H Cation Lewis Structure
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A hydrogen ion is created when a hydrogen atom loses or gains an electron. A positively charged hydrogen ion (or proton) can readily combine with other particles and therefore is only seen isolated when it is in a gaseous state or a nearly particle-free space.[1] Due to its extremely high charge density of approximately 2×1010 times that of a sodium ion, the bare hydrogen ion cannot exist freely in solution as it readily hydrates, i.e., bonds quickly.[2] The hydrogen ion is recommended by IUPAC as a general term for all ions of hydrogen and its isotopes.[3] Depending on the charge of the ion, two different classes can be distinguished: positively charged ions and negatively charged ions.
*H Cation Lewis Structure Pogil
*H Cation Lewis Structure Definition
*H Cation Lewis Structure Worksheet AnswersCation (positively charged)[edit]Zundel cation
Electron dot diagrams for ions are the same as for atoms, except that some electrons have been removed for cations, while some electrons have been added for anions. Thus in comparing the electron configurations and electron dot diagrams for the Na atom and the Na + ion, we note that the Na atom has a single valence electron in its Lewis diagram.
A hydrogen atom is made up of a nucleus with charge +1, and a single electron. Therefore, the only positively charged ion possible has charge +1. It is noted H+.
Depending on the isotope in question, the hydrogen cation has different names:
*Hydron: general name referring to the positive ion of any hydrogen isotope (H+)
*Proton: 1H+ (i.e. the cation of protium)
*Deuteron: 2H+, D+
*Triton: 3H+, T+
Model Ill: Complex Lewis Dot Structures / Polyatomic Ions This! Use the atom cards and Cheerios to build the polyatomic ion before drawing it. Polyatomic ions are a group of atoms that are covalently bonded and act as a single unit with a charge. The Lewis dot structure(s) for NH20H. Rules for Writing Lewis Structures. Count the total number of valence electrons in the molecule or polyatomic ion. (For example, H 2 O has 2x1 + 6 = 8 valence electrons, CCl 4 has 4 + 4x7 = 32 valence electrons.) For anions, add one valence electron for each unit of negative charge; for cations, subtract one electron for each unit of positive charge. Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure.
In addition, the ions produced by the reaction of these cations with water as well as their hydrates are called hydrogen ions:
*Hydronium ion: H3O+
*Zundel cation: H5O2+ (named for Georg Zundel)
*Eigen cation: H9O4+ (or H3O+ •3H2O) (named for Manfred Eigen)
Zundel cations and Eigen cations play an important role in proton diffusion according to the Grotthuss mechanism.
In connection with acids, ’hydrogen ions’ typically refers to hydrons.
In the image at left the hydrogen atom (center) contains a single proton and a single electron. Removal of the electron gives a cation (left), whereas addition of an electron gives an anion (right). The hydrogen anion, with its loosely held two-electron cloud, has a larger radius than the neutral atom, which in turn is much larger than the bare proton of the cation. Hydrogen forms the only cation that has no electrons, but even cations that (unlike hydrogen) still retain one or more electrons are still smaller than the neutral atoms or molecules from which they are derived.Anion (negatively charged)[edit]
Hydrogen anions are formed when additional electrons are acquired:
*Hydride: general name referring to the negative ion of any hydrogen isotope (H−)
*Protide: 1H−
*Deuteride: 2H−, D −
*Tritide: 3H−, T −Uses[edit]
Hydrogen ions drive ATP synthase in photosynthesis. This happens when hydrogen ions get pushed across the membrane creating a high concentration inside the thylakoid membrane and a low concentration in the cytoplasm. However, because of osmosis, the H+ will force itself out of the membrane through ATP synthase. Using their kinetic energy to escape, the protons will spin the ATP synthase which in turn will create ATP. This happens in cellular respiration as well though the concentrated membrane will instead be the inner membrane of the mitochondria.
Hydrogen ions concentration, measured as pH, is also responsible for the acidic or basic nature of a compound. Water molecules split to form H+ and hydroxide anions. This process is referred to as the self-ionization of water.Ocean acidification[edit]
The concentration of hydrogen ions and pH are inversely proportional; in an aqueous solution, an increased concentration of hydrogen ions yields a low pH, and subsequently, an acidic product. By definition, an acid is an ion or molecule that can donate a proton, and when introduced to a solution it will react with water molecules (H2O) to form a hydronium ion (H3O+), a conjugate acid of water.[4] For simplistic reasoning, the hydrogen ion (H+) is often used to abbreviate the hydronium ion.
