Lewis Structures & Valence Bond Theory路易斯结构、价键理论23页PPT
IBChemistry
The Octet Rule
• Atoms tend to gain, lose, or share e- until they are surrounded by 8 valence e- and are thus energetically stable.
• Exceptions do occur (and will be discussed later.)
2
Valence Electrons
• Valence electrons are the electrons in the outer shell that are on the same period or row. These are the electrons that are involved in bonding atoms together.
2. Check to see if charges are balanced.
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
– Ionic compounds are neutral, the positive charges balance out the negative charges.
– Ionic bonds form a lattice structure
20
Ionic Bonding
• Results as atoms lose or gain e- to achieve 8 electrons in the highest shell
Lewis_Structure_Rules
Rules for Lewis Structures1. Write the skeletal structure.-Hydrogen cannot be central.-Atoms that occur once are usually central.Example: CBr4-The least electronegative atom is central.-Use chemical intuition.2. Sum the total number of valence electrons of all atoms.Special - Polyatomic ions:Add one electron for each negative charge.Subtract one electron for each positive charge.In the following steps subtract electrons used from this total.3. Make single bonds between atoms in skeletal structure.4. Make octets on outer atoms by placing lone pairs. (Except hydrogen)5. Place remaining electrons on central atom(s) to make octets.6. If the central atom(s) do not have an octet then make double or triple bonds.7. Look for ambiguities in structure such:Hydrogen with more than 2 valence e-Atoms without typical number of bondsAtoms with more than an octet8. For ions place brackets: [ ] around the structure and the charge as a superscript following. For example: [Mg]2+9. Sometimes you can consider more than one possible structure. When this is the case consider:-Formal charges-Electronegativity-Are there multiple resonance structures?Examples:1. N22. CO23. Cl2COResonance StructuresIn some cases there are multiple Lewis structures that are equally plausible.Each of the Lewis structures is a resonance structure.The real molecule is a combination of each of these resonance structures and is referred to as a resonance hybrid.Resonance forms differ in placement of electrons but not atom arrangements. Example: O3Formal ChargeThe formal charge of an atom in a molecule is the difference between the number of valence electrons in the free atom and the number of valence electrons "owned" by the atom in the molecule.Formal Charge = C f = E v - (E u + 1/2E b)E v = number of valence electrons in free atom.E u = number of unshared electrons owned by the atom.E b = number of bonding electrons owned by the atom.Uses of formal charge:1. To determine the correct Lewis structure from multiple possibilities2. To determine when conventional Lewis structures based on the octet rule are incorrect.Another way to determine formal charge on each atom is to count valence electrons “owned” by an atom and compare to the number of valence electrons in the atom by itself. There are two type of counting valence electrons and each method is for a different purpose:1. Counting for the octet rule: the two electrons in a bonding pair both count when determining whether the octet rule is satisfied.2. Counting for formal charge: All non-bonding “lone pairs” count as “owned” when determining formal charge. However only one electron in a bonded pair is considered “owned” when counting for formal charge determination.Exceptions to the octet rule1. Radicals - odd electron speciesIn some cases there is no way to draw a Lewis structure in which all atoms have an octet. Example: NO, NO22. Incomplete octets - Compounds with central atoms with fewer tan eight valence electrons.Example: BF3, BeCl23. Expanded valence - Compounds with a central atom that has more than eight valence electrons. The central atom is a nonmetal of the third, fourth or fifth period.- Elements in the third period or higher can have more than eight valence electrons and so are capable of expanded valence.- Because of expanded valence some noble gases are capable of forming compounds.- Use the formal charge concept to see that expanded valence still allows for stable Lewis structures.Examples: PCl5 , SF6 , XeF4 , ClF3Note: some people use expanded valence for molecules such as SO42-Shapes of Molecules from Lewis Structures (VSEPR)When we speak of the shape, or molecular geometry, of a small molecule we mean the three dimensional arrangement of atoms about the central atom. For larger molecules, the same idea applies but the molecular geometry refers to the arrangement of atoms about any particular atom, rather than the geometry of the entire molecule.To consider the molecular geometry about an atom it must be connected to more than one other atom. Diatomic molecules must be linear by definition.With three or more atoms attached the assignment of a molecular geometry is not so clear. For example consider the two molecules ammonia, NH3, and formaldehyde,CH2O. Both have a central atom connected to three other atoms and yet different molecular geometries. Formaldehyde is described as having a trigonal planar molecular shape with all four atoms in the same plane. Ammonia is not a flat molecule. Instead the three hydrogen atoms are in a plane with the nitrogen atom above (or below depending on your perspective) the plane and the molecular geometry is described as