Chapter 9 Electronic Structure and Periodic Trends

The Structure of Atoms

Plum Pudding Model Chapter 9

Tootsie Pop Model 2

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The Dual Nature of Light: The Particle and The Wave • From the time of the ancient Greeks, people have thought of light as a stream of tiny particles - like marbles or billiard balls • Thomas Young (in 1807) performed the now classic double slit experiment to test this theory. For more info about these topics: http://www.colorado.edu/physics/2000/index.pl?Type=TOC

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A great animation of the idea of light as waves and particles can be found here: http://video.google.com/videoplay?docid=-4237751840526284618&q=quantum

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The Dual Nature of Light: The Wave • As evidenced by the double slit experiment, light travels through space as a wave, similar to an ocean wave. • Wavelength (λ) is the distance light travels in one cycle. • Frequency (ν) is the number of wave cycles completed each second. • Amplitude is the height measured from the center of the wave. The square of this value gives the Intensity.

Velocity (c) = λ • ν • Light has a constant velocity (c) of 3.00 × 108 m/s. Chapter 9

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The Dual Nature of Light: The Particle • In 1900, Max Planck proposed that radiant energy is not continuous, but is emitted in small bundles of energy called photons. – This is the quantum concept. • It was found that when light is emitted, it behaves like a stream of small particles that move in a wave-like fashion. • The energy of an emitted photon is directly proportional to the frequency of the emitted light

E = hν Chapter 9

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The Electromagnetic Spectrum • The complete electromagnetic spectrum (all possible wavelengths and frequencies) is an un-interrupted band, or continuous spectrum. • The radiant energy spectrum includes most types of radiation, most of which are invisible to the human eye. ROY G BIV

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Atomic Spectra • When a particular element is heated and the light emitted is passed through a prism, the resulting breakdown results in a non-continuous spectrum or a discrete line spectrum • This spectrum is called the element’s atomic spectrum • Each element has its own, unique spectrum • This discovery is one important reason for Bohr’s hypothesis regarding hydrogen. Chapter 9

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Quantized versus Continuous • The quantum concept means that the energy of the electron and its radius of orbit around the nucleus is limited to specific values (and cannot be anywhere in between!). • If these values were continuous, they would be free to have any value. Energy E3 hv = E3 – E2

E2 hv = E3 – E1 hv = E2 – E1 Chapter 9

E1

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The Bohr Model of the Atom • In 1913, Niels Bohr suggested a new model of the atom that explained why hydrogen had a discrete line spectrum rather than a continuous spectrum. • Bohr's basic theory: electrons in atoms can only be at certain energy levels, and they can give off or absorb radiation only when they jump from one level to another. • In his model that an atom consists of an extremely dense nucleus that is surrounded by electrons that travel in set orbits around the nucleus. • He hypothesized that the energy possessed by these electrons and the radius of the orbits are quantized, meaning it is limited to specific values and is never between those values. • These “orbits” were of varying energies, The dependent on their distance from the nucleus

Gobstopper Model

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The Bohr Theory: Problems • Bohr’s theory only works for hydrogen atoms. • Once you have more than one electron, the calculations for the electron energy and the orbit radii breakdown. • Therefore, a new theory needed to be developed for multi-electron elements. • Bohr’s theory did make two important contributions: – It suggested a reasonable explanation for the discrete line spectrum of the elements – It introduced the idea of quantized electron energy levels (orbits!) Chapter 9

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The Quantum Mechanical Model of the Atom • In the 1920’s Erwin Schrödinger applied the principles of wave mechanics to atoms and developed the Quantum Mechanical Model of the Atom • This theory builds on Bohr’s idea of quantized energy levels (orbits) and adds additional requirements for electron location and energy. • These requirements are summarized by the four Quantum Numbers • These four numbers give a type of “address” for each electron within an atom. Chapter 9

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Quantum Numbers • In order to identify the probable location of an electron within an atom, four Quantum Numbers were established. • There are four quantum numbers: – – – –

n is the principal energy level (Bohr’s Orbits!) l is the sublevel ml is the orbital ms is the electron spin

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Principal Energy Level (n) • Principal Energy Level (n): provides info about the distance of the electron from the nucleus – the higher n is, the further away the electron is from the nucleus – n also helps to determine the energy of the electron because the further you are away from the nucleus, the higher the energy

Allowed Values: n = 1, 2, 3, ... Chapter 9

never 0 13

Sublevel (l) • It was later shown that the principal energy levels were split into discrete sublevels • l identifies the sublevel location of an electron • The number of sublevels within an energy level is equal to the principal quantum number (n)

l Value: Letter Used:

0 s

1 p

2 d

3 f

increasing energy Chapter 9

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Orbitals (ml) • Each sublevel is further split into orbitals (not orbits!) • The location of the electron within the sublevel is designated by the orbital (ml) – The orbital is not as discrete as Bohr’s Orbit. It is a region in space around a nucleus where there is a high probability of finding the electron. – This “probability” idea is based on the Uncertainty Principle which states that you cannot know both the position of the electron and its velocity at the same time.

