Transcript Chapter

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 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 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 not stable;
 2.the electron would
quickly spiral into the
nucleus.
 3. Not is the line
spectra of atoms
Continuous spectrum
Atomic Line Spectra
Na
(H、He、Li、Na、Ba、Hg、Ne light emission)
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 unit
quantum
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
(1879 -1955)
h Planck’s constant
= 6.623×10-34J s.
E = hv =
h
c

(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 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.
 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
h
Ln
2
量子化条件:
L
P
r
m
v
代表电子运动轨道的角动量(L= p ·r =mv r )
是电子轨道运动动量,
是轨道半径,
是电子的质量,
是电子的运动速度。
电子在任意轨道做圆周运动的角动量mv r
h
必须是 2 的整数倍, n = 1, 2, 3,
n=3
n=2
n=1
+
 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 = - (Z2/n2) ×2.180 × 10-18J (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 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
• Orbitof– emitted photon
ΔE = Ei - Ef = hv
Excited state
a higherstate
energy
•EiGround
– 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 spectrum of
hydrogen could be reproduced by a simple
formula.
1
1
1
--- = 1.097 × 107m-1 ( ---- - -----)
λ
22
n2
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  spectrum are called the
Low n
High E
Short 
High n
BALMER series.
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)
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)
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.
h
h


mn p
 Mass:
mn >> h ,
(7-4)

is short
wave properties ignored
 Particle:
mn
<<h,  can not ignored
wave properties
[例7-1]
分别计算m=2.5×10-2kg,v = 300m·s-1的子弹
和me=9.1×10-31kg v =1.5×106 m·s-1的电子的
波长,并加以比较。
 解: 按(7-4)式,子弹的波长为:
6.626 10 34
35
 23


8
.
8

10
(
m
)

8
.
8

10
( pm)
2
2.5 10  300
电子的波长为:

6.626  10 34
9.1 10
31
 1.5  10
6
 500 ( pm)
计算结果表明,子弹的波长很短,完全可以不予考虑。
电子的波粒二象性(Davisson和Germer实验 )
1927年美国物理学家Davisson C和Germer L根据电子的波长
与X射线波长相近,用电子束代替X射线,用镍晶体薄层
作为光栅进行衍射实验,得到与X射线衍射类似的图像,
证实了电子的波动性。
X-diffracted
electron diffracted
7-1.3 The Heisenberg Uncertainty principle
 1927 ,He recognized :
Heisenberg German physicist
(1901-1971)
no experimental
method can be
devised that will
measure
simultaneously the
precise position(x)
as well us the
momentum (mv) of
an object.
Uncertainty principle formula
h
h
px 
或 x 
4
4mν
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p148_)
•Example
 Suppose Δx=1.0 ×10- 4 m for a substance with
mass of 0.01kg
h
6.626  10 34
v 

4mx
4  3.14  1.0  10  2  1.0  10  4
 5.3  10  29 (m  s 1 )
In hydrogen atom, for an electron, v =106m/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
h
6.626 10 34
v 

31
10
4mx
4  3.14  9.1110 10
 5.8 105 (m  s 1 )
Quantum or Wave Mechanics
Schrodinger applied idea of ebehaving as a wave to the
problem of electrons in atoms.
 2   2   2  8 2 m
 2  2  2 ( E  V )  0
2
x
y
z
h
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)
 2   2   2  8 2 m
 2  2  2 ( E  V )  0
2
x
y
z
h
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 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 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 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 n ,
there are n different kinds of orbital, each
with a distinctive shape, denoted by the
l quantum number.
 subshell
l
s p d f g...
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
orientation, or direction, in space.
m = (2 l + 1) or 0  1  2  3  l...
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:
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
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 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 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
 n.l .m ( r . . )  Rn.l ( r )Yl .m ( . )
Radial wave
function
angular
wave
function
 n.l .m ( r . . )  Rn.l ( r )Yl .m ( . )
 .0.30
Spherical coordinates
x = r sin cos
y = r sin sin
z = r cos
Shapes of the orbitals
 Shapes of the
orbitals for:
 (a) an s subshell
?
 (b) a p subsell
 (c) a d subshell
如:氢原子的角度部分
【s轨道】
1
Ys ( ,  ) 
4
Ys是一常数与(,)无关,半径为:
【pz轨道】
1
4
3
Ypz ( ,  ) 
cos
4
节面:当cos  0时,Y0,  90°
我们下来试做一下函数在Pz平面的图形。
z
0
30° 30
+ θ
60°
60
90
x,y
-

