CHEMISTRY 1: EXAM 3
Chapter: 6 Electronic Structure of Atoms
Chapter 7: Periodic Properties of Elements
Chapter 8: Basic Concepts of Chemical Bonding
Chapter 9: Molecular Geometry and Bonding Theories
CHAPTER 6: ELECTRONIC
STRUCTURE OF ATOMS
• 6.1: The Wave Nature of Light
• 6.2: Quantized Energy and Photons
• 6.3: Line Spectra and the Bohr Model
• 6.4: The Wave Behavior of Matter
• 6.5: Quantum Mechanics and Atomic Orbitals
• 6.6: Representation of Orbitals
• 6.7: Many Electron Atoms
• 6.8: Electron Configurations
• 6.9: Electron Configurations and the Periodic Table
6.1 THE WAVE NATURE OF LIGHT
• Electromagnetic Radiation:
• Carries energy through space as waves Radiant Energy
• All electromagnetic radiation moves at the speed of light
•
•
•
•
𝑐𝑐 = 𝜆𝜆𝜈𝜈
c = 3.00 * 10^8 m/s
λ = wavelength (nm)
ν = frequency (Hz or 1/s)
6.2: QUANTIZED ENERGY AND
PHOTONS
• Wave model cannot explain all behavior or energy
• Black Body Radiation: light emitted from hot objects
• Photoelectric Effect: electrons being emitted from metal surfaces
• Emission Spectra: emission of light from excited gas atoms
• Quantum: Amount of energy that can be emitted or absorbed as
electromagnetic radiation
• Energy can only be emitted in strict quantum
• Work Function: energy required to overcome attractive forces holding
it in a metal form
• 𝐸𝐸 = ℎ𝜈𝜈
• E = Energy of a single quantum or photon
• ℎ = 6.626 ∗ 10−34 J-s
6.3: LINE SPECTRA AND
THE BOHR MODEL
• Continuous Spectrum: contains light of all wavelengths
• Line Spectrum: Single wavelength
• 𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅 𝐸𝐸𝐸𝐸𝐸𝐸𝐸𝐸𝐸𝐸𝐸𝐸𝐸𝐸𝐸𝐸:
1
𝜆𝜆
= 𝑅𝑅ℎ
• Rh = 1.096776*10^7 1/m
1
𝑛𝑛12
−
1
𝑛𝑛22
where n2>n1 and positive
6.3: LINE SPECTRA AND
THE BOHR MODEL
• Bohr Model
1. Distinct Orbitals with different Radii
2. Can exist only in specific orbitals with allowed energy states
3. Emission and Absorption only occurs when energy states change
• Only explains Hydrogen
6.4: THE WAVE BEHAVIOR OF
MATTER
• 𝜆𝜆 =
ℎ
𝑚𝑚𝑚𝑚
• mv = momentum
• 𝑃𝑃𝑃𝑃𝑃𝑃𝑃𝑃𝑃𝑃𝑘𝑘 ′ 𝑠𝑠 𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶: ℎ = 6.626 ∗ 10−34 𝐽𝐽 − 𝑠𝑠
• Uncertainty Principle
• Δ𝑥𝑥 ∗ Δ(𝑚𝑚𝑚𝑚) ≥
ℎ
4𝜋𝜋
6.5: QUANTUM MECHANICS AND
ATOMIC ORBITALS
• Orbital: Distribution of electrons
1. Principal Quantum Number: increase in n = higher energy =less tightly bound to
nucleus
• Same Orbital =Electron Shell
2. Angular Momentum Quantum Number = s=0, p=1, d=2, f=3
3. Magnetic Quantum Number: Describes orientation
6.6 REPRESENTATION OF ORBITALS
• S Orbital
• Lowest Energy Orbital
• Spherically Symmetric
• Radial probability Function
n=1
6.6 REPRESENTATION OF ORBITALS
n=2
n=3 or more
6.7: MANY ELECTRONS ATOMS
Electron Spin
1
1
• + ,−
2
2
Pauli Exclusion Principle
• Filled in order of increasing energy
6.8: ELECTRON CONFIGURATIONS
• Hund’s Rule: Put one electron in
each orbital until the full
6.9: ELECTRON CONFIGURATION
AND THE PERIODIC TABLE
6.9: ELECTRON CONFIGURATION
AND THE PERIODIC TABLE
• Example:
• Ga
[Ar] 3d10 4s2 4p1
• Pb
[Xe] 4f14 5d10 6s2 6p2
CHAPTER 7: PERIODIC PROPERTIES
OF THE ELEMENTS
• 7.1: Development of the Periodic Table
• 7.2: Effective Nuclear Charge
• 7.3: Sizes of Atoms and Ions
• 7.4: Ionization Energy
• 7.5: Electron Affinities
• 7.6: Metals, Nonmetals and Metalloids
• 7.7: Group Trends and the Active Metals