being trigonal pyramidal.Even though the molecular geometry about an atom only describes the arrangement of connected atoms, according to VSEPR theory the lone pairs of electrons must be considered to correctly identify the shape.VSEPR (valence shell electron pair repulsion) theory basically states that all regions of electron density directed out and away from an atom repel each other. The lowest energy state is one in which all regions of electron density are as far apart from each other as possible.The geometry about an atom is found by looking at the Lewis structure. What are regions of electron density directed away from an atom in VSEPR theory? There are two types: bonds and lone pairs. The molecular geometry can be found by counting bonded atoms and lone pairs about a central atom.Sometimes VSEPR is explained as a consideration of the number of “electron pairs” about an atom with a “pair” being a bond or lone pair. Multiple bonds can be hard to deal with because a double bond for instance is two pairs but both pairs must be directed away from the central atom in the same direction (towards the other atom). So in terms of VSEPR a multiple bond is only one “pair”. This is why it is better to count “bonded atoms plus lone pairs” instead of “pairs”. A bonded atom counts once regardless of bond order (single, double or triple bond).。
美国普通化学笔记5
ChemistryChapter 8∙In 1864, English chemists john newlands noticed that when the known elements were arranged in order of atomic mass, every eighth element hadsimilar properties. Newlands referred to this peculiar relationship as thelaw of octaves. Howe ver, this “law” turned out to be inadequate forelements beyond calcium, and newland’s work was not accepted by thescientific community.∙Representative elements are the elements in groups 1A through 7A, all of which have incompletely filled s or p subshells of the highest principalquantum number. With the exception of helium, the noble gases (thegroup 8A elements) all have a completely filled p subshell. The transitionmetals are the elements in groups 1B and 3B through 8B, which haveincompletely filled d subshells or readily produce cations withincompletely filled d subshells (these metals are sometimes referred to asthe d-blok transition elements). The group 2B elements are Zn, Cd, andHg, which are neither representative elements nor transition metals. Thelanthanides and actinides are sometimes called f-block transition elementsbecause they have incompletely filled f subshells∙The outer electrons of an atom, which are those involved in chemical bonding are often called the valence electrons. Having the same number ofvalence electrons accounts for similarities in chemical behavior among theelements within each of these groups.∙Ions, or atoms and ions, that have the same number of electrons and hence the same ground-state electron configuration are said to be isoelectronic.∙Atomic radius of a metal is one-half the distance between the two-nuclei in two adjacent atoms. For elements that exist as diatomic molecules, theatomic radius is one-half the distance between the nuclei of the two atomsin a particular molecule.∙When looking at a periodic table:o The elements are increasing as in atomic radius as you go fromright to left, and from up to down. ****∙Ionic radius is the radius of a cation or an anion. Ionic radius affects the physical and chemical properties of an ionic compound.∙When a neutral atom is converted to an ion, we expect a change in size. If the atoms forms an anion, its size increases, because the nuclear chargeremains the same but the repulsion resulting from the additional electronenlarges the domain of the electron cloud. On the other hand, a cation issmaller than the neutral atom, because removing one or more electronsreduces electron-electron repulsion but the nuclear charge remains thesame, so the electron cloud shrinks.∙Focusing on isoelectronic cations, we see that the radii of tripostive ions (that is, ions that bear three positive charges) are smaller than those ofdipositive ions (that is, ions that bear two positive charges) which in turnare smaller than unipositve ions (that is, ions that bear one positive charge).∙Ionization energy – is the minimum energy required to remove an electron from a gaseous atom in its ground state. The magnitude of ionizationenergy is a measure of the effort required to force an atom to give up anelectron, or of how “tightly” the electron is held in the atom., the higherthe ionization energy the more difficult it is to remove the electron.∙For a many-electron atom, the amount of energy required to remove the first electron, from the atom in its ground state:o Energy + X(g) -> X+(g) + e-o Is called the first ionization energy (I1). In this equation Xrepresent a gaseous atom of any element and e- is an electron.Unlike an atom in the condensed liquid and solid phases, an atomis the gaseous phase is virtually uninfluenced by its neighbors.o Energy + X+(g) -> X2+(g) + e- Second ionizationo Energy + X2+(g) -> X3+(g) + e- Third Ionization∙When a electron is removed from a neutral atom, the repulsion among the remains electrons decreases. Because the nuclear charge remains constant,more energy is needed to remove another electron from the positivelycharged ions. Thus for the same element ionization energies alwaysincrease in this order:o I1<I2<I3<….