0
30
60 90 120
150
180
cos 1
0.866 0.5 0 - 0.5
- 0.866
-1
YP z 0.489 0.423 0.244 0 - 0.244 - 0.423 - 0.489
节面:当θ = 90° cosθ= 0 Y=0时
波函数的角度分布图
由图可知,原子轨道的角度分布图有正负之分,
这对于讨论分子的化学键及对称性十分重要。
同样地,可以画出其它原子轨道的角度分布图。
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 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.
2pz
2px
f orbitals
电子云的径向分布图
 n.l .m ( r . . )  Rn.l ( r )Yl .m ( . )
|Ψn,l,m(r,θ,φ)
2
| =
2
R
n,l(r)
•
2
Y
l,m(θ,φ)
Probability density
Probability
2
P=|Ψ| •
dV
几率(dP)=几率密度(|ψ|2)×体积(dV)
电子云的径向分布图
考虑离核距离为r,厚度为dr的薄层球壳内发现电
子的几率.
1s球壳微体积: dV = 4πr2dr
D(r) =4πr2dr •R2(r)
-----壳层几率(球壳层
内发现电子的几率)
Probability
2
P=|Ψ| •
dV=
2
=4πr dr
2
•R
2
2
|Ψ| •4πr dr
(r)
= D(r) •
•
2
Y
2
Y
l,m(θ,φ)
l,m(θ,φ)
Radial distribution Angular distribution
function diagram
function diagram
P= |Ψ|2 •4πr2dr
 离核越近:
r值越小,体积越小,|ψ|2越大,D(r)不是最大,
 离核越远:
r值越大,体积越大,|ψ|2越小,D(r)亦不是最大,
 在ao处:
|ψ|2不是最大的, 但体积较大,使D(r)可达最大。
ao=52.9pm处。
当r=2ao时, D(r)=0,出现第一个节面。
当r=4ao时, D(r)又出现最大值,此即2s电子云
当r=2ao时, D(r)=0,出现第一个节面。
当r=4ao时, D(r)又出现最大值,此即2s电子云
电子云的径向分布图 峰数= n-l
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 quantum n and l .
 E1s E2sE 2p E3sE3p E3d E4sE 4p E 4d E4f E5s…
1s 2s2p 3s3p 4s3d4p 5s4d5p 6s4f5d6p 7s…
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. The penetrating effect of an electron can
decrease the energy of orbital.
The penetrating effect
D(r)
D(r)
1s
2s
3d
3s
3s
图1 l 相同, n不同时的比较
3p
r
图2 n 相同, l 不同时的比较
从上图可以看出:
(1) l相同,n不同: 1s<2s<3s . n 增大时,电子离核的距离
(主峰)将增加。
(2) n相同,l不同 3s<3p<3d. l 值大,峰个数减少。
l 值小,电子在核附近出现的机会(钻穿峰)较多。
r
(3) n,l都不同时,将出现能级交错 :
4s<3d<4p
钻穿效应: 外层电子向内层穿透,导致内层电
子对它的屏蔽作用减弱的效应叫钻穿效应
Question
为什么 2s 价电子比 2p 价电
子受到较小的屏蔽?
2s电子云径向分布
曲线除主峰外,还有
一个距核更近的小峰.
这暗示, 部分电子云
钻至离核更近的空间,
从而部分回避了其他
电子的屏蔽.
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 in the
lowest energy orbitals possible as long as
permission of Pauli exclusion principle .
The electrons filling order is:
1s, 2s2p, 3s3p, 4s3d4p, 5s4d5p, 6s4f5d6p, 7s5f……
6p
6s
5p
5s
4p
4s
3s
2s
1s
3p
2p
5d
4d
3d
4f
2. Pauli Exclusion Principle
(2n2)
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 lowest energy is attained when the
number of electrons with the same spin is
maximized.
N Ô-×ÓÐòÊýΪ7
按洪特规则的基态电子构型
而不是
是
1s
1s
2s
2p
2s
2p
Example
Boron(atomic number =5)