• 7.8: Group Trends for Selected Nonmetals
Valence
Orbitals
(Columns)
7.1: DEVELOPMENT OF THE
PERIODIC TABLE
Important People
• Mendeleev
• Meyer
• Rutherford
• Moseley
7.2: EFFECTIVE NUCLEAR CHARGE
• 𝑍𝑍𝑒𝑒𝑒𝑒𝑒𝑒 = 𝑍𝑍 − 𝑆𝑆
• Zeff = Effective Nuclear
Charge
• Net attraction to the Nucleus
• Z=Number of Protons in the
Nucleus
• S=Screening Constant
Number of core electrons
GENERAL PERIODIC TRENDS
7.4 IONIZATION ENERGY
• Minimum Energy required to remove an electron from the ground
state
• Greater the Ionization Energy the harder it is to remove an electron
7.4 IONIZATION ENERGY
7.5: ELECTRON AFFINITIES
• Energy change that occurs when an electron is added to a gaseous
atom
• Attraction Energy released when added
7.6: METALS, NONMETALS , AND
METALLOIDS
7.6: METALS, NONMETALS,
METALLOIDS
• Metals
• Shiny luster
• Low ionization energy likes positive ions
• Nonmetals
•
•
•
•
Vary in appearance
Poor conductors of heat and electricity
Low melting points
High electron affinity likes to be negative
• Metalloids
• Properties are an intermediate between metals and nonmetals
Group 1A:
Alkali Metals
Group 2A:
Alkaline Earth
Metals
7.7: GROUP TRENDS FOR THE
ACTIVE METALS
Group 7A:
Halogens
Group 8A:
Noble Gases
CHAPTER 8: BASIC CONCEPTS OF
CHEMICAL BONDING
• 8.1: Chemical Bonds, Lewis Symbols and the Octet Rule
• 8.2: Ionic Bonding
• 8.3: Covalent Bonding
• 8.4: Bond Polarity and Electronegativity
• 8.5: Drawing Lewis Structures
• 8.6: Resonance Structures
• 8.7: Exceptions to the Octet Rule
• 8.8: Strengths of Covalent Bonds
8.1: CHEMICAL BONDS, LEWIS
STRUCTURES AND THE OCTET RULE
• Chemical Bonds: Two atoms or ions are strongly attached
• Ionic Bond: metal and nonmetal bonded
• Covalent Bond: two non metals bonded, share electrons
• Metallic Bonds: two metals bonded, electrons are free to move
• Lewis Symbols
• Dots represent valence electrons
• Octet Rule
• Atoms will gain or lose electrons until they have 8 valence electrons
8.2: IONIC BONDING
• Metal and Nonmetal
• Electron Transfer
• Losing electrons is
endothermic
• Gaining electrons is
exothermic
• Lattice Energy: Energy
required to separate a mole of
solid ionic compound into ions
• 𝐸𝐸𝑒𝑒𝑒𝑒 =
𝜅𝜅𝑄𝑄1 𝑄𝑄2
𝑑𝑑
• k=8.99*10^9 J*m/C^2
8.3: COVALENT BONDING
• Two nonmetals bonded
• Shared electrons
• Brittle, high melting points, crystalline
• Can have double and triple bonds
8.4: BOND POLARITY AND
ELECTRONEGATIVITY
• Polar and Nonpolar Bonds depend on
electronegativity
• Electronegativity: Ability to attract Electrons
• Dipole Moment = 𝜇𝜇 = 𝑄𝑄𝑄𝑄
• Q= charge, r = distance
8.5: DRAWING LEWIS STRUCTURES
8.6: RESONANCE STRUCTURES
8.7: EXCEPTIONS TO THE OCTET
RULE
• BF3
• PCl5
• SF4
• PO4
8.8: STRENGTHS OF COVALENT
BONDS
∆𝐻𝐻𝑟𝑟𝑟𝑟𝑟𝑟 = � 𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏 𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒 𝑜𝑜𝑜𝑜 𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏 𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏 − � 𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏 𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒𝑒 𝑜𝑜𝑜𝑜 𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏𝑏 𝑓𝑓𝑓𝑓𝑓𝑓𝑓𝑓𝑓𝑓𝑓𝑓
Increase in number in bonds = bonds grow shorter
CHAPTER 9: MOLECULAR
GEOMETRY AND BONDING
THEORY
THE VSEPR MODEL
CONSTRUCTIVE AND DESTRUCTIVE
HOMONUCLEAR
BONDING
HETERONUCLEAR
BONDING