Ocean acidification is the direct consequence of elevated concentrations of hydrogen ions and carbonate saturation from significant absorption of carbon dioxide (CO2) by the world’s oceans.[5] The pre-industrial state of the ocean’s carbonate chemistry has been notably stable, including the balance of its pH.[6] Following the industrial revolution, anthropogenic emissions of burning fossil fuels, cement production, and changes in land use, have increased the oceans uptake of carbon dioxide from the atmosphere by 30%.[7] In the ocean, the absorption capacity of this greenhouse gas is 59 times higher than in the atmosphere;[8] the ocean acts as the largest carbon sink on the planet, playing a significant role in climate regulation.[9] In addition to carbon fluxes, the natural process of carbon sequestration from the atmosphere into the deep ocean is facilitated by two systems, the biological pump and the solubility pump. The solubility pump is a physico-chemical process that transfers CO2 at the air-sea interface.[10] Based on Henry’s Law, the amount of dissolved CO2 in an aqueous solution is directly proportional to the partial pressure of CO2 in the atmosphere.[11] To maintain equilibrium, a state of high atmospheric partial pressure of CO2 leads to an increased oceanic exchange of this gas by molecular diffusion.
In the surface waters, dissolved atmospheric carbon dioxide (CO2(aq)) reacts with water molecules to form carbonic acid (H2CO3), a weak diprotic acid. Diprotic acids consist of two ionizable hydrogen atoms in each molecule.[12] In an aqueous solution, partial dissociation of carbonic acid releases a hydrogen proton (H+) and a bicarbonate ion (HCO3-), and subsequently, the bicarbonate ion dissociates into an additional hydrogen proton and a carbonate ion (CO32-).[13] The dissolving and dissociating of these inorganic carbon species generate an increase in the concentration of hydrogen ions and inversely lowers ambient surface ocean pH. The carbonate buffering system governs the acidity of seawater by maintaining dissolved inorganic carbon species in chemical equilibrium.
The chemical equation consists of reactants and products that may react in either direction. More reactants added to a system yield more product production (the chemical reaction shifts to the right) and if more product is added, additional reactants will form, shifting the chemical reaction to the left. Therefore, in this model, a high concentration of the beginning reactant, carbon dioxide, produces an increased amount of end-product (H+ and CO32-), thus lowering pH and creating a more acidic solution. The natural buffering system of the ocean resist the change in pH by producing more bicarbonate ions generated by free acid protons reacting with carbonate ions to form an alkaline character.[14] However, increasing atmospheric CO2 concentrations may exceed the buffering capacity threshold, consequently resulting in higher rates of ocean acidification. Shifts in the ocean’s carbonate chemistry has the potential to manipulate ocean biogeochemical cycles for many elements and compounds causing profound impacts on marine ecosystems. Furthermore, the solubility of CO2 is temperature dependent; elevated surface water temperatures reduce CO2 solubility. A continual rise in atmospheric partial pressure of CO2 could potentially convert the ocean from acting as sink (the vertical transport of carbon to the depths of the ocean) to becoming a source (CO2 degassing from the ocean), further increasing global temperatures.[15]See also[edit]References[edit]
*^’Hydrogen ion - chemistry’. britannica.com. Retrieved 18 March 2018.
*^due to its extremely high charge density of approximately 2×1010 times that of a sodium ion
*^Compendium of Chemical Terminology, 2nd edition McNaught, A.D. and Wilkinson, A. Blackwell Science, 1997 ISBN0-86542-684-8, also onlineArchived 2005-12-12 at the Wayback Machine
*^OpenStax, Chemistry. OpenStax CNX. Jun 20, 2016 http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@9.311.
