Examples of electron in the following topics:
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- Therefore, these electrons are not as strongly bound as electrons closer to the nucleus.
- Ne has 10 electrons.
- Thus the number of nonvalence electrons is 2 (10 total electrons - 8 valence).
- Flourine has 9 electrons but F- has gained an electron and thus has 10.
- Sodium has 11 electrons but the Na+ ion has lost an electron and thus has 10.
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- Such charges are produced by removing (or adding) electrons from (or to) an object.
- Electron deficient species, which may or may not be positively charged, are attracted to electron rich species, which may or may not be negatively charged.
- Electrophiles: Electron deficient atoms, molecules or ions that seek electron rich reaction partners.
- Nucleophiles: Electron rich atoms, molecules or ions that seek electron deficient reaction partners.
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- An atom's electrons exist in discrete atomic orbitals, and the atom's electron configuration can be determined using a set of guidelines.
- This nucleus is surrounded by electrons.
- An atom's electron shell can accommodate 2n2 electrons, where n is the energy level.
- An element's electron configuration is the arrangement of the electrons in the shells.
- Electrons that occur together in an orbital are called an electron pair.
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- The Aufbau principle determines an atom's electron configuration by adding electrons to atomic orbitals following a defined set of rules.
- As electrons are added, they assume the most stable shells with respect to the nucleus and the electrons already present.
- When there are two electrons in an orbital, the electrons are called an electron pair.
- If the orbital only has one electron, this electron is called an unpaired electron.
- Put one electron in each of the three p orbitals in the second energy level (the 2p orbitals) and then if there are still electrons remaining, go back and place a second electron in each of the 2p orbitals to complete the electron pairs.
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- The electron affinity (Eea) of a neutral atom or molecule is defined as the amount of energy released when an electron is added to it to form a negative ion, as demonstrated by the following equation:
- Mulliken used a list of electron affinities to develop an electronegativity scale for atoms by finding the average of the electron affinity and ionization potential.
- A molecule or atom that has a more positive electron affinity value is often called an electron acceptor; one with a less positive electron affinity is called an electron donor.
- To use electron affinities properly, it is essential to keep track of the sign.
- Electron affinity follows the trend of electronegativity: fluorine (F) has a higher electron affinity than oxygen (O), and so on.
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- The total number of electrons represented in a Lewis structure is equal to the sum of the numbers of valence electrons in each individual atom.
- Non-valence electrons are not represented in Lewis structures.
- After the total number of available electrons has been determined, electrons must be placed into the structure.
- When counting electrons, negative ions should have extra electrons placed in their Lewis structures; positive ions should have fewer electrons than an uncharged molecule.
- The hypochlorite ion, ClO−, contains 13 + 1 = 14 electrons.
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- Diamagnetic atoms have only paired electrons, whereas paramagnetic atoms, which can be made magnetic, have at least one unpaired electron.
- In other words, one of the electrons has to be "spin-up," with $m_s = +\frac{1}{2}$, while the other electron is "spin-down," with $m_s = -\frac{1}{2}$.
- Whenever two electrons are paired together in an orbital, or their total spin is 0, they are called diamagnetic electrons.
- Electrons that are alone in an orbital are called paramagnetic electrons.
- Remember that if an electron is alone in an orbital, the orbital has a net spin, because the spin of the lone electron does not get canceled out.
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- Electrons tend to minimize repulsion by occupying their own orbital, rather than sharing an orbital with another electron.
- To avoid confusion, scientists always draw the first electron, and any other unpaired electron, in an orbital as "spin-up."
- For example, take the electron configuration for carbon: 2 electrons will pair up in the 1s orbital, 2 electrons pair up in the 2s orbital, and the remaining 2 electrons will be placed into the 2p orbitals.
- As another example, oxygen has 8 electrons.
- Therefore, two p orbitals will each get 1 electron and one will get 2 electrons.
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- Eight electrons fill the valence level for all noble gases, except helium, which has two electrons in its full valence level.
- This arrangement of shared electrons between O and H results in the oxygen atom having an octet of electrons, and each H atom having two valence electrons.
- These are 'lone pairs' of electrons.
- Each O atom starts out with six (red) electrons and C with four (black) electrons, and each bond behind an O atom and the C atom consists of two electrons from the O and two of the four electrons from the C.
- Two H atoms, each contributing an electron, share a pair of electrons.
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- Lewis symbols (also known as Lewis dot diagrams or electron dot diagrams) are diagrams that represent the valence electrons of an atom.
- Lewis structures (also known as Lewis dot structures or electron dot structures) are diagrams that represent the valence electrons of atoms within a molecule.
- For example, the Lewis symbol of carbon depicts a "C' surrounded by 4 valence electrons because carbon has an electron configuration of 1s22s22p2.
- Notice that the outermost level has only one electron.
- Each of the four valence electrons is represented as a dot.