 10 -7 m

SAMPLE PROBLEM #3 � What is the energy of light of frequency 2. 13 x 1012 Hz? Plank’s constant = 6. 626 x 10 -34 J/Hz. � v = 2. 13 x 1012 Hz � h = 6. 626 x 10 -34 J/Hz �E = ? �E = hv = (6. 626 x 10 -34 J/Hz)(2. 13 x 1012 Hz) � E = 1. 41 x 10 -21 J

ATOMIC SPECTRA � � If a current is passed through a gas at low pressure, the electrons of the atoms become energized. � Ground state – lowest energy state of an atom � Excited state – condition in which the atom has more energy than its ground state. When the atom returns to ground state, it gives off electromagnetic radiation (light). � Example: Neon Signs

HYDROGEN’S LINE EMISSION SPECTRUM This diagram is called a line-emission spectrum which shows the frequencies of light emitted by hydrogen when its electrons return to ground state.

BOHR’S MODEL � Linked the atom’s electron with photon emission. � Electrons are allowed to circle the nucleus only in fixed paths (orbits). � Atom is in the lowest energy level when the electron is closest to the nucleus. � An electron can move to a higher orbital by gaining energy. � When an electron falls from an excited state, a photon is emitted � Bohr’s model mathematically explained the lineemission spectrum of hydrogen.

QUANTUM THEORY

QUANTUM THEORY � Electrons, like light, are found to have a dual wave-particle nature. But where are they? � Heisenberg uncertainty principle states that it is impossible to know the velocity and position of a particle at the same time. � This and Schrödinger’s wave equation lead to modern quantum mechanics, which describes the motions of subatomic particles as waves.

ORBITALS � Electrons are not in neat fixed orbits like Bohr proposed but in orbitals. � 3 -D region around the nucleus that indicates the probable location (90%) of an electron. � Each electron has an “address” consisting of four Quantum numbers.

QUANTUM NUMBERS Quantum Number Type Abbre v. Meaning Principle Quantum Number n Main energy level occupied by electrons (1, 2, 3, 4, etc) i. e. which row it’s in Angular Momentum Quantum Number l Indicates the shape of the orbital (s, p, d, f) Magnetic Quantum Number m Indicates the orientation of an orbital around the nucleus (For p, -1, 0, +1) Spin Quantum Number s Two possible values (+ ½ or – ½) which indicate the spin value of an electron in an orbital.

S-ORBITALS �S orbital has one shape. � It can hold 2 electrons � For s, m=0

P-ORBITALS � 3 different orientations of the p orbitals. � Each orientation can hold two electrons. � Each corresponds to an m value m = -1, 0, or +1

D-ORBITALS 5 different orientations � Each can hold two electrons � m = -2, -1, 0, +1, or +2 �

HOW DO WE USE THIS INFO? � The periodic table is organized based on electron configuration.

WRITING ELECTRON CONFIGURATIONS

AUFBAU RULE: Add electrons to the lowest energy levels and sublevels first and then go to the next level until all of the electrons have been accounted for. Diagonal rule: the order of filling

Noble gases He Examples 1 H 2 He n = 1 s 1 (1 electron, close to the nucleus) Ne 1 s 2 (2 electrons, both in the s-sublevel) Ar Kr 3 Li 1 s 22 s 1 Xe Rn 10 Ne 28 Ni 1 s 22 p 6 [He]2 s 22 p 6 1 s 22 p 63 s 23 p 6 4 s 23 d 8 Argon’s configuration 17 Cl 1 s 22 p 6 3 s 23 p 5 Neon’s configuration Short-Cuts 18 e 2 [Ar]4 s 3 d 8 Ar Row above element given 10 e - [Ne]3 s Ne 23 p 5 Same row (n) as element given Short-cut electron configuration: build configuration on the noble gas that ends the previous row on the periodic table. Begin with the “n’s” (1, 2, 3 …) , always followed by the s clouds first (1 s, 2 s, 3 s etc. )and continue to fill electrons until you are at your given element.

HUND’S RULE: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results. 1 s 2 2 s 2 nitrogen 2 p 3

Pauli Exclusion Principle: Each orbital can hold a maximum of 2 electrons. To occupy the same orbital, 2 electrons must spin in opposite directions. [Ne] 2 s 2 2 p 6 When 2 electrons share an orbital, they are called “paired”.