Name: drawing lewis structures with valence electron counts
Uploaded: 2020-09-03T13:38:28.000Z
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Chemistry

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Lewis Structures of Molecular Compounds and Polyatomic Ions

هياكل لويس هي تمثيلات مبسطة من الروابط الكيميائية بين الذرات. كتابة تراكيب لويس،للمركبات المحايدة مثل الإيثين أو الأمونيا،أو للأيونات متعددة الذرات مثل الهيدرونيوم أو الفوسفات،يتضمن التسلسل التالي من الخطوات. أولاًاحسب العدد الإجمالي لإلكترونات التكافؤ في الجزيء.ضع في اعتبارك الإيثين،على سبيل المثال. تحتوي كل ذرة كربون على 4 إلكترونات تكافؤ،وكل ذرة هيدروجين 1 إلكترون تكافؤ مع 2 ذرة كربون و 4 ذرات هيدروجين،يحتوي الإيثين على إجمالي 12 إلكترونًتكافؤ. بعد ذلك،ارسم الهيكل العظمي للجزيء.بشكل عام،الذرة الأولى في الصيغة الكيميائية أو الذرة الأقل سالبيةكهربيًا،يتم وضعها مركزيًا. هذه الذرة تسمى أيضًا الذرة المركزية ويجب أن تكون قادرة على تكوين رابطين على الأقل. ذرات الهيدروجين،أو الذرات الاكثر سالبيةكهربية،يتم وضعها بشكل نهائي حول الذرة المركزية.هنا،تحتل ذرات الكربون المواقع المركزية مع ذرات الهيدروجين المحيطة بهم. الخطوة التالية هي إكمال الثمانية. أولاًضع زوجًا من الإلكترونات بين كل ذرتين لتمثيل أزواج الترابط.يتم توزيع الإلكترونات المتبقية كأزواج مفردة إلى الذرات الطرفية والمركزية لتحقيق التكوين الثنائي أو الثماني للذرات. في هذه الحالة،منذ أن وصل الهيدروجين إلى دويتو،يتم نقل الإلكترونات المتبقية لذرات الكربون. الآن،استخدم الأزواج المفردة لتكوين روابط متعددة بين الذرات التي تفتقر إلى اكتمال الثمانية.في الإيثين،لم تصل ذرات الكربون إلى ثمانية بعد. لذلك،يتم نقل الزوج المفرد إلى منطقة الترابط لتشكيل رابطة مزدوجة تسمح لجميع الذرات للوصول إلى تكوين إلكتروني مستقر. لكتابة بنية لويس للأيونات متعددة الذرات المشحونة،مثل الهيدرونيوم أو الفوسفات،نفس النهج يتبع،مع التعديل لاستيعاب الشحنة الأيونية.لرسم بنية لويس للهيدرونيوم،أولاًاحسب عدد إلكترونات التكافؤ على الأيون. بما أن أيون الهيدرونيوم له شحنة موجبة،نُزيل إلكترون من إجمالي إلكترونات التكافؤ،لتقليل العدد إلى ثمانية. بعد ذلك،وضِح الجزيء من خلال تركيبه الهيكلي.الآن ضع زوجًا من الإلكترونات بين كل ذرة،متبوعًا بزوج الإلكترونات المفردة على الذرة المركزية. بهذه الطريقة،ذرات الهيدروجين والأكسجين تصل للتكوين الثنائي والثماني،على التوالي. أخيرًا،اكتب بنية لويس بين قوسين و شحنة موجبة في الزاوية العلوية اليمنى.

