Topic 2.2: Electrons Honors Chemistry 2014-15 Mrs. Peters 1.

Topic 2.2: Topic 2.2: ElectronsElectrons

Honors Chemistry 2014-15Mrs. Peters

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2.2: Electron 2.2: Electron

ConfigurationConfiguration

Essential Idea: The electron configuration of an atom can be deduced from its atomic number.Nature of Science: •Developments in scientific research follow improvements in apparatus – the use of electricity and magnetism in Thomson’s cathode rays. (1.8)•Theories being superseded – quantum mechanics is among the most current models of the atom (1.9) •Use theories to explain natural phenomena – line spectra explained by the Bohr model of the atom (2.2)

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ConfigurationConfiguration Understandings: •Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level.•The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies.•The main energy level or shell is given an integer number, n, and can hold a maximum number of electrons, 2n2.

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ConfigurationConfiguration Understandings: (Continued)•A more detailed model of the atom describes the division of the main energy level into s, p, d, and f sub-levels of successively higher energies.•Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron.•Each orbital has a defined energy state for a given electronic configuration and chemical environment and can hold two electrons of opposite spin.

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ConfigurationConfiguration Applications and Skills:•Description of the relationship between colour, wavelength, frequency, and energy across the electromagnetic spectrum.•Distinction between a continuous spectrum and a line spectrum.

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ConfigurationConfiguration Applications and Skills: (Continued)•Description of the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels.•Recognition of the shape of an s atomic orbital and the px, py, and pz atomic orbitals.

•Application of the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z=36.

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Review of Topic 2.1Review of Topic 2.1

Let’s Review!•List how many protons, neutrons and electrons the following elements have:

o Lio Co Oo Mgo Po Aro Ca

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• Let’s Review!• List how many protons, neutrons and electrons

the following elements have: o Li: p = 3, n = 4, e = 3o C: p = 6, n = 6, e = 6o O: p = 8, n = 8, e = 8o Mg: p = 12, n = 12, e = 12o P: p = 15, n = 16, e = 15o Ar: p = 18, n = 22, e = 18o Ca: p = 20, n = 20, e = 20

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Atomic Structure Review:

Protons and Neutrons are located in the nucleus.Electrons are found in the electron cloud outside the nucleus.

This unit will focus on the electron cloud and where to find electrons.

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Bohr ModelBohr Model

Bohr Model for the Atom:•Nucleus in the center with protons and neutrons•Electrons in layers or levels around nucleus

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Bohr Model for the Atom:•Electrons are arranged in energy levels (layers)•Shows the number of electrons in each energy level.•Electron orbits are circular paths

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Bohr Model for the Atom:

•Which element is this the electron arrangement for?•How do you know?

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Bohr ModelBohr ModelBohr Model for the Atom:

•Useful for explaining and predicting chemical properties•Based on the fundamental idea that electrons exist in definite, discrete energy levels•Electrons can move from one energy level to another

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Bohr ModelBohr ModelBohr Model for the Atom:

Limitations of this model:•Assumes all orbits are fixed •Assumes all energy levels are circular•Suggests incorrect scale for atom

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U4. Quantum Mechanical U4. Quantum Mechanical ModelModel

Quantum Mechanical Model:•Sophisticated mathematical theory that incorporates wave-like nature of electrons•Based on 2 key ideas:

o Schrodinger’s Equationo Heisenberg’s uncertainty principle

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Heisenberg’s Uncertainty PrincipalIt is impossible to determine accurately both the momentum and the position of a particle simultaneously.

It is not possible to state precisely the location of an electron and its exact momentum, we can calculate the probability of finding an electron in a given region of space

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Schrodinger’s EquationoFormulated in 1926 by Austrian physicist Erwin SchrodingeroEquation integrates the dual wave-like and particle nature of the electronoDescribe atomic orbitals: a region in space where there is a high probability of finding an electron.

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U4. Quantum Mechanical U4. Quantum Mechanical

ModelModel

SublevelsThere are 4 types:s, p, d, and f

Each type has a characteristic shape, specific number of orbitals and associated energy.

Each orbital holds a maximum of 2 electrons

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U5. Sublevels & A4. S U5. Sublevels & A4. S

SublevelsSublevels

Sublevels

s: spherical shape, 1 orbital, holds 2 e-

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U5. Sublevels & A4. P U5. Sublevels & A4. P

SublevelsSublevels

Sublevels

P: dumbbell shaped, 3 orbitals, holds 6 e-

Draw the three sublevel shapes

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U5. SublevelsU5. Sublevels

Sublevelsd: double dumbbell shaped, 5 orbitals, holds 10 e-f: funky shaped, 7 orbitals, holds 14 e-

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The Quantum Mechanical ModelElectrons are not found at certain distances from the

nucleus but are located in a region in space that is described by a set of 4 quantum numbers. The exact location and path of the electron can’t be determined.

