Atomic structure
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Atomic structure Atomic structure Document Transcript

  • ATOMIC STRUCTURE Summary: Atomic Structure Atoms consist of protons and neutrons in the nucleus, surrounded by electrons that reside in orbitals. Orbitals are classified according to the four quantum numbers that represent any one particular orbital's energy, shape, orientation, and the spin of the occupying electron. The first section of this SparkNote on Atomic Structure will focus on the electron and the mechanism of describing electrons and their orbitals. As we shall see in the second section, electrons fill up orbitals in a systematic fashion, following the rules of the Aufbau principle. The configuration of electrons in an atom play a vital role in chemistry. Virtually every chemical process relies on the interactions of electrons between atoms, most particularly on the tendency of atoms to follow the octet rule, the tendency to gain a full valence shell electrons. In the second section of this SparkNote, we will discuss the properties of electrons, distinguishing between valence electrons and inner electrons, then broadening the discussion into an examination of the properties of electron conifgurations. Unsurprisingly, given the importance electron configurations play in determining the chemical and physical characteristics of an atom, atoms with similar electron configurations also display similar characteristics. In other words, much of the periodicity of the Periodic Table arises from electron configuration. To see the periodic table, click here. Once the window appears, roll your mouse over the elements to see their specific information. You can also access the periodic table by going into the SparkNotes reference section that resides at the top of every SparkNotes page. We discussed a number of periodic trends in the SparkNote on the Periodic Table table. In the third section we will quickly discuss those previous trends again, and then move to a description of the periodic trends relating to atomic size, ionization energy, electron affinity, and electronegativity. Terms Anion - An ion with a net negative charge. Atomic orbital - An orbital, associated with only one particular atom, in which electrons reside. Though they are called orbitals, atomic orbitals should not be conceived as akin to the orbits of planets rather around a star. Instead, orbitals describe a locus of space in which an electron is likely to reside. Each orbital can hold up to two electrons. Aufbau principle - German for "building up", a systematic procedure for determining the electron configuration of any atom. Incorporates the Pauli Exclusion Principle and Hund's Rule. Cation - An ion with a net positive charge. degenerate orbitals - Orbitals with identical energies. Electron - A negatively charged elementary particle of mass 9.109390x10- 31 . Electrons of an unbonded atom move around the atomic nucleus in orbitals. Those electrons in the orbitals furthest from the nucleus are the highest in energy, play a crucial role in chemical processes such as bonding, and are called valence electrons. Electron affinity - The energy change in an atom when it gains an electron. Electronegativity - A measure of the ability of an atom to attract electrons to itself. Incorporates the atom's ionization energy and electron affinity.
  • Hund's Rule - A rule which says that, when choosing between orbitals, electrons prefer to go in separate orbitals of the same energy. In this way, every orbital within a particular shell (or subshell when the orbitals are not degenerate) will be ha lf-filled before any single one orbital becomes completely filled. Ion - Any atom or molecule with a net charge. Ionization energy - The energy it takes to remove an electron from an atom. Isoelectronic - Description for two elemental species with the same electronic configuration. Isotope - Atoms with the same number of protons (i.e. same atomic number) but a different number of neutrons. Neutron - An uncharged atomic particle of mass 1.67493x10- 27 . It resides in the nucleus. Nucleus - The small, dense central region of an atom around which electrons orbit. The nucleus is made up of protons and neutrons. Octet rule - The cardinal rule of bonding. The octet rule states that atoms gain stability when they have a full complement of 8 electrons in their valence shells. Pauli Exclusion Principle - States that no two electrons in an atom or molecule can have the same set of four quantum numbers. Proton - A positively charged particle of mass 1.6726x10- 27 . Protons reside in the nucleus. Quantum Numbers - The four numbers that define each particular electron of an atom. The Principle Quantum Number (n) describes the electrons' energy and distance from the nucleus. The Angular Momentum Quantum Number (l) describes the shape of the orbital in which the electron resides. The Magnetic Quantum Number (m describes the orientation of the orbital in space. The Spin Quantum number describes whether the spin of the electron is positive or negative. Shell - A group of subshells of similar energy levels. For example, 2s and 2p subshells occupy the same shell. Indicated by the principle quantum number. Shielding - When the attraction from the nucleus felt by one electron is lessened or blocked by intermediate electrons. Shielding can split degenerate orbitals. For example, since s-orbital electrons shield for p-orbital electrons and rece ive little shielding themselves, s-orbitals are usually of lower energy level than p-orbitals of the same shell. Splitting - Through shielding, the breaking of degenerate orbitals within a shell in multi-electron atoms. Subshell - Orbitals of the same subshell are of the same shape and energy. p-orbitals are of the same subshell, while s-orbitals are of a separate subshell. Indicated by the angular momentum quantum number. Uncertainty Principle - A tenet of quantum mechanics that says that the position and momentum of any particle cannot both be known precisely at the same time. Valence electrons - The electrons in the outermost energy shell of an atom. The configuration of these electrons determine the chemical properties of the element.
