Download Chapter 1 Atomic Structure and Periodic Properties of the Elements

Chapter 1 Atomic Structure and Periodic
Properties of the Elements (原子结构与元素周期表)
1.1 Light, electromagnetic radiation and spectrum
(光、电磁辐射和光谱)
1.2 Hydrogen spectrum and Bohr’s Model
(氢光谱与波尔模型)
1.3 Electronic Structure of Atoms
(原子的电子结构)
1.4 Electron Configurations and the Periodic Table
(核外电子排布与元素周期表)
1-1
1.1 Light, electromagnetic radiation and spectrum
(光、电磁辐射和光谱)
Wave-particle duality of Light (光的波粒二象性)
When we say “light,” we generally are referring to visible
light(可见光)(λ= 400 ~ 750 nm) — a type of electromagnetic
radiation(电磁辐射). Visible light constitutes a very small
segment(片段) of the electromagnetic spectrum(电磁波谱).
Other types of electromagnetic radiation in this spectrum include
microwave(微波), infrared(红外), and ultraviolet radiation(紫外).
光的波动性主要体现在：
Properties :
Wavelength(波长), l (nm). Frequency(频率), n (s-1, Hz)
Amplitude(振幅), A.
constant speed. c (3.00×108 m.s-1)
1-2
• All waves have:
frequency(频率)
and wavelength(波长)
• symbol: n (Greek letter “nu”)
l (Greek “lambda”)
• units:
“cycles per sec” = Hertz
“distance” (nm)
• All radiation: l • n = c
Where v is the frequency of the radiation in reciprocal(倒数)
seconds (s–1), l is the wavelength in meters, and c is the speed
of light in meters per second.(3.00 × 108 m/sec)
1-3
Note: Long wavelength
 small frequency
Short wavelength
 high frequency
Example: Red light has l = 700 nm.
Calculate the frequency, n.
c
=
n=
l
3.00 x 108 m/s
7.00 x 10-7 m
increasing
frequency
increasing
wavelength
= 4.29 x 1014 Hz
• Wave nature of light is shown by classical wave
properties such as
• interference(干涉)
• diffraction(衍射)
1-4
光的粒子性：Quantized Energy and Photons
(能量量子化与光子)
Classical physics(经典物理学),energy was considered to be
“continuous” (连续的). In 1900 Max Planck (普朗克) proposed
that, like matter, energy is quantized(量子化的). He termed the
smallest “particle” of energy that can be emitted or absorbed as
electromagnetic radiation the quantum(量子).
• Planck’s hypothesis(假设): An object can only gain or lose
energy by absorbing or emitting radiant energy in QUANTA.
According to Planck’s theory, the energy of radiation is directly
proportional to its frequency. E is the energy of a single
quantum, and h is Planck’s constant(普朗克常数)(6.63×10-34
J ·s). Energy is always emitted or absorbed in whole-number
multiples(整数倍) of hv; for example, hv, 2 hv, 3 hv, and so on.
1-5
Light with a large l (small n) has a small E.
E = hv
Light with a short l (large n) has a large E.
Energy of Radiation
Problem: Calculate the energy of 1.00 mol of photons of red light.
l = 700 nm n = 4.29 x 1014 sec-1
E = h•n
= (6.63 × 10-34 J ·s)(4.29 × 1014 sec-1)
= 2.85×10-19 J per photon
E per mol = (2.85 × 10-19 J/ph)(6.02 × 1023 ph/mol)
= 171.6 kJ/mol
- the range of energies that can break bonds.
1-6
Photoelectric Effect(光电效应) Albert Einstein (1879-1955)
Photoelectric effect demonstrates the particle nature of light.
No e- observed until light of a
certain minimum E is used.
Number of e- ejected does NOT
depend on frequency, rather it
depends on light intensity.
• Classical theory said that E of ejected electron should
increase with increase in light intensity — not observed!
• Experimental observations can be explained if light
consists of particles called PHOTONS of discrete(不连续
的) energy.
1-7
1.2 Hydrogen spectrum and Bohr’s Model
(氢光谱与波尔模型)
Sunlight is a continuous spectrum(连续光谱), atom is a line
spectrum (线状光谱), which consists of only a few specific
wavelengths.