∙Another property that greatly influences the chemical behavior of atoms is their ability to accept one or more electrons. This ability is called electronaffinity, which is the negative of the energy change that occurs when anelectron is accepted by an atom of an element in the gaseous stateo X(g) + e- -> X-(g) deltaH = -XXXkJ▪If delta h has a positive value (ie. 390 kj/mol) means thatthe process is exothermic▪If delta h has a negative value, that means that the processis endothermic∙Another trend in chemical behavior of the representative elements is the diagonal relationship. Diagonal relationship refers to similarities that existbetween pairs of elements in different groups and period of the periodictable. Specifically the first three members of the second period (Li, Be andB) exhibit many similarities to the elements located diagonally belowthem in the periodic table.If you would like to further understand this chapter, I suggested reading the summary. Or if you would like to learn more about the individual group elements, then I suggest reading the last few pages of this chapter.Chapter 9∙Lewis dot symbol – consists of the symbol of an element and one dot for each valence electron in an atom of the element.∙Covalent bond – a bond in which two electrons are shared by two atoms.Covalent compounds are compounds that contain only covalent bonds.∙Lone pairs – pairs of valence electrons that are not involved in covalent bond formation (ie. F2)∙Lewis structures is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms. Only valence electrons are shown in a Lewis structure.∙Octet rule – an atom other than hydrogen tends to from bonds until it is surrounded by eight valence electrons. In other words, a covalent b ond forms when there are not enough electrons for each individual atom tohave a complete octet. By sharing electrons in a covalent bond, theindividual atoms can complete their octets. The requirement for hydrogen is that it attains the electron configuration of helium, or a total of twoelectrons.o The octet rule works mainly for elements in the second period of the periodic table.∙Atoms can form different types of covalent bonds. In a single bond – two atoms are held together by one electron pair. Many compounds are held together by multiple bonds, that is, bonds formed when two atoms shre two or more pairs of electrons. If two atoms share two pairs of electrons, the covalent bond is called a double bond.∙ A triple bond arises when two atoms share three pairs of electrons, (N2) ∙Bond length – is defined as the distance between the nuclei of two covalently bonded atoms in a molecule.∙The bond HF is called a polar covalent bond, or simply a polar bond, because the electrons spend more time in the vicinity of one atom than the other. The HF bond and other polar bonds can be though of as beingintermediate between a (nonpolar) covalent bond, in which the sharing of electrons is exactly equal, and an ionic bond, in which the transfer of the electron(s) is nearly complete.∙ A property that helps us distinguish a nonpolar covalent bond from a polar covalent bond is electronegativity, the ability of an atom to attract toward itself the electrons in a chemical bond. Elements with highelectronegativity have a greater tendency to attract electrons than doelements with low electronegativity.o Electronegativity is related to electron affinity and ionization energy.o Electronegativity is a relative concept, mea ign tha t an element’s ectronegativity can be measured only in relation theelectronegativity of other elements.o Linus Pauling devised a method for calculating relativeelectronegativities of most elements.∙There is no sharp distinction between a polar bond and an ionic bond, but the following rule is helpful in distinguishing between them. An ionicbond forms when the electronegativity difference between the twobonding atoms is 2.0 more. This rule applies to most but not all ioniccompounds. Sometimes chemists use the quantity percent ionic characterto describe the nature of a bond. A purely ionic bond would have 100percent ionic character, although no such bond is known, whereas anonpolar or purely covalent bond has 0 percent ionic character.∙Electronegativity and electron affinity are related but different concepts.Both indicate the tendency of an atom to attract electrons. However,electron affinity refers to an isolated atom’s attraction for an additionalelectron, whereas electronegativity signifies the ability of an atom in achemical bond (with another atom) to attract the shared electron.Furthermore, the electron affinity is an experimentally measurablequantity, whereas electronegativity is an estimated number that cannot be measured.∙An atom’s formal charge is the electrical charge difference between the valence electrons in an isolated atom and the number of electrons assigned to an atom in a lewis structure.∙To assign the number of electrons on an atom in a lewis structure, we proceed as:o All the ato m’s nonbonding electrons are assigned to the atomo We break the bond(s) between the atom and other atom(s) and assign half of the bonding electrons to the atom∙When you write formal charges, these rules are helpful:o For molecules, the sum of the formal charges must add up to zero because they are electrically neutral species.o For cations, the sum of the formal charges must equal the positive chargeo For anions, the sum of the formal charges must equal the negative charge∙Keep in mind, that formal charges do not represent actual charge separation within the molecule.