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AP Chemistry Electronic Structures Worksheet
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Electronic Structures of Atoms / Periodic Trends / Ionic Bonding / Solids / Phase Changes
H Advanced ChemistryUnit 3
Objectives #1-3 Atomic Theory
*review of electromagnetic radiation characteristics:
(diagrams)
Examples of Electromagnetic Radiation
Objectives #1-3 Atomic Theory
frequency, wavelength, energyfrequency vs. wavelength (inverse
relationship)frequency vs. energy (direct relationship)wavelength vs. energy (inverse
relationship)c=fλ (c = speed of light in m/s, f =
frequency in Hz (1/s), λ = wavelength in m)
Max Planck (1858-1947)
Objectives #1-3 Atomic Theory
E = hf or hc/λ h = Planck’s Constant (energy for waves)
Albert Einstein (1879-1955)
Objectives #1-3 Atomic Theory
E = mc2 (energy for particles)*Wave particle-dualityMatter has wave and particle
characteristics; acts as particle when interacting with matter; acts as wave when travelling through space
Louