Chapter Seven Atomic Structure atoms neutrons protons positive

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 Chapter Seven Atomic Structure § atoms neutrons protons (positive charge ) electrons (negative

Chapter Seven Atomic Structure § atoms neutrons protons (positive charge ) electrons (negative charge)

7 -1 Changing Ideas about Atomic Structure 7 -2 The Quantum Mechanical Description of

7 -1 Changing Ideas about Atomic Structure 7 -2 The Quantum Mechanical Description of Electron in Hydrogen Atoms 7 -3 Electron Configuration of Manyelectron Atoms 7 -4 The Periodic Table and Periodic Law

7 -1. 1 The Bohr theory of Hydrogen Atom § 1805 dolton proposed atom

7 -1. 1 The Bohr theory of Hydrogen Atom § 1805 dolton proposed atom theory, proved exist of atom § 1900 electron were discovered § 1911 Ruthrford proposed the atomic nucleus by α-ray scatting § 1931 neutron were discovered

Ruthrford’s nuclear model § Figure 7 -1: In classical theory: 1. atoms constructed are

Ruthrford’s nuclear model § Figure 7 -1: In classical theory: 1. atoms constructed are not stable; § 2. the electron would quickly spiral into the nucleus. § 3. Not is the line spectra of atoms

Continuous spectrum

Continuous spectrum

Atomic Line Spectra Na

Atomic Line Spectra Na

(H、He、Li、Na、Ba、Hg、Ne light emission)

(H、He、Li、Na、Ba、Hg、Ne light emission)

In 1913, Niels Bohr(1885 -1962), founded Bohr theory by using the work of Planck

In 1913, Niels Bohr(1885 -1962), founded Bohr theory by using the work of Planck and Einstein

Quantum of concept no continuum emission § Atom a copy of energy absord Least

Quantum of concept no continuum emission § Atom a copy of energy absord Least unit quantum

l. The Photoelectric Effect Einstein used the quantum theory to explain the photoelectric effect

l. The Photoelectric Effect Einstein used the quantum theory to explain the photoelectric effect : Each energy packet called photon, is a quantum of energy E=h v Physicist Albert Einstein h Planck’s constant (1879 -1955) = 6. 623× 10 -34 J s.

E = hv = (波粒二象性) Photons of high frequency radiation have high energies, whereas

E = hv = (波粒二象性) Photons of high frequency radiation have high energies, whereas photons of lower frequency radiation have lower energy.

7 -1. 1 The Bohr theory of Hydrogen Atom § Bohr set down the

7 -1. 1 The Bohr theory of Hydrogen Atom § Bohr set down the following § two postulates to account for: (1) the stability of the hydrogen atom (that the atom exists and its electron does not continuously radiate energy and spiral into the nucleus) (2) the line spectrum of the atom.

l Bohr theory of Hydrogen Atom § Bohr assumed that: 1. Energy-level postulate an

l Bohr theory of Hydrogen Atom § Bohr assumed that: 1. Energy-level postulate an atom looked something like the solar system: 1) a nucleus at the center 2) the electron could have only certain orbits L 代表电子运动轨道的角动量(L= p ·r =mv r ) P 是电子轨道运动动量, r 是轨道半径, m 是电子的质量, v 是电子的运动速度。 量子化条件: 电子在任意轨道做圆周运动的角动量mv r 必须是 的整数倍, n = 1, 2, 3,

n=3 n=2 n=1 5 = r m p 9 2. +

n=3 n=2 n=1 5 = r m p 9 2. +

l Bohr theory of Hydrogen Atom § 3) energy levels: an electron in an

l Bohr theory of Hydrogen Atom § 3) energy levels: an electron in an atom can have only specific energy values, which are called the energy levels of the electron in the atom En = - (Z 2/n 2) × 2. 180 × 10 -18 J (for H atom) Z : 核电荷数 n : 能级数 1, 2, 3, --- ∞ n值越大,表示电子运动轨道离核越远,能量越高。