تركيبات لويس للمركبات الجزيئية والأيونات متعددة الذرات

لرسم هياكل لويس للجزيئات المعقدة والأيونات الجزيئية، من المفيد اتباع الإجراء خطوة بخطوة كما هو موضح: <ol><li> حدد العدد الإجمالي لإلكترونات التكافؤ (الغلاف الخارجي). بالنسبة للكاتيونات، اطرح إلكتروناً واحداً لكل شحنة موجبة. بالنسبة للأنيونات، أضف إلكتروناً واحداً لكل شحنة سالبة. </li><li> ارسم شكلاً هيكلياً للجزيء أو الأيون، مع ترتيب الذرات حول ذرة مركزية. (بشكل عام ، يجب وضع أقل عنصر كهرسلبي في المركز.) قم بتوصيل كل ذرة بالذرة المركزية برابطة واحدة (زوج إلكترون واحد). </li><li> وزع الإلكترونات المتبقية كأزواج وحيدة على الذرات الطرفية (باستثناء الهيدروجين)، مع استكمال مجموعة ثمانية حول كل ذرة. </li><li> ضع كل الإلكترونات المتبقية على الذرة المركزية. </li><li> أعد ترتيب إلكترونات الذرات الخارجية لعمل روابط متعددة مع الذرة المركزية للحصول على مجموعة ثمانية كلما أمكن ذلك. </li></ol> على سبيل المثال، ضع في الاعتبار SiH4 و CHO2– ، NO+ و OF2 كأمثلة يمكن من أجلها تطبيق هذا المبدأ التوجيهي العام لتحديد هياكل لويس الخاصة بهم. <ol><li> تحديد العدد الإجمالي لإلكترونات التكافؤ (الغلاف الخارجي) في الجزيء أو الأيون. & nbsp؛ </li></ol> بالنسبة لجزيء مثل SiH4، يتم إضافة عدد إلكترونات التكافؤ على كل ذرة في الجزيء: & nbsp؛ = [4 تكافؤ e−/Si atom × 1 Si atom] + [ 1 تكافؤ e−/H atom × 4 H atoms] = 8 تكافؤ e−<h4>CHO2–</h4> بالنسبة للأيون السالب، مثل  CHO2–، تتم إضافة عدد إلكترونات التكافؤ على الذرات إلى عدد الشحنات السالبة على الأيون (يتم اكتساب إلكترون واحد لكل شحنة سالبة): = [4 تكافؤ e−/C atom × 1 C atom] + [1 تكافؤ e−/H atom × 1 H atom] + [6 تكافؤ e−/O atom × 2 O atoms] + [1 additional e−] = 18 تكافؤ e−<h4>NO+</h4> بالنسبة إلى أيون موجب ، مثل NO+، يتم إضافة عدد إلكترونات التكافؤ على الذرات في الأيون، متبوعاً بطرح عدد الشحنات الموجبة على الأيون (يُفقد إلكترون واحد مقابل كل شحنة موجبة واحدة) من العدد الإجمالي لإلكترونات التكافؤ: = [5 تكافؤ e−/N atom × 1 N atom] + [6 تكافؤ e−/O atom × 1 O atom] + [−1 e−] = 10 تكافؤ e−<h4>OF2 </h4>OF2 نظراً لكونه جزيئاً محايداً، يتم ببساطة إضافة عدد إلكترونات التكافؤ: = [6 تكافؤ e−/O atom × 1 O atom] + [7 تكافؤ e−/F atom × 2 F atoms] = 20 تكافؤ e−<ol start="2"><li> ارسم جدولاً هيكلياً للجزيء أو الأيون، ورتب الذرات حول ذرة مركزية وربط كل ذرة بالذرة المركزية برابطة واحدة (زوج إلكترون واحد). (لاحظ أنه يتم الإشارة إلى الأيونات بأقواس حول الهيكل، والشحنة الأيونية خارج الأقواس:) <img alt="Figure1" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_1.svg" style="height: 80px;" /> في الحالات التي تكون فيها عدة ترتيبات للذرات ممكنة ، كما هو الحال في CHO2−، يتم استخدام الدليل التجريبي لاختيار الترتيب الصحيح. بشكل عام، من المرجح أن تكون العناصر الأقل كهرسلبية هي ذرات مركزية. في  CHO2−، تحتل ذرة الكربون  الأقل كهرسلبيّة الموقع المركزي مع ذرات الأكسجين والهيدروجين المحيطة بها. تشمل الأمثلة الأخرى P في POCl3, S in SO2، و Cl في  ClO4−. استثناء هو أن الهيدروجين لا يكاد يكون ذرة مركزية. باعتباره العنصر الأكثر كهرسلبية، لا يمكن أن يكون الفلور أيضاً ذرة مركزية. </li><li> وزع الإلكترونات المتبقية كأزواج وحيدة على الذرات الطرفية (باستثناء الهيدروجين) لإكمال غلاف التكافؤ بثمانية إلكترونات. (مع عدم وجود إلكترونات متبقية على SiH4، لم يتغير هيكلها.) <<img alt="Figure2" src="https://www.jovecore.com/CDNSource/lessons/1063/1063image_2.svg" style="height: 80px;" /></li><li> ضع كل الإلكترونات المتبقية على الذرة المركزية. & nbsp؛<ul><li> بالنسبة إلى SiH4,  CHO2−,  و NO+ ، هناك لا توجد إلكترونات متبقية. بالنسبة لـ OF2، من بين 16 إلكتروناً متبقياً، يتم وضع 12 إلكتروناً، وبالتالي يتم وضع 4 إلكترونات على الذرة المركزية: <img alt="Figure3" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_3.svg" style="height: 40px;" /></li></ul></li></ul></li>
 
