Chapter 4 Arrangement of Electrons in Atoms
Jan 17, 2016
Chapter 4
Arrangement of Electrons in Atoms
4-1 The Development of a New Atomic Model Rutherford’s model did not explain where
electrons were – what prevented electrons from being drawn into nucleus?
New model arose from experiments involving absorption and emission of light by matter
4-1 Properties of Light Light can behave as a wave Visible light is a kind of electromagnetic radiation (energy that
exhibits wavelike behavior as it travels through space) EM radiation includes X rays, UV and IR light, microwaves,
radiowaves
4-1 Properties of Light
All EM radiation moves at the same speed in a vacuum: 3.0 x 108 m/s
Wave motion is repetitive Wavelength (λ): distance between
corresponding points on adjacent waves (m, cm, nm)
Frequency (v): number of waves that pass a given point in a specific time, usually one second (1/s, Hz)
4-1 Properties of Light
4-1 Properties of Light
Since speed is constant, frequency and wavelength are related to each other mathematically
c = λv
Wavelength and frequency are INVERSELY proportional because their product is a constant.
4-1 Sample Problem
Determine the frequency of light with wavelength 550 nm. Convert nm to mUse formula c=λν to determine v
4-1 The Photoelectric Effect (Light as a Particle) 1900s – an experiment that
cannot be explained by the wave theory of light
Photoelectric effect – refers to the emission of electrons from a metal surface when light shines on the metal
For a given metal, no electrons are emitted if the light’s frequency is below a certain minimum, regardless of how intense the light or how long it is shone on the metal
4-1 The Photoelectric Effect
Wave theory predicts that ANY frequency of light could supply enough energy to eject an electron from the metal surface
Wave theory can’t explain why light must be of certain minimum frequency
4-1 The Particle Description of Light 1900 – Max Planck –
studying emission of light by hot objects
Proposed matter does not emit energy continuously but in small, specific amounts called quanta
Idea is called Quantum Theory. Planck wins Nobel prize in 1918 for his work.
4-1 The Particle Description of Light Quantum – minimum quantity of energy
that can be lost or gained by an atom Energy of a quantum is related to
frequency
E = hv
4-1 The Particle Description of Light 1905 – Einstein – light
has a dual nature – sometimes it acts like a wave, sometimes it acts like a particle
Light has wave properties Light is also like a stream
of particles, each particle carries a quantum of energy
Einstein called these particles photons
4-1 Explanation for the Photoelectric Effect Electrons are bound to the atom with a certain
amount of energy. Metal surface must be struck by a photon of light
carrying at least this amount of energy to knock the electron loose.
Energy and frequency are directly proportional. (E=hν)
Only frequencies equal to or greater than the threshold frequency will knock an electron off an atom.
4-1 Sample Problem
Calculate the energy associated with a photon of light of frequency 4.1 x 1014 Hz.
4-1 Hydrogen-Atom Line-Emission Spectrum Spectrum – a pattern of
energy observed when matter absorbs and emits energy
Ground state – lowest energy state of an atom or molecule
Excited state –state in which atom or molecule has higher PE than ground state
4-1 Hydrogen-Atom Line-Emission Spectrum
Current passed through vacuum tube with hydrogen gas inside
Pink light passed through prism to separate into specific frequencies of light
4-1 Hydrogen-Atom Line-Emission Spectrum Why does hydrogen give
off only specific frequencies of light?
1913 – Niels Bohr proposed a model for hydrogen atom that linked the atom’s electron with photon emission
Ties line emission spectrum to quantum theory.
4-1 Bohr Model of the Hydrogen Atom Electron can circle nucleus only in allowed
paths, or orbits Orbit closest to nucleus has lowest energy
(ground state) Orbits farther from nucleus have higher energy
(excited states) When electron absorbs energy, it jumps to
higher orbit When electron emits energy, it drops to lower
orbit
4-1 Bohr Model of the Hydrogen Atom Electron can only
exist in certain allowed orbits.
Can only absorb and emit amounts of energy that correspond to energy differences between orbits.
4-1 Bohr Model of the Hydrogen Atom Bohr’s model did not explain the spectra of
atoms with more than one electron Bohr’s theory did not explain the chemical
behavior of atoms
4-2 Electrons as Waves
It was already known that light can behave as a particle or a wave.
1924 – Louis deBroglie asked if electrons could also have dual wave-particle nature
4-2 deBroglie’s Hypothesis
Electrons are particles but they can act like waves
A wave confined to a space can only have certain frequencies – seems to correspond to Bohr’s quantized electron orbits
The electron-wave is confined to a certain space – the region around the nucleus – so electron-waves can only have certain frequencies, which correspond to certain energies (E = hv)
4-2 Wave-Particle Duality of Nature
Particles can have wave properties.
Waves can have particle properties.
