Sodium Electrons

Posted on  by admin

Sodium fluoride NaF or FNa CID 5235 - structure, chemical names, physical and chemical properties, classification, patents, literature, biological activities. Atomic number of sodium is 11. Atomic number signifies, number of electrons or protons present in the system. Since, atomic number of sodium is 11, the electrons present in the system is 11. But, the charge on sodium is +! This suggest that, one electron is lost by sodium. Hence, number of electrons present in sodium ion is 10. Sodium(1+) is a monoatomic monocation obtained from sodium.It has a role as a human metabolite and a cofactor. It is an alkali metal cation, an elemental sodium, a monovalent inorganic cation and a monoatomic monocation.

Sodium electron configuration

Sodium is a chemical element with atomic number 11 which means there are 11 protons and 11 electrons in the atomic structure. The chemical symbol for Sodium is Na. The atom consist of a small but massive nucleus surrounded by a cloud of rapidly moving electrons. The nucleus is composed of protons and neutrons. If sodium loses an electron, it now has 11 protons, 11 neutrons, and only 10 electrons, leaving it with an overall charge of +1. It is now referred to as a sodium ion. Therefore, it tends to gain an electron to create an ion with 17 protons, 17 neutrons, and 18 electrons, giving it a net negative (–1) charge.

How many valence electrons does sodium have?

3 Answers

Matt N. · mrpauller.weebly.com

Sodium Electrons Per Shell

Sodium, like all the group 1 alkali metals, has one valence electron.

Explanation:

Valence electrons are the outermost electrons, and are the ones involved in bonding. Sodium has 11 electrons: its atomic number is 11, so it has 11 protons; atoms are neutral, so this means sodium also has 11 electrons.

Electrons are arranged in 'shells' or energy levels. Depending on your level of Chemistry, it is probably easier to think of them as particles orbiting the nucleus. The first 'shell' can have 2 electrons. The second 'shell' can have up to 8 electrons. The third 'shell' is a bit more complicated but let's just say that it takes up to 8 electrons as well (for now...).

So, sodium's 11 electrons are arranged this way: 2 electrons in the first 'shell', 8 electrons in the second 'shell'; and 1 electron (the valence electron) in the third 'shell'. We write this as 2.8.1. The last number is how we know the number of valence electrons.

Protons

Aluminium has the electron arrangement 2.8.3. It has 3 valence electrons. Fluorine has the electron arrangement 2.7. It has 7 valence electrons. This trend only gets broken at Sc (#21).

Here is a video which gives further explanation of this topic.
video from: Noel Pauller

Ari · Media Owl

Sodium has one valence electron. Valence electrons are electrons found in the outermost shell of an atom. The shell number representing the valence shell will differ depending on the atom in question. For sodium, which is in the 3rd row of the periodic table, the valence electrons will be found in the 3rd shell. For fluorine, which is in the second row, the valence electrons will be found in the second shell. (Quick note: In the periodic table, rows are horizontal lines, rows are vertical lines.)

There are a number of ways to approach this question. The easiest way is to use the periodic table as your guide. If we start by looking at the second row, and follow along horizontally, we can assign valence electron numbers to each element. Starting with lithium (Li), which has one valence electron, we can move across horizontally and add one valence electron each time as follows,

Element : Number of Valence Electrons
Li : 1
Be : 2
B : 3
C : 4
N : 5
O : 6
F : 7
Ne : 8 (or zero)

Furthermore, every atom on the periodic table below lithium (Li) will also have one valence electron (i.e. Na, K, Rb, Cs, Fr). The same goes for every other element listed, as well. For instance, Cl, Br, I and At will each have seven valence electrons like fluorine (F). Don't worry too much about elements found in between the Be column and the B column. For the so-called transition metals, valence electrons are not as well defined.

Explanation:

Valence electrons are found in the highest energy s and p orbitals in the main group of elements. The electron configuration for a neutral sodium is #1's'^2'2s'^2'2p'^6'3s'^1'#. The highest energy 3s orbital contains one electron, therefore, neutral sodium atoms have one valence electron.

Actually, all of the elements in Group1/IA have one valence electron.

Related questions

A solvated electron is a freeelectron in (solvated in) a solution, and is the smallest possible anion. Solvated electrons occur widely, although it is difficult to observe them directly because their lifetimes are so short.[1] The deep color of solutions of alkali metals in liquid ammonia arises from the presence of solvated electrons: blue when dilute and copper-colored when more concentrated (> 3 molar).[2] Classically, discussions of solvated electrons focus on their solutions in ammonia, which are stable for days, but solvated electrons also occur in water and other solvents – in fact, in any solvent that mediates outer-sphere electron transfer. The real hydration energy of the solvated electron can be estimated by using the hydration energy of a proton in water combined with kinetic data from pulse radiolysis experiments. The solvated electron forms an acid–base pair with atomic hydrogen.