It estimates the probability of finding an electron within a certain volume of space surrounding the nucleus. Electron positions can be represented by a fuzzy cloud surrounding the nucleus (electron cloud).

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U3. Quantum Mechanical ModelU3. Quantum Mechanical Model

The Quantum Mechanical Model4 Quantum Numbers: n: energy level (called the principal quantum

number) l: sublevels ml: orbital

ms: spin

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U3. Quantum Mechanical ModelU3. Quantum Mechanical Model

The Quantum Mechanical Model

Each energy level can hold a maximum of electrons based on 2n2

Ex: if n is 3, there can be up to 2(3)2 electrons =18 e-

Argon has 18e- and is at the end of energy level 3

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A & S 5: Electron A & S 5: Electron ConfigurationConfiguration

Three principles (rules) must be followed when representing electron configurations:1. Aufbau Principle: electrons fill the lowest energy orbital first 2. Pauli Exclusion Principle: any orbital can hold a maximum of 2 electrons and those electrons have opposite spin.3. Hund’s rule of maximum multiplicity: when filling orbitals of equal energy, electrons fill all orbitals singly before occupying in pairs.

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Order of Energy Levels for ElectronsOrder of Energy Levels for Electrons

This order must be followed every time!Each level must be filled

before moving to the next level

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

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A&S 5. Electron A&S 5. Electron


Diagramming Electron Arrangement

There are two methods for diagramming electronarrangement

• Orbital Filling Diagram• Electron Configuration

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Orbital Filling Diagrams

1. Draw a box or line for each orbital and sublevel.

(s = 1 line, p = 3 lines, d= 5 lines)2. Place arrows to denote electrons. Maximum of 2

electrons per box. The first arrow is pointing up, the second arrow is pointing down to represent opposite spins.

3. Within a sublevel, each space must get an electron before the second electron is added ( ie: each p sublevel gets one before doubling up)

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A&S 5. Electron ConfigurationA&S 5. Electron Configuration

Orbital Filling

Draw orbital filling diagrams for the following atoms.

• H• Be• O

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Orbital Filling


• H __1s

• Be __ __ 1s 2s

• O __ __ __ __ __ 1s 2s 2p

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Orbital Filling


• H __1s

• Be __ __ 1s 2s

• O __ __ __ __ __ 1s 2s 2p

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Orbital Filling


• Al

• Ca

•

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• Al __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p

• Ca __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s

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Electron Configuration

Start with 1s and follow the order, filling each orbital with the maximum number of electrons until all the electrons in the atom have a place.

Write electron configurations for the following elements:H 1s1

Be 1s22s2

O 1s22s22p4

AlCa

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Electron Configuration

Start with 1s and follow the order, filling each orbital with the maximum number of electrons until all the electrons in the atom have a place.

Write electron configurations for the following elements:H 1s1

Be 1s22s2

O 1s22s22p4

Al 1s22s22p63s23p1 Ca 1s22s22p63s23p64s2

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• Let’s Review Ions!• What is an ion?• What are the two types of ions?• List how many protons, neutrons and electrons do

the following ions have: o Li+1

o O-2

o Mg+2

o P -3

o Al+3

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• List how many protons, neutrons and electrons do the following ions have: o Li+1 p = 3, n = 4, e = 2o O-2 p = 8, n = 8, e = 10o Mg+2 p = 12, n = 12, e = 10o P -3 p = 15, n = 16, e = 18o Al+3 p = 13, n = 14, e = 10

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A & S 5. Ion ConfigurationA & S 5. Ion Configuration

Cations lose electrons, Anions gain electrons

For orbital filling: add or subtract the number of electrons in the charge, draw in electrons

For Electron Configuration: add or subtract the number of electrons in the charge, fill in sublevels and orbitals

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A & S 5. Ion ConfigurationA & S 5. Ion Configuration

Practice:Compare Mg and Mg+2

Number of electrons:Electron configuration:

Compare O and O-2

Number of electrons:Electron configuration:

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A & S 5. Condensed A & S 5. Condensed


This is a short cut!•Full electron configurations become lengthy and cumbersome with increasing atomic number

•Condensed Configurations are more convenient

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Core Electrons: electrons that are in the inner energy levelsValence Electrons: electrons that are in the outer energy level

Condensed = [nearest noble gas] + valence electrons

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Condensed = [nearest noble gas core] + valence electrons

Ex: Oxygen: has 8 e-, 2 in core and 6 valence, nearest noble gas is He (which has 2 e-)[He] 2s2 2p4

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Condensed = [nearest noble gas core] + valence electrons

Ex: Cobalt: [Ar} 4s2 3d7

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Condensed = [nearest nobel gas core] + valence electrons

Practice: Write the condensed configurationCl:Zn:Mn:

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Periodic Table NotesPeriodic Table Notes

Let’s Label the periodic table!