*^W.S. Broecker, T. Takahashi (1997) Neutralization of fossil fuel CO2 by marine calcium carbonate
*^P.N. Pearson, M.R. Palmer (2000) Atmospheric carbon dioxide concentrations over the past 60 million years Nature, 406, pp. 695-699
*^C.L. Sabine, et al. (2004). The oceanic sink for anthropogenic CO2Science, 305 (5682), pp. 367-371
*^Lal R. (2008). Carbon sequestration. Philosophical transactions of the Royal Society of London. Series B, Biological sciences, 363(1492), 815–830. https://doi.org/10.1098/rstb.2007.2185
*^Ben I. Mcneil & Richard J. Matear (2007). Climate change feedbacks on future oceanic acidification, Tellus B: Chemical and Physical Meteorology, 59:2, 191-198
*^Hessen, D., Ågren, G., Anderson, T., Elser, J., & De Ruiter, P. (2004). Carbon Sequestration in Ecosystems: The Role of Stoichiometry. Ecology, 85(5), 1179-1192. Retrieved November 22, 2020, from http://www.jstor.org/stable/3450161
*^Avishay DM, Tenny KM. Henry’s Law. [Updated 2020 Sep 7]. In: StatPearls [Internet]. Treasure Island (FL): StatPearls Publishing; 2020 Jan-. Available from: https://www.ncbi.nlm.nih.gov/books/NBK544301/
*^OpenStax, Chemistry. OpenStax CNX. Jun 20, 2016 http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@9.311.
*^OpenStax, Chemistry. OpenStax CNX. Jun 20, 2016 http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@9.311.
*^Middelburg, J. J., Soetaert, K., & Hagens, M. (2020). Ocean Alkalinity, Buffering and Biogeochemical Processes. Reviews of geophysics (Washington, D.C. : 1985), 58(3), e2019RG000681. https://doi.org/10.1029/2019RG000681
*^Matsumoto, K. (2007). Biology-mediated temperature control on atmosphericpCO2and ocean biogeochemistry. Geophysical Research Letters, 34(20). doi:10.1029/2007gl031301Retrieved from ’https://en.wikipedia.org/w/index.php?title=Hydrogen_ion&oldid=993122316’
GENERAL CHEMISTRY TOPICSLewis structures Examples of how to draw Lewis structures: Water (H2O), Dinitrogen monoxide (Nitrous oxide, N2O), acetic acid (C2H4O2). General rules for drawing Lewis structures.
Lewis structures are structural formulas for molecules and polyatomic ions that represent all valence electrons. Since valence electrons are typically represented as dots, these structural formulas sometimes are called Lewis dot stutctures. Here we present some examples of how to draw Lewis structures. The general rules for drawing Lewis structures are given below.
Example 1. Water. The formula for water is H2O. Counting valence electrons yields eight total (six from oxygen, one each from the two hydrogens). Hydrogen is special since it can only accommodate a duet, and therefore can form at most one bond. Atoms that can only form one bond must be terminal (or peripheral) atoms in the structure.
The skeletal structure for water must be H O H, and not H H O, even though hydrogen has lower electronegativity than oxygen (compare this case to that of N2O, which follows a general rule that less electronegative atoms tend to be central atoms in Lewis structures). H Cation Lewis Structure Pogil
In panel (a), the skeletal structure for water is shown with the atoms represented by their Lewis symbols. The eight total valence electrons are explicitly represented. In (b) we allow single electrons - one each from hydrogen and oxygen - to form a bonding pair between the nuclei. In (c), we have replaced both bonding pairs with a line (or dash) to symbolize the covalent bond formed between the atoms by the bonding electron pair. Note that oxygen is surrounded by an octet of electrons, satisfying the octet rule. We have drawn a valid Lewis structure for water.
Example 2. Dinitrogen monoxide (Nitrous oxide, N2O). Since either N or O can serve as a central atom, in choosing a skeletal structure we use the rule that the least electronegative atom be placed in the middle.
Having the appropriate skeletal structure and correct count of valence electrons, the goal is to place the electrons in bonds or lone pairs so that each atom has an octet of electrons. The figure at left shows a procedure for generating candidate structures by moving electrons singly, or in pairs, starting with the Lewis symbols for the elements. Since N and O are both Period 2 elements, no atom can exceed an octet. In this case, it is possible to draw three valid Lewis structures.
These three valid Lewis structures for dinitrogen monoxide are known as non-equivalent resonance structures. As an assessment tool, formal charge assignments can be used to predict the relative contributions of the resonance forms to the resonance hybrid, which represents a more realistic conception of the electron distribution within the molecule.