B
Nitrogen (atomic number =7) N
1s22s2 2p1
1s22s2 2p3
Magnesium (atomic number =12) Mg
1s22s2 2p63s2
or [Ne]3s2
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 for other elements.
Example:
chromium (Z = 24) [Ar]4s13d 5
or [Ar]4s23d4
half-filled (s1 p3 d5)
Subshells completely empty(s0p0d0) stability
completely filled (s2 p6 d10)
电子层结构式要与原子的电子排布式区别开,
以29号元素铜为例:
电 子 排 布 式: 29Cu: 1s2 2s2 2p6 3s2 3p6 4s1 3d10
电子层结构式: 29Cu: 1s2 2s2 2p6 3s2 3p6 3d10 4s1
(或电子构型式)
SPECTROSCOPIC NOTATION
for H, atomic number = 1
1
1s
value of n
no. of
electrons
value of l
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 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).
维尔纳长式周期表
IA
1 H
2
1 氢
IIA
Li 4 Be
3
IIIA IVA
5 B 6 C
2 锂 铍
11 Na 12 Mg
3 钠 镁 IIIB
VA VIA VIIA
7 N 8 O 9 F
He
氦
10 Ne
硼 碳 氮 氧 氟 氖
13
IVB VB VIB VIIB
19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn
VIII
IB IIB
26 Fe 27 Co 28 Ni 29 Cu 30 Zn
Al
14
Si
15
P
16
S
17
Cl
18 Ar
铝 硅 磷 硫 氯 氩
31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr
4 钾 钙 钪 钛 钒 铬 锰 铁 钴 镍 铜 锌 镓 锗 砷 硒 溴 氪
37
Rb
38 Sr 39
Y
40 Zr 41 Nb 42 Mo 43
Tc
44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49
In
50 Sn 51 Sb 52 Te 53
I
54 Xe
5 铷 锶 钇 锆 铌 钼 锝 钌 铑 钯 银 镉 铟 锡 锑 碲 碘 氙
6
75Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn
铯 钡 La-Lu 铪 钽 钨 铼 锇 铱 铂 金 汞 铊 铅 铋 钋 砹 氡
55
Cs
56 Ba 57 71 72 Hf 73 Ta 74 W
87 Fr 88 Ra 89-103 104 Rf 105Db 106Sg 107Bh 108 Hs 109Mt 110
7 钫 镭
Ac-Lr
111
112
钅
卢 钅杜 钅喜 钅波 钅
麦 Uun Uuu Uub
黑 钅
57 La 58 Ce 59
Pr
60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65
Tb
114
66 Dy 67 Ho 68
Er
118
116
69Tm 70 Yb 71
Lu
镧系 镧 铈 镨 钕 钷 钐 铕 钆 铽 镝 钬 铒 铥 镱 镥
89 Ac 90 Th 91 Pa 92 U 93 Np 94 Pu 95Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101Md 102No 103 Lr
锕系 锕 钍 镤 铀 镎 钚 镅 锔 锫 锎 锿 镄 钔 锘 铹
periods
 short period
long periods
First
(2 element)
second (8 element)
(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).
 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 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 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
(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 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.