4-2 Heisenberg Uncertainty Principle
If the electron is both a particle and a wave, where is it?
Werner Heisenberg, German physicist, 1927
Electrons are detected by hitting them with photons, but hitting them changes their position
It is impossible to determine simultaneously the position and velocity of an electron
4-2 The Schrodinger Wave Equation 1926 – Erwin
Schrodinger uses assumption that electron behaves as a wave to describe mathematically the wave properties of electrons and other very small particles (Quantum theory)
4-2 What does it mean?
Solutions to the Schrodinger equation are called wave functions
Wave functions can give probability of finding an electron at a particular position in the space around the nucleus
An orbital is a 3D region around the nucleus that indicates the probable location of an electron
4-2 Atomic Orbitals and Quantum Numbers Quantum numbers specify the properties
of atomic orbitals and the properties of electrons in orbitals.
Each electron in an atom can be assigned a set of four quantum numbers.
4-2 The Principal Quantum Number
Symbolized by n Indicates the main energy level
occupied by an electron Values of n are positive
integers (ex. n = 1 is the first energy level)
Principal quantum number also gives approximate distance from nucleus/size of energy level or shell
Total number of electrons that can exist in a given energy level, n, is equal to 2n2.
Energy level, n Maximum number of
electrons, 2n2
1 2
2 8
3 18
4 32
5 50
6 72
7 98
4-2 Angular Momentum Quantum Number
Symbolized by l Indicates the shape of the orbital “sublevels” For each energy level, n, the number of
orbital shapes possible is equal to n The first four shapes are given letter
symbols (s, p, d and f)
4-2 Magnetic Quantum Number
Symbolized by m Indicates the orientation of an orbital
around the nucleus
s orbital (1 orientation)
p orbital (3 orientations)
d orbitals (5 orientations)
f orbitals (7 orientations)
sublevel number of orbitals available
number of electrons sublevel can hold
s 1 2
p 3 6
d 5 10
f 7 14
4-2 Spin Quantum Numbers
Electrons in orbitals spin on internal axes. When charged bodies spin, they induce a
magnetic field. An electron can spin in one of two possible
directions. The spin quantum number has two possible
values, + ½ and – ½ A single orbital can hold a total of two electrons,
which MUST have opposite spins.
4-3 Electron Configuration
The arrangement of electrons in an atom Assigns an energy level and sublevel to
each electron in an atom.
4-3 Rules Governing Electron Configurations The Aufbau Principle – an electron
occupies the lowest-energy orbital available. (aufbau is German for “building up”
Electrons fill low energy orbitals before filling higher energy orbitals.
4-3 Electron Configuration
1s has the lowest energy.
Energies of sublevels in different main energy levels begin to overlap in n=3
Use orbital filling diagram to determine order in which sublevels are filled.
start
4-3 Rules Governing Electron Configurations Pauli Exclusion Principle –
no two electrons in the same atom can have the same set of four quantum numbers
In other words, if two electrons are going to occupy the same orbital, they must have opposite spin.
-let horizontal line represent orbital
-an up arrow and a down arrow represent two electrons of opposite spin
4-3 Rules Governing Electron Configurations Hund’s rule – orbitals of equal energy
(degenerate orbitals) are occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin
Bus seat rule
4-3 Ways to Represent Electron Configuration Electron
Configuration Notation
Assigns each electron to an energy level and a sublevel.
Examples:
Na
P
Br
Rb
K
Ar
4-3 Electron Configuration Sample Problems Name the elements indicated by the
following electron configurations:1s22s22p63s23p5
1s22s22p63s23p64s23d5
4-3 Electron Configuration Sample Problems Write the electron configuration for an
element that has the following number of electrons:7141933
4-3 Ways to Represent Electron Configuration Orbital Notation – uses lines and arrows to
represent orbitals and electrons Example: Write the orbital notations for
nitrogen and oxygen.
N
O
4-3 Ways to Represent Electron Configuration Noble Gas Notation – to simplify an
element’s electron configuration, use the preceding noble gas as shorthand to indicate all the electrons possessed by that noble gas
Example – Ne and Na
4-3 Valence Electrons
Valence electrons are electrons in the outermost energy level of an atom, farthest from the nucleus
They are important because they are the electrons that are usually involved in chemical reactions.
How many valence electrons does sodium have?
Bromine?
Silicon?
4-3 Electron Configurations with Special Stability Octet – the outer energy level is
considered filled when the s and p sublevels are completely filled with 8 electrons
A filled outer energy level (8 electrons) is a very stable electron configuration.
The noble gases have filled outer energy levels. This is why they are unreactive.
4-3 Electron Configurations with Special Stability Filled and half-
filled sublevels have special stability (especially d).
This fact sometimes results in electron configurations that deviate from the Aufbau principle.
Chromium
Copper
Molybdenum
Silver
Pig boots!