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Rutherford’s AtomElectromagnetic Radiation

Emission of Energy by AtomsEnergy Levels of Hydrogen

Atomic ModelsHydrogen Orbitals

Electron ArrangementsElectron Configuration

Atomic Properties and the Periodic Table pages 278-315

Unit 10 – Atomic Theory

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Atomic Theory• The concept of atoms explains many important

observations, such as why compounds always have the same composition (a specific compound always contains the same types and numbers of atoms) and how chemical reactions occur (they involve a rearrangement of atoms).

• We learned to picture the atom as a positively charged nucleus composed of protons and neutrons at its center and electrons moving around the nucleus in a space very large compared to the nucleus.

• In this unit, we develop a picture of the electron arrangements in atoms.

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Rutherford’s Atom

• Replay Video 3 - Ernest Rutherford and the nucleus

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Rutherford’s Atom• What are the electrons doing? How are they arranged

and how do they move?• Rutherford suggested the electrons might revolve

around the nucleus like the planets revolve around the sun in our solar system.

• He couldn’t explain why the negatively charged electrons aren’t attracted into the positive nucleus, causing the atom to collapse.

• More observations of the properties of atoms would be needed to understand the structure of the atom.

• To help understand these observations, we need to discuss the nature of light and how it transmits energy.

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Electromagnetic Radiation• Energy is transmitted by light, more properly called

electromagnetic radiation. Think of a light bulb, the flame of a Bunsen burner or the sun.

• Many different forms of this energy exist, including x-rays, microwaves and radio waves as well as light.

• A wave is characterized by three properties: wavelength, frequency and speed.

• Wavelength (symbolized by the Greek letter lambda, λ) is the distance between consecutive wave peaks.

• Frequency (symbolized by the Greek letter nu, ν) indicates how many waves pass a certain point in a given period of time.

• The speed of a wave indicates how fast a peak travels.

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Electromagnetic Radiation

Examples using ocean waves.

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Electromagnetic Radiation• Light (electromagnetic radiation) also travels as waves.• The various types of electromagnetic radiation (x-rays,

microwaves) differ in their wavelengths. X-rays have very short wavelength, while radio waves have very long wavelengths.

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Electromagnetic Radiation• Radiation provides an important means of energy

transfer. Energy from the sun reaches the earth mainly as visible and ultraviolet radiation.

• We visualize radiation (light) as a wave that carries energy through space. Sometimes, though, light does not act as a wave.

• Electromagnetic radiation can sometimes have properties that are characteristic of particles. Another way to think of light is as a stream of packets of energy called photons.

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Electromagnetic Radiation• Is light a wave or a stream

of particles of energy?• It appears to both. This is

called the wave-particle nature of light.

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Electromagnetic Radiation

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Emission of Energy by Atoms• When compounds are burned, they emit a color

characteristic of the cation. Li+, for example, emits a red flame when burned. Na+ burns with a yellow flame, Cu2+ with a green flame.

• The colors of the flames result from atoms releasing energy in the form of visible light of specific wavelengths, or colors.

• The heat from the flame causes the atom to absorb energy. The atom becomes excited. Some of the excess energy is released as light. The atom moves to a lower energy state as it emits a photon of light.

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Emission of Energy by Atoms

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Emission of Energy by Atoms• When atoms receive energy, they become excited.

They can release the energy by emitting light. The emitted energy is carried away by a photon.

• The energy of the photon corresponds exactly to the energy change of the emitting atom.

• High energy photons correspond to short wavelength light. Low energy photons correspond to long wavelength light.

• The photons of red light have less energy than the photons of blue light because red light has a longer wavelength than blue light.

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Energy Levels of Hydrogen• An atom with excess energy is said to be in an excited

state.• An excited atom can release some or all of its excess

energy by emitting a photon and thus move to a lower energy state. The lowest possible energy state of an atom is called its ground state.

• Different wavelengths of light carry different amounts of energy per photon.

• When a photon is emitted, the energy contained in the photon correspond to the change in energy that the atom experiences in going from the excited state to the lower state.

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Energy Levels of Hydrogen

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Energy Levels of Hydrogen

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Energy Levels of Hydrogen• When we study the photons of visible light emitted,

we see only certain colors.• Only certain types of photons are produced.• Because only certain photons are emitted, only

certain energy changes are occurring.• So, hydrogen atoms must have certain discrete

energy levels.• We say the energy levels of hydrogen are quantized,

that is, only certain values are allowed.• Energy levels of all atoms are quantized.