Example 3. Acetic acid (C2H4O2). There are cases where even for very simple molecules there are several chemically plausible skeletal structures. For structures of larger molecules, in which more than one atom is joined to two or more atoms, there is no longer a unique central atom, and the probability of encountering isomers dramatically increases. In order to draw the Lewis structure for a given isomer, more information about the skeletal structure is necessary. Such is the case for most organic molecules, those containing primarily carbon, hydrogen, and oxygen, with other elements such as nitrogen, sulfur, and phosphorous also common in molecules of biological importance. An example is acetic acid, an important example of a weak acid. For acetic acid, the skeletal structure is centered on a chain of atoms bonded together as C C O. The first carbon atom uses its three remaining valence electrons to form bonds to three of the four hydrogens. To the second carbon atom is attached another oxygen atom. Finally, the remaining hydrogen atom attaches to one of the oxygen atoms.
In this example, once the skeletal structure is specified, a Lewis structure follows fairly readily. Drawing bonds in place of pairs of electrons composed of one electron from each atom yields a bonded skeletal structure. The octet rule can be satisfied for all non-hydrogen atoms when the remaining unpaired electrons are moved in to form a double bond between carbon and oxygen.
Rules for drawing Lewis structures
Goal: Given a chemical formula corresponding to a molecule or molecular ion, draw a Lewis structure.
1. First of all, a correct count of all valence electrons is essential. One way to do this is to write the Lewis symbols for all of the atoms in the formula, and count up all the ’dots’. For a molecule (uncharged), that count is the correct number of valence electrons. For polyatomic ions, total the valence electrons for all atoms in the formula and subtract one electron for each positive charge of a cation, and add one electron for each unit negative charge of an anion.
2. Draw a skeletal structure. What this means is that we decide how the atoms are to be bonded. Choose a central atom (we’ll start with small molecule examples for which there is only one central atom, and the other atoms - the peripheral atoms - are all bonded to the central atom). Hydrogen (H) and fluorine (F) each have valence of 1, and generally these will not be central atoms (bonded to more than one atom). Given a formula, the central atom is typically the first atom (ClF4 example), although this convention is not always followed (e.g. HNO3). Another good way to choose is to pick the least electronegative atom. Inevitably, there will be cases where it is possible to draw more than one skeletal structure.
3. Draw bonds as lines between atoms. Each bond counts as 2 e−.
4. Add electrons as non-bonding lone pairs around the peripheral atoms so that they have octets (eight electrons total). Note that this does not apply to H, which can only accommodate a duet (2 e−).
5. Add remaining pair(s) of electrons to the central atom so that its octet is complete (if not already). Never exceed an octet for a period 2 atom! For periods 3 and greater, atoms are large enough to accommodate more than an octet in their valence bonding shell. If there are no further electrons available and the central still does not have a complete octet, a lone pair on a peripheral atom may be pushed into a second (or third) bond with the central atom. Carbon and nitrogen are 2nd period elements that commonly form double and triple bonds as central atoms, and oxygen as a peripheral atom is often found in a double bond with the central atom. H Cation Lewis Structure Definition
6. If all atoms from the 2nd period and greater have at least an octet, and no 2nd period atom exceeds an octet, and the total number of electrons in bonds and lone pairs is equal to the total number of valence electrons available, then a valid Lewis structure has been produced. The convention for ions is to enclose the structure in brackets, and indicate the net charge at the upper right corner.
7. Evaluation, exceptions, and use. Note that there are a few cases where the best Lewis structure has an incomplete octet on a central atom. Since it is often possible to draw more than one valid Lewis structure for a molecule or molecular ion, we will need evaluate which one(s) are more plausible or make better chemical sense. As noted above, formal charge is used as a guide in that total number of formal charges zero or a minimum is generally best, and the formal charge of an atom is considered in relation to its electronegativity. Remember, a Lewis structure is not the molecule, but only a shorthand symbolism that is meant to convey some information about it. Such information informs prediction of the likely physical or chemical properties of the real molecule or the bulk substance that is made up of those molecules. One of the most common uses we make of a valid Lewis structure (and for this we do not need the best Lewis structu
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A hydrogen ion is created when a hydrogen atom loses or gains an electron. A positively charged hydrogen ion (or proton) can readily combine with other particles and therefore is only seen isolated when it is in a gaseous state or a nearly particle-free space.[1] Due to its extremely high charge density of approximately 2×1010 times that of a sodium ion, the bare hydrogen ion cannot exist freely in solution as it readily hydrates, i.e., bonds quickly.[2] The hydrogen ion is recommended by IUPAC as a general term for all ions of hydrogen and its isotopes.[3] Depending on the charge of the ion, two different classes can be distinguished: positively charged ions and negatively charged ions.