The line spectrum of hydrogen contains four visible lines
Niels Bohr(波尔) applied quantum theory(量子论) to explain the
line spectrum of hydrogen in terms of the behavior of the
electron in a hydrogen atom.
1-8
Bohr’s Model
When an electron resides(居住) in the orbit designated by n = 1, it is
said to be in the ground state(基态). This is the lowest possible energy
level in which hydrogen’s electron can exist. If hydrogen’s electron is
in a higher energy orbit, with n greater than 1, the atom is said to be in
an excite state(激发态). Bohr assumed that the electron could “jump”
from one allowed energy state to another by absorbing(吸收) or
emitting(发射) photons of radiant energy of certain frequencies. He
described the lines in the hydrogen spectrum as the energy given off
when an electron in an excited state returns to the ground state. A flame
or the application of high voltage imparts(给予) energy to the electron
in a hydrogen atom and promotes(促使) it to an orbit of higher n value.
When the excited state electron returns to the ground state, it releases
the excess(过剩) energy in the form of visible light. The frequency of
this emitted radiant energy corresponds exactly to the energy difference
between the two states (△E ).
1-9
excite state
n=2
visible light
energy
ground state
n=1
Bohr proposed that the electron in a hydrogen atom could circle
the nucleus(原子核) only in specific orbits designated by a
quantum number(量子数) n. The quantum number can have
integer values(整数值), with n = 1 corresponding to the orbit
closest to the nucleus. He showed the relationship between the
value of n and the energy of an electron is
RH is the Rydberg constant (2.18×10–18J).(雷德堡常数)
1-10
Infrared
Visible
Ultra Violet
△E = hv
The visible lines in the hydrogen line spectrum
1-11
Atomic Line Spectra
• Bohr’s greatest contribution to science was in
building a simple model of the atom.
• It was based on understanding the Sharp
Line Spectra(锐线光谱) of excited atoms.
Niels Bohr (1885-1962)
(Nobel Prize, 1922)
• Excited atoms emit light of only certain
wavelengths
• The wavelengths of emitted light depend
on the element.
H
Hg
Ne
1-12
Summary：Atomic Spectra and Bohr Model
Problems(Classical):
1. Classically any orbit should be
possible and so is any energy.
2. But a charged particle moving
in an electric field should emit energy.
+
Electron
orbit
End result should be destruction!
Bohr said :
• Classical view is wrong.
• Need a new theory — now called QUANTUM(量子论)
or WAVE MECHANICS(波动力学).
• e- can only exist in certain discrete orbits (不连续的轨道)
— called stationary states(静态).
• e- is restricted(受限于) to QUANTIZED energy states.
1-13
The Wave Behavior of Matter (物质波)
From Bohr model to Quantum mechanics (量子力学)
Problems existed with Bohr theory. (1) theory only successful
for the H atom. (2) introduced quantum idea artificially(人为
的).So, we go on to QUANTUM or WAVE MECHANICS
Quantum or Wave Mechanics
• Light has both wave & particle properties
• de Broglie (1924) proposed that all moving objects have
wave properties.
• For light: E = hn = hc / l
• For particles: E = mc2 (Einstein)
and for particles
(mass)x(velocity) = h / l
Therefore, mc = h / l
1-14
WAVE properties of matter
l for particles is called the de Broglie
wavelength
L. de Broglie
(1892-1987)
Electron diffraction with
electrons of 5-200 kev
- Al metal
Davisson & Germer 1927
Na Atom Laser
beams(衍射条纹)
l = 15
micometers (mm)
Andrews, Mewes,
Ketterle
M.I.T. Nov 1996
1-15
Question
What is the de Broglie wavelength of a 350 g object moving at a
speed of 5.00 m/s?
（1） 1.26 × 10–18 m；（2） 3.79 × 10–34 m；
（3） 3.79 × 10–37 m；（4） 1.89 × 10–33 m
Inspection of this equation reveals that only extremely small
objects, such as subatomic(亚原子) particles, have wavelengths
sufficiently large as to be observable. In other words, the
wavelength associated with a golf ball, for example, is so tiny as
to be completely out of the range of any visible observation.
Scientists proved experimentally that electrons are diffracted by
crystalline solids(晶体). Diffraction(衍射) is a wavelike
behavior.