∙Resonance structure – one of two or more lewis structures for a single molecule that cannot be represented accurately by only one lewis structure.The double-headed arrow indicates that the structures shown areresonance structures.∙The term resonance itself means the use of two or more lewis structures to represent a particular molecule.∙Exceptions to the octet rule:o The incomplete octet:▪In some compounds the number of electrons surround thecentral atom in a stable molecule is fewer than eight.▪Elements in group 3A, particularly boron and aluminum,also tend to form compounds in which they are surroundedby fewer than eight electrons.∙ A resonance structure with a double bond betweenB and F can be drawn that satisfies the octet rule forB.▪The B-N bond is different from the covalent bondsdiscussed so far in the sense that both electrons arecontributed by the N atom. A covalent bond in which oneof the atoms donated both electrons is called a coordinatecovalent bond. Although the properties of a coordinatecovalent bond do not differ from those of a normal covalentbond (because all electrons are alike no matter what theirsource), the distinction is useful for keeping tack of valenceelectrons and assigning formal charges)o Odd-Electron Molecules▪Some molecules contain an odd number of electrons.Among them are nitric oxide (NO) and nitrogen dioxide(NO2)▪Because we need an even number of electrons for completepairing (to reach eight) the octet rule clearly cannot besatisfied for all the atoms in any molecule that has an oddnumber of electronso The expanded octet:▪In a number of compounds there are more than eightvalence electrons around an atom. These expanded octetsare needed only for atoms of elements in and beyond thethird period of the periodic table.∙ A measure of the stability of a molecule is its bond energy, which is the enthalpy change required to break a particular bond in 1 mole of gaseousmolecules. (bond energies in solids and liquids are affected byneighboring molecules.)∙In many cases, it is possible to predict the approximate enthalpy ofreaction by using the average bond energies. Because energy is alwaysrequired to break chemical bonds and chemical bond formation is alwaysaccompanied by a release of energy, we can estimate the enthalpy of areaction by counting the total number of bonds broken and formed in thereaction and recording all the corresponding energy changes. The enthalpyof reaction in the gas phase is given by:o deltaH o = sigma(BE(reactants)) – sigma(BE(products))o where be stands for average bond energy and sigma is thesummation signTo further understand Bond energies, and Lewis dot structures and resonance I suggest taking a deeper look into the textbook.。
国外教材有机化学的英文讲义知识交流
Atomic Structure
• protons, neutrons, and electrons • isotopes
C 12
6
C 14
6
Atomic Orbitals
2s orbital (spherical)
2p orbital
Electronic Configurations
➢ Aufbau principle: electrons occupy the orbitals with the lowest energy first ➢ Hund’s rule: electrons will occupy empty degenerated orbitals before pairing
Major Contributor?
H +H CN
H
H
major
H+ H
CN
H
H
minor, carbon does not have octet.
Resonance
• Delocalization of the negative charge on the conjugate base will stabilize the anion, so the substance is a stronger acid.
• More resonance structures usually mean greater stabilization.
H H
HCO H
Satisfy the octet rule!
Practice: write the Lewis structures for the below molecules
chemicalBonding化学键解析
• Metallic bond: attractive force holding pure metals together.
Figure 8.3: Ionic Bonding
+ and the negative pole -.
HyperChem
Drawing Lewis Structures
Follow Step by Step Method (See Ng Web-site)
1. Total all valence electrons. [Consider Charge]
Strengths of Covalent Bonds
Chemical Bonds, Lewis Symbols, and the Octet Rule Lewis Symbols
Chemical Bonds, Lewis Symbols, and the Octet Rule
The Octet Rule • All noble gases except He has an s2p6 configuration. • Octet rule: atoms tend to gain, lose, or share electrons
2. Write symbols for the atoms and guess skeleton structure [ define a central atom ].
3. Place a pair of electrons in each bond.
4. Complete octets of surrounding atoms. [ H = 2 only ]
有机化学结构与功能第一章
The structure of the molecule is related to the reactions that it can undergo
What is organic synthesis? Organic synthesis is the making of new organic molecules
有 机化
H
C Br
Cl F
1
学
Chapter 1 Structure and Bonding In Organic Molecules
1-1 1-2 1-3 1-4 1-5 1-6 1-7 1-8 1-9 The Scope of Organic Chemistry: An Overview Coulomb Forces: A Simplified View of Bonding Ionic and Covalent Bonds: The Octet Rule Electron-Dot Model of Bonding: Lewis Structures Resonance Forms Atomic Orbitals: A Quantum Mechanical Description Molecular Orbitals and Covalent Bonding Hybrid Orbitals: Bonding in Complex Molecules Structures and Formulas of Organic Molecules
14
Elements tend to form molecules in such a way
organic chemistry有机化学 第五版 LGWade JR 答案
-
. .-
+
+
(d)
..
•
H - N - C = C - C - N - H .. 1 1 1 1 1 H H H H H mmor
• •
(g)
:0: :0 : :0 : :0 : :0 : :0 : II II 1 II II 1 H - C - C- C - H----- H - C = C - C - H ----- H - C - C = C - H 1 1 1 H H H major mITIor major these two have equivalent energy and are major because the negative charge is on the more electronegative oxygen atom
(c) :O-N=O (d) (e)
I I
+
O=N-O: ..
+
. .
-
H-C=C-C-H
H H H
I
.. H-C-C=C-H
H H H
I
I
I
da
: 0:
II
w.