*H Cation Lewis Structure Pogil
*H Cation Lewis Structure Definition
*H Cation Lewis Structure Worksheet AnswersCation (positively charged)[edit]Zundel cation
Electron dot diagrams for ions are the same as for atoms, except that some electrons have been removed for cations, while some electrons have been added for anions. Thus in comparing the electron configurations and electron dot diagrams for the Na atom and the Na + ion, we note that the Na atom has a single valence electron in its Lewis diagram.
A hydrogen atom is made up of a nucleus with charge +1, and a single electron. Therefore, the only positively charged ion possible has charge +1. It is noted H+.
Depending on the isotope in question, the hydrogen cation has different names:
*Hydron: general name referring to the positive ion of any hydrogen isotope (H+)
*Proton: 1H+ (i.e. the cation of protium)
*Deuteron: 2H+, D+
*Triton: 3H+, T+
Model Ill: Complex Lewis Dot Structures / Polyatomic Ions This! Use the atom cards and Cheerios to build the polyatomic ion before drawing it. Polyatomic ions are a group of atoms that are covalently bonded and act as a single unit with a charge. The Lewis dot structure(s) for NH20H. Rules for Writing Lewis Structures. Count the total number of valence electrons in the molecule or polyatomic ion. (For example, H 2 O has 2x1 + 6 = 8 valence electrons, CCl 4 has 4 + 4x7 = 32 valence electrons.) For anions, add one valence electron for each unit of negative charge; for cations, subtract one electron for each unit of positive charge. Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure.
In addition, the ions produced by the reaction of these cations with water as well as their hydrates are called hydrogen ions:
*Hydronium ion: H3O+
*Zundel cation: H5O2+ (named for Georg Zundel)
*Eigen cation: H9O4+ (or H3O+ •3H2O) (named for Manfred Eigen)
Zundel cations and Eigen cations play an important role in proton diffusion according to the Grotthuss mechanism.
In connection with acids, ’hydrogen ions’ typically refers to hydrons.
In the image at left the hydrogen atom (center) contains a single proton and a single electron. Removal of the electron gives a cation (left), whereas addition of an electron gives an anion (right). The hydrogen anion, with its loosely held two-electron cloud, has a larger radius than the neutral atom, which in turn is much larger than the bare proton of the cation. Hydrogen forms the only cation that has no electrons, but even cations that (unlike hydrogen) still retain one or more electrons are still smaller than the neutral atoms or molecules from which they are derived.Anion (negatively charged)[edit]
Hydrogen anions are formed when additional electrons are acquired:
*Hydride: general name referring to the negative ion of any hydrogen isotope (H−)
*Protide: 1H−
*Deuteride: 2H−, D −
*Tritide: 3H−, T −Uses[edit]
Hydrogen ions drive ATP synthase in photosynthesis. This happens when hydrogen ions get pushed across the membrane creating a high concentration inside the thylakoid membrane and a low concentration in the cytoplasm. However, because of osmosis, the H+ will force itself out of the membrane through ATP synthase. Using their kinetic energy to escape, the protons will spin the ATP synthase which in turn will create ATP. This happens in cellular respiration as well though the concentrated membrane will instead be the inner membrane of the mitochondria.
Hydrogen ions concentration, measured as pH, is also responsible for the acidic or basic nature of a compound. Water molecules split to form H+ and hydroxide anions. This process is referred to as the self-ionization of water.Ocean acidification[edit]
The concentration of hydrogen ions and pH are inversely proportional; in an aqueous solution, an increased concentration of hydrogen ions yields a low pH, and subsequently, an acidic product. By definition, an acid is an ion or molecule that can donate a proton, and when introduced to a solution it will react with water molecules (H2O) to form a hydronium ion (H3O+), a conjugate acid of water.[4] For simplistic reasoning, the hydrogen ion (H+) is often used to abbreviate the hydronium ion.