The solvated electron is responsible for a great deal of radiation chemistry.

Alkali metals dissolve in liquid ammonia giving deep blue solutions, which conduct electricity. The blue colour of the solution is due to ammoniated electrons, which absorb energy in the visible region of light. Alkali metals also dissolve in some small primary amines, such as methylamine and ethylamine[3] and hexamethylphosphoramide, forming blue solutions. Solvated electron solutions of the alkaline earth metals magnesium, calcium, strontium and barium in ethylenediamine have been used to intercalate graphite with these metals.[4]

Sodium

History[edit]

The observation of the color of metal-electride solutions is generally attributed to Humphry Davy. In 1807–1809, he examined the addition of grains of potassium to gaseous ammonia (liquefaction of ammonia was invented in 1823). James Ballantyne Hannay and J. Hogarth repeated the experiments with sodium in 1879–1880. W. Weyl in 1844 and C. A. Seely in 1871 used liquid ammonia while Hamilton Cady in 1897 related the ionizing properties of ammonia to that of water. Charles A. Kraus measured the electrical conductance of metal ammonia solutions and in 1907 attributed it to the electrons liberated from the metal.[5][6] In 1918, G. E. Gibson and W. L. Argo introduced the solvated electron concept.[7] They noted based on absorption spectra that different metals and different solvents (methylamine, ethylamine) produce the same blue color, attributed to a common species, the solvated electron. In the 1970s, solid salts containing electrons as the anion were characterized.[8]

Properties[edit]

Focusing on solutions in ammonia, liquid ammonia will dissolve all of the alkali metals and other electropositive metals such as Ca,[9]Sr, Ba, Eu, and Yb (also Mg using an electrolytic process[10]), giving characteristic blue solutions.

Solutions obtained by dissolution of lithium in liquid ammonia. The solution at the top has a dark blue color and the lower one a golden color. The colors are characteristic of solvated electrons at electronically insulating and metallic concentrations, respectively.

A lithium–ammonia solution at −60 °C is saturated at about 15 mol% metal (MPM). When the concentration is increased in this range electrical conductivity increases from 10−2 to 104ohm−1cm−1 (larger than liquid mercury). At around 8 MPM, a 'transition to the metallic state' (TMS) takes place (also called a 'metal-to-nonmetal transition' (MNMT)). At 4 MPM a liquid-liquid phase separation takes place: the less dense gold-colored phase becomes immiscible from a denser blue phase. Above 8 MPM the solution is bronze/gold-colored. In the same concentration range the overall density decreases by 30%.

Dilute solutions are paramagnetic and at around 0.5 MPM all electrons are paired up and the solution becomes diamagnetic. Several models exist to describe the spin-paired species: as an ion trimer; as an ion-triple—a cluster of two single-electron solvated-electron species in association with a cation; or as a cluster of two solvated electrons and two solvated cations.

Solvated electrons produced by dissolution of reducing metals in ammonia and amines are the anions of salts called electrides. Such salts can be isolated by the addition of macrocyclicligands such as crown ether and cryptands. These ligands bind strongly the cations and prevent their re-reduction by the electron.

In neutral of partially-oxidized metal-ammonia or metal-aqua complexes diffuse solvated electrons are present. These species are recognized as 'Solvated electron precursors' (SEPs). Simply a SEP is a metal complex that bear diffuse electrons in the periphery of the ligands.[11] The diffuse solvated electron cloud occupies a quasi-spherical atomic s-type orbital and populate higher angular momentum p-, d-, f-, g-type orbitals in excited states.[12][13][14][15]

Its standard electrode potential value is -2.77 V.[16] Equivalent conductivity 177 Mho cm2 is similar to that of hydroxide ion. This value of equivalent conductivity corresponds to a diffusivity of 4,75*10−5 cm2s−1.[17]

Some thermodynamic properties of the solvated electron have been investigated by Joshua Jortner and Richard M. Noyes (1966)[18]

Alkaline aqueous solutions above pH = 9.6 regenerate the hydrated electron through the reaction of hydrated atomic hydrogen with hydroxide ion giving water beside hydrated electrons.

Below pH = 9.6 the hydrated electron reacts with the hydronium ion giving atomic hydrogen, which in turn can react with the hydrated electron giving hydroxide ion and usual molecular hydrogen H2.

The properties of solvated electron can be investigated using the rotating ring-disk electrode.

