Name: how ka and kb show acid-base strength
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 2 min 29 s
Channel: JoVE

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relative-strengths-of-conjugate-acid-base-pairs

Relative Strengths of Conjugate Acid-Base Pairs

Zależność między stałymi dysocjacji sprzężonej pary kwas-zasada można wyrazić ilościowo.
W przypadku słabego kwasu, HA, jego stała dysocjacji kwasu, Ka, jest wyrażona jako stężenie jonów hydroniowych razy stężenie jonów A podzielone przez stężenie HA.
Stała dysocjacji zasady, Kb, dla sprzężonej zasady, jonu A, jest wyrażona jako stężenie HA razy stężenie jonów wodorotlenkowych podzielone przez stężenie jonu A.
Jeżeli wyrażenia dla Ka i Kb zostaną pomnożone, otrzymane równanie jest wyrażeniem równowagi dla Kw. Równanie to pokazuje, że Ka dla słabego kwasu i Kb dla jego sprzężonej zasady są odwrotnie proporcjonalne. Mocniejszy kwas o wyższej wartości Ka zawsze ma proporcjonalnie słabszą sprzężoną zasadę o niższym Kb i odwrotnie.
Ponieważ wartość KW jest stała, równanie to można wykorzystać do obliczenia Ka lub Kb sprzężonej pary kwas-zasada, jeśli którakolwiek z nich jest znana. Na przykład, jeśli Kb dla zasady wynosi 1 × 10−6, to Ka dla jego sprzężonego kwasu można obliczyć jako 1 × 10−8.
Związek między Ka, Kb i KW można również wyrazić w postaci ich ujemnych logarytmów, pKa, pKb i pKW.
Dla równania Ka razy Kb jest równe Kw, gdy weźmie się ujemny logarytm obu stron, otrzymane równanie to pKa plus pKb jest równe pKW, które wynosi 14 w temperaturze 25 °C.
pKa i pKb można również wykorzystać do porównania mocy słabych kwasów i zasad. Im niższa wartość pKa, tym silniejszy kwas. W ten sam sposób, im niższa wartość pKb, tym silniejsza podstawa. Na przykład kwas o pKa 2,1 jest silniejszy niż kwas o pKa 4,6.

względne siły sprzężonych par kwas-zasada

Chemia kwasowo-zasadowa Brønsteda-Lowry'ego polega na przenoszeniu protonów; zatem logika sugeruje związek między względną mocą sprzężonych par kwas-zasada. Siłę kwasu lub zasady określa się ilościowo poprzez stałą jonizacji Ka lub Kb, która reprezentuje stopień reakcji jonizacji kwasu lub zasady. Dla sprzężonej pary kwas-zasada HA / A−, równania równowagi jonizacji i wyrażenia na stałą jonizacji są następujące:

<img alt="Eq1" src="/files/ftp_upload/11398/11398_Equation_1.png" style="height: 90px;" />

<img alt="Eq2" src="/files/ftp_upload/11398/11398_Equation_2v2.png" style="height: 70px;" />

Dodanie tych dwóch równań chemicznych daje równanie autojonizacji wody:

<img alt="Eq3" src="/files/ftp_upload/11398/11398_Equation_3.png" style="height: 50px;" />

Jak wspomniano wcześniej, stała równowagi dla zsumowanej reakcji jest równa iloczynowi matematycznemu stałych równowagi dla dodanych reakcji, a zatem

<img alt="Eq4" src="/files/ftp_upload/11398/11398_Equation_4.png" style="height: 90px;" />

Równanie to określa związek między stałymi jonizacji dla dowolnej sprzężonej pary kwas-zasada, a mianowicie ich iloczyn matematyczny jest równy iloczynowi jonowemu wody, KW. Przekształcając to równanie, staje się oczywista wzajemna zależność między mocami sprzężonej pary kwas-zasada:

<img alt="Eq5" src="/files/ftp_upload/11398/11398_Equation_5.png" style="height: 90px;" />

Odwrotna proporcjonalność pomiędzy Ka i Kb oznacza, że im silniejszy jest kwas lub zasada, tym słabszy jest jego sprzężony partner.

