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Transcript
Chapter 1 Remembering General Chemistry: Electronic Structure and Bonding Paula Yurkanis Bruice University of California, Santa Barbara What makes carbon so special? • Atoms to the left of carbon give up electrons. • Atoms to the right of carbon accept electrons. • Carbon shares electrons. 1.1 The Structure of an Atom Atoms have an internal structure consisting of one or more subatomic particles: protons, neutrons, and electrons. proton positive charge mass = 1.673 x 10-27 kg neutron no charge mass = 1.675 x 10-27 kg electron negative charge mass = 9.109 x 10-31 kg 1.1 The Structure of an Atom Most of the mass of an atom is concentrated in the nucleus. The nucleus contains one or more positively charged protons, and one or more neutrons with no electrical charge. 1.1 The Structure of an Atom One or more negatively charged electrons are in constant motion somewhere outside the nucleus. The number of electrons is equal to the number of protons; the atom has no overall electrical charge. 1.1 The Structure of an Atom An atom is mostly free space because the volume of the nucleus and the electrons outside the nucleus are extremely small compared to the overall volume of the atom. 1.1 The Structure of an Atom Most of the mass of an atom is in its nucleus. Most of the volume of an atom is occupied by its electrons, and that is where our focus will be because it is the electrons that form chemical bonds. 1.1 The Structure of an Atom 1.1 The Structure of an Atom Atomic number • equals the number of protons in its nucleus, and is also the number of electrons that surround the nucleus. Mass number • the sum of its protons and neutrons. Isotopes • have the same atomic number, but different mass numbers because they have different numbers of neutrons (i.e., 12C, 13 C, and 14C). Atomic weight (Atomic mass) • of a naturally occurring element is the average mass of its atoms. Molecular weight • of a compound is the sum of the atomic weights of all the atoms in the molecule. The Structure of an Atom Protons are positively charged. Neutrons have no charge. Electrons are negatively charged. Atomic number = # of protons Atomic number of carbon = 6 Neutral carbon has six protons and six electrons. Isotopes 1.2 How the Electrons in an Atom are Distributed The electrons in an atom can be thought of as occupying a set of shells that surround the nucleus. The first shell is the smallest and the one closest to the nucleus. Atomic orbitals • The probability distribution about one atomic nucleus (each shell contains subshells known as atomic orbitals). The Distribution of Electrons in an Atom • The first shell is closest to the nucleus. • The closer the atomic orbital is to the nucleus, the lower its energy. • Within a shell, s is lower in energy than p. Atomic Orbitals s orbital p orbitals px py pz -Aufbau principle: An electron goes into the available atomic orbital with the lowest energy. -Pauli exclusion principle: No more than two electrons can occupy each atomic orbital. -Hund’s rule: An electron goes into an empty degenerate orbital rather than pairing up. Carbon Atomic Number = 6 2p 2s 1s Oxygen Atomic Number = 8 2p 2s 1s Electronic configuration Core electrons • Electrons in inner shells (those below the outermost shell) Valence electrons • Electrons in the outermost shell. • Elements in the same column of the periodic table have the same number of valence electrons, thus having similar chemical properties. The chemical behavior of an element depends on its electronic configuration. 