2. Transitions(跃迁)between energy levels § photons are given off or absorbed when an electron

2. Transitions(跃迁)between energy levels § photons are given off or absorbed when an electron moves from one orbit to another. ground state a lower energy state ( if n = 1, is called ground state ) excited state a high energy state ( if n = 2、3……, is called ground state)

Ground state Energy of emitted photon ΔE = Ei - Ef = hv Excited

Ground state Energy of emitted photon ΔE = Ei - Ef = hv Excited state Ei a higher energy level (initial energy level) Ef a lower energy level (final energy level )

§ In 1885, J. J. Balmer showed that the wavelengths, λ, in the visible

§ In 1885, J. J. Balmer showed that the wavelengths, λ, in the visible spectrum of hydrogen could be reproduced by a simple formula. 1 1 1 --- = 1. 097 × 107 m-1 ( ---- - -----) λ 22 n 2 postulate from level n = 4 to level n = 2 light of wavelength 486 nm (blue green ) is emitted

Hydrogen atom spectra Low E Visible lines in H atom Long l spectrum are

Hydrogen atom spectra Low E Visible lines in H atom Long l spectrum are called the Low n BALMER series. High E Short l High n 6 5 4 Energy 3 2 1 Ultra Violet Lyman Visible Balmer Infrared Paschen n

Bohr’s theory § Successful 1. established the concept of atomic energy levels (atomic orbit)

Bohr’s theory § Successful 1. established the concept of atomic energy levels (atomic orbit) 2. explaining the spectrum of hydrogen § Unsuccessful 1. atomic orbit is fastness 2. can’t explain minuteness the spectrum of hydrogen atom

7 -1. 2 De Broglie Waves (Matter Waves) Louis-Victor de Broglie, (1892 -1987, France)

7 -1. 2 De Broglie Waves (Matter Waves) Louis-Victor de Broglie, (1892 -1987, France) In 1929, he was awarded the Nobel Prize for Physics for his research on quantum theory and his discovery of the wave nature of electrons. He showed that the wavelength of moving particles is equal to Planck's constant divided by the momentum.

(7 -4) § Mass: is short >> h , wave properties ignored § Particle:

(7 -4) § Mass: is short >> h , wave properties ignored § Particle: <<h, can not ignored wave properties

7 -1. 3 The Heisenberg Uncertainty principle § 1927 , He recognized : no

7 -1. 3 The Heisenberg Uncertainty principle § 1927 , He recognized : no experimental method can be devised that will measure simultaneously the precise position(x) as well us the Heisenberg German physicist momentum (mv) of (1901 -1971) an object.

Uncertainty principle formula Δp uncertainty of the momentum Δx uncertainty of the position h

Uncertainty principle formula Δp uncertainty of the momentum Δx uncertainty of the position h Planck's constant The more precisely one knows Δp, the less precisely Δx is known, and vice versa.

(中文p 148_) • Example § Suppose Δx=1. 0 × 10 - 4 m for

(中文p 148_) • Example § Suppose Δx=1. 0 × 10 - 4 m for a substance with mass of 0. 01 kg § In hydrogen atom, for an electron, v =106 m/s , suppose Δx=1. 0 × 10 - 10 m, 电子速度的不准确量 与其速度本身十分接近

Quantum or Wave Mechanics Schrodinger applied idea of ebehaving as a wave to the

Quantum or Wave Mechanics Schrodinger applied idea of ebehaving as a wave to the problem of electrons in atoms. E. Schrodinger 1887 -1961 1933 received the Nobel Prize E the total energy V the potential energy m a particle in terms of its mass x y z respect to its position in three dimensions

7 -1. 4 Schrődinger Equation (wave function) Solution to WAVE EQUATION gives set of

7 -1. 4 Schrődinger Equation (wave function) Solution to WAVE EQUATION gives set of mathematical expressions called WAVE FUNCTIONS ψ (psi) § wave function ψ has an amplitude(振幅) at each position in space (just as for a water wave or a classical electromagnetic wave).