Biorąc logarytm ujemny obu stron równania, otrzymujemy Ka × Kb = KW

<img alt="Eq6" src="/files/ftp_upload/11398/11398_Equation_6.png" style="height: 50px;" />

Następnie

<img alt="Eq7" src="/files/ftp_upload/11398/11398_Equation_7.png" style="height: 50px;" />

Ponieważ pKW wynosi 14 w temperaturze 25 ° C, równanie to można również zapisać jako

<img alt="Eq1" src="/files/ftp_upload/11398/11398_Equation_8.png" style="height: 50px;" />

Wartości pKa i pKb reprezentują także odpowiednio moc kwasów i zasad. Podobnie jak pH i pOH, im wyższa wartość pKa lub pKb, tym słabszy odpowiednio kwas lub zasada.

<div class="overflow">
<table style="margin-left: auto; margin-right: auto;">
	<tbody>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Zasada</td>
			<td style="color: black; border: 1px solid black; text-align: center;">kwasowa</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Kwas nadchlorowy (HClO4)*</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon nadchloranowy (ClO4−)**</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Kwas siarkowy (H2SO4)*</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon wodorosiarczanowy (HSO4−)**</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Jodowodór (HI)*</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon jodkowy (I−)**</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Bromowodór (HBr)*</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon bromkowy (Br−)**</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Chlorowodór (HCl)*</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon chlorkowy (Cl−)**</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Kwas azotowy (HNO3)*</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon azotanowy (NO3−)**</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon hydroniowy (H3O+)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Woda (H2O)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon wodorosiarczanowy (HSO4−)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon siarczanowy (SO42−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Kwas fosforowy (H3PO4)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon diwodorofosforanowy (H2PO4−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Fluorowodór (HF)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon fluorkowy (F−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Kwas azotawy (HNO2)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon azotynowy (NO2−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Kwas octowy (CH3CO2H)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon octanowy (CH3CO2−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Kwas węglowy (H2CO3)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon wodorowęglanowy (HCO3−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Siarkowodór (H2S)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon siarkowodoru (HS−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon amonowy (NH4+)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Amoniak (NH3)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Cyjanowodór (HCN)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon cyjankowy (CN−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon wodorowęglanowy (HCO3−)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon węglanowy (CO32−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Woda (H2O)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon wodorotlenkowy (OH−)</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon siarkowodoru (HS−)†</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon siarczkowy (S2−)‡</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Etanol (C2H5OH)†</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon etanolanowy (C2H5O−)‡</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Amoniak (NH3)†</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon amidowy (NH2−)‡</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Wodór (H2)†</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon wodorkowy (H−)‡</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Metan (CH4)†</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Jon metidowy (CH3−)‡</td>
		</tr>
		<tr>
			<td colspan="2" style="color: black; border: 1px solid black; ">*Poddaje się całkowitej jonizacji kwasowej w wodzie</td>
		</tr>
		<tr>
			<td colspan="2" style="color: black; border: 1px solid black; ">† Nie ulega jonizacji kwasowej w wodzie</td>
		</tr>
		<tr>
			<td colspan="2" style="color: black; border: 1px solid black; ">**Nie ulega zasadowej jonizacji w wodzie</td>
		</tr>
		<tr>
			<td colspan="2" style="color: black; border: 1px solid black; ">‡ Poddaje się całkowitej jonizacji zasady w wodzie</td>
		</tr>
	</tbody>
</table>
</div>

Przedstawiona lista sprzężonych par kwas-zasada jest ułożona tak, aby pokazać względną siłę każdej substancji w porównaniu z wodą. W kolumnie kwasowej gatunki wymienione poniżej wody są słabszymi kwasami niż woda. Gatunki te nie ulegają jonizacji kwasowej w wodzie. Nie są to kwasy Brønsteda-Lowry'ego. Wszystkie wymienione powyżej gatunki wody są silniejszymi kwasami, przenoszącymi w pewnym stopniu protony do wody po rozpuszczeniu w roztworze wodnym w celu wytworzenia jonów hydroniowych. Gatunki znajdujące się nad wodą, ale poniżej jonu hydroniowego, to słabe kwasy ulegające częściowej jonizacji kwasowej, podczas gdy te powyżej jonu hydroniowego to mocne kwasy, które są całkowicie zjonizowane w roztworze wodnym.