1.3 Ionic and Covalent Bonds Octet rule (proposed by G. N. Lewis) • An atom is most stable if its outer shell is either filled or contains eight electrons and it has no electrons of higher energy. Atoms on the Left Side of the Periodic Table Easily Lose an Electron Atoms on the Right Side of the Periodic Table Easily Gain an Electron A Hydrogen Atom can Easily Lose or Gain an Electron Chemical Bonding Chemical Bonds • The forces holding atoms together in compounds. Only valence electrons are used in bonding. Lewis Dot Representation of Atoms • Dots around the chemical symbol of an atom represent the valence electrons. Examples Atom Electronic Structure Electronic Configuration Lewis Dot Structure 3p 3s Boron 2 2 1 1s 2s 2p 2p 2s B 1s 3p 3s Phosphorus [Ne] 3s23p3 2p 2s 1s P 內殼層電子也能形成化學鍵 鍵結規則出現例外,教科書即將改寫。撰文/莫斯柯維茨(Clara Moskowitz) 翻譯/甘錫安 我們在高中化學裡學到,電子在不同的能階環繞原子核運行。能量較低 的能階稱為「內電子殼層」,能量最高的能階構成外殼層;原子必須共 用或交換最外殼層的電子,才能形成化學鍵。不過,有位化學家可能發 現我們熟悉的鍵結規則出現了例外。在極高的壓力下,原子內殼層的電 子似乎也能形成化學鍵。 美國加州大學聖巴巴拉分校及中國北京計算科學研究中心的化學家苗茂 生表示:「這種現象打破了內殼層電子不會參與反應、不會進入化學領 域的學說。」根據苗茂生的計算,在極高的壓力下,銫原子和氟原子可 能形成具有內殼層鍵結的特異分子。 這兩種原子通常只會形成比較簡單的鍵結。銫是鹼金屬,最外殼層只有 一個孤立的價電子。另一方面,氟是鹵素,最外殼層只差一個電子就完 全填滿,正好可以接收銫多出來的一個電子。 但是苗茂生卻發現,在高壓下,有兩種分子的形成和銫的內殼層電子有 關。為了形成三氟化銫(CsF3),銫原子必須提供一個價電子以及兩個 內殼層電子,和三個氟原子共用。若有四個內殼層電子參與,則可形成 五氟化銫(CsF5)。苗茂生表示:「這種分子的形狀類似海星,相當漂 亮。」他在《自然.化學》發表這項發現。1981年諾貝爾化學獎得主、 美國康乃爾大學榮譽教授霍夫曼(Roald Hoffmann)表示,這些分子的 形狀和形成機率都「相當令人驚奇」。 【摘自科學人2014年第145期3月號】 An Ionic Compound is Formed by the Attraction Between Ions of Opposite Charge The Ionic Bond Transfer of Electrons from One Atom to Another Example: • Sodium chloride Na Cl + 1 e Na + 1e Cl Covalent Bonds are Formed by Sharing Electrons Nonpolar covalent bond = bonded atoms are the same Polar covalent bond = bonded atoms are different The Covalent Bond Some atoms do not transfer electrons from one atom to another to form ions. Instead they form a chemical bond by sharing pairs of electrons between them. A covalent bond consists of a pair of electrons shared between two atoms. How Many Bonds Does an Atom Form? Bond Polarity Depends on Electronegativity Differences Electronegativity (EN) Electronegativity is a measure of the ability of an atom to attract a bonding pair of electrons toward itself. Electronegativity differences in covalently bonded atoms result in polar covalent bonds. Polar Covalent Bonds The polarity of a bond is determined by the difference in electronegativity values. If the electronegativities are the same the bond is nonpolar and the electrons are shared equally. If the atoms have greatly differing electronegativities the bond is very polar. Polar Covalent Bonds Dipole Moments the greater the difference in electronegativity, the greater the dipole moment, and the more polar the bond Electrostatic Potential Maps 1.4 How The Structure of A Compound Is Represented Nonbonding or lone-pair electrons: • Electrons not used in bonding. Formal charge: • The difference between the number of valence electrons an atom has when it is not bonded to any other atoms and the number of electrons it “owns” when it is bonded. Formal charge = number of valence electrons – (number of lone-pair electrons + ½ number of bonding electrons) 1.4 Representation of Structure Neutral Carbon Forms Four Bonds if carbon does not form four bonds, it has a charge (or it is a radical) Formal Charge: +1 –1 0 Neutral Nitrogen Forms Three Bonds Nitrogen has one lone pair. If nitrogen does not form three bonds, it is charged. Neutral Oxygen Forms Two Bonds Oxygen has two lone pairs. If oxygen does not form two bonds, it is charged. Hydrogen and the Halogens Form One Bond A halogen has three lone pairs. if hydrogen or halogen does not form one bond, it has a charge (or it is a radical) The Number of Bonds and Lone Pairs • Halogen = 3 lone pairs • Oxygen = 2 lone pairs • Nitrogen = 1 lone pair Drawing Lewis Dot Structures Determine the total number of valence electrons supplied by all of the atoms in the structure. Write down the skeletal arrangement of the atoms and connect them with a single covalent bond (two electrons). • “Normal” bonding: H: one covalent bond halogens: one covalent bond O: two covalent bonds N: three covalent bonds C: four covalent bonds Drawing Lewis Dot Structures Distribute pairs of electrons (pairs of dots) around each atom (except H) to give each atom eight electrons around it (the octet rule). If there are not enough electrons to give these atoms eight electrons, change single bonds between atoms to double or triple bonds by shifting non-bonded pairs of electrons as needed. Assign formal charges and evaluate the structure. Lewis Dot Structures How to Draw a Lewis Structure NO3– Determine the total number of valence electrons (5 + 6 + 6 + 6 = 23). Because they are negatively charged, add another electron = 24. Avoid O–O bonds. Check for formal charges. 1.5 Atomic Orbitals s orbital p orbitals 1.6 How Atoms form Covalent Bonds Sigma (σ) bond The covalent bond that is formed when the two orbitals overlap 1.6 How Atoms form Covalent Bonds 1.7 How Single Bonds Are Formed in Organic Compounds Representations of Methane In Order to Form Four Bonds, Carbon Must Promote an Electron Four Orbitals are Mixed to Form Four Hybrid Orbitals An sp3 orbital has a large lobe and a small lobe. Hybridization However, mixing the carbon 2s orbital and the 3 carbon 2p orbitals gives 4 sp3 hybrid orbitals: + 2s + 2p + 2p + 2p 2p 2s sp 3 + + sp 3 sp 3 sp 3 sp3 The Carbon in Methane is sp3 Carbon is tetrahedral. The tetrahedral bond angle is 109.5°. The Bonding in Ethane Representations of Ethane 1.8 How A Double Bond Is Formed: The Bond in Ethene An sp2 Carbon Has Three sp2 Orbitals and One p Orbital The Carbons in Ethene are sp2 Representations of Ethene 1.9 How A Triple Bond Is Formed: The Bonds in Ethyne The Two sp Orbitals Point in Opposite Directions The Two p Orbitals are Perpendicular The Carbons in Ethyne are sp Representations of Ethyne 1.10 The Bonds in the Methyl Cation, the Methyl Radical, and the Methyl Anion The Methyl Cation The Methyl Radical The Methyl Anion 1.11 The Bonds in Ammonia and in the Ammonium Ion Nitrogen Has Three Unpaired Valence Electrons and Forms Three Bonds in NH3 Nitrogen does not have to promote an electron. The Bonds in Ammonia (NH3) The Ammonium Ion (+NH4) 1.12 The Bonds in Water Oxygen Has Two Unpaired Valence Electrons and Forms Two Bonds in H2O Oxygen does not have to promote an electron. The Bonds in Water (H2O) 1.13 The Bond in a Hydrogen Halide Overlap of an s Orbital with an sp3 Orbital The Length and Strength of a Hydrogen Halide Bond 1.14 Summary: Hybridization, Bond lengths, Bond strengths, and Bond angles Summary of Hybridization The orbitals used in bond formation determine the bond angle. Single Bond: 1 σ Double bond: 1 σ + 1 π Triple Bond: 1 σ + 2 π Hybridization of C, N, and O Bond Strength and Bond Length The shorter the bond, the stronger it is. A π Bond is Weaker Than a σ Bond Hybridization, Bond Angle, Bond Length, Bond Strength 1.15 The Dipole Moments of Molecules 1.15 The Dipole Moments of Molecules