  • Valence shell - The highest energy shell in an atom, containing valence electrons. All interactions between atoms take place through the electrons of the valence shell. Atoms and Atomic Orbitals Fundamentals of the Atom An atom consists of a nucleus of protons and neutrons, surrounded by electrons. Each of the elements in the periodic table is classified according to its atomic number, which is the number of protons in that element's nucleus. Protons have a charge of +1, electrons have a charge of -1, and neutrons have no charge. Neutral atoms have the same number of electrons and protons, but they can have a varying number of neutrons. Within a given element, atoms with different numbers of neutrons are isotopes of that element. Isotopes typically exhibit similar chemical behavior to each other. Electrons have such little mass that they exhibit properties of both particles and waves; in. We further know from Heisenberg's Uncertainty Principle that it is impossible to know the precise location of an electron. Despite this limitation, there are regions around the atom where the electron has a high probability of being found. Such regions are referred to as atomic orbitals. Atomic Orbitals and Quantum Numbers The relation of a particular electron to the nucleus can be described through a series of four numbers, called the Quantum Numbers. The first three of these numbers describe the energy (Principle quantum number), shape (Angular momentum quantum number), and orientation of the orbital (magnetic quantum number). The fourth number represents the "spin" of the electron (spin quantum number). The four quantum numbers are described below. Principle Quantum Number (n) The principle quantum number indicates how the distance of the orbital from the nucleus. Electrons are farther away for higher values of n . Electrons are negatively charged, so electrons that are closer to the positively charged nucleus are more powerfully attracted and tightly bound than those that are farther away. Electrons that are closer to the nucleus are thus more stable, and less likely to be lost by the atom. In other words, as n increases, so does the energy of the electron and the likelihood of that electron being lost by the atom. In a given atom, all the atomic orbitals with the same n are collectively known as a shell. n can take on integer values of 1 or higher (ex. 1, 2, 3, etc.). Angular Momentum Quantum Number (l) The angular momentum quantum number describes the shape of the orbital. The angular momentum number (or subshell) can be represented either by a number (any integer from 0 up to n-1) or by a letter (s,p,d,f,g, and then up the alphabet), with 0 corresponding to s, 1 to p, 2 to d, and so on. For example: when n = 1, l can only equal 0; meaning that shell n = 1 has only an s orbital (l = 0). when n = 3, l can equal 0, 1, or 2; meaning that shell n = 3 has s,p, and d orbitals. s orbitals are spherical, whereas p orbitals are dumbbell-shaped. d orbitals and beyond are much harder to visually represent.
  • Figure %: s and p atomic orbital shapes Magnetic Quantum Number (m) Gives the orientation of the orbital in space; in other words, the value of m describes whether an orbital lies along the x-, y-, or z-axis on a three-dimensional graph, with the nucleus of the atom at the origin. m can take on any value from -l to l. For our purposes, it is only important that this quantum number tells us that for each value of n there may be up to one s -orbital, three p -orbitals, five d -orbitals, and so on. For example: The s orbital (l = 0) has one orbital, since m can only equal 0. That orbital is spherically symmetrical about the nucleus. Figure %: s orbital The p orbital (l = 1) has three orbitals, since m = -1, 0, and 1. These three orbitals lie along the x -, y -, and z -axes.