Ocean acidification is the direct consequence of elevated concentrations of hydrogen ions and carbonate saturation from significant absorption of carbon dioxide (CO2) by the world’s oceans.[5] The pre-industrial state of the ocean’s carbonate chemistry has been notably stable, including the balance of its pH.[6] Following the industrial revolution, anthropogenic emissions of burning fossil fuels, cement production, and changes in land use, have increased the oceans uptake of carbon dioxide from the atmosphere by 30%.[7] In the ocean, the absorption capacity of this greenhouse gas is 59 times higher than in the atmosphere;[8] the ocean acts as the largest carbon sink on the planet, playing a significant role in climate regulation.[9] In addition to carbon fluxes, the natural process of carbon sequestration from the atmosphere into the deep ocean is facilitated by two systems, the biological pump and the solubility pump. The solubility pump is a physico-chemical process that transfers CO2 at the air-sea interface.[10] Based on Henry’s Law, the amount of dissolved CO2 in an aqueous solution is directly proportional to the partial pressure of CO2 in the atmosphere.[11] To maintain equilibrium, a state of high atmospheric partial pressure of CO2 leads to an increased oceanic exchange of this gas by molecular diffusion.
In the surface waters, dissolved atmospheric carbon dioxide (CO2(aq)) reacts with water molecules to form carbonic acid (H2CO3), a weak diprotic acid. Diprotic acids consist of two ionizable hydrogen atoms in each molecule.[12] In an aqueous solution, partial dissociation of carbonic acid releases a hydrogen proton (H+) and a bicarbonate ion (HCO3-), and subsequently, the bicarbonate ion dissociates into an additional hydrogen proton and a carbonate ion (CO32-).[13] The dissolving and dissociating of these inorganic carbon species generate an increase in the concentration of hydrogen ions and inversely lowers ambient surface ocean pH. The carbonate buffering system governs the acidity of seawater by maintaining dissolved inorganic carbon species in chemical equilibrium.
The chemical equation consists of reactants and products that may react in either direction. More reactants added to a system yield more product production (the chemical reaction shifts to the right) and if more product is added, additional reactants will form, shifting the chemical reaction to the left. Therefore, in this model, a high concentration of the beginning reactant, carbon dioxide, produces an increased amount of end-product (H+ and CO32-), thus lowering pH and creating a more acidic solution. The natural buffering system of the ocean resist the change in pH by producing more bicarbonate ions generated by free acid protons reacting with carbonate ions to form an alkaline character.[14] However, increasing atmospheric CO2 concentrations may exceed the buffering capacity threshold, consequently resulting in higher rates of ocean acidification. Shifts in the ocean’s carbonate chemistry has the potential to manipulate ocean biogeochemical cycles for many elements and compounds causing profound impacts on marine ecosystems. Furthermore, the solubility of CO2 is temperature dependent; elevated surface water temperatures reduce CO2 solubility. A continual rise in atmospheric partial pressure of CO2 could potentially convert the ocean from acting as sink (the vertical transport of carbon to the depths of the ocean) to becoming a source (CO2 degassing from the ocean), further increasing global temperatures.[15]See also[edit]References[edit]
*^’Hydrogen ion - chemistry’. britannica.com. Retrieved 18 March 2018.
*^due to its extremely high charge density of approximately 2×1010 times that of a sodium ion
*^Compendium of Chemical Terminology, 2nd edition McNaught, A.D. and Wilkinson, A. Blackwell Science, 1997 ISBN0-86542-684-8, also onlineArchived 2005-12-12 at the Wayback Machine
*^OpenStax, Chemistry. OpenStax CNX. Jun 20, 2016 http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@9.311.