Reactivity and applications[edit]

The solvated electron reacts with oxygen to form a superoxideradical (O2.−).[19] With nitrous oxide, solvated electrons react to form hydroxyl radicals (HO.).[20] The solvated electrons can be scavenged from both aqueous and organic systems with nitrobenzene or sulfur hexafluoride[citation needed].

A common use of sodium dissolved in liquid ammonia is the Birch reduction. Other reactions where sodium is used as a reducing agent also are assumed to involve solvated electrons, e.g. the use of sodium in ethanol as in the Bouveault–Blanc reduction.

Sodium

Solvated electrons are involved in the reaction of sodium metal with water.[21] Two solvated electrons combine to form molecular hydrogen and hydroxide ion.

Solvated electrons are also involved in electrode processes.[22]

Diffusion[edit]

The diffusivity of the solvated electron in liquid ammonia can be determined using potential-step chronoamperometry.[23]

In gas phase and upper atmosphere of Earth[edit]

Solvated electrons can be found even in the gas phase. This implies their possible existence in the upper atmosphere of Earth and involvement in nucleation and aerosol formation.[24]

References[edit]

  1. ^Schindewolf, U. (1968). 'Formation and Properties of Solvated Electrons'. Angewandte Chemie International Edition in English. 7 (3): 190–203. doi:10.1002/anie.196801901.
  2. ^Cotton, F. A.; Wilkinson, G. (1972). Advanced Inorganic Chemistry. John Wiley and Sons Inc. ISBN978-0-471-17560-5.
  3. ^Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN978-0-08-037941-8.
  4. ^W. Xu and M. M. Lerner, 'A New and Facile Route Using Electride Solutions To Intercalate Alkaline Earth Ions into Graphite', Chem. Mater. 2018, 30, 19, 6930–6935. https://doi.org/10.1021/acs.chemmater.8b03421
  5. ^Kraus, Charles A. (1907). 'Solutions of Metals in Non-Metallic Solvents; I. General Properties of Solutions of Metals in Liquid Ammonia'. J. Am. Chem. Soc.29 (11): 1557–1571. doi:10.1021/ja01965a003.
  6. ^Zurek, Eva (2009). 'A Molecular Perspective on Lithium–Ammonia Solutions'. Angew. Chem. Int. Ed.48 (44): 8198–8232. doi:10.1002/anie.200900373. PMID19821473.
  7. ^Gibson, G. E.; Argo, W. L. (1918). 'The Absorption Spectra of the Blue Solutions of Certain Alkali and Alkaline Earth Metals in Liquid Ammonia and Methylamine'. J. Am. Chem. Soc.40 (9): 1327–1361. doi:10.1021/ja02242a003.
  8. ^Dye, J. L. (2003). 'Electrons as Anions'. Science. 301 (5633): 607–608. doi:10.1126/science.1088103. PMID12893933.
  9. ^Edwin M. Kaiser (2001). 'Calcium-Ammonia'. Calcium–Ammonia. Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rc003. ISBN978-0471936237.
  10. ^Combellas, C; Kanoufi, F; Thiébault, A (2001). 'Solutions of solvated electrons in liquid ammonia'. Journal of Electroanalytical Chemistry. 499: 144–151. doi:10.1016/S0022-0728(00)00504-0.
  11. ^Ariyarathna, Isuru R.; Pawłowski, Filip; Ortiz, Joseph Vincent; Miliordos, Evangelos (2018). 'Molecules mimicking atoms: monomers and dimers of alkali metal solvated electron precursors'. Physical Chemistry Chemical Physics. 20 (37): 24186–24191. doi:10.1039/c8cp05497e.
  12. ^Ariyarathna, Isuru R.; Khan, Shahriar N.; Pawłowski, Filip; Ortiz, Joseph Vincent; Miliordos, Evangelos (4 January 2018). 'Aufbau Rules for Solvated Electron Precursors: Be(NH 3 ) 4 0,± Complexes and Beyond'. The Journal of Physical Chemistry Letters. 9 (1): 84–88. doi:10.1021/acs.jpclett.7b03000.
  13. ^Ariyarathna, Isuru R.; Miliordos, Evangelos (2019). 'Superatomic nature of alkaline earth metal–water complexes: the cases of Be(H 2 O)0,+4 and Mg(H 2 O)0,+6'. Physical Chemistry Chemical Physics. 21 (28): 15861–15870. doi:10.1039/c9cp01897b.
  14. ^Ariyarathna, Isuru R.; Miliordos, Evangelos (2020). 'Geometric and electronic structure analysis of calcium water complexes with one and two solvation shells'. Physical Chemistry Chemical Physics. 22 (39): 22426–22435. doi:10.1039/d0cp04309e.
  15. ^Ariyarathna, Isuru (1 March 2021). 'First Principle Studies on Ground and Excited Electronic States: Chemical Bonding in Main-Group Molecules, Molecular Systems with Diffuse Electrons, and Water Activation using Transition Metal Monoxides'.
  16. ^Baxendale, J. H. (1964), Radiation Res. Suppl., 114 and 139
  17. ^Hart, Edwin J. (1969). 'The Hydrated Electron'. Survey of Progress in Chemistry. 5: 129–184. doi:10.1016/B978-0-12-395706-1.50010-8. ISBN9780123957061.
  18. ^Jortner, Joshua; Noyes, Richard M. (1966). 'Some Thermodynamic Properties of the Hydrated Electron'. The Journal of Physical Chemistry. 70 (3): 770–774. doi:10.1021/j100875a026.
  19. ^Hayyan, Maan; Hashim, Mohd Ali; Alnashef, Inas M. (2016). 'Superoxide Ion: Generation and Chemical Implications'. Chemical Reviews. 116 (5): 3029–3085. doi:10.1021/acs.chemrev.5b00407. PMID26875845.
  20. ^Janata, Eberhard; Schuler, Robert H. (1982). 'Rate constant for scavenging eaq- in nitrous oxide-saturated solutions'. The Journal of Physical Chemistry. 86 (11): 2078–2084. doi:10.1021/j100208a035.
  21. ^Walker, D.C. (1966). 'Production of hydrated electron'. Canadian Journal of Chemistry. 44 (18): 2226–. doi:10.1139/v66-336.
  22. ^B. E. Conway, D. J. MacKinnon, J. Phys. Chem., 74, 3663, 1970
  23. ^Harima, Yutaka; Aoyagui, Shigeru (1980). 'The diffusion coefficient of solvated electrons in liquid ammonia'. Journal of Electroanalytical Chemistry and Interfacial Electrochemistry. 109 (1–3): 167–177. doi:10.1016/S0022-0728(80)80115-X.
  24. ^F. Arnold, Nature 294, 732-733, (1981)

