Electrons in Atoms

Electrons in Atoms • Rutherford’s model has some limitations • It did not explain the chemical properties of the elements • It did not address the electrons

Low energy ( = 700 nm) High energy ( = 380 nm) Frequency  (s-1) 3 x 106 3 x 1012 3 x 1022 10-8 10-14 102 Electrons in Atoms According to the wave model, light consists of electromagnetic waves.

Electrons in Atoms • Visible light of different wavelengths can be separated into a spectrum of colors. • In the visible spectrum, red light has the lowest energy. • Violet light has the highest energy.

Electrons in Atoms • In 1913, Niels Bohr develops a new atomic model • Bohr stated that the electrons orbit the nucleus like the planets orbit the sun.

Electrons in Atoms • Each possible electron orbit in Bohr’s model has a fixed energy. • The fixed energies an electron can have are called energy levels. • Each energy level further from the nucleus is of greater energy

Electrons in Atoms • The rungs on this ladder are like the energy levels in Bohr’s model. • A person on a ladder cannot stand between the rungs. • Similarly, the electrons in an atom cannot exist between energy levels.

Electrons in Atoms • The rungs on this ladder are like the energy levels in Bohr’s model. • The energy levels in atoms are unequally spaced, like the rungs in this unusual ladder. • The higher energy levels are closer together.

Electrons in Atoms • The Rutherford model could not explain why elements that have been heated to higher temperatures give off different colors of light. • The Bohr model explains how the energy levels of electrons in an atom change when the atom emits light.

Electrons in Atoms • When an electron occupies the lowest possible energy level – it is said to be in its ground state • An electron can absorb energy from an external source • Sun • Fire • Electricity

Electrons in Atoms • When the electron absorbs energy it can jump up to higher energy levels – this is called the excited state • However, the electron cannot stay in an excited state • When it returns to its ground state it must release the energy it has absorbed

Electrons in Atoms • It releases the energy in the form of light • This is called the emission line spectrum • No two elements have the same line emission spectrum

Electrons in Atoms • Bohr’s model was based on the emission line spectrum

Electrons in Atoms • Unfortunately, Bohr’s model did not apply to other atoms • That led scientists to question his model • They wondered why the electron had to be located in a precise orbit

Electrons in Atoms • That led to the Heisenburg uncertainty principle • states that it’s impossible to determine both the position and velocity of an electron

Electrons in Atoms • Further developments led to Schrodinger developing the quantum mechanical model • This model describes mathematically the position of electrons in an atom • It is based on the allowed energies an electron can have • It shows how likely it is to find an electron in a particular location around the nucleus of an atom.

Electron cloud Electrons in Atoms • The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloudlike region called an orbital. • The cloud is more dense where the probability of finding the electron is high.

Electrons in Atoms • To specify the probable location of an electron in an atom, chemists use quantum numbers • Principle quantum number • Angular momentum quantum number • Magnetic quantum number • Spin quantum number

Electrons in Atoms • The principle quantum number indicates the main energy level occupied by the electron • Often referred to as the energy level • The energy level corresponds to the periods of the periodic table • 1st energy level = period 1

Electrons in Atoms • The angular momentum quantum number indicates the shape of the orbital • l = s, p, d, f • The s orbitals are spherical. • The p orbitals are dumbbell shaped.

Electrons in Atoms • The magnetic quantum number indicates the orientation of the orbital • m = x, y, z

Electrons in Atoms • Parts of the periodic table corresponds to each orbital shape • Groups 1 & 2 – s block • Groups 13-18 – p block • Groups 3-12 – d block • Bottom two rows – f block

Electrons in Atoms • The spin quantum number indicates the spin of the electron • Itmay be thought of as clockwise or counterclockwise. • A vertical arrow indicates an electron and its direction of spin ( or ).

Electrons in Atoms • The numbers and types of atomic orbitals depend on the principal energy level. (each orbital can hold up to 2 electrons)

Electrons in Atoms • To describe the arrangement of the electrons in an atom we use electron configuration • To describe spin we use orbital notation

Electrons in Atoms • To determine electron configuration, follow three simple rules • The Aufbau principle states that an electron occupies the lowest energy orbital that can receive it • We fill lowest to highest energy

6p 5d 6s 4f 5p 4d 5s 4p 4s 3d 3p 3s Increasing energy 2p 2s The aufbau diagram shows the relative energy levels of the various atomic orbitals. Orbitals of greater energy are higher on the diagram. 1s Electrons in Atoms

Electrons in Atoms • To determine electron configuration, follow three simple rules • The Pauli Exclusion principle states that no two electrons in the same atom can have the same set of 4 quantum numbers • Hunds Rule states that orbitals of the same energy must be occupied by one electron before it can be occupied by a second electron.

Electrons in Atoms

Electrons in Atoms • For larger atoms, the electron configuration can be tedious • We use a shorter method called the noble gas electron configuration

Electrons in Atoms