teori atom

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1. ATOMIC THEORY 2. IONIC BONDING3. COVALENT BONDING4. SPECTROSCOPY5. SPECIAL TOPICS

Dr. Akhmad Syoufian

Department of ChemistryFaculty of Mathematics and Natural SciencesGadjah Mada University

ATOMIC THEORY

Dr. Akhmad Syoufian


Modern atomic theory:

In the early years of the 19th century, John Daltondeveloped his atomic theory in which he proposed that each chemical element is composed of atoms of a single, unique type, and that though they are both immutable and indestructible, they can combine to form more complex structures (chemical compounds). How precisely Dalton arrived at his theory is not entirely clear, but nonetheless it allowed him to explain various new discoveries in chemistry that he and his contemporaries made.

Modern atomic theory (continued):

Atoms were thought to be the smallest possible division of matter until 1897 when J.J. Thomsondiscovered the electron through his work on cathode rays. He thus concluded that atoms were divisible, and that the “corpuscles” (they would later be renamed electrons by other scientists) were their building blocks. To explain the overall neutral charge of the atom, he proposed that the corpuscles were distributed in a uniform sea or cloud of positive charge; this was the plum pudding model


Thomson's plum pudding model was disproved in 1909 by one of his students, Ernest Rutherford, who discovered that most of the mass and positive charge of an atom is concentrated in a very small fraction of its volume, which he assumed to be at the very center.


The gold foil experimentTop: Expected results: alpha particles passing through the plum pudding model of the atom with negligible deflection.Bottom: Observed results: a small portion of the particles were deflected, indicating a small, concentrated positive charge.

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Quantum theory revolutionized physics at the beginning of the 20th century, when Max Planckand Albert Einstein postulated that light energy is emitted or absorbed in discrete amounts known as quanta (singular, quantum).

In 1913, Niels Bohr incorporated this idea into his Bohr model of the atom, in which the electrons could only orbit the nucleus in particular circular orbits with fixed angular momentum and energy, their distances from the nucleus being proportional to their respective energies.


The Bohr model of the atom; Under this model electrons could not spiral into the nucleus because they could not lose energy in a continuous manner; instead, they could only make instantaneous "quantum leaps" between the fixed energy levels. When this occurred, light was emitted or absorbed at a frequency proportional to the change in energy.


Bohr's model was only able to predict the spectral lines of hydrogen; it couldn't predict those of multi-electron atoms. Worse still, as spectrographic technology improved, additional spectral lines in hydrogen were observed which Bohr's model couldn't explain.

In 1916, Arnold Sommerfeld added elliptical orbits to the Bohr model to explain the extra emission lines, but this made the model very difficult to use, and it still couldn't explain complex atoms.


In 1924, Louis de Broglie proposed that all moving particles-particularly subatomic particles such as electrons-exhibit a degree of wave-like behavior.

Erwin Schrödinger, fascinated by this idea, explored whether or not the movement of an electron in an atom could be better explained as a wave rather than as a particle.

Schrödinger's equation, published in 1926, describes an electron as a wave-function instead of as a point particle, and it elegantly predicted many of the spectral phenomena Bohr's model failed to explain.


The five atomic orbitals of a neon atom, separated and arranged in order of increasing energy from left to right, with the last three orbitals being equal in energy. Each orbital holds up to two electrons, which exist for most of the time in the zones represented by the colored bubbles. Each electron is equally in both orbital zones, shown here by color only to highlight the different wave phase.


Since a wave-function incorporates time as well as position, it is impossible to simultaneously derive precise values for both the position and momentum of a particle for any given point in time; this became known as the uncertainty principle.

This invalidated Bohr's model, with its neat, clearly defined circular orbits.

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The modern model of the atom describes the positions of electrons in an atom in terms of probabilities. An electron can potentially be found at any distance from the nucleus, but - depending on its energy level - tends to exist more frequently in certain regions around the nucleus than others; this pattern is referred to as its atomic orbital. THANK YOU FOR YOUR ATTENDANCE

@ 2008


ENERGIES AND ORBITALS IN MANY-ELECTRON ATOMS

Dr. Akhmad Syoufian


Many-Electron Atoms: The First Two Rules:

The quantum state of an electron is specified by the orbital quantum numbers, n, l, and m, plus an electron spin quantum number s. So far we have neglected this property of electrons, and we will not say any more about it at this stage except to note that scan have one of two values, +½ or -½.

The Pauli Exclusion Principle says that no two electrons in an atom may have be in the same quantum state.

Many-Electron Atoms: The First Two Rules:

That is, no two electrons can have the same four quantum numbers, n, l, m, and s.

This is equivalent to saying that no orbital (specified by n, l, and m) can be occupied by more than two electrons.

The second rule is that electrons in atoms (and molecules) generally exist in their lowest possible energy state. This is called the ground state.

This is enough to begin to handle multi-electron atoms, at least He.

Worked Example: The Ground State Electronic Configuration of He: Using the atomic orbitals obtained for the hydrogen atom, we

fill orbitals beginning with the lowest energy. (To do this we are pretty much ignoring the interactions between electrons, and treating them as two independent waves bound to the same (2+) nucleus.)

Electron 1 goes into the 1s orbital (n=1, l=0, m=0) with s = +½

Electron 2 goes into the 1s orbital (n=1, l=0, m=0) with s = -½

The ground state electron configuration of He is written as 1s2. (For H it is written 1s1.)

What happens to the next electron? What is the ground state configuration of Li?

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Filling the n = 2 orbitals: Rule 3:

After He, the n=1 (1s) orbital is full. According to the wave equation for the hydrogen atom, the 2s and three 2p orbitalsall have the same energy, so the next electron could go into any of the four n=2 orbitals.