1-16
物质波的特点：The Uncertainty Principle(测不准原理)
On the heels of de Broglie’s theory, Werner Heisenberg(海森堡)
concluded that there is a fundamental limitation on how precisely
we can simultaneously(同时地) measure the location (位置) and
the momentum(动量) of an object small enough to have an
observable wavelength. This limitation is known as the
Heisenberg uncertainty principle(海森堡测不准原理). When
applied to the electrons in an atom, this principle states that it is
inherently impossible to know simultaneously both the exact
momentum of the electron and its exact location in space.
• Cannot simultaneously define the position and momentum (=
m ·v) of an electron. (Dx ·Dp = h)
• At best we can describe the position and velocity of an electron
by a Probability density(几率密度), which is given by Y 2 .
1-17
1.3 Electronic Structure of Atoms
(原子的电子结构)
1.3.1 Wave functions and Schrödinger equation
(波函数与薛定鄂方程)
Erwin Schrödinger(薛定鄂) developed an equation to
incorporate(结合) both the wave and particle properties of the
electron. The solution to the Schrödinger equation(薛定鄂方程)
yields a series of wave functions(波函数,Ψ). The square of a
wave function(Ψ2) is the probability density(几率密度), or the
probability that an electron will be found at a given point in
space (also called electron density -电子密度). Regions where
there is a high probability of finding the electron are regions of
high electron density. The result of Schrödinger‘s work is a
more sophisticated(复杂的,深奥微妙的) view of the atom.
1-18
E. Schrodinger
1887-1961
• Schrodinger applied idea of e- behaving as a
wave to the problem of electrons in atoms.
• Solution to WAVE EQUATION gives set of
mathematical expressions called WAVE
FUNCTIONS, Y
• Each describes an allowed energy state of
an e- Quantization introduced naturally.
• Y is a function of distance and two angles —Ψ(r,θ,φ)
• For 1 electron, Y corresponds to an ORBITAL — the region
of space within which an electron is found.
• Y does NOT describe the exact location of the electron.
• Y 2 is proportional to the probability(几率) of finding an eat a given point.
1-19
1.3.2 Wave Function and Electron Atmosphere
(波函数与电子云图形)
Y 2 is proportional to the probability(几率) of finding an e- at
a given point. The diagram of Y 2 is Electron Atmosphere
(电子云). The value of Y 2 is electron cloud density(电子
云密度).
53 pm
基态氢原子的电子云
1-20
Angular distribution function (ADF)
电子云角度分布函数
Ψ(r,θ,φ) = R(r) · Y(θ,φ)
Y(θ,φ) --angular
distribution
function(ADF)
(角度分布函数)
1-21
Radial distribution function (RDF)
电子云径向分布函数
Ψ(r,θ,φ) = R(r) · Y(θ,φ)
R(r) --radial distribution
function(RDF)(径向分布函数)
表示电子在离核
半径为r的单位厚
度的薄球壳内出
现的概率
R(r) = r2R2，
r越小，R越大；
r越大，R越小。
1-22
1.3.3 Orbital Quantum Numbers(轨道量子数)
In Bohr’s model of the atom, a single quantum number
described an orbit in which an electron could exist. In the
quantum mechanical model of the atom, three quantum numbers
are required to describe what is now called an orbital(轨道)
1. The principal (major) quantum number (主量子数）n
The principal(major) quantum number n, can have positive
integer values (正整数).The principal quantum number
determines the size of the orbital. The larger the value of n, the
larger the orbital.
n
1
2
3
4
Symbol
K
L
M
N
1-23
2. The second or angular quantum number(副量子数或
角量子数) l
The second or angular quantum number l, can have integer
values from 0 to n –1. The value of l determines the shape of
the orbital(原子轨道形状). Each value of l has a letter
associated with it to designate orbital shape.
l = 0, 1, 2, 3,…(n -1)
l
0
1
2
3
Symbol
s
p
d
f
3. The magnetic quantum number(磁量子数) m
The magnetic quantum number m, can have integer values from
–l , via 0, to +l. The magnetic quantum number determines the
orbital’s orientation in space(轨道空间伸展方向).
ml = 0, ±1, ±2, … ±l
1-24
Summary
Each value of n defines an electron shell (电子层).
Within a shell, each value of l defines an electron
subshell (电子亚层). Within a subshell, each value of m
defines an individual orbital (每一个单独的轨道.)
1-25
How to use the three quantum numbers ?
1. The shell with principal quantum number n will consist of
exactly n subshells. Each subshell corresponds to a different
allowed value of l from 0 to n–1. Thus, the first shell (n = 1)
consists of only one subshell, the 1s (l = 0); the second shell (n =
2) consist of two subshells, the 2s (l = 0) and 2p (l = 1); the third
shell (n = 3) consists of three subshells, 3s (l = 0), 3p (l = 1), 3d
(l = 2), and so forth.
2. Each subshell consists of a specific number of orbitals. Each
orbital corresponds to a different allowed value of ml . For a
given value of l, there are 2l + 1 allowed values of ml, ranging
from –l to l. Thus, each s (l = 0) subshell consists of one orbital;
each p (l = 1) subshell consists of three orbitals; each d (l = 2)
subshell consists of five orbitals, and so forth.
1-26
3. The total number of orbitals in a shell is n2, where n is the
principal quantum number of the shell. The resulting number of
orbitals for the shells—1, 4, 9, 16—has a special significance
with regard to the periodic table(周期表): We see that the
number of elements in the rows(排)of the periodic table—2, 8,
18, and 32—are equal to twice these numbers.
• How many orbitals
are there in the n = 2
shell?
n
2
l
0
ml
0
orbital
s
1
+1
0
-1
three p
1-27
4. Electron Spin Quantum Number ms(自旋量子数)
In the quantum mechanical model of the atom, each orbital can
accommodate(容纳) two electrons. Electrons possess an intrinsic
(内在的,本质的)property known as electron spin(电子自旋). In
order to distinguish the two electrons in a single orbital, we must
use a fourth quantum number, the electron spin quantum
number ms.
The electron spin quantum number, ms, can have values of
+ ½ and – ½.
Diamagnetic(反磁性): NOT attracted to a
magnetic field,All electrons are paired--N2
Paramagnetic(顺磁性): attracted to a
magnetic field.Substance has unpaired
electrons --O2
1-28
QUANTUM
NUMBERS
n (shell)
1, 2, 3, 4, ...
l (subshell)
0, 1, 2, ... n - 1
ml (orbital)
- l ... 0 ... + l
ms (electron spin) +1/2, -1/2
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 ml
1-29
1.3.4 Representations of Orbitals (轨道的表示方法）
s Orbitals
All s orbitals are spherical(球形) in shape.
1-30
p Orbitals
For n = 2, l = 0 and 1, there are 2
types of orbitals — 2 subshells.
For l = 0, ml = 0, this is a s subshell
For l = 1, ml = -1, 0, +1, this is a p
subshell with 3 orbitals
Typical p orbital
When l = 1, there is a planar
node through the nucleus.
pz
90 o
A p orbital
px
py
The three p orbitals lie
90o apart in space
1-31
The p orbitals have electron density distributed in dumbbellshaped(哑铃形) regions. The three p orbitals lie along the x, y,
and z axes(轴) and are distinguished by the labels px, py, and pz.
1-32
d Orbitals
For n = 3, what are the values of l ?
l = 0, 1, 2 -- and so there are 3 subshells in the shell.
For l = 0, m = 0  s subshell with single orbital. s orbitals
(l = 0) are spherical.
For l = 1, m = -1, 0, +1  p subshell with 3 orbitals. p orbitals
have l = 1, are “dumbbell” shaped.
For l = 2, m = -2, -1, 0, +1, +2  d subshell with 5 orbitals.
By analogy (以此类推), For l = 3, m = -3, -2, -1, 0, +1, +2, +3
 f subshell with 7 orbitals.
1-33
Four of the d orbitals have double dumbbell shapes(双哑铃形
或纺锤形) that lie in planes(平面) defined by the x, y, and z
axes. The fifth d orbital has two distinct components: a
dumbbell shape lying along the z axis and a doughnutshape (圆
环形) that encircles the z axis.
typical d orbital
1-34
all orbitals of the n = 1, n = 2 and n = 3 shells
n=
3d
3
2
There are
n2
orbitals in
the nth SHELL
1
1-35
1.4 Electron Configurations and the Periodic Table
(核外电子排布与元素周期表)
1.4.1 Effective Nuclear Charge(有效核电荷) Zeff (or Z*)
In a many-electron atom, repulsions (斥力) between electrons
result in differences in orbital energies within an electron shell. In
effect, electrons are shielded(被屏蔽) from the nucleus by other
electrons. The attraction between an electron and the nucleus is
diminished by the presence of other electrons. Instead of the
magnitude(量) of attraction being determined by the nuclear
charge, in a many-electron system the magnitude of attraction is
determined by the effective nuclear charge(有效核电荷) Zeff (or
Z*). This is the diminished nuclear charge that is felt by the
electrons in orbitals with principal quantum number 2 and higher.