• The number of orbitals per sublevel is determined by: 2l + 1

l

0 (s) 1 (p) 2 (d)

3 (f)

2l + 1 Chapter 9

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Shapes of the Orbitals • Each orbital has a specific shape. • As you increase the principle energy level, the orbitals increase in size!

s p

• Remember, these orbitals represent a region in space where there is a high probability of finding the electron. These are not discrete locations! Chapter 9

d

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Electron Occupancy in Sublevels • The Pauli Exclusion Principle states that an orbital can hold up to two electrons • The maximum number of electrons in each of the energy sublevels depends on the number of orbitals within that sublevel: – The s sublevel holds a maximum of 2 electrons (1 orbital). – The p sublevel holds a maximum of 6 electrons (3 orbitals). – The d sublevel holds a maximum of 10 electrons (5 orbitals). – The f sublevel holds a maximum of 14 electrons (7 orbitals)

• The maximum electrons per energy level (n) is obtained by adding the maximum number of electrons in each sublevel. Chapter 9

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Electrons per Energy Level

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Electron Configurations • Many of an element’s chemical properties depend on its electron configuration • The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. • The principal energy level (n) is written first, followed by the letter designation of the sublevel (l) then a superscript with the number of electrons in the sublevel. – Electrons are arranged about the nucleus in a regular manner. – The first electrons fill the energy sublevels closest to the nucleus. – Electrons continue filling each sublevel until it is full and start filling the next closest sublevel. Chapter 9

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Electron Configurations: Aufbau Principle Pauli Exclusion Principle: No two electrons in an atom can have the same quantum numbers (n, l, ml, ms). Hund’s Rule: When filling orbitals in the same sublevel, maximize the number of parallel spins (so fill then pair!). Rules of Aufbau Principle: 1. Lower n orbitals fill first. 2. Each orbital holds two electrons; each with different ms. 3. Half-fill degenerate orbitals before pairing electrons. Chapter 9

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Using the Periodic Table for Electron Configurations • The periodic table took on its current shape once the quantum model of the atom was developed. • You can use it to fill up your sublevels and orbitals to build your electron configurations.

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Writing Electron Configurations • Step 1: Locate the element on the periodic table. • Step 2: Determine the number of electrons the element Iron has 26 electrons has: • Step 3: Starting at the beginning of the Periodic Table, move left to right across the periods, filling each sublevel with electrons until you reach the location of your element:

• Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6 • Step 4: Check that the sum of the superscripts equals the atomic number of iron (26). Chapter 9

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An Alternative Method

Increasing Energy

Core [He] [Ne] [Ar] [Kr] [Xe] [Rn]

1s 2s 3s 4s 5s 6s 7s

2p 3p 4p 5p 6p 7p

3d 4d 4f 5d 5f 6d

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Writing Electron Configurations • Give the ground-state electron configurations for: Ne (Z = 10)

Mn (Z = 25)

Zn (Z = 30)

Eu (Z = 63)

W (Z = 74)

Lw (Z = 103)

• Identify elements with ground-state configurations: 1s2 2s2 2p4

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d6

1s2 2s2 2p6

[Ar] 4s2 3d1

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[Xe] 6s2 4f14 5d10 6p5

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Exceptions to the Filling Order • When filling the 3d sublevel, exceptions occur for the chromium (Cr) and copper (Cu) families:

Cr

Cu

4s

3d

4p

4s

3d

4p

4s

3d

4p

4s

3d

4p

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Noble Gas Core Electron Configurations • Recall, the electron configuration for Na is: Na: 1s2 2s2 2p6 3s1 • We can abbreviate the electron configuration by indicating the innermost electrons with the symbol of the preceding noble gas. • The preceding noble gas with an atomic number less than sodium is neon, Ne. We rewrite the electron configuration: Na: [Ne] 3s1 Chapter 9

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Valence Electrons • When an atom undergoes a chemical reaction, only the outermost electrons are involved. • These electrons are of the highest energy and are furthest away from the nucleus. These are the valence electrons. • The valence electrons are the s and p electrons beyond the noble gas core. Chapter 9

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Predicting Valence Electrons • The Group number indicates the number of valence electrons.

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Periodic Trends • The arrangement of the periodic table means that the physical properties of the elements follow a regular pattern. • Some trends include: – Atomic Radius – Ionization Energy – Metallic Character – Electronegativity (Chapter 6) Chapter 9

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Atomic Radius • An atom’s atomic radius is the distance from the nucleus to the outermost electrons.

Why do you think the radius increases in this way? Chapter 9

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Ionization Energy • The ionization energy of an atom is the amount of energy required to remove an electron in the gaseous state.

• The closer the electron is to the nucleus, the harder it is to remove. Chapter 9

Why??

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Metallic Character • Metallic character is the degree of metal character of an element. • A metal is defined as an element that can lose electrons and form positively charged ions (cations). – The easier the loss of these electrons, the more metallic the element.

• Metallic character decreases left to right across a period and from bottom to top in a group.

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