 <li> إعادة ترتيب إلكترونات الذرات الخارجية لعمل روابط متعددة مع الذرة المركزية، للحصول على مجموعة ثمانية كلما أمكن ذلك. &nbsp؛<ul><li> SiH4: يحتوي Si بالفعل على مجموعة ثمانية، لذا لا يلزم فعل أي شيء. & nbsp؛ <img alt="Figure4" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_4.svg" style="height: 70px;" /></li><li>CHO2−: nbsp؛ يتم توزيع إلكترونات  التكافؤ كأزواج وحيدة على ذرات الأكسجين، لكن ذرة الكربون تفتقر إلى مجموعة ثمانية. & nbsp؛ <img alt="Figure5" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_5.svg" style="height: 80px;" /></li><li> ومن ثم، يتم التبرع بزوج واحد من الإلكترونات من أحد الأكسجين إلى ذرة الكربون لتشكيل رابطة مزدوجة. اعتماداً على ذرة الأكسجين التي تبرعت بالإلكترونات، يمكن أن يكون هناك هيكلان محتملان، يطلق عليهما هياكل الرنين. </li><li> NO + : بالنسبة لهذا الأيون، تتم إضافة ثمانية إلكترونات تكافؤ، لكن لا تحتوي أي من الذرة على مجموعة ثمانية. لا يمكن إضافة إلكترونات إضافية نظراً لاستهلاك إجمالي الإلكترونات بالفعل. في هذا السيناريو، يجب تحريك الإلكترونات لتشكيل روابط متعددة. تحتوي ذرة النيتروجين على زوجين منفردين من الإلكترونات وذرة الأكسجين بها زوج واحد. & nbsp ؛ <img alt="Figure6" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_6.svg" style="height: 55px;" /></li><li> لا يزال هذا لا ينتج مجموعة ثمانية، لذلك يجب تحريك زوج آخر لتكوين رابطة ثلاثية. & nbsp؛ <img alt = "Figure7" src = "https://www.jovecore.com/CDNSource/lessons/1063/1063Image_7.svg" style = "height: 50px؛" /> </li><li> في OF2، لا شيء يتغير لأن كل ذرة تحتوي بالفعل على مجموعة ثمانية. <img alt="Figure8" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_8.svg" style="height: 40px;" /></li></ul></li></ol>هذا النص مقتبس من <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chemistry 2e, Section 7.3: Lewis Symbols and Structures.</a>

لرسم هياكل لويس للجزيئات المعقدة والأيونات الجزيئية ، من المفيد اتباع إجراء خطوة بخطوة كما هو موضح:

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drawing lewis structures with valence electron counts

Drawing Lewis structures starts with counting valence electrons and building the atom arrangement step by step. Valence electrons are the outer-shell electrons that help show how atoms bond. For ions, a positive charge means one electron is removed for each plus sign, and a negative charge means one electron is added for each minus sign.

The next step is to sketch the skeleton structure. The least electronegative atom usually goes in the center, while hydrogen is almost never central. Fluorine also cannot be a central atom because it is the most electronegative element. Ions are written in brackets, with the charge shown outside the brackets.

After the skeleton is drawn, connect the atoms to the central atom with single bonds. Each single bond represents one electron pair. Then place the remaining electrons as lone pairs on the terminal atoms, except hydrogen, until those atoms have octets when possible. Any electrons still left go on the central atom.

If the central atom still does not have an octet, move lone pairs to make double or triple bonds. This can create resonance structures, which are different valid arrangements of electrons. For the formate ion, CHO2−, one oxygen can donate a lone pair to make a double bond with carbon. That gives two possible resonance forms.

