Download Electronic Structure and Covalent Bonding

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