Z* = Z – S. S is the screening constant(屏蔽常数)
1-36
1.4.2 Shielding Effect and Penetration Effect
(屏蔽效应与穿透效应)
Shielding Effect
Shielding Effect is:
K>L>M> N>…
And
ns > np > nd > nf
So the Energy is:
K< L < M < N < …
And
ns < np < nd < nf
1-37
Penetration Effect
能级交错现象
So the Energy is:
4s < 3d < 4p
5s < 4d < 5p
6s < 4f < 5d < 6p
电子云径向密度分布图
ns < (n-2)f < (n-1)d < np
1-38
1.4.3 Atomic Energy Level Diagrams
(原子轨道能级图)
美国化学家Pauling经过计算，将原子轨道分为七个能级组。
第一组：1s
第二组：2s2p
第三组：3s3p
第四组：4s3d4p
第五组：5s4d5p
第六组：6s4f5d6p
第七组：7s5f6d7p
特点：
1、能级能量由低到高。
2、组与组之间能量差大，组内各轨道间能量差小，随n逐
渐增大，这两种能量差逐渐减小。
3、第一能级组只有1s一个轨道，其余均有两个或两个以
上，且以ns开始np结束。
4、能级组与元素周期相对应。
1-39
Pauling近似能级图
1-40
1.4.4 核外电子排布规律
1. Pauli Exclusion Principle(保里不相容原理)
The Pauli exclusion principle states that no two electrons in an
atom can have the same set of four quantum numbers: n, l, ml, 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 “region of
space” (= orbital)
Consequences:
• No orbital can have more than 2 electrons
• No two electrons in the same atom can have the same
set of 4 quantum numbers (n, l, ml , ms)
• “Each electron has a unique address.”
1-41
2. Lowest Energy Principle(能量最低原理)
All of the electrons in an atom reside in the lowest energy orbitals
possible as long as permission of Pauli exclusion principle .
Electron
Filling
Order
1-42
3. Hund’s rule (洪特规则)
Hund‘s Rule of Maximum Multiplicity (多样性) is an
observational rule which states that a greater total spin state usually
makes the resulting atom more stable. Accordingly, it can be taken
that if two or more orbitals of equal energy are available, electrons
will occupy them singly before filling them in pairs.
np2
true
false
np3
true
false
The orbitals with same energy are called degenerate orbitals (简
并轨道). e.g. p sub-shell has 3 degenerate orbitals namely px, py
and pz. The d and f sub-shell has 5 and 7 degenerate orbitals,
respectively.
Specially,all-full (p6, d10 , f14 ) or half-full(p3 , d5 , f 7 ) or empty
(p0 , d0 , f 0 ) is lower energy and more stable relatively.
1-43
Note that when electrons begin to occupy the 2p subshell, they
occupy all three p orbitals singly before pairing in a single orbital.
This is the phenomenon that is described by Hund’s rule.
1-44
Writing Atomic Electron Configurations
Two ways of writing configurations.
One is called the
spectroscopic notation.
（光谱型表示方法）
A second way is called the
orbital box notation.