*^W.S. Broecker, T. Takahashi (1997) Neutralization of fossil fuel CO2 by marine calcium carbonate
*^P.N. Pearson, M.R. Palmer (2000) Atmospheric carbon dioxide concentrations over the past 60 million years Nature, 406, pp. 695-699
*^C.L. Sabine, et al. (2004). The oceanic sink for anthropogenic CO2Science, 305 (5682), pp. 367-371
*^Lal R. (2008). Carbon sequestration. Philosophical transactions of the Royal Society of London. Series B, Biological sciences, 363(1492), 815–830. https://doi.org/10.1098/rstb.2007.2185
*^Ben I. Mcneil & Richard J. Matear (2007). Climate change feedbacks on future oceanic acidification, Tellus B: Chemical and Physical Meteorology, 59:2, 191-198
*^Hessen, D., Ågren, G., Anderson, T., Elser, J., & De Ruiter, P. (2004). Carbon Sequestration in Ecosystems: The Role of Stoichiometry. Ecology, 85(5), 1179-1192. Retrieved November 22, 2020, from http://www.jstor.org/stable/3450161
*^Avishay DM, Tenny KM. Henry’s Law. [Updated 2020 Sep 7]. In: StatPearls [Internet]. Treasure Island (FL): StatPearls Publishing; 2020 Jan-. Available from: https://www.ncbi.nlm.nih.gov/books/NBK544301/
*^OpenStax, Chemistry. OpenStax CNX. Jun 20, 2016 http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@9.311.
*^OpenStax, Chemistry. OpenStax CNX. Jun 20, 2016 http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@9.311.
*^Middelburg, J. J., Soetaert, K., & Hagens, M. (2020). Ocean Alkalinity, Buffering and Biogeochemical Processes. Reviews of geophysics (Washington, D.C. : 1985), 58(3), e2019RG000681. https://doi.org/10.1029/2019RG000681
*^Matsumoto, K. (2007). Biology-mediated temperature control on atmosphericpCO2and ocean biogeochemistry. Geophysical Research Letters, 34(20). doi:10.1029/2007gl031301Retrieved from ’https://en.wikipedia.org/w/index.php?title=Hydrogen_ion&oldid=993122316’
GENERAL CHEMISTRY TOPICSLewis structures Examples of how to draw Lewis structures: Water (H2O), Dinitrogen monoxide (Nitrous oxide, N2O), acetic acid (C2H4O2). General rules for drawing Lewis structures.
Lewis structures are structural formulas for molecules and polyatomic ions that represent all valence electrons. Since valence electrons are typically represented as dots, these structural formulas sometimes are called Lewis dot stutctures. Here we present some examples of how to draw Lewis structures. The general rules for drawing Lewis structures are given below.
Example 1. Water. The formula for water is H2O. Counting valence electrons yields eight total (six from oxygen, one each from the two hydrogens). Hydrogen is special since it can only accommodate a duet, and therefore can form at most one bond. Atoms that can only form one bond must be terminal (or peripheral) atoms in the structure.
The skeletal structure for water must be H O H, and not H H O, even though hydrogen has lower electronegativity than oxygen (compare this case to that of N2O, which follows a general rule that less electronegative atoms tend to be central atoms in Lewis structures). H Cation Lewis Structure Pogil
In panel (a), the skeletal structure for water is shown with the atoms represented by their Lewis symbols. The eight total valence electrons are explicitly represented. In (b) we allow single electrons - one each from hydrogen and oxygen - to form a bonding pair between the nuclei. In (c), we have replaced both bonding pairs with a line (or dash) to symbolize the covalent bond formed between the atoms by the bonding electron pair. Note that oxygen is surrounded by an octet of electrons, satisfying the octet rule. We have drawn a valid Lewis structure for water.
Example 2. Dinitrogen monoxide (Nitrous oxide, N2O). Since either N or O can serve as a central atom, in choosing a skeletal structure we use the rule that the least electronegative atom be placed in the middle.
Having the appropriate skeletal structure and correct count of valence electrons, the goal is to place the electrons in bonds or lone pairs so that each atom has an octet of electrons. The figure at left shows a procedure for generating candidate structures by moving electrons singly, or in pairs, starting with the Lewis symbols for the elements. Since N and O are both Period 2 elements, no atom can exceed an octet. In this case, it is possible to draw three valid Lewis structures.
These three valid Lewis structures for dinitrogen monoxide are known as non-equivalent resonance structures. As an assessment tool, formal charge assignments can be used to predict the relative contributions of the resonance forms to the resonance hybrid, which represents a more realistic conception of the electron distribution within the molecule.