Further reading[edit]

  • Sagar, D. M.; Colin; Bain, D.; Verlet, Jan R. R. (2010). 'Hydrated Electrons at the Water/Air Interface'. J. Am. Chem. Soc. 132 (20): 6917–6919. doi:10.1021/ja101176r. PMID20433171.
  • Martyna, Glenn (1993). 'Electronic states in metal-ammonia solutions'. Physical Review Letters. 71 (2): 267–270. Bibcode:1993PhRvL..71..267D. doi:10.1103/physrevlett.71.267. PMID10054906.
  • Martyna, Glenn (1993). 'Quantum simulation studies of singlet and triplet bipolarons in liquid ammonia'. Journal of Chemical Physics. 98 (1): 555–563. Bibcode:1993JChPh..98..555M. doi:10.1063/1.464650.
  • Solvated Electron. Advances in Chemistry. 50. 1965. doi:10.1021/ba-1965-0050. ISBN978-0-8412-0051-7.
  • Anbar, Michael (1965). 'Reactions of the Hydrated Electron'. Solvated Electron. Advances in Chemistry. 50. pp. 55–81. doi:10.1021/ba-1965-0050.ch006. ISBN978-0-8412-0051-7.
  • Abel, B.; Buck, U.; Sobolewski, A. L.; Domcke, W. (2012). 'On the nature and signatures of the solvated electron in water'. Phys. Chem. Chem. Phys. 14 (1): 22–34. Bibcode:2012PCCP...14...22A. doi:10.1039/C1CP21803D. PMID22075842.
  • Harima, Y.; Aoyagui, S. (1981). 'Determination of the chemical solvation energy of the solvated electron'. Journal of Electroanalytical Chemistry and Interfacial Electrochemistry. 129 (1–2): 349–352. doi:10.1016/S0022-0728(81)80027-7.
  • Hart, Edwin J. (1969). 'The Hydrated Electron'. Survey of Progress in Chemistry Volume 5. Survey of Progress in Chemistry. 5. pp. 129–184. doi:10.1016/B978-0-12-395706-1.50010-8. ISBN9780123957061.
  • The electrochemistry of the solvated electron. Technische Universiteit Eindhoven.
  • IAEA On the Electrolytic Generation of Hydrated Electron.
  • Fundamentals of Radiation Chemistry, chapter 6, p. 145–198, Academic Press, 1999.
  • Tables of bimolecular rate constants of hydrated electrons, hydrogen atoms and hydroxyl radicals with inorganic and organic compounds International Journal of Applied Radiation and Isotopes Anbar, Neta
Retrieved from 'https://en.wikipedia.org/w/index.php?title=Solvated_electron&oldid=1012855156'