7 -2. 1 Wave Function, Atomic Orbital and Electron Cloud ψ is a function

7 -2. 1 Wave Function, Atomic Orbital and Electron Cloud ψ is a function of distance and two angles. ——— Ψ(r, θ, φ)、 For 1 electron, ψ corresponds to an ORBITAL — the region of space within which an electron is found. ψ does NOT describe the exact location of the electron.

7 -2. 2 Atomic Orbital ____ Quantum Numbers n the principal quantum number l

7 -2. 2 Atomic Orbital ____ Quantum Numbers n the principal quantum number l the angular momentum quantum number m the magnetic quantum number. they will be used to describe atomic orbitals and to label electrons that reside in them.

1. Principal quantum number (n): § Shell K L M N. . . n

1. Principal quantum number (n): § Shell K L M N. . . n 1 2 3 4. . . As n increases, the orbitals extend farther from the nucleus, average position of an electron in these orbitals is farther from the nucleus Energies: K<L<M<N<O< … 1<2< 3< 4< 5 < …

2. Angular momentum quantum number (l ) § Within each shell of quantum number

2. Angular momentum quantum number (l ) § Within each shell of quantum number n , there are n different kinds of orbital, each with a distinctive shape, denoted by the l quantum number. § subshell s p d f g. . . l 0 1 2 3 4. . . (n-l) Energies: s<p < d < f < g…

3. Magnetic quantum number (m): A subshell has the same shape, but a different

3. Magnetic quantum number (m): A subshell has the same shape, but a different orientation, or direction, in space. m = (2 l + 1) or Each orbital of a particular subshell (no matter how it is oriented in space) has the same energy. Example: p orbit have 3 different orientation p x. p y p z

About Quantum Numbers —— Orbital An atomic orbital is defined by 3 quantum numbers:

About Quantum Numbers —— Orbital An atomic orbital is defined by 3 quantum numbers: n l m Electrons are arranged in shells and subshells of RBITALS. n shell l subshell m designates an orbital within a subshell

Table 7 -1: The allowed sets of quantum numbers for atomic orbitals

Table 7 -1: The allowed sets of quantum numbers for atomic orbitals

4. Spin quantum number (ms) : ms the spin quantum number refers to a

4. Spin quantum number (ms) : ms the spin quantum number refers to a magnetic property of electrons called spin. Values for the spin quantum number are +1/2 and – 1/2.

A fourth quantum number Note: n. l. m. ms they will be used to

A fourth quantum number Note: n. l. m. ms they will be used to describe electrons that reside in them

QUANTUM NUMBERS 1. Which of the following is not a valid set(有效的组合) of four

QUANTUM NUMBERS 1. Which of the following is not a valid set(有效的组合) of four quantum numbers to describe an electron in an atom? (1) 1, 0, 0, +½ (2) 2, 1, 1, +½ (3) 2, 0, 0, –½ (4) 1, 1, 0, +½ 2. The energy of an orbital in a many-electron atom depends on (1) the value of n only (2) the value of l only (3) the values of n and l (4) the values of n, l, and m

7 -2. 3 Sizes and Shapes of Atomic Orbitals Radial wave function angular wave

7 -2. 3 Sizes and Shapes of Atomic Orbitals Radial wave function angular wave function

Spherical coordinates x = r sin cos y = r sin z = r

Spherical coordinates x = r sin cos y = r sin z = r cos

Shapes of the orbitals § Shapes of the orbitals for: § (a) an s

Shapes of the orbitals § Shapes of the orbitals for: § (a) an s subshell ? § (b) a p subsell § (c) a d subshell

The Probability Function (ψ2) —— Electron Cloud ψ2 is related to the probability per

The Probability Function (ψ2) —— Electron Cloud ψ2 is related to the probability per unit volume such that the product of ψ 2 and a small volume (called a volume element) yields the probability of finding the electron within that volume.

1. Electron Cloud § The total probability of locating the electron in a given

1. Electron Cloud § The total probability of locating the electron in a given volume (for example, around the nucleus of an atom) is then given by the sum of all the products of ψ2 and the corresponding volume elements.