 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 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
“电子仁”或“电子实”
1 氢H
2 氦He 1s2
3 锂Li [He] 2s1
4 铍Be [He] 2s2
5硼B [He] 2s22p1
6 碳C [He] 2s22p2
7 氮N [He] 2s22p3
8 氧O [He] 2s22p4
9 氟F [He] 2s22p5
10氖Ne 1s2 2s22p6
11钠Na [Ne] 3s1
12镁Mg [Ne] 3s2
13铝Al [Ne] 3s23p1
14硅Si [Ne] 3s23p2
1s1
价电子层
价层电子
15磷P [Ne] 3s23p3
16硫S [Ne] 3s23p4
17氯Cl [Ne] 3s23p5
18氩Ar 1s22s22p63s23p6
19钾K [Ar] 4s1
20钙Ca [Ar] 4s2
21钪Sc [Ar] 3d14s2
22钛Ti [Ar] 3d24s2
23钒V [Ar] 3d34s2
24铬Cr* [Ar] 3d54s1
25锰Mn [Ar] 3d54s2
26铁Fe
[Ar] 3d64s2
27钴Co [Ar] 3d74s2
28镍Ni
[Ar] 3d84s2
不
符
合
构
造
原
理
1-48号元素的核外电子层结构
1
H
1s1
17
Cl
[Ne]3s23p5
33
As
[Ar]3d104s24p3
2
He
1s2
18
Ar
[Ne]3s23p6
34
Se
[Ar]3d104s24p4
3
Li
[He]2s1
19
K
[Ar]4s1
35
Br
[Ar]3d104s24p5
4
Be
[He]2s2
20
Ca
[Ar]4s2
36
Kr
[Ar]3d104s24p6
5
B
[He]2s22p1
21
Sc
[Ar]3d14s2
37
Rb
[Kr]5s1
6
C
[He]2s22p2
22
Ti
[Ar]3d24s2
38
Sr
[Kr]5s2
7
N
[He]2s22p3
23
V
[Ar]3d34s2
39
Y
[Kr]4d15s2
8
O
[He]2s22p4
24
Cr
[Ar]3d54s1
40
Zr
[Kr]4d25s2
9
F
[He]2s22p5
25
Mn
[Ar]3d54s2
41
Nb
[Kr]4d45s1
10
Ne
[He]2s22p6
26
Fe
[Ar]3d64s2
42
Mo
[Kr]4d55s1
11
Na
[Ne]3s1
27
Co
[Ar]3d74s2
43
Tc
[Kr]4d55s2
12
Mg
[Ne]3s2
28
Ni
[Ar]3d84s2
44
Ru
[Kr]4d75s1
13
Al
[Ne]3s23p1
29
Cu
[Ar]3d104s1
45
Rh
[Kr]4d85s1
14
Si
[Ne]3s23p2
30
Zn
[Ar]3d104s2
46
Pd
[Kr]4d10
15
P
[Ne]3s23p3
31
Ga
[Ar]3d104s24p1
47
Ag
[Kr]4d105s1
16
S
[Ne]3s23p4
32
Ge
[Ar]3d104s24p2
48
Cd
[Kr]4d105s2
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 shell
= greatest oxidation number
Example:
17Cl
15P
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 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.
电负性的应用
1.判断元素的金属性和非金属性 金属性
元素的电负性一般在2.0以下,非金属性性
元素一般在2.0以上。电负性最大的元素是
位于右上方的F,电负性最小的元素是位于
左下方的Fr(Fr是放射性元素)
2.估计化学键的类型 在化合物中,可以根据电负
性的差值大小,估计化学键的类型。
电负性差越大,离子性越强,一般说来,电负性差
大于1.7时,可认为是离子键,小于1.7时为共价键。
原子半径及其周期性变化
原子半径:有三种不同的定义
(1)共价半径:同种元素的两原子以共价单键相连时,两原子
核间距离的一半叫共价半径.
(2)金属半径:金属元素的两原子以密堆积方式(金属键)结
合成金属晶体时,两原子核间距离的一半叫金属半径.
(3)范氏半径:同种元素的两原子不能以共价键相连,只靠分
子间力作用时,两原子核间距离的一半叫范德华半径.
1、原子半径及其变化规律
2、主族元素原子半径的变化规律
主族元素:
(1)同周期 从左到右
逐渐减小.但到希
有气体元素时又有
增大,原因是半径
定义不同.
(2)同一族 从上到下
逐渐增大.从左到
右,核电荷增加是
主要因素,但从上
到下,电子层增加
是主要因素.
副族元素原子半径的变化规律
副族元素:
(1)同周期 从左到右逐渐减小.但减小幅度不如主
族元素,这是由于最后一个电子是填在(n-1)层d
轨道上对核电荷的抵消作用造成的.
(2)同一族 从上到下逐渐增
大。但增加幅度较小,甚
至第五、第六周期基本没
有增加,这是由于镧系收
缩的原因造成的.