Effective Nuclear Charge

6 Key excerpts on "Effective Nuclear Charge"

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  eBook - ePub

  BIOS Instant Notes in Inorganic Chemistry

  • Tony Cox(Author)
  • 2004(Publication Date)
  • Taylor & Francis

    (Publisher)

  A general increase of radius and decrease in IE down most groups is dominated by the increasing principal quantum number of outer orbitals. Effective Nuclear Charge also increases, and can give rise to irregularities in the IE trends.

    States of ionization IEs for positive ions always increase with the charge. Electron affinities are the IEs of negative ions and are always less than IEs for neutral atoms.   Relativistic effects Deviations from the nonrelativistic predictions become significant for heavy atoms, and contribute to especially high IEs for later elements in the sixth period.  

  Related topics

  Many-electron atoms (A3 )

  The periodic table (A4 )

  Chemical periodicity (B2 )

  Energies and sizes

  The first ionization energy (IE) of an atom (M) is the energy required to form the positive ion M:

  M → M+ + e–

  The IE value reflects the energy of the orbital from which the electron is removed, and so depends on the principal quantum number (n) and Effective Nuclear Charge (Zeff ; see Topic A3 ):

  The average radius of an orbital depends on the same factors (see Topic A2 ):

  Smaller orbitals generally have more tightly bound electrons with higher ionization energies.

  It is sometimes useful to assume that the distance between two neighboring atoms in a molecule or solid can be expressed as the sum of atomic or ionic radii. Metallic, covalent or ionic radii can be defined according to the type of bonding between atoms, and van der Waals’ radii for atoms in contact but not bonded. Such empirically derived radii are all different and are not easily related to any simple predictions based on isolated atoms. They are, however, qualitatively related to orbital radii and all follow the general trends discussed below (see, e.g. Topic D4 , Table 1 , for ionic radii).

  Horizontal trends

  Increasing nuclear charge is accompanied by correspondingly more electrons in neutral atoms. Moving from left to right in the periodic table, the increase of nuclear charge has an effect that generally outweighs the screening from additional electrons. Increasing Zeff leads to an increase of IE across each period, which is the most important single trend in the periodic table (see Topic B2

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  eBook - ePub

  Essentials of Nuclear Medicine Physics and Instrumentation

  • Rachel A. Powsner, Matthew R. Palmer, Edward R. Powsner(Authors)
  • 2013(Publication Date)
  • Wiley-Blackwell

    (Publisher)

  The electrical charges of the atom are balanced, that is, the total negative charge of the electrons equals the total positive charge of the nucleus. This is simply another way of pointing out that the number of orbital electrons equals the number of nuclear protons. Furthermore, the electrons must fill the shells with the highest binding energy first. At least in the elements of low-atomic-number electrons, the inner shells have the highest binding energy.

  If the arrangement of the electrons in the shells is not in the stable state, they will undergo rearrangement in order to become stable, a process often referred to as de-excitation . Because the stable configuration of the shells always has less energy than any unstable configuration, the de-excitation releases energy as photons, often as X-rays .

  Nucleus

  Like the atom itself, the atomic nucleus also has an inner structure (Figure 1.8 ). Experiments have shown that the nucleus consists of two types of particles: protons , which carry a positive charge, and neutrons , which carry no charge. The general term for protons and neutrons is nucleons . The nucleons, as shown in the first two rows of Table 1.1 , have a much greater mass than electrons. Like electrons, nucleons have quantum properties, including spin. The nucleus has a spin value equal to the sum of the nucleon spin values.

  Figure 1.8 The nucleus of an atom is composed of protons and neutrons.