（轨道型表示方法）
One electron has
n = 1, l = 0, ml = 0, ms = + 1/2
Other electron has
n = 1, l = 0, ml = 0, ms = - 1/2
SPECTROSCOPIC NOTATION
for H, atomic number = 1
1
1s
value of n
no. of
electrons
value of l
ORBITAL BOX NOTATION
for He, atomic number = 2
Arrows
2
depict
electron
spin
1s
1s
1-45
Example
Li — 1s22s1 or [He] 2s1
Na — 1s22s22p63s1 or [Ne] 3s1
Fe — 1s22s22p63s23p63d64s2 or [Ar] 3d64s2
Fe2+ — 1s22s22p63s23p63d6 or [Ar] 3d6
Fe3+ — 1s22s22p63s23p63d5 or [Ar] 3d5
Cr — 1s22s22p63s23p63d54s1 or [Ar] 3d54s1
Cu — 1s22s22p63s23p63d104s1 or [Ar] 3d104s1
1-46
1.4.5 Electron Configurations and the Periodic Table
(电子结构与周期表)
Determining electron configurations can be greatly simplified by
viewing the periodic table in terms of orbitals. Figure below
shows the division of the periodic table into the s block(区), the p
block, the d block, the ds block, and the f block. In each valence
electrons(价电子) are in an orbital designated by the block. For
example, all elements in the s block have their valence electrons
in s orbitals. The d block elements are the transition elements
(过渡元素). Their valence electrons are in d orbitals. The f block
elements are the lanthanides (镧系) and the actinide elements
(锕系).
1-47
Relationship of Electron Configuration
and Regions of the Periodic Table
d block
s block
ds
p block
f block
1-48
Electron Configuration and Periodic Table
1-49
1.4.6 Periodic Properties of the Elements
(元素性质的周期性)
Development of the Periodic Table(周期表的发展)
In 1869 when two scientists, Dmitri Mendeleev(门捷列夫) and
Lothar Meyer, put forth independent but very similar classification
schemes for the elements. Both Mendeleev and Meyer organized the
elements according to physical and chemical properties. Further, they
both noticed that certain properties were repeated periodically if the
elements were arranged in order of increasing atomic weight.
Mendeleev went so far as to leave blank spaces in his table where no
known element had the properties that would place it there logically.
He called the missing elements eka-aluminum(类铝) and eka-silicon
(类硅), and he predicted their physical and chemical properties based
on his theory of periodic properties. The subsequent discovery of
these missing elements and the fact that their properties so closely
match Mendeleev‘s predictions lent tremendous weight(具有重要地
位) to Mendeleev's theory.
1-50
In 1913 Henry Moseley improved on the periodic table by
developing the concept of atomic numbers. Arranging the
elements in order of increasing atomic number eliminated(排除)
some of the inconsistencies(不一致) in the periodic table, which
has been based on atomic weights. For example, in a table ordered
according to atomic mass, potassium(K) would appear below
neon(Ne,氖) in group 8A. Clearly the properties of potassium,
which is a solid at room temperature and highly reactive, would
not logically place it in that group. When the elements are instead
arranged in order of increasing atomic number, argon(Ar,氩) and
potassium appear in their correct places. The modern periodic
table arranges elements in order of increasing atomic number.
1-51
1. Atomic radius (原子半径)
According to the quantum mechanical model of the atom, radial
(径向) electron density — the probability of finding an electron
at a particular distance from the nucleus — falls off (减少) as
distance from the nucleus increases, but never really reaches
zero. For the noble gases, which exist as isolated atoms, this
complicates the problem of defining atomic "size."
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可以假定原子是球形，来测定相邻原子的核间距，根据原子
的不同存在形式可将原子半径分为：
金属半径：金属晶体中，金属原子被看为刚性球体，彼此相
切，其核间距的一半，为金属半径。
共价半径：同种元素的两个原子以共价单键结合时，原子核
间距离的一半。
van der Waals 半径：以分子间力结合的两个同种原子核间距
的一半。
d (I2) =2.66 Å, so r (I) = 1.33 Å
r+
r-
d (H2)=0.74 Å, so r (H)= 0.37 Å
d
So, d (HI)= 1.33 + 0.37 = 1.7 Å
d
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• Size goes UP on going down
a GROUP
• Because each successive step
down with a group adds
another shell of electrons.
• Size goes DOWN on going across a PERIOD.
• Size decreases due to increase in Z*.
• Each added electron feels a greater and greater + charge.
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Atomic Radii
Atomic radii increase
from top to bottom
within a group
Atomic radii decrease
from left to right
within a period
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Trends in Atomic Size
Radius (pm)
250
K
1st transition
se rie s
3rd pe riod
200
Na
2nd pe riod
Li
150
Kr
100
Ar
Ne
50
He
0
0
5
10
15
20
25
30
35
40
Atomic Numbe r
Question: Which series of elements is arranged correctly in order of
increasing atomic radius?