2 pz 2 px

2 pz 2 px

f orbitals

f orbitals

Probability 2 P=|Ψ| • d. V= 2 =4πr dr = D(r) 2 2 |Ψ|

Probability 2 P=|Ψ| • d. V= 2 =4πr dr = D(r) 2 2 |Ψ| • 4πr dr 2 • R (r) • • 2 Y l, m(θ, φ) Radial distribution Angular distribution function diagram

7 -3 Electron Configuration of Many-electron Atoms § 1. An electron configuration describes the

7 -3 Electron Configuration of Many-electron Atoms § 1. An electron configuration describes the arrangement of electrons in the subshells of an atom. § 2. The chemical properties of elements are related to these configurations. § 3. The four quantum numbers n, l, m, and ms enable us to label completely an electron in any orbital in any atom.

Order of filling orbitals § Generally, the energy of an orbital depends on the

Order of filling orbitals § Generally, the energy of an orbital depends on the quantum n and l. § E 1 s E 2 s. E 2 p E 3 s. E 3 p E 3 d E 4 s. E 4 p E 4 d E 4 f E 5 s… 1 s 2 s 2 p 3 s 3 p 4 s 3 d 4 p 5 s 4 d 5 p 6 s 4 f 5 d 6 p 7 s…

Why? This phenomenon can be explained by shielding effect (screening effect) and penetrating effect.

Why? This phenomenon can be explained by shielding effect (screening effect) and penetrating effect. 1. The shielding effect is that it reduces the electrostatic attraction between protons in the nucleus and the electron in outside orbital. 2. 2. The penetrating effect of an electron can decrease the energy of orbital.

The electron fill law 1. principle of energy levels lowest Electrons in an atom

The electron fill law 1. principle of energy levels lowest Electrons in an atom occupy the lowest possible energy levels, or orbitals. 2. The Pauli exclusion principle: No two electrons in the same atom can have the same set of four quantum numbers. 3. Hund's rule: Every orbital in a subshell is singly occupied (filled) with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin;

1. principle of energy levels lowest All of the electrons in an atom reside

1. principle of energy levels lowest All of the electrons in an atom reside in the lowest energy orbitals possible as long as permission of Pauli exclusion principle. The electrons filling order is: 1 s, 2 s 2 p, 3 s 3 p, 4 s 3 d 4 p, 5 s 4 d 5 p, 6 s 4 f 5 d 6 p, 7 s 5 f…… 6 p 6 s 5 p 5 s 4 p 4 s 3 s 2 s 1 s 3 p 2 p 5 d 4 d 3 d 4 f

2. Pauli Exclusion Principle (2 n 2) The Pauli exclusion principle states that no

2. Pauli Exclusion Principle (2 n 2) The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers: n l m and ms. Thus, for two electrons to occupy the same orbital, one must have ms = + ½ and the other must have ms = – ½. • electrons with the same spin keep as far apart as possible • electrons of opposite spin may occupy the same orbital

3. Hund’s rule(洪特规则) This rule states that for orbitals with the same energy, the

3. Hund’s rule(洪特规则) This rule states that for orbitals with the same energy, the lowest energy is attained when the number of electrons with the same spin is maximized.

Example Boron(atomic number =5) Nitrogen (atomic number =7) B 1 s 22 s 2

Example Boron(atomic number =5) Nitrogen (atomic number =7) B 1 s 22 s 2 2 p 1 N 1 s 22 s 2 2 p 3 Magnesium (atomic number =12) Mg 1 s 22 s 2 2 p 63 s 2 or [Ne]3 s 2 Chromium (atomic number =24) Copper (atomic number =29) ? Lanthanum (atomic number =57)

According to Hund’s rule and Pauli exclusion principle, we can writing the electron configurations

According to Hund’s rule and Pauli exclusion principle, we can writing the electron configurations for other elements. Example: chromium (Z = 24) [Ar]4 s 13 d 5 or [Ar]4 s 23 d 4 half-filled (s 1 p 3 d 5) Subshells completely empty(s 0 p 0 d 0) stability completely filled (s 2 p 6 d 10)

7 - 4 The Periodic Table and Periodic Law § Then in 1869, Russian

7 - 4 The Periodic Table and Periodic Law § Then in 1869, Russian chemist Dimitri Mendeleev (1834 -1907) proposed arranging elements by atomic weights and properties (Lothar Meyer independently reached similar conclusion but published results after Mendeleev). Mendeleev's periodic table of 1869 contained 17 columns with two partial periods of seven elements each (Li-F & Na -Cl) followed by two nearly complete periods (K-Br & Rb-I).