  Table 1.1 The subatomic particles

  A simple but useful model of the nucleus is a tightly bound cluster of protons and neutrons. Protons naturally repel each other, since they are positively charged; however, there is a powerful binding force called the nuclear force that holds the nucleons together very tightly (Figure 1.9 ). The work (energy) required to overcome the nuclear force, the work to remove a nucleon from the nucleus, is called the nuclear binding energy . Typical binding energies are in the range of 6 million to 9 million electron volts (MeV) (approximately one thousand to one million times the electron binding force). The magnitude of the binding energy is related to another fact of nature: the measured mass of a nucleus is always less than the mass expected from the sum of the masses of its neutrons and protons. The “missing” mass is called the mass defect , the energy equivalent of which is equal to the nuclear binding energy. This interchangeability of mass and energy was immortalized in Einstein’s equation E  = mc 2

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  Introductory Electrical Engineering With Math Explained in Accessible Language

  • Magno Urbano(Author)
  • 2019(Publication Date)
  • Wiley

    (Publisher)

  −19 C.

  In the early days, science pictured electrons spinning around atomic nucleus like planets spinning around the sun. Today we know that this image is false. Electrons oscillate frenetically around the atomic nucleus, and their position cannot be determined,2 like they were in all positions at the same time, a kind of “energetic” cloud.

  The idea of an electron orbiting an atomic nucleus like a planet around a sun is wrong because it makes us think about electrons as a solid and discrete element, like a particle. In real life, electrons behave both as particles and waves.3

  Physicists knows that electrons occupy orbits around the nucleus at a certain distance, like they were shells. Each orbit is at a discrete level of energy and can only be occupied by electrons having that particular level of energy. If an electron gains energy, it may pass to a superior orbit. If it loses energy, it will decay to a lower orbit and release the extra energy by emitting a photon.4

  Figure 3.2

  Atom representation.

  A representation of an atom is shown in Figure 3.2 . The spheres at the center represent the nucleus, nothing more than an agglomeration of protons and neutrons. The circles represent the energetic orbits. Electrons are shown in the diagram as “e” inside the orbits but in real life they can be at any point and in all points at the same time inside a particular orbit.

  3.4 Strong Force and Weak Force

  Atoms and their protons and neutrons are kept together by powerful attraction forces. The laws of physics say that elements with the same electric charge repel each other and that elements with opposite electric charge attract each other.

  Atomic nuclei may be formed by several protons and neutrons. These protons could never be kept together with other protons, at the required close proximity, to form these atomic nuclei, simply because they would repel each other. A bigger force must exist to overpower the electric repulsion force. This force is called the strong nuclear force or strong interaction.

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  eBook - ePub

  Radiation Detection

  Concepts, Methods, and Devices

  • Douglas McGregor, J. Kenneth Shultis(Authors)
  • 2020(Publication Date)
  • CRC Press

    (Publisher)

  The number of electrons in an atom equals its atomic number Z and determines its position in the Periodic Table. The chemical properties are determined by the number and arrangement of the electrons. Each element in the table is formed by adding one electron to that of the preceding element in the Periodic Table in such a way that the electron is most tightly bound to the atom. The arrangement of the electrons for the elements with electrons in only the first four shells is shown in Table 3.3. Table 3.3. Electron shell arrangement for the lightest elements. 3.4.7 Success of Quantum Mechanics Quantum mechanics has been an extremely powerful tool for describing the energy levels and the distributions of atomic electrons around a nucleus. Each energy level and configuration is uniquely defined by four quantum numbers: n the principal quantum number, ℓ the orbital angular momentum quantum number, m ℓ the z -component of the angular momentum, and m s = ±1/2, the electron spin number. These numbers arise naturally from the analytical solution of the wave equation (as modified by Dirac to include special relativity effects) and thus avoid the ad hoc introduction of orbital quantum numbers required in earlier atomic models. Inside the nucleus, quantum mechanics is also thought to govern. However, the nuclear forces holding the neutrons and protons together are much more complicated than the electromagnetic forces binding electrons to the nucleus. Consequently, much work continues in the application of quantum mechanics (and its more general successor quantum electrodynamics) to predicting energy and configuration states of nucleons. Nonetheless, the fact that electronic energy levels of an atom and nuclear excited states are discrete with very specific configurations is a key concept in modern physics

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  Foundations for Teaching Chemistry

  Chemical Knowledge for Teaching

  • Keith S. Taber(Author)
  • 2019(Publication Date)
  • Routledge

    (Publisher)

  Even here, the basic explanation can involve an explanatory sleight-of-hand. A simple explanation might be that as one moves across a period the nuclear charge becomes greater, so there is a greater attraction from the nucleus acting on the outer-shell electrons, pulling them towards the nucleus. As a result, the electrons are held closer to the nucleus as one moves across the period.