  • Figure %: p orbitals The d orbital (l = 2) has five orbitals, since m = -2, -1, 0, 1, and 2. It is far more difficult to describe the orientation of d orbitals, as you can see:
  • Figure %: d orbitals Spin Quantum Number (s): The spin quantum number tells whether a given electron is spin up (+1/2) or spin down (-1/2). An orbital contains two electrons, and each of those electrons must have different spins. Orbital Energy Diagrams It is often convenient to depict orbitals in an orbital energy diagram, as seen below in . Such diagrams show the orbitals and their electron occupancies, as well as any orbital interactions that exist. In this case we have the orbitals of the hydrogen atom with electrons omitted. The first electron shell (n = 1) contains just the 1s orbital. The second shell (n = 2) holds a 2s orbital and three 2p orbitals. The third shell (n = 3) holds one 3s orbital, three 3p orbitals, and five 3d orbitals, and so forth. Note that the relative spacing between orbitals becomes smaller for larger n. In fact, as n gets large the spacing becomes infinitesimally small.
  • Figure %: Energy diagram of the unoccupied atomic orbitals of hydrogen. Potential energy is on the y-axis. You will see such energy diagrams quite often in your continuing study of chemistry. Notice that all orbitals with the same n have the same energy. Orbitals with identical energies are said to be degenerate (not in the moral sense!). Electrons in higher-level orbitals have more potential energy and are more reactive, i.e. more likely to undergo chemical reactions. Multi-electron atoms When an atom only contains a single electron, its orbital energies depend only on the principle quantum numbers: a 2s orbital would be degenerate with a 2p orbital. However, this degeneracy is broken when an atom has more than one electron. This is due to the fact that the attractive nuclear force any electron feels is shielded by the other electrons. s-orbitals tend to be closer to the nucleus than p-orbitals and don't get as much shielding, and hence become lower in energy. This process of breaking degeneracies within a shell is known as splitting. In general s orbitals are lowest in energy, followed by p orbitals, d orbitals, and so forth. Figure %: Splitting of orbital energies in multi-electron systems Electron Energy The energy diagram of imply a further fact about the energy of electrons. Note that the energy levels in these diagrams do not follow a continuous line: an atom is either in one energy subshell or it is in another. There is no in between. In this way, the diagram perfectly represents the quantized nature of electrons, meaning that electrons can only exist at specific and defined energy levels. The energy level of an electron in a particular energy shell can be determined according to the following equation:
  • En = /frac-2.178x10- 18joulesn2 where n is the principal quantum number and En is the energy level at that quantum number. When an electron absorbs a specific quanta of energy it can jump to a higher energy level. It can also give off a specific quanta and fall back to a lower energy level. An atom whose electrons are at their lowest energy levels is said to be in the ground state. The discovery of the quantum nature of energy and electrons, first formulated by Max Planck in 1900, led to the creation of an entirely new field, quantum mechanics. Electron Configuration and Valence Electrons Electron Configuration The electrons in an atom fill up its atomic orbitals according to the Aufbau Principle; "Aufbau," in German, means "building up." The Aufbau Principle, which incorporates the Pauli Exclusion Principle and Hund's Rule prescribes a few simple rules to determine the order in which electrons fill atomic orbitals: 1. Electrons always fill orbitals of lower energy first. 1s is filled before 2s, and 2s before 2p. 2. The Pauli Exclusion Principle states no two electrons within a particular atom can have identical quantum numbers. In function, this principle means that if two electrons occupy the same orbital, they must have opposite spin. 3. Hund's Rule states that when an electron joins an atom and has to choose between two or more orbitals of the same energy, the electron will prefer to enter an empty orbital rather than one already occupied. As more electrons are added to the atom, these electrons tend to half-fill orbitals of the same energy before pairing with existing electrons to fill orbitals. Figure %: The ground state electron configuration of carbon, which has a total of six electrons. The configuration is determined by applying the rules of the Aufbau Principle. Valency and Valence Electrons The outermost orbital shell of an atom is called its valence shell, and the electrons in the valence shell are valence electrons. Valence electrons are the highest energy electrons in an atom and are therefore the most reactive. While
  • inner electrons (those not in the valence shell) typically don't participate in chemical bonding and reactions, valence electrons can be gained, lost, or shared to form chemical bonds. For this reason, elements with the same number of valence electrons tend to have similar chemical properties, since they tend to gain, lose, or share valence electrons in the same way. The Periodic Table was designed with this feature in mind. Each element has a number of valence electrons equal to its group number on the Periodic Table. Figure %: The periodicity of valence electrons This table illustrates a number of interesting, and complicating, features of electron configuration. First, as electrons become higher in energy, a shift takes place. Up until now, we have said that as the principle quantum number, increases, so does the energy level of the orbital. And, as we stated above in the Aufbau principle, electrons fill lower energy orbitals before filling higher energy orbitals. However, the diagram above clearly shows that the 4s orbital is filled before the 3d orbital. In other words, once we get to principle quantum number 3, the highest subshells of the lower quantum numbers eclipse in energy the lowest subshells of higher quantum numbers: 3d is of higher energy than 4s. Second, the above indicates a method of describing an element according to its electron configuration. As you move from left to right across the periodic table, the above diagram shows the order in which orbitals are filled. If we were the actually break down the above diagram into groups rather than the blocks we have, it would show how exactly how many electrons each element has. For example, the element of hydrogen, located in the uppermost left-hand corner of the periodic table, is described as 1s1 , with the s describing which orbital contains electrons and the 1 describing how many electrons reside in that orbital. Lithium, which resides on the periodic table just below hydrogen, would be described as 1s2 2s1 . The electron configurations of the first ten elements are shown below (note that the valence electrons are the electron in highest energy shell, not just the electrons in the highest energy subshell).