However we have already seen that the wave-functions for the s and p orbitals are different.

s orbitals have their maximum amplitude at the nucleus. This means that electrons in s orbitals are bound by the true nuclear charge (3+ for Li, etc.)

p orbitals have a node at the nucleus. Their interaction with the nucleus is screened by electrons closer in, so electrons in 2p orbitals are bound by a lower effective charge.


The different effective nuclear charges lower the energy of the ns orbital relative to np, so the s orbital fills first with up to 2 electrons. The ground state configuration of Li is 1s22s1, and for Be it is 1s22s2.


After Be, the 1s and 2s orbitals are full. The 2p orbitals are next to fill.

Three 2p orbitals can accommodate a total of six electrons, which gives the configurations of elements B through to Ne.


This is summarized in Hund’s Rule, that the lowest energy electron configuration in orbitals of equal energy is the one with the maximum number of unpaired electrons with parallel spins.

Hund’s Rule:

Electron configurations are often represented in an orbital diagram, which explicitly shows the number and spin of electrons in various atomic orbitals.

Filling higher orbitals:

The same rules apply for the order of orbital filling as we deduced for n=2. First the 3s orbitals fill (Na & Mg), and then 3p (Al-Ar).

As the angular momentum quantum number, l, increases, the orbitals extend further from the nucleus, and all orbitalsexcept s have nodes at the nucleus. This means that the energy of an orbital increases with l for a given n.

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Filling higher orbitals:

This effect is big enough that the energy of the 4s orbital is lower than 3d. The order of increasing energies and of filling is shown in the diagram at right.

An important consequence of this is in atomic spectroscopy is that of the energy level spacing or differences ΔE are unique to each atom, which means that we can identify atoms by their atomic absorbance or emission spectra.

Multi-Electron Configurations:

Pauli Exclusion Principle. No two electrons in an atom may be in the same quantum state {n, l, m, s}

Aufbau Principle. Electrons adopt the lowest possible energy configuration.

Penetration. Orbitals of equal n nearest the nucleus have lowest energy: s< p< d< f…

Hund’s Rule. Maximiseunpaired electron spins in degenerate orbitals.

Orbitals table: Worked example: Multi-Electron Configurations:

What are the electron configurations of atomic Ca and Ge?

Ca has 20 electrons, which we fill as follows 1s22s22p63s23p64s2 or [Ar]4s2

Ge has 32 electrons, which we fill as follows 1s22s22p63s23p64s23d104p2 or [Ar]4s23d104p2

Multi-Electron Configurations:

Orbitals with the same principal quantum number but different azimuthal quantum numbers have different energies in multi-electron atoms.

?What gives

Orbitals and Electron Shells:

Periodic trends are related to electron configurations. The classical model of the atom included the concept of electron “shells” derived from the row lengths in the periodic table.

Noble gases are unreactive because they contain filled electron shells. This emerges from quantum theory as a natural consequence of the allowed orbital structure.

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Structure of the Periodic Table:

Atoms with the same outer shell configuration are expected to have similar chemical properties. Outer shell or valenceelectrons are important in the formation of chemical bonds (as we shall see later). They will lie in the same group in the periodic table, and form compounds with the same stoichiometry.

Structure of the Periodic Table:

The periodic table can be regarded in terms of electron configurations, denoted by orbital angular momentum quantum number. The periodic table may thus be divided into s, p, d, and f blocks according to which orbital is being filled.

THANK YOU FOR YOUR ATTENDANCE

@ 2008


PERIODIC ATOMIC PROPERTIES AND QUANTUM THEORY

Dr. Akhmad Syoufian


1. Atomic Radius:

The atomic radius is determined by the electronic configuration, and particularly by how far the electron density extends from the nucleus. The wave-functions and potential energy help make sense of the observed trends.

1. Atomic Radius (continued):

Radius increases down a group as electrons add to new “shells.”

Across a row the radius decreases as the nuclear charge increases.

From group 8 (noble gas) to the group 1 (alkali metal). The one additional electron goes into the next s-orbital, increasing the radius markedly.

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1. Atomic Radius (continued): 2. Ionization Energy:

Ionization energy is related to the energy required to remove an electron from an atom.

Clearly a multi-electron atom would have many ionization energies. So by definition the first ionization energy is defined as the energy required to remove the outer most electron from a neutral atom in the gas phase.

Quantum theory also helps make sense of ionization energy trends.

2. Ionization Energy (continued):

Stepping down a group, the outer electrons of each element is another

shell further away from the nucleus. Inner electrons screen the

nuclear attraction that binds the electron, so ionization becomes easier.

Across a row, electrons are added to the same shell. The increase in nuclear charge without additional screening holds the electrons more tightly to the nucleus.

2. Ionization Energy (continued):

The figure below shows the 1st ionization energies.

Take look more closely! At the trend in ionization energies we see two deviations. Boron's ionization energy is lower than beryllium's and oxygen's is lower than nitrogen's.

What gives ?

3. Electron Affinity:

Electron affinity is the energy change associated with the addition of an electron to a neutral atom in the gas phase to form the single charged anion.

Except for Group 2 and Noble Gases which have filled sub-shells, forming the anion is exothermic - i.e. EA is negative.

The halides have the largest electron affinity with the greatest amount of energy released on anion formation.

In general, the size of the energy change (+ or -) decreases as you go down the group.

3. Electron Affinity (continued):

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3. Electron Affinity (continued):

Fluorine (atomic number 9) has a very negative electron affinity, energy is released when is gains an electron. Meanwhile, lithium and beryllium do not want to gain electrons. Beryllium is particularly uninterested.

?What gives

THANK YOU FOR YOUR ATTENDANCE

@ 2008


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