Względne siły sprzężonych par kwas-zasada

Chemia kwasowo-zasadowa Brønsteda-Lowry'ego to przenoszenie protonów; Tak więc logika sugeruje związek między względnymi siłami sprzężonego kwasowo-zasadowego ...

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how ka and kb show acid-base strength

Brønsted-Lowry acid-base chemistry is based on proton transfer. That idea helps show how the strength of an acid is linked to the strength of its conjugate base, and vice versa. Acid and base strength is measured with ionization constants, Ka and Kb, which tell how far an ionization reaction goes.

For the conjugate pair HA and A−, the acid ionizes in water and the base reaction can be written for the partner species. When the two reactions are added, they give the autoionization of water. The equilibrium constant for the summed reaction is the product of the two individual constants, so Ka × Kb = Kw for any conjugate acid-base pair.

This relationship shows a reciprocal pattern. If one member of the pair is stronger, its conjugate partner is weaker. The stronger the acid, the weaker its conjugate base, and the stronger the base, the weaker its conjugate acid.

Taking the negative log of Ka × Kb = Kw gives the pKa and pKb form of the same idea. At 25 °C, pKw is 14, so the relationship can also be written as pKa + pKb = 14. Like pH and pOH, a larger pKa or pKb means a weaker acid or base.

The conjugate acid-base list is arranged to compare each species with water. Species below water in the acid column are weaker acids than water, so they do not undergo acid ionization in water and are not Brønsted-Lowry acids. Species above water transfer protons to water to some extent and form hydronium ions.

Among those acids, the ones above water but below hydronium ion are weak acids, so they ionize only partially in aqueous solution. The acids above hydronium ion are strong acids, and they are completely ionized in water. The list includes examples such as perchloric acid, sulfuric acid, hydrogen chloride, nitric acid, acetic acid, ammonia, ethanol, hydride, and methane with their conjugate partners.

How Ka and Kb Show Acid-Base Strength

bronsted-lowry acids and bases

Proton Transfer in Acids and Bases

Acids and bases can be understood through proton transfer. Early scientists described acids by their sour taste, their effect on dyes, and their reaction with limestone, which releases carbon dioxide. Robert Boyle also noted that acid solutions can lose these properties after contact with alkali, or base, solutions. The modern view developed step by step. Humphry Davy showed that hydrogen is the essential part of acids. Joseph Louis Gay-Lussac then described acids as substances that neutralize...

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Water as an Amphiprotic Substance

Water can act as both a Brønsted-Lowry acid and a Brønsted-Lowry base. When hydrogen fluoride dissolves in water, it donates protons to water. This is acid ionization. The products are hydronium ions and fluoride ions. Water can also accept protons from another substance. When pyridine dissolves in water, it accepts a proton from water molecules. This is base ionization. The products are hydroxide ions and pyridinium ions. These reactions show that water can play either role. A substance that...

ph scale

pH and pOH in Water

pH and pOH show how hydronium and hydroxide ions behave in water and other aqueous solutions. These ion concentrations are linked by the ion product of water, K w , so when one goes up, the other goes down. Because ion concentrations can differ by huge amounts, a logarithmic scale is used to describe them. pH is defined from the hydronium ion concentration, [H3O+], and pOH is the matching p-function for hydroxide ion concentration. These two values are also connected through the K w...

relative strengths of conjugate acid-base pairs

Brønsted-Lowry acid-base chemistry is based on proton transfer. That idea helps show how the strength of an acid is linked to the strength of its conjugate base, and vice versa. Acid and base strength is measured with ionization constants, Ka and Kb, which tell how far an ionization reaction goes. For the conjugate pair HA and A−, the acid ionizes in water and the base reaction can be written for the partner species. When the two reactions are added, they give the autoionization of water. The...

strong acid and base solutions

Strong Acids, Strong Bases, and pH

Strong acids and strong bases change the pH of a solution by dissociating completely in water. A strong acid produces hydronium ions, which are H3O+ ions, and a strong base produces hydroxide ions, OH-. Because the dissociation is complete, the ion concentration can be found from the starting concentration and the molar ratio. For example, 0.20 M hydrobromic acid dissociates fully in water. It produces 0.20 M hydronium ions and 0.20 M bromide ions. A 0.015 M KOH solution also dissociates...