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A & S 1: Wavelength, Frequency and A & S 1: Wavelength, Frequency and

EnergyEnergy

Light consists of electromagnetic waves that can travel through space and matter

All electromagnetic waves travel in a vacuum at the speed of light (c)

C = 3.0 x 108 m/s

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EnergyEnergy

Three components of electromagnetic waves

• Amplitude (y): Height from the origin to the crest

• Wavelength (λ) : Distance between the crests

• Frequency (ν): Number of wave cycles to pass a given point per unit time

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EnergyEnergy

Wavelength is related to the frequency of the radiation by the equation

c = λν

Draw and label the wave diagram.

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A & S 1. Electromagnetic A & S 1. Electromagnetic

spectrumspectrum

The electromagnetic spectrum is an arrangement of all of the types of electromagnetic radiation in increasing order of wavelength or decreasing frequency

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spectrumspectrum

The higher the frequency, the shorter the wavelength and the higher the energy

E = hfE = energy (Joules)

h = Planck’s constant (6.63 x 10-34 J s)

f = frequency (s-1)

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spectrumspectrum

Electromagnetic Spectrum Range:

• Radiowaves: long wavelength, low energy radiation

• Microwaves• Infrared radiation (IR)

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spectrumspectrum

Electromagnetic Spectrum Range:

• Visible Light: what we can see (ROYGBIV)

• Ultraviolet waves (UV)• X-Rays• Gamma Rays: high

energy radiation, short wavelength

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SpectrumSpectrum

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A & S 2. Continuous and Line SpectrumsA & S 2. Continuous and Line Spectrums

Continuous SpectrumEmission showing a

continuous range of wavelengths and frequencies; all the colors together without any space between them

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A & S 2. Continuous and Line SpectrumsA & S 2. Continuous and Line Spectrums

Line SpectrumEmission of specific

elements showing a series of discrete lines; individual lines of color with space between each line

On your paper, draw the difference between a continuous and line spectrum

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U1. Emission spectrumU1. Emission spectrum

Emission Spectrum: a series of lines against a black background (type of line spectrum)

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U1. Emission spectrumU1. Emission spectrum

Absorption Spectrum: a pattern of dark lines against a colored background

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A & S2. Continuous and line A & S2. Continuous and line spectrumspectrum

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A & S 2. continuous and line A & S 2. continuous and line spectrumspectrum

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We use an instrument called a spectroscope to detect the emission spectrum for a given source of

light.

U2. Line Emission Spectrum for U2. Line Emission Spectrum for HydrogenHydrogen

Why do we see different colors of light in line spectrums?• When electrons of a gaseous atom get excited, they are

raised to a higher energy level. • The extra energy is released as light when they drop back

down to lower energy levels.

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U2. Line Emission Spectrum for U2. Line Emission Spectrum for HydrogenHydrogen

Why do we see different colors of light in line spectrums?• The energy is provided by thermal or electrical energy• Each element has its own unique line spectrum

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U2. Line Emission Spectrum U2. Line Emission Spectrum for Hydrogenfor Hydrogen

• Must know the Hydrogen line series.• Draw this in your notes or on the back of the

Electromagnetic Spectrum paper. Indicate the colors and the wavelengths, notice the space between colors.

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A & S 3 Emission Spectrum of A & S 3 Emission Spectrum of HydrogenHydrogen

The line spectrum that we see from visible light is called the Balmer series.

Similar sets of lines can be seen in ultraviolet (Lyman Series) and Infrared (Paschen Series)

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Electrons move in orbits around the nucleus of the atom. Each orbit has a fixed amount of potential energy. The farther from the nucleus the orbit is, the more potential energy it has.

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• When electrons absorb energy they can move out to higher energy levels (the excited state).

• When they fall back to the lower energy level (the ground state) they emit a photon, a discrete amount of energy.

• Photon energy is seen as light.

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A & S 3 Emission Spectrum of HydrogenA & S 3 Emission Spectrum of Hydrogen

• Depending on how far they fall, different colors of light are given off.• Red = short fall• Violet = long fall

• The Balmer series of lines (visible light) is formed when the electrons fall back to the second energy level (n=2)

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A & S 3 Emission Spectrum of HydrogenA & S 3 Emission Spectrum of Hydrogen

Balmer Series: VisibleTransition of electrons from outer levels to n=2.

Spectral lines converge at increased values of n due to closer spacing of energy levels

Red: n=3 to n=2Blue-green: n=4 to n=2Blue: n=5 to n=2Violet: n=6 to n=2

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STOP HERE!STOP HERE!

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Topic 2.2: Electrons Honors Chemistry 2014-15 Mrs. Peters 1.

Documents

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electron configurations

electron orbits

electron arrangement

energy levels layersshows

discrete energy levels

energy transitions

existence of electrons