(1) C < N < O < F
(2) F < S < Al < Rb
(3) Sr < Ca < Mg < Be (4) Li < Na < K < Ca
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Ion Sizes - CATIONS
Does the size go up or down when an atom loses an electron to
form a cation?
+
Li, 152 pm
Li+, 60 pm
CATIONS are SMALLER than the parent atoms.
Ion Sizes - ANIONS
Does the size go up or down when gaining an electron to form
an anion?
F, 64 pm
-
F-, 136 pm
ANIONS are LARGER than the parent atoms.
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Trends in Ion Sizes
CATIONS
ANIONS
(59 pm)
(207 pm)
Trends in relative ion sizes are the same as atom sizes.
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2. Ionization Energy(电离能)--I
The minimum amount of energy required to remove an electron
from the ground state of an atom or ion in the gas phase is called
the ionization energy(电离能)(I). The first ionization energy of
magnesium(Mg) is represented by the equation:
Mg(g)
Mg+(g) + e
I1 = 738 kJ/mol
Magnesium's second and third ionization energies are
Mg+(g)
Mg2+(g) + e
I2 = 1450 kJ/mol
Mg2+(g)
Mg3+(g) + e
I3 = 7730 kJ/mol
Note that the amount of energy required increased for the second
ionization. Because of coulombic (库仑) attraction, it is more
difficult to remove an electron from a positively charged ion than
it is to remove one from a neutral atom.
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This third ionization energy of magnesium is 7730 kJ/mol. This
represents a significant increase over the first two. The reason
for this jump(突跃) in ionization energy is that magnesium has
only two valence electrons. The third ionization of magnesium
removes an inner-shell, or “core” electron(核电子).
Note that the first column of Table above shows a general
upward trend in first ionization energies from left to right across
the third period of the periodic table.
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•The I1 generally increase from left to right across the
periodic table. The I1 decrease from top to bottom with a
group.
•Energy ‘cost’ is very high to remove an INNER SHELL e.
This is why oxidation. No. = Group No.
•Metals lose electrons more easily than nonmetals. Metals
are good reducing agents.
•Nonmetals lose electrons with difficulty.
Question
1. Which pair is not correctly arranged in order of increasing I1?
(1) O < F
(2) Ar < Ne
(3) Mg < N
(4) Br < Se
2. Which do you expect to have the highest third ionization energy?
(1) N (2) C (3) Be (4) P
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3. Electron Affinities (电子亲合能)--E
Essentially the opposite of ionization, electron affinity(电子亲
合能)(E) is the energy released for the addition of an electron to
an atom in the gas phase.
Cl(g) + e-
Cl-(g)
E1 = +349 kJ/mol
ATTENTION: In here, EXOTHERMIC(放热) -- E is positive
(+) and ENDOTHERMIC(吸热) -- E is negative (-).
•A few elements GAIN electrons to form anions.
•Electron affinity is the energy released when an atom gains
an electron.
A(g) + e-  A-(g)
E1 = - (EA- - EA )
•If EA- < EA then the anion(A-) is more stable and there is an
exothermic reaction.
•The greater the DE, the easier the element gain electron.
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Trends in Electron Affinity
•Affinity for electron increases across a period.
•Affinity decreases down a group.
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4. Electronegativity (电负性) --X
X is a measure of the ability of an atom in a
molecule to attract electrons to itself.
Concept proposed by Linus Pauling (190194).Nobel prizes:Chemistry (54), Peace (63)
4
N
3.5
F
O
Cl
C
3
2.5
Linus Pauling
(1901-1994)
H
Si
2
P
S
1.5
1
• F has maximum X.
• 同一周期：从左到
右电负性依次增大。
• 同一主族：从上到
下电负性依次变小，
0.5
0
1
2
3
4
5
6
7
8
9
10 11 12 13 14 15 16 17 18
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Electronegativity and Bond Polarity
(电负性与键的极性)
Which bond is more polar ? (has larger bond DIPOLE偶极)
O—H
O—F
X(A) - X(B)
DX
3.5 - 2.1
1.4
3.5 - 4.0
0.5
H
O F
X 2.1 3.5 4.0
DX (O-H) > DX (O-F)
Therefore OH is more polar than OF
Also note that polarity is “reversed.”
O—H
-δ + δ
O—F
+δ -δ
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