weak acid solutions

Weak Acid Equilibrium in Water

Weak acid equilibrium in water explains why many acids only partially react when dissolved. Most weak acid molecules stay in their original form, and only a small amount of hydronium ions forms. Weak acids are common in nature. They help create the tangy taste of citrus fruits, the stinging sensation from insect bites, and the smells linked to body odor. Acetic acid is a familiar weak acid. It is the main ingredient in vinegar. Its equation uses a double arrow to show that the reaction with...

weak base solutions

Weak Base pH and Kb Calculations

Weak base pH and Kb calculations show how partially reacting bases change water chemistry. Weak bases produce hydroxide ions when they dissolve by reacting with water molecules. Because the reaction only happens partly, these compounds are classified as weak bases. Weak bases are common in nature and in technology. Ammonia is one major example. Global production of ammonia is typically well over 100 metric tons each year. It is used as an agricultural fertilizer, a starting material for...

mixtures of acids

pH of Acid Mixtures

The pH of acid mixtures depends on the acids present, their dissociation constants, and their starting concentrations. When a solution contains two different acids, the pH can be found by choosing the method that fits their relative strengths. This makes acid mixtures a useful case for applying acid-base rules in chemistry. In a mixture of a strong acid and a weak acid, the strong acid dissociates completely and supplies almost all of the hydronium ions in the solution. The weak acid...

ions as acids and bases

Salt Hydrolysis and pH of Solutions

Salt hydrolysis explains why some salt solutions are acidic, basic, or neutral. Salts are ionic compounds made of cations and anions. In water, either ion may react with water through an acid or base ionization reaction. The pH of the solution depends on the relative acid-base strength of the ions in the salt. Ammonium chloride is a good example of an acidic salt. When it dissolves, it separates into ammonium ions and chloride ions. The ammonium ion is the conjugate acid of ammonia, so it can...

determining the ph of salt solutions

Salt Hydrolysis and pH in Solutions

Salt hydrolysis and pH in solutions depend on the ions released when a salt dissolves. Some ions react with water, while others do not. The balance between the cation and anion helps show whether the solution becomes acidic, basic, or neutral. Salts with pH-neutral anions and cations that produce hydronium ions give acidic solutions with pH below 7. Ammonium nitrate is one example. In NH4NO3, NO3− does not react with water, but NH4+ forms hydronium ions, so the solution is acidic. Other salts...

acid/base strengths and dissociation constants

acidbase-strengths-and-dissociation-constants

Acid and Base Strength in Water

Acid and base strength in water depends on how much a substance ionizes after it dissolves. If ionization is nearly complete, the substance is called strong. If only a small amount ionizes, it is called weak. Weak acids and weak bases are much more common than strong ones. The most common strong acids are HClO4, HCl, HBr, HI, HNO3, and H2SO4. The most common strong bases are LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2. These strong acids and bases ionize essentially completely in water. <br...

polyprotic acids

Stepwise Ionization of Polyprotic Acids

Polyprotic acids lose hydrogen ions in steps. The number of ionizable protons in a molecule determines whether an acid is monoprotic, diprotic, or triprotic. This video also shows that polyprotic bases can accept more than one proton. Monoprotic acids have one ionizable hydrogen atom per molecule. Examples include HCl, HNO3, and HCN. Their reactions with water form hydronium ions and the matching anion. Acetic acid, CH3CO2H, is also monoprotic because only the hydrogen in the carboxyl group...

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How Bonding Affects Acid Strength

Acid strength depends on bonding, electronegativity, and oxidation number. These factors help explain why some hydrogen compounds act as strong acids, while others behave as weak acids or bases. For binary acids, which are hydrogen compounds with one other nonmetal, acidity increases as the H-A bond gets weaker down a group. That is why the order for group 17 is HF &lt; HCl &lt; HBr &lt; HI, and for group 16 it is H2O &lt; H2S &lt; H2Se &lt; H2Te. Across a period, acid strength increases as...

lewis acids and bases

Electron-Pair Reactions in Acid-Base Chemistry

Electron-pair reactions explain Lewis acid-base chemistry. In 1923, G. N. Lewis described acids and bases by how they accept or donate a pair of electrons. When that electron pair forms a coordinate covalent bond, one atom supplies both bonding electrons. This bond is also called a dative bond. A simple example is the reaction of water with a hydrogen ion to make a hydronium ion. Ammonia can also combine with a hydrogen ion to form an ammonium ion. In both cases, the reaction creates a Lewis...

key terms and concepts

See how Ka, Kb, pKa, and pKb reveal the strength of conjugate acid-base pairs, including why stronger acids have weaker conjugate bases.