7 - 4 The Periodic Table and Periodic Law § The modem Periodic Table

7 - 4 The Periodic Table and Periodic Law § The modem Periodic Table consists of 7 horizontal(水平) rows of elements (often referred to as periods or series) and 32 vertical(垂直) columns of elements (referred to as families or groups).

periods § short period long periods First (2 element) second (8 element) third fourth

periods § short period long periods First (2 element) second (8 element) third fourth 18 elements fifth 18 elements sixth 32 elements seventh 32 elements

periods or series § The first short period contains two elements hydrogen (H)and helium(He).

periods or series § The first short period contains two elements hydrogen (H)and helium(He). § The second short period contains eight elements, beginning with lithium (Li) and ending with neon (Ne). § The third short period also contains eight elements, beginning with sodium (Na)and ending with argon (Ar).

The two long periods, § The fourth period and the fifth period are two

The two long periods, § The fourth period and the fifth period are two long periods, each of which contains 18 elements. The fourth period includes the elements from potassium (K)through krypton (kr). Within this period are the elements from scandium (Sc)through copper(Cu), which are known as the first transition series.

§ The fifth period is begins with rubidium (Rb)and ends with xenon (Xe). Within

§ The fifth period is begins with rubidium (Rb)and ends with xenon (Xe). Within this period are the elements yttrium (Y) through silver (Ag), which comprise the second transition series.

The sixth period § The sixth period, beginning with cesium (Cs)and ending with radon

The sixth period § The sixth period, beginning with cesium (Cs)and ending with radon (Rn), contains 32 elements. § The third transition series, made up of lanthanum (La)and the elements hafnium (Hf)through gold (Au)

The sixth period § The third transition series is split: between lanthanum and hafnium

The sixth period § The third transition series is split: between lanthanum and hafnium is a series of 14 elements, cerium (Ce) through lutetium (Lu), called the first inner transition series, or the lanthanide series or the rare earth elements.

The seventh period § The seventh period extends from francium through element number 118.

The seventh period § The seventh period extends from francium through element number 118. § However, no elements after element 109 have been characterized. § The known elements in this period include a part of the fourth transition series (actinium, and elements 104 through 109).

Electronic Structure and the Periodic Law § the periodicity with respect to the number

Electronic Structure and the Periodic Law § the periodicity with respect to the number of valence electrons; § valence electrons that is, electrons in the outermost shell. § the Periodic Table is simply an arrangement of atoms that puts elements with the same number of valence electrons in the same group.

表:基态电中性原子的电子组态 价层电子 “电子仁”或“电子实 价电子层 ” 23 p 3 1 15磷P [Ne] 3 s 1

表:基态电中性原子的电子组态 价层电子 “电子仁”或“电子实 价电子层 ” 23 p 3 1 15磷P [Ne] 3 s 1 氢H 1 s 23 p 4 2 16硫S [Ne] 3 s 2 氦He 1 s 17氯Cl [Ne] 3 s 23 p 5 3 锂Li [He] 2 s 1 22 s 22 p 63 s 23 p 6 2 18氩Ar 1 s 4 铍Be [He] 2 s 不 19钾K [Ar] 4 s 1 5硼B [He] 2 s 22 p 1 符 2 20钙Ca [Ar] 4 s 6 碳C [He] 2 s 22 p 2 合 1 2 2 3 21钪Sc [Ar] 3 d 4 s 7 氮N [He] 2 s 2 p 构 24 s 2 2 4 22钛Ti [Ar] 3 d 8 氧O [He] 2 s 2 p 造 34 s 2 2 5 23钒V [Ar] 3 d 9 氟F [He] 2 s 2 p 54 s 1 原 2 2 6 24铬Cr* [Ar] 3 d 10氖Ne 1 s 2 s 2 p 54 s 2 理 1 25锰Mn [Ar] 3 d 11钠Na [Ne] 3 s 26铁Fe [Ar] 3 d 64 s 2 12镁Mg [Ne] 3 s 2 74 s 2 2 1 27钴Co [Ar] 3 d 13铝Al [Ne] 3 s 3 p 28镍Ni [Ar] 3 d 84 s 2 14硅Si [Ne] 3 s 23 p 2