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Energy Levels of Hydrogen

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Energy Levels of Hydrogen

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Energy Levels of Hydrogen

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The Bohr Model of the Atom

This model proposed by Niels Bohr worked well to explain the hydrogen atom, but the model did not explain other atoms.

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Wave Mechanical Model

Louis de Broglie and Erwin Schrodinger developed the wave mechanical model. The model gives no information about when the electron occupies a certain point or how the electron moves.

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The Hydrogen OrbitalsThe probability map is called an orbital. The orbital shown in Figure 10.20 is called the 1s orbital and describes the ground (lowest) state of energy for hydrogen.

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The Hydrogen OrbitalsHydrogen has discrete energy levels. They are called principal energy levels and labelled with an integer. Each principal energy level has sublevels.

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The Hydrogen Orbitals

Principal level 2 has 2 sublevels. They are called 2s and 2p. Principal level 3 has 3 sublevels called 3s, 3p and 3d. Principal level 4 has 4 sublevels called 4s, 4p, 4d and 4f.

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The Hydrogen Orbitals

The principal levels describe size and shape. The s orbital is spherical. Level 1 is smaller than level 2, which is smaller than level 3.

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The Hydrogen OrbitalsThe three 2p orbitals are lobed, not spherical. They are oriented along the x, y or z axis.

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The Hydrogen Orbitals

The shapes of the five 3d orbitals are shown below.

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Electron Arrangements

• An atom has as many electrons as it does protons, so all atoms beyond hydrogen have more than one electron.

• Each electron appears to spin like a top on its axis. It can only spin in one direction. We represent spin with an arrow, ↑ or ↓. Electrons in the same orbital must have opposite spins.

• This leads to the Pauli exclusion principle: an atomic orbital can hold a maximum of two electrons and those two electrons must have opposite spins.

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Electron Arrangements• Hydrogen has an atomic number of 1 (Z =1) and

therefore a single electron to have a net charge of zero. To show its electron configuration, we write the principal energy level followed by the sublevel, 1s. The number of electron in the orbital is placed as a superscript, 1s1.

• The electron configuration can also be shown using an orbital diagram, or box diagram, as below.

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Electron Arrangements• Hydrogen (Z=1) 1s1

• Helium (Z=2) 1s2

• Lithium (Z=3) 1s2 2s1

• Berylium (Z=4) 1s2 2s2

• Boron (Z=5) 1s2 2s2 2p1

• Carbon (Z=6) 1s2 2s2 2p2

• Nitrogen (Z=7) 1s2 2s2 2p3

• Oxygen (Z=8) 1s2 2s2 2p4

• Fluorine (Z=9) 1s2 2s2 2p5

• Neon (Z=10) 1s2 2s2 2p6

The orbital diagram for nitrogen is below.

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Electron Arrangements

• Sodium (Z=11) 1s2 2s2 2p6 3s1 or [Ne] 3s1

• Magnesium (Z=12) [Ne] 3s2

• Aluminum (Z=13) [Ne] 3s2 3p1

• Silicon (Z=14) [Ne] 3s2 3p2

• Phosphorous (Z=15) [Ne] 3s2 3p3

• Sulfur (Z=16) [Ne] 3s2 3p4

• Chlorine (Z=17) [Ne] 3s2 3p5

• Argon (Z=18) [Ne] 3s2 3p6

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Electron Arrangements

• Valence electrons are the electrons in the outermost (highest) principal energy level of an atom. These are the electrons involved in bonding of atoms to each other.

• Also note that the atoms of elements in the same group have the same number of electrons in a given type of orbital, except that the orbitals are in different principal energy levels. Elements with the same valence electron arrangement show very similar chemical behavior.

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Electron Arrangements

• The order of filling orbitals changes for Z=19. Experiments show that the chemical properties of potassium are very similar to lithium and sodium. We predict that the 4s orbital will fill before the 3d orbital. This means that principal energy level 4 begins to fill before level 3 is full.

• Potassium (Z=19) [Ar] 3s2 3p6 3d 1

• Potassium (Z=19) [Ar] 3s2 3p6 4s1

• Calcium (Z=20) [Ar] 3s2 3p6 4s2

• Scandium (Z=21) [Ar] 3s2 3p6 4s2 3d1

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Electron Arrangements

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Electron Arrangements

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Periodic TrendsAs you go down a group, metals are more likely to lose an electron.

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Periodic TrendsIonization energy is the energy required to remove an individual atom in the gas phase. Metals have relatively low ionization energies.

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Periodic Trends

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Electron Configuration and Periodic Trends

HomeworkPages 311-312; problems 50, 54, 60, 80 and 82