The examples in the model show how the procedure works in real cases. SiH4 has 8 valence electrons total, so its structure does not need more changes once the single bonds are drawn. CHO2− has 18 valence electrons. NO+ has 10 valence electrons, and its atoms must be rearranged into a triple bond to satisfy octets as much as possible. OF2 has 20 valence electrons, and after the lone pairs are placed, its structure stays unchanged because each atom already has an octet.

Drawing Lewis Structures with Valence Electron Counts

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types of chemical bonds

Chemical Bonding and Bond Properties

Chemical bonding explains how atoms and ions join to form molecules and compounds. American chemist Gilbert N. Lewis pioneered chemical bonding theory with the Lewis model. This model helps explain why bonds form, why some are stronger than others, and why compounds have fixed formulas like H2O instead of H4O. Ionic bonding involves ions, which are atoms or molecules with an electrical charge. A cation is a positive ion formed when a neutral atom loses one or more electrons from its valence...

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Predicting Ionic Stability with Lattice Energy

Lattice energy helps explain why ionic compounds are stable. It measures the strength of attraction between positive and negative ions in a crystal. In one common definition, lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions, so it has a positive value. Another definition uses the opposite sign and describes the energy released when gaseous ions form a lattice. When checking data, it is important to confirm which convention is being used. In either...

lewis symbols and the octet rule

Valence Electrons, Lewis Symbols, and Bonding

Chemical bonds form when atoms or ions interact in ways that lower the potential energy of a molecule. Gilbert N. Lewis created a simple model to show bond formation using valence electrons, the electrons in the outer shell. In this model, electrons are shared or transferred to help atoms reach a stable electron configuration. The Lewis model helps describe two common bond types. An ionic bond forms when electrons move from a metal to a nonmetal. A covalent bond forms when two nonmetals share...

electronegativity

Electronegativity and Polar Bonds

Electronegativity helps explain why some covalent bonds are nonpolar and others are polar covalent. It is a measure of how strongly an atom attracts shared electrons toward itself. That pull controls how electron density is distributed in a bond. Linus Pauling proposed the first electronegativity values by studying bond energies in heteronuclear molecules such as hydrogen fluoride. He expected the bond energy of HF to be about the average of the bond energies of H2 and F2, or 296 kJ/mol. The...

ionic bonding and electron transfer

Ionic Compounds and Crystal Lattices

Ionic compounds form when positive and negative ions attract each other. A cation is a positive ion, and it forms when an atom loses one or more valence electrons. An anion is a negative ion, and it forms when an atom gains one or more valence electrons. Compounds made of these ions are called ionic compounds, or salts, and the force between them is an ionic bond. The properties of ionic solids help show how strong this attraction is. They usually form rigid, brittle crystals and have high...

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Lattice Energy Patterns in Ionic Solids

Lattice energy shows how strongly ions hold an ionic solid together. It is the energy required to separate one mole of a solid into its gaseous ions. For sodium chloride, this process requires 769 kJ per mole. The opposite convention may also be used, where lattice energy is the energy released when gaseous ions form the solid. In either case, a larger magnitude means a more stable ionic compound. For ionic crystals, lattice energy follows a pattern based on charge and ion size. The...

covalent bonding and lewis structures

Shared Electrons in Covalent Molecules

Covalent molecules form when atoms share electrons instead of transferring them. This bonding usually happens between nonmetal atoms with similar tendencies to attract electrons. Because the atoms pull on the electrons in a shared way, the bond is held together by mutual attraction. Covalent compounds often have physical properties that differ from ionic compounds. The attraction between neutral molecules is weaker than the attraction between charged ions. As a result, many covalent compounds...

bond polarity, dipole moment, and percent ionic character

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How Bond Polarity Shapes Molecular Dipoles

Bond polarity depends on electronegativity, the ability of an atom to attract shared electrons. The absolute value of the electronegativity difference, or ΔEN, gives a rough guide to bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is larger, the bond may be polar covalent or ionic. Examples show how this works in real compounds. H–H has a ΔEN of 0, so it is nonpolar. H–Cl has a ΔEN of 0.9, so it is polar covalent. Na–Cl has a ΔEN of 2.1, so it...