1 -48号元素的核外电子层结构 1 H 1 s 1 17 Cl [Ne]3 s 23 p 5

1 -48号元素的核外电子层结构 1 H 1 s 1 17 Cl [Ne]3 s 23 p 5 33 As [Ar]3 d 104 s 24 p 3 2 He 1 s 2 18 Ar [Ne]3 s 23 p 6 34 Se [Ar]3 d 104 s 24 p 4 3 Li [He]2 s 1 19 K [Ar]4 s 1 35 Br [Ar]3 d 104 s 24 p 5 4 Be [He]2 s 2 20 Ca [Ar]4 s 2 36 Kr [Ar]3 d 104 s 24 p 6 5 B [He]2 s 22 p 1 21 Sc [Ar]3 d 14 s 2 37 Rb [Kr]5 s 1 6 C [He]2 s 22 p 2 22 Ti [Ar]3 d 24 s 2 38 Sr [Kr]5 s 2 7 N [He]2 s 22 p 3 23 V [Ar]3 d 34 s 2 39 Y [Kr]4 d 15 s 2 8 O [He]2 s 22 p 4 24 Cr [Ar]3 d 54 s 1 40 Zr [Kr]4 d 25 s 2 9 F [He]2 s 22 p 5 25 Mn [Ar]3 d 54 s 2 41 Nb [Kr]4 d 45 s 1 10 Ne [He]2 s 22 p 6 26 Fe [Ar]3 d 64 s 2 42 Mo [Kr]4 d 55 s 1 11 Na [Ne]3 s 1 27 Co [Ar]3 d 74 s 2 43 Tc [Kr]4 d 55 s 2 12 Mg [Ne]3 s 2 28 Ni [Ar]3 d 84 s 2 44 Ru [Kr]4 d 75 s 1 13 Al [Ne]3 s 23 p 1 29 Cu [Ar]3 d 104 s 1 45 Rh [Kr]4 d 85 s 1 14 Si [Ne]3 s 23 p 2 30 Zn [Ar]3 d 104 s 2 46 Pd [Kr]4 d 10 15 P [Ne]3 s 23 p 3 31 Ga [Ar]3 d 104 s 24 p 1 47 Ag [Kr]4 d 105 s 1 16 S [Ne]3 s 23 p 4 32 Ge [Ar]3 d 104 s 24 p 2 48 Cd [Kr]4 d 105 s 2

families or groups § 1. Elements in any one group have the same number

families or groups § 1. Elements in any one group have the same number of electrons in their outermost shell § 2. The similarity in chemical properties among elements of the same group occurs because they have the same numbers of valence electrons § 3. The number of electrons in the valence shell of an atom determines its chemical properties. § 4. It is the loss, gain, or sharing of valence electrons that determines how elements react.

families or groups § 1. A number of groups = electron number of outmost

families or groups § 1. A number of groups = electron number of outmost shell = greatest oxidation number Example: 17 Cl 15 P 2. B number of groups =lose electron number [(n-1)d+ns] (except ⅧB) =greatest oxidation number(but it can be changed ) Example: Cr +2, +3, +6 Mn +2 , +3, +4, +6, +7

Electronegativity § The electronegativity of an atom is a measure of the ability of

Electronegativity § The electronegativity of an atom is a measure of the ability of an atom to draw bonding electrons to itself when chemically combined with another atom § In general, electronegativity increases in any row of the periodic table from left to right, and it decreases in going from the top of a column to the bottom.