  • The Octet Rule Our discussion of valence electron configurations leads us to one of the cardinal tenets of chemical bonding, the octet rule. The octet rule states that atoms become especially stable when their valence shells gain a full complement of valence electrons. For example, in above, Helium (He) and Neon (Ne) have outer valence shells that are completely filled, so neither has a tendency to gain or lose electrons. Therefore, Helium and Neon, two of the so-called Noble gases, exist in free atomic form and do not usually form chemical bonds with other atoms. Most elements, however, do not have a full outer shell and are too unstable to exist as free atoms. Instead they seek to fill their outer electron shells by forming chemical bonds with other atoms and thereby attain Noble Gas configuration. An element will tend to take the shortest path to achieving Noble Gas configuration, whether that means gaining or losing one electron. For example, sodium (Na), which has a single electron in its outer 3s orbital, can lose that electron to attain the electron configuration of neon. Chlorine, with seven valence electrons, can gain one electron to attain the configuration of argon. When two different elements have the same electron configuration, they are called isoelectronic. Diamagnetism and Paramagnetism The electron configuration of an atom also has consequences on its behavior in relation to magnetic fields. Such behavior is dependent on the number of electrons an atom has that are spin paired. Remember that Hund's Rule and the Pauli Exclusion Principle combine to dictate that an atom's orbitals will all half-fill before beginning to completely fill, and that when they completely fill with two electrons, those two electrons will have opposite spins. An atom with all of its orbitals filled, and therefore all of its electrons paired with an electron of opposite spin, will be very little affected by magnetic fields. Such atoms are called diagmetic. Conversely, paramagnetic atoms do not have all of their electrons spin-paired and are affected by magnetic fields. There are degrees of paramagnetism, since an atom might have one unpaired electron, or it might have four. Periodic Trends In the SparkNote on the Periodic table we discussed a number of simple periodic trends. In this section we will discuss a number of more complex trends, the understanding of which relies on knowledge of atomic structure. Before getting into these trends, we should engage a quick review and establish some terminology. As seen in the previous section on the octet rule, atoms tend to lose or gain electrons in order to attain a full valence shell and the stability a full valence shell imparts. Because electrons are negatively charged, an atom becomes positively or negatively charged as it loses or gains an electron, respectively. Any atom or group of atoms with a net charge
  • (whether positive or negative) is called an ion. A positively charged ion is a cation while a negatively charged ion is an anion. Now we are ready to discuss the periodic trends of atomic size, ionization energy, electron affinity, and electronnegativity. Atomic Size (Atomic Radius) The atomic size of an atom, also called the atomic radius, refers to the distance between an atom's nucleus and its valence electrons. Remember, the closer an electron is to the nucleus, the lower its energy and the more tightly it is held. Moving Across a Period Moving from left to right across a period, the atomic radius decreases. The nucleus of the atom gains protons moving from left to right, increasing the positive charge of the nucleus and increasing the attractive force of the nucleus upon the electrons. True, electrons are also added as the elements move from left to right across a period, but these electrons reside in the same energy shell and do not offer increased shielding. Moving Down a Group The atomic radius increases moving down a group. Once again protons are added moving down a group, but so are new energy shells of electrons. The new energy shells provide shielding, allowing the valence electrons to experience only a minimal amount of the protons' positive charge. Cations and Anions Cations and anions do not actually represent a periodic trend in terms of atomic radius, but they do affect atomic radius, and so we will discuss them here. A cation is positively charged, meaning that it is an atom that has lost an electron or electrons. The positive charge of the nucleus is thus distributed over a smaller number of electrons and electron-electron repulsion is decreased, meaning that the electrons are held more tightly and the atomic radius is smaller than in the normal neutral atom. Anions, conversely, are negatively charged ions: atoms that have gained electrons. In anions, electron-electron repulsion increases and the positive charge of the nucleus is distributed over a large number of electrons. Anions have a greater atomic radius than the neutral atom from which they derive.