How Ka and Kb Show Acid-Base Strength | JoVE K12

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The Pauli Exclusion Principle

The arrangement of electrons in the orbitals of an atom is called its electron configuration. We describe an electron configuration with a symbol that ...

The Energies of Atomic Orbitals

In an atom, the negatively charged electrons are attracted to the positively charged nucleus. In a multielectron atom, electron-electron repulsions are ...

The Aufbau Principle and Hund's Rule

To determine the electron configuration for any particular atom, we can build the structures in the order of atomic numbers. Beginning with hydrogen, and ...

electron-configuration-of-multielectron-atoms

Electron Configuration of Multielectron Atoms

The alkali metal sodium (atomic number 11) has one more electron than the neon atom. This electron must go into the lowest-energy subshell available, the ...

Use models to explain how electrons are distributed in atoms.

PS.Chm.1.3

Periodic Classification of the Elements

The periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their ...

The Periodic Table

As early chemists discovered more elements, they realized that various elements could be grouped by their similar chemical behaviors. One such grouping ...

Atomic Mass

Atoms &mdash; and the protons, neutrons, and electrons that compose them &mdash; are extremely small. For example, a carbon atom weighs less than 2 ...

Molar Mass

The identity of a substance is defined not only by the types of atoms or ions it contains but by the quantity of each type of atom or ion. For example, ...

Use the Periodic Table as a model to predict the relative properties of elements based on the pattern of valence electrons in the outermost energy levels of atoms.

PS.Chm.2.1

atomic-radii-and-effective-nuclear-charge

Atomic Radii and Effective Nuclear Charge

The elements in groups of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number ...

Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it ...

Ionization Energy

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy ...

Electron Affinity

The electron affinity (EA) is the energy change for adding an electron to a gaseous atom to form an anion (negative ion).



This process can be either ...

Construct an explanation to infer the atomic size, reactivity, electronegativity, and ionization energy of an element based on its position in the Periodic Table.

PS.Chm.2.2

Types of Chemical Bonds

Chemical bonding theories were pioneered by American chemist Gilbert&#xA0;N. Lewis. He developed a model called the Lewis model to explain the type and ...

Lewis Symbols and the Octet Rule

Chemical bonds are complex interactions between two or more atoms or ions, which reduce the potential energy of the molecule. Gilbert N. Lewis developed a ...

Ionic Bonding and Electron Transfer

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence ...

The Born-Haber Cycle

Lattice Energy&#xA0;

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a ...

trends-in-lattice-energy-ion-size-and-charge

Trends in Lattice Energy: Ion Size and Charge

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure ...

Covalent Bonding and Lewis Structures

Compared to ionic bonds, which results from the transfer of electrons between metallic and nonmetallic atoms, covalent bonds result from the mutual ...

Electronegativity

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity.&#xA0;

Electronegativity values of ...

bond-polarity-dipole-moment-and-percent-ionic-character

Bond Polarity, Dipole Moment, and Percent Ionic Character

Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

lewis-structures-of-molecular-compounds-and-polyatomic-ions

Lewis Structures of Molecular Compounds and Polyatomic Ions

To draw Lewis structures for complicated molecules and molecular ions, it is helpful to follow a step-by-step procedure as outlined:


	Determine the ...

Resonance

The Lewis structure of a nitrite anion (NO2&minus;) may actually be drawn in two different ways, distinguished by the locations of the N-O and N=O ...

Formal Charges

In some cases, there are seemingly more than one valid Lewis structures for molecules and polyatomic ions. The concept of formal charges can be used to ...

Exceptions to the Octet Rule

Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:


	...

Bond Energies and Bond Lengths

Stable molecules exist because covalent bonds hold the atoms together. The strength of a covalent bond is measured by the energy required to break it, ...

Analyze and interpret data to explain the mechanisms and properties of the two main types of intramolecular (ionic and covalent) bonds.

PS.Chm.3.1

Analyze the structure of atoms and isotopes.

How Ka and Kb Show Acid-Base Strength

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