lewis structures of molecular compounds and polyatomic ions

Drawing Lewis structures starts with counting valence electrons and building the atom arrangement step by step. Valence electrons are the outer-shell electrons that help show how atoms bond. For ions, a positive charge means one electron is removed for each plus sign, and a negative charge means one electron is added for each minus sign. The next step is to sketch the skeleton structure. The least electronegative atom usually goes in the center, while hydrogen is almost never central. Fluorine...

resonance

Resonance Forms in Lewis Structures

Resonance forms in Lewis structures help explain molecules and ions that cannot be shown correctly with just one drawing. In nitrite ion, NO2−, the Lewis structure can be drawn in two ways. The N-O and N=O bonds switch positions in the two drawings, but the atoms stay in the same arrangement. If nitrite really had one single bond and one double bond, the two bond lengths should be different. A double bond is shorter and stronger than a single bond between the same atoms. Experiments show that...

formal charges

Lewis Structures with Formal Charge

Lewis structures can sometimes show more than one reasonable arrangement for a molecule or polyatomic ion. Formal charge helps chemists choose the best Lewis structure when that happens. It is a useful way to compare possible structures in chemistry. Formal charge is the hypothetical charge an atom would have if bonding electrons were shared evenly between the atoms. It can also be found by starting with the valence electrons of the neutral atom, subtracting the nonbonding electrons, and then...

exceptions to the octet rule

Octet Rule Exceptions in Lewis Structures

Some covalent molecules do not follow the octet rule in their Lewis structures. Their central atoms can fall into three groups: odd-electron molecules, electron-deficient molecules, and hypervalent molecules. These cases show why a Lewis structure may not always give every atom eight valence electrons. Odd-electron molecules have an odd number of valence electrons, so one electron remains unpaired. Molecules with an odd number of electrons are called radicals. Nitric oxide, NO, is one example,...

bond energies and bond lengths

How Bond Energy Predicts Reaction Heat

Bond energy helps explain why molecules stay together and how much heat a reaction may release or absorb. A covalent bond is stronger when more energy is needed to separate the atoms. The energy needed to break one specific covalent bond in one mole of gaseous molecules is called bond energy, or bond dissociation energy. For a diatomic molecule, bond energy is the standard enthalpy change for an endothermic reaction that breaks that bond. The H-H bond is a common example. Its bond energy is...

bonding in metals

Electron Sea Model in Metals

Electron sea model explains how metallic bonding holds metals together. In this model, most metal atoms do not have enough valence electrons to form ionic or covalent bonds in the usual way. Their valence electrons are held loosely because metals have low electronegativity, and their ionization energy is low, so electrons are removed easily. When a metal atom loses these outer electrons, it becomes a positively charged metal ion. The freed electrons form a delocalized electron cloud around...

key terms and concepts

Learn the step-by-step method for drawing Lewis structures, counting valence electrons, placing lone pairs, and handling ions, resonance, and multiple bonds.

Drawing Lewis Structures with Valence Electron Counts | JoVE K12

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Nuclear Binding Energy

The difference between the calculated and experimentally measured masses is known as the mass defect of the atom. In the case of helium-4, the mass defect ...

Radioactive Decay and Radiometric Dating

Radioactivity is a spontaneous disintegration of an unstable nuclide and is a random process, as all the nuclei in the sample do not decay simultaneously. ...

Nuclear Fission

Many heavier elements with smaller binding energies per nucleon can decompose into more stable elements that have intermediate mass numbers and larger ...

Nuclear Power

Controlled nuclear fission reactions are used to generate electricity. Any nuclear reactor that produces power via the fission of uranium or plutonium by ...

Nuclear Fusion

The process of converting very light nuclei into heavier nuclei is also accompanied by the conversion of mass into large amounts of energy, a process ...

Nuclear Transmutation

Nuclear transmutation is the conversion of one nuclide into another. It can occur by the radioactive decay of a nucleus, or the reaction of a nucleus with ...

Use models to compare nuclear reactions including alpha decay, beta decay and gamma decay; nuclear fusion and nuclear fission.

PS.Chm.1.2

Atomic Orbitals

An atomic orbital represents the three-dimensional regions in an atom where an electron has the highest probability to reside. The radial distribution ...

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

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

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

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

Analyze the structure of atoms and isotopes.

Drawing Lewis Structures with Valence Electron Counts

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