Example 3. Acetic acid (C2H4O2). There are cases where even for very simple molecules there are several chemically plausible skeletal structures. For structures of larger molecules, in which more than one atom is joined to two or more atoms, there is no longer a unique central atom, and the probability of encountering isomers dramatically increases. In order to draw the Lewis structure for a given isomer, more information about the skeletal structure is necessary. Such is the case for most organic molecules, those containing primarily carbon, hydrogen, and oxygen, with other elements such as nitrogen, sulfur, and phosphorous also common in molecules of biological importance. An example is acetic acid, an important example of a weak acid. For acetic acid, the skeletal structure is centered on a chain of atoms bonded together as C C O. The first carbon atom uses its three remaining valence electrons to form bonds to three of the four hydrogens. To the second carbon atom is attached another oxygen atom. Finally, the remaining hydrogen atom attaches to one of the oxygen atoms.
In this example, once the skeletal structure is specified, a Lewis structure follows fairly readily. Drawing bonds in place of pairs of electrons composed of one electron from each atom yields a bonded skeletal structure. The octet rule can be satisfied for all non-hydrogen atoms when the remaining unpaired electrons are moved in to form a double bond between carbon and oxygen.
Rules for drawing Lewis structures
Goal: Given a chemical formula corresponding to a molecule or molecular ion, draw a Lewis structure.
1. First of all, a correct count of all valence electrons is essential. One way to do this is to write the Lewis symbols for all of the atoms in the formula, and count up all the ’dots’. For a molecule (uncharged), that count is the correct number of valence electrons. For polyatomic ions, total the valence electrons for all atoms in the formula and subtract one electron for each positive charge of a cation, and add one electron for each unit negative charge of an anion.
2. Draw a skeletal structure. What this means is that we decide how the atoms are to be bonded. Choose a central atom (we’ll start with small molecule examples for which there is only one central atom, and the other atoms - the peripheral atoms - are all bonded to the central atom). Hydrogen (H) and fluorine (F) each have valence of 1, and generally these will not be central atoms (bonded to more than one atom). Given a formula, the central atom is typically the first atom (ClF4 example), although this convention is not always followed (e.g. HNO3). Another good way to choose is to pick the least electronegative atom. Inevitably, there will be cases where it is possible to draw more than one skeletal structure.
3. Draw bonds as lines between atoms. Each bond counts as 2 e−.
4. Add electrons as non-bonding lone pairs around the peripheral atoms so that they have octets (eight electrons total). Note that this does not apply to H, which can only accommodate a duet (2 e−).
5. Add remaining pair(s) of electrons to the central atom so that its octet is complete (if not already). Never exceed an octet for a period 2 atom! For periods 3 and greater, atoms are large enough to accommodate more than an octet in their valence bonding shell. If there are no further electrons available and the central still does not have a complete octet, a lone pair on a peripheral atom may be pushed into a second (or third) bond with the central atom. Carbon and nitrogen are 2nd period elements that commonly form double and triple bonds as central atoms, and oxygen as a peripheral atom is often found in a double bond with the central atom. H Cation Lewis Structure Definition
6. If all atoms from the 2nd period and greater have at least an octet, and no 2nd period atom exceeds an octet, and the total number of electrons in bonds and lone pairs is equal to the total number of valence electrons available, then a valid Lewis structure has been produced. The convention for ions is to enclose the structure in brackets, and indicate the net charge at the upper right corner.
7. Evaluation, exceptions, and use. Note that there are a few cases where the best Lewis structure has an incomplete octet on a central atom. Since it is often possible to draw more than one valid Lewis structure for a molecule or molecular ion, we will need evaluate which one(s) are more plausible or make better chemical sense. As noted above, formal charge is used as a guide in that total number of formal charges zero or a minimum is generally best, and the formal charge of an atom is considered in relation to its electronegativity. Remember, a Lewis structure is not the molecule, but only a shorthand symbolism that is meant to convey some information about it. Such information informs prediction of the likely physical or chemical properties of the real molecule or the bulk substance that is made up of those molecules. One of the most common uses we make of a valid Lewis structure (and for this we do not need the best Lewis structu
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