H - C = C - C - H .. .. H - C - C=C - H I I I I I I H H H H H H (f) Sulfur can have up to 12 electrons around it because it has d orbitals accessible . :0 : : 0: : 0: I I II O==S-O: .. :O - S - O: .. .. :O - S=O
·8 8
H - C == C - C leD I I 8 H H
分子结构
，共用电子对（A:B），表示为A-B; 同
样地，双键，两个共享电子对（A::B） 表示为A=B。 没有共用的电子（A :）称为孤对电子， 它们对分子的形状发挥重要作用。
Gilbert N. Lewis
2017化学竞赛
2017化学竞赛
（2） QF的计算公式：
QF=实际成键数-理论成键数 例：判断N2O和HN3的QF
2017化学竞赛
（3） Lewis结构式的稳定性 a. 在Lewis结构式中，QF应尽可能小，若共价分子中所有
原子的形式电荷均为零，则是最稳定的Lewis结构式 b. 两相邻原子之间的形式电荷尽可能避免同号
XeO4：8 + 2 4 = 16
2017化学竞赛 问题2：有些富电子化合物为什么可以不修正呢？
例如：OPCl3
当配位原子数小于或等于键数时，可以不修正，因为连 接配位原子的单键已够了。但中心原子周围的配位原子
数目超过4，必须要修正
8e
10e
2017化学竞赛 例题：假设SNF、NSF和SFN都存在，试比较他们的 稳定性顺序。
AX4E0
CH4、NH4+、 四面体 SO42-
4
AX3E1
三角锥
NH3、PCl3
AX2E2
V形
H2O、H2S
2017化学竞赛
m+n AXmEn 分子构型 举例
AX5E0
三角双锥
PCl5
5
AX4E1
变形四面 体（跷跷 板）
SF4
2017化学竞赛
m+n AXmEn 分子构型 举例
AX3E2
有机化学英文课件chapter1
•• ••
••
Na+ F• • ••
• in forming Na+F-, the single 3s electron from Na is transferred to the partially filled valence shell of F
N a(1s22s 22p63s1) + F(1s 22s22p5)
1-7
Electron Configuration of Atoms
The pairing of electron spins
1-8
Electron Configuration of Atoms
Table 1.3 The Ground-State Electron Configuration of Elements 1-18
1-6
Electron Configuration of Atoms
Aufbau Principle:
• orbitals fill in order of increasing energy from lowest energy to highest energy
Pauli Exclusion Principle:
• an atom that gains electrons becomes an anion • an atom that loses electrons becomes a cation • the attraction of anions and cations leads to the
Chemical Bonding
Lewis Structures
Elements with 2 valence electrons are called the Alkaline Earth metals (group 2) (Label)
Lewis Structures
Elements with 7 valence electrons are called the Halogens (group 7) (Label)
+1
-2 0
Nitrogen? -3
Charges
Charges
A few exceptions!
Metals are always a positive charge!! Non metals are always negative!! Metalloids can go either way (you are not responsible for choosing – I will tell you) According to rule of 4’s….its a +3 charge But since it’s a nonmetal, we use -5!
Ionic Bonds
Usually formed by members of the Alkali group (ones with +1 electron) Electrons are donated to another molecule Between elements from opposite sides of the chart Forms crystals (salts) & most dissolve in water
Formal charge Draw a Lewis dot structure for ClO Which
•
Polyatomic ions Lewis dot structures do not show a molecule’s shape.
18
Readings for Next Class
10.7
19
Resonance
How many Lewis structures can you draw for ozone?
What is the average bond order for an O-O bond in ozone?
-½ ••
O O O •• ••
••+ •• -½
If more than one valid Lewis structure can be drawn for a molecule, the true electron distribution is a hybrid of the possible distributions. This is called resonance.
1
Formal charge
The sum of formal charges on individual atoms is the total charge of the molecule or ion
Examples
◦ Ammonium ion (NH4)+
2
EXAMPLE 10-7
Using the Formal Charge in Writing Lewis Structures. Write the most plausible Lewis structure of nitrosyl chloride, NOCl, one of the oxidizing agents present in aqua regia.
Lewis Structure
Shared Electrons are considered to contribute to the electron requirements of both atoms involved; thus, the electron pairs shared by H and O in the water molecule are counted toward both the 8-electron requirement of oxygen and 2-electron requirement of hydrogen. Double Bonds: 4 electrons contained(Carbon Dioxide);
Examples：
Cyanic acid, Isocyanic acid and fulminic acid
Applications of Lewis structure（Lewis structure的 应用）.