  One issue with the language here is that forces are interactions between pairs of bodies and act on both bodies: the nucleus is attracted to the electrons as much as vice versa. That is a complication we may not feel is important here in providing an accessible account. It is worth bearing in mind, however, as the way we teach can have repercussions later (perhaps when teaching about bonding, see Chapter 11 ). For example, when atomic structure is first introduced, students are often told that that despite repelling each other, negatively charged electrons are held in an atom because of the attraction from the positively charged protons in the nucleus of the atom, because ‘opposite charges attract and similar charges repel’. (It is customary to refer to ‘similar’ charges, but we should be aware that in everyday language ‘similar’ means somewhat alike, and here we mean of the same type.) This is reasonable: but often there is no attempt to explain how the ‘similar’ positive charges are located so closely together in the nucleus given they also repel each other.

  Sometimes if teachers do not provide explanations, students invent their own. One reason given by students for the stability of the nucleus is that the protons are forced together because they are pushed into the nucleus due to a force from the atomic electrons (Taber, 2000a). This alternative conception reflects a kind of symmetry that some students find appealing: the protons pull the electrons, so the electrons push the protons. This explanation is not only counter to Newton’s third law (that forces are interactions that act on the bodies in opposite directions, so either they are both attractions – towards the other body – or both repulsions – away from the other body), but also ignores the inverse square law (Coulomb’s law) that the magnitude of the force falls off with distance. The repulsion between protons packed together in a nucleus must be much greater than any force due to the relatively

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  eBook - ePub

  AP® Chemistry All Access Book + Online + Mobile

  • Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
  • 2014(Publication Date)
  • Research & Education Association

    (Publisher)

  The atomic radii of atoms increase when moving from right to left across the periodic table. The atomic radius also increases as you move down the periodic table. Cations have smaller radii than their corresponding neutral atoms. This is due to a greater Effective Nuclear Charge (greater number of protons in the nucleus compared to electrons) and thus the positively charged nucleus pulls the valence electrons closer. Anions have larger radii than their corresponding neutral atoms. This can be attributed to the addition of valence electrons and their repulsion, thus increasing the atomic radius.

  Figure 5.5.

  TEST TIP For the AP Chemistry exam, be able to describe why the trends in the atomic radius exist.

  Ionization Energy

  The ionization energy is the energy required to remove an electron from an atom, resulting in the formation of a cation. The ionization energy decreases when moving from right to left across the periodic table. It also decreases when moving down a group within the periodic table. It should be noted that more than 1 electron may be removed to form ions of greater charge, but the energy required to remove successive electrons will increase exponentially. Cations that have electron configurations containing filled octets are extremely stable and require an enormous amount of energy to remove an electron.

  Figure 5.6.

  Electron Affinity

  Electron affinity is the ability of an atom to gain electrons in order to form anions. Nonmetal atoms have a much higher electron affinity, because metals will not form anions. The electron affinity increases moving from left to right across the periodic table (see Figure 5.7 ). It also increases when moving up a group within the periodic table.

  Figure 5.7.

  TEST TIP An explanation of most of these trends on the periodic table can be linked to the concept of the atomic radius.

  Electronegativity

  Electronegativity is the ability of an atom to attract electron density to it when forming a covalent bond. The electronegativity of an atom exhibits the same trend as electron affinity in the periodic table. You should know the values from Pauling’s scale of electronegativity for all nonmetallic atoms. This will enable you to determine if the covalent bond is polar or if a molecule is polar. The chart below shows the electronegativity values for the most common nonmetals.

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