  • Ionization Energy and Electron Affinity The process of gaining or losing an electron requires energy. There are two common ways to measure this energy change: ionization energy and electron affinity. Ionization Energy The ionization energy is the energy it takes to fully remove an electron from the atom. When several electrons are removed from an atom, the energy that it takes to remove the first electron is called the first ionization energy, the energy it takes to remove the second electron is the second ionization energy, and so on. In general, the second ionization energy is greater than first ionization energy. This is because the first electron removed feels the effect of shielding by the second electron and is therefore less strongly attracted to the nucleus. If a particular ionization energy follows a previous electron loss that emptied a subshell, the next ionization energy will take a rather large leap, rather than follow its normal gently increasing trend. This fact helps to show that just as electrons are more stable when they have a full valence shell, they are also relatively more stable when they at least have a full subshell. Ionization Energy Across a Period Ionization energy predictably increases moving across the periodic table from left to right. Just as we described in the case of atomic size, moving from left to right, the number of protons increases. The electrons also increase in number, but without adding new shells or shielding. From left to right, the electrons therefore become more tightly held meaning it takes more energy to pry them loose. This fact gives a physical basis to the octet rule, which states that elements with few valence electrons (those on the left of the periodic table) readily give those electrons up in order to attain a full octet within their inner shells, while those with many valence electrons tend to gain electrons. The electrons on the left tend to lose electrons since their ionization energy is so low (it takes such little energy to remove an electron) while those on the right tend to gain electrons since their nucleus has a powerful positive force and their ionization energy is high. Note that ionization energy does show a sensitivity to the filling of subshells; in moving from group 12 to group 13 for example, after the d shell has been filled, ionization energy actually drops. In general, though, the trend is of increasing ionziation energy from left to right. Ionization Energy Down a Group Ionization energy decreases moving down a group for the same reason atomic size increases: electrons add new shells creating extra shielding that supersedes the addition of protons. The atomic radius increases, as does the energy of the valence electrons. This means it takes less energy to remove an electron, which is what ionization energy measures. Electron Affinity An atom's electron affinity is the energy change in an atom when that atom gains an electron. The sign of the electron affinity can be confusing. When an atom gains an electron and becomes more stable, its potential energy decreases: upon gaining an electron the atom gives off energy and the electron affinity is negative. When an atom becomes less stable upon gaining an electron, its potential energy increases, which implies that the atom gains energy as it acquires the electron. In such a case, the atom's electron affinity is positive. An atom with a negative electron affinity is far more likely to gain electrons. Electron Affinities Across a Period Electron affinities becoming increasingly negative from left to right. Just as in ionization energy, this trend conforms to and helps explain the octet rule. The octet rule states that atoms with close to full valence shells will tend to gain
  • electrons. Such atoms are located on the right of the periodic table and have very negative electron affinities, meaning they give off a great deal of energy upon gaining an electron and become more stable. Be careful, though: the nobel gases, located in the extreme right hand column of the periodic table do not conform to this trend. Noble gases have full valence shells, are very stable, and do not want to add more electrons: noble gas electron affinities are positive. Similarly, atoms with full subshells also have more positive electron affinities (are less attractive of electrons) than the elements around them. Electron Affinities Down a Group Electron affinities change little moving down a group, though they do generally become slightly more positive (less attractive toward electrons). The biggest exception to this rule are the third period elements, which often have more negative electron affinities than the corresponding elements in the second period. For this reason, Chlorine, Cl, (group VIIa and period 3) has the most negative electron affinity. Electronegativity Electronegativity refers to the ability of an atom to attract the electrons of another atom to it when those two atoms are associated through a bond. Electronegativity is based on an atom's ionization energy and electron affinity. For that reason, electronegativity follows similar trends as its two constituent measures. Electronegativity generally increases moving across a period and decreases moving down a group. Flourine (F), in group VIIa and period 2, is the most powerfully electronegative of the elements. Electronegativity plays a very large role in the processes of Chemical Bonding.