1. Judge the bonding stability of atoms in molecules.（判断分子中原子连接方式的 稳定性）。 N2O:
计算分子中原子之间的键级进而可以判断键能的大小和键长的长短
Yuwen Cui
CONTENTS
How to show Lewis structure（Lewis structure的表 示方法）? Determination of the numbers of bonds in covalent molecules（共价分子中成键数的确定）. Drawing Lewis structure（Lewis structure的画法）. Applications of Lewis structure（Lewis structure的 应用）. Special Cases（特殊情况）
4-分子结构-化学原理 PPT课件
共价键类型
Multiple Bonds
double bonds triple bonds
配位健（Coordination covalent bond）
CO
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键参数 (Bond Parameters)
键长(bond length) 键角(bond angle) 键能(bond energy)
O HH
S HH
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6.2 Covalent Bonding
价键理论
Valence bond theory (VB)
杂化轨道理论 Hybrid orbital theory
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离子键 （ionic bond）
Walther Kossel 1916
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1s
x
1s
3px
(a)
x 3px
(b)
x 3px
(c) 2019/10/17 24
共价键类型 σ键
s-s s-px px-px
π键
py-py pz-pz
x
x
轨道的重叠部分沿键轴呈圆 柱形对称分布
键轴（internuclear axis）
轨道的重叠部分垂直于键 轴并呈镜面反对称分布
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键的极性
(Polarity of bond)
电负性
分子的空间构型
极性共价键
(Polar covalent bond)
非极性共价键
(nonpolar covalent bond)
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e- poor e- rich
d+ d-
lewis-structures
[
]
• A resonance structure can be drawn for any molecule in which a double bond can be formed from two or more identical choices • Resonance structures can be drawn 2 – – ways… O 1 H C H C O – O 2 H C • Resonance implies that the bond flips back and forth. Really, it lies between extremes • Read 7.7 (pg. 242). Do PE 8. O
Placing electrons around atoms
Compound 4) Octet for O Cl O peripheral 16 - 12 = 4 atoms 1) Skeletal 5) RemainO Cl O O Cl O –s on Structure ing e center atom 4 - 4 = 0 2) Count 7x1 + 6x2 6) Create electrons multiple +1 = 20 No need bonds? 3) Electron O Cl O Final O Cl O pairs in structure or 20 - 4 = 16 O Cl O bonds ClO2–
2 peripheral Lewis’ bonded to a central Lewis
Skeletal structures
• Because there are exceptions to the octet rule, we need a set of rules to determine how many electrons surround atoms • The first step is to determine how the atoms are bonded in a molecule • Generally, if there is only one of one element and multiple copies of another element, the unique element is central • Commonly, H is peripheral, bonded to O • Read 7.6 (pg. 236) up to PE5. Do PE5.
思维导图 8
Chemical Bonds bond length The distance between two bounded atoms (where the bonding energy is minimal) The zero point of energy is defined with the atoms at infinite separa=on Electronega=vity The different affini=es of atoms for the electrons in a bond in a bond are described by a property called electronega=vity. When two atoms with very different electronega=vi=es interact, electron transfer can occur to form the ions that make up an ionic substance. Intermediate cases give polar covalent bonds with unequal electron sharing.Bond Polarity and Dipole Moments A molecule such as HF that has a center of posi=ve charge and a center of nega=ve charge is said to be dipolar, or to have a dipole moment When a nonmetal and a representa=ve-group metal react to form a binary ionic compound, the ions form so that the valence electron configura=on of the nonmetal achieves the electron configura=on of the next noble gas atom and the valence orbitals of the metal are emp=ed. In this way both ions achieve noble gas electron configura=ons.Energy Effects in Binary Ionic Compounds aPce energy: —the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid breaking bonds absorb energy while form a new bond releas energy Covalent Bond Energies and Chemical Reac=ons Hp – Hr = H The Localized Electron Bonding Model Descrip=on of the valence electron arrangement in the molecule using Lewis structures (will be discussed in the next sec=on). 2. Predic=on of the geometry of the molecule using the valence shell electron-pairrepulsion (VSEPR) model (will be discussed in Sec=on 8.13). 3. Descrip=on of the type of atomic orbitals used by the atoms to share electrons or hold lone pairs (will be discussed in Chapter 9).Lewis Structures The Lewis structure of a molecule shows how the valence electrons are arranged among the atoms in the molecule.Steps for Wri=ng Lewis Structures 1.Sum the valence electrons from all the atoms. Do not worry about keeping track of which electrons come from which atoms. It is the total number of electrons that is important. e a pair of electrons to form a bond between each pair of bound atoms. 3.Arrange the remaining electrons to sa=sfy the duet rule for hydrogen and the octet rule for the second-row elements. Lewis Structures: Comments About the Octet Rule ›T h e s e c o n d -r o w elements C, N, O, and F should always be assumed to obey the octet rule. ›The second-row elements B and Be oben have fewer than eight electrons around them in their compounds. These electron-deficient compounds are very reac=ve. ›The second-row elements never exceed the octet rule, since their valence orbitals (2s and 2p) can accommodate only eight electrons. ›Third-row and heavier elements oben sa=sfy the octet rule but can exceed the octet rule by using their empty valence d orbitals. ›When wri=ng the Lewis structure for a molecule, sa=sfy the octet rule for the atoms first. If electrons remain aber the octet rule has been sa=sfied, then place them on the elements having available d orbitals (elements in Period 3 or beyond)Resonance Resonance is invoked when more than one valid Lewis structure can be wricen for a par=cular molecule. The resul=ng electron structure of the molecule is given by the average of these resonance structures Atoms in molecules try to achieve formal charges as close to zero as possible. 2. Any nega=ve formal charges are expected to reside on the most electronega=ve atoms.Lone pair electrons belong en=rely to the atom in ques=on. 2. Shared electrons are divided equally between the two sharing atoms The number of valence electrons on the free neutral atom (which has zero net charge because the number of electrons equals the number of protons) 2. The number of valence electrons “belonging” to the atom in a molecule Molecular Structure: The VSEPR Model The main postulate of this model is that the structure around a given atom is determined principally by minimizing electron-pair repulsions Steps to Apply the VSEPR Model 1.D r a w t h e L e w i s structure for the molecule. 2.Count the electron pairs and arrange them in the way that minimizes repulsion (that is, put the pairs as far apart as possible). 3.D e t e r m i n e t h e posi=ons of the atoms from the way the electron pairs are shared. 4.Determine the name of the molecular structure from the posi=ons of the atoms.。
高等有机化学PPT
The ionic bonding force arises from the electrostatic attraction between ions of opposite charge.
The covalent bonding force arises from sharing of electron pairs between atoms.
O
CH3 C NH2
O = CH3 C + NH2
more stable major contributor
less stable minor contributor
Draw resonance structures for the following species. If the structures are not equivalent, indicate which would be the major contributor.
O3
OOO
OOO
Rules for Resonance Structures:
3. All of the structures must be proper Lewis structures.
O3
OOO
OOO
Rules for Resonance Structures:
4. The energy of the actual molecule is lower than the energy that might be estimated for any contributed structure.
高等有机化学 反应、机理与结构（原著 第五版 修订），[美]史密斯（Smith M.B） 马奇（March J），化学工业出版社
covalent-bonding-AP化学电子教案
Lewis structure VSEPR model Hybridization
Molecular orbital model
Localized Electron Model
• Simple model, easily applied. • A molecule is composed of atoms that
are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
• Three Parts
1) Valence electrons using Lewis structures 2) Prediction of geometry using VSEPR 3) Description of the types of orbitals
• In a nonpolar bond, the bonding electrons are shared equally:
H2, Cl2:
• In a polar bond, electrons are shared unequally because of the difference in electronegativity
H Cl
2.1 3.0
= 0.9
Electronegativity Difference
Bond Type
0 - 0.5
Nonpolar
0.5 - 1.7
Polar
1.7
Ionic
Localized Electron Model
化学结构知识点英语总结
化学结构知识点英语总结Atomic StructureThe foundation of chemical structures lies in the understanding of atomic structure. Atoms are the basic building blocks of matter and consist of three major subatomic particles: protons, neutrons, and electrons. Protons are positively charged particles found in the nucleus of an atom, while neutrons are neutral particles also located in the nucleus. Electrons, which have a negative charge, orbit the nucleus in specific energy levels called atomic orbitals.The number of protons in an atom determines its identity as an element, and this is known as the atomic number. For example, an atom with six protons is the element carbon, and an atom with eight protons is the element oxygen. The total number of protons and neutrons in an atom is referred to as the atomic mass, which is often a whole number for most elements. Isotopes are atoms of the same element that have different numbers of neutrons and therefore different atomic masses.BondingChemical bonds are the forces that hold atoms together in a compound. There are three main types of chemical bonds: ionic, covalent, and metallic. Ionic bonding occurs when atoms transfer electrons to achieve a full outer shell and form positively and negatively charged ions. The attraction between these ions creates an ionic bond. Covalent bonding occurs when atoms share electrons to complete their outer shells, resulting in the formation of a molecule. Metallic bonding is found in metals, where the outer electrons are delocalized and free to move throughout the structure, creating a "sea of electrons" that holds the metal atoms together.Molecular GeometryMolecular geometry refers to the three-dimensional arrangement of atoms in a molecule. The shape of a molecule is determined by the arrangement of its atoms and the type of bonds between them. This has important implications for the physical and chemical properties of the molecule. For example, the shape of a molecule can affect its polarity, which in turn influences its solubility, reactivity, and biological activity.VSEPR TheoryThe Valence Shell Electron Pair Repulsion (VSEPR) theory is a useful tool for predicting the shape of molecules. According to this theory, electron pairs in the valence shell of an atom repel each other, and the molecule adopts a shape that minimizes this repulsion. For example, molecules with two electron pairs around the central atom have a linear shape, while molecules with four electron pairs adopt a tetrahedral shape. The VSEPR theory is also helpful in understanding the concept of molecular polarity and the behavior of molecules in chemical reactions.Lewis StructuresLewis structures are diagrams that show the bonding and non-bonding electron pairs in a molecule. These structures are a helpful way to visualize the arrangement of atoms and electrons in a compound. By drawing Lewis structures, chemists can predict the shape and bonding in a molecule, as well as the reactivity of the compound. For example, a molecule with a double bond between two atoms will have a different shape and reactivity than a molecule with a single bond between the same atoms.ResonanceIn some cases, a molecule can have multiple valid Lewis structures, and the actual structure of the molecule is a hybrid of these possibilities. This phenomenon is called resonance. Resonance occurs when electrons can be delocalized and are not confined to a single bond or atom. This concept is particularly important in understanding the properties of aromatic compounds and the stability of certain molecules.HybridizationHybridization refers to the mixing of atomic orbitals to form new hybrid orbitals in a molecule. These hybrid orbitals have different shapes and energies than the original atomic orbitals and are used to explain the geometry and bonding in molecules. For example, in methane (CH4), the carbon atom undergoes sp^3 hybridization, resulting in the formation of four equivalent sp^3 hybrid orbitals that point towards the corners of a tetrahedron.Molecular Orbital TheoryThe molecular orbital theory is another approach to understanding the bonding in molecules. According to this theory, electrons in a molecule are delocalized and are described by molecular orbitals, which extend over the entire molecule rather than being associated with a specific bond or atom. The combination of atomic orbitals gives rise to molecular orbitals, which can be bonding, antibonding, or nonbonding in nature. This theory provides a more comprehensive explanation of the electronic structure and properties of molecules than the simpler valence bond theory.ConclusionIn summary, chemical structures are the foundation of chemistry, and understanding them is crucial for making sense of the properties and behavior of substances. The concepts of atomic structure, bonding, and molecular geometry provide valuable insights into the nature of matter and the interactions between atoms and molecules. By applying these principles, chemists can design new materials, drugs, and technologies, and gain a deeper understanding of the world around us.。
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Consideration of effective nulear charge:
Ψ1s = N(z/a0)3/2exp(-zr/a0), z = 1.17 ― curve c
The z-modification gives good bond length, but energy is still substantially off.
What is the difference between f. c. & ox. no
Formal charge is for the comparison of the elctron change between the bonded and the non-bonded atomic molds. Oxidation number is for the comparison of the electron change between the oxidized and the reduced molds. Formal charge and oxidation numbers are the same for the ionic compounds.
Disadvantage of Lewis Structure
1 Octet rule is no good for d orbital or d electrons involved species – use 18 electron rule
2 no good for electron delocalization – use resonance 3 Octet rule is no good for odd electron species 4 no information for geometry- use VSEPR 5 It does not explain the bonding character – use
If g is set as 0.1, which gives 99% (1/1.01) of 1s and
1% (0.01/1.01) of 2pz. Note that 1% of contribution of pz accounts for about 5% energy stabilization, mainly due to the enhancement of orbital overlap efficiency.
Heitler –London functions: Attractive ΨII+ = ψ1sa(1)ψ1sb (2) + ψ1sa(2)ψ1sb (1) ― curve b
Repulsive ΨII− = ψ1sa(1)ψ1sb (2) − ψ1sa(2)ψ1sb (1) ― curve g
ΨII– = ψ1sa(1) ψ1sb(2) – ψ1sa(2) ψ1sb(1) ΨI = ψ1sa(1) ψ1sb (2)
ΨII– = ψ1sa(1) ψ1sb(2) – ψ1sa(2) ψ1sb(1)
ΨI = ψ1sa(1) ψ1sb (2)
z = 1.17
ΨII+ = ψ1sa(1) ψ1sb(2) + ψ1sa(2) ψ1sb(1)
experimental
Consideration of polarizability or hybridization:
VBT or MOT
Valence Bond Theory (VBT)
Atomic orbitals that involve in chemical bonding are the major concern. ψiψj ≠ 0, wherein ψi and ψj represent the orbital wave functions of different atoms.
A lone pair is that the electron pair is localized in an atom.
Octot (bar) formula 2 valence eletrons & net charge 3 multiple bond 4 dative bond (arrow to the acceptor) 5 formal charge 6 What is the difference between f. c. & ox. 6 gnoenerally good thru the 3rd period
Lewis Structure
Chem istry
Localized Electron Bonding Model
A molecule is composed of atoms that are bound together by sharing electron pairs.
A bonding pair is of which two electrons are localized between two adjacent nuclei.
ψa = N (1sa + g 2pza) ψb = N (1sb + g 2pzb), assuming z is the nuclear axis.
Normalization condition gives N = 1/(1 + g2)1/2 Contribution of 1sa to ψa (or 1sb to ψb ) is 1/(1 + g2) Contribution of 2pza to ψa (or 2pzb to ψb ) is g2/(1 + g2)
VBT for H2
Consideration of simple orbital coulomb overlap: ΨI = ψ1sa(1)ψ1sb (2) ― curve a
ΨI = ψ1sa(1) ψ1sb (2) experimental
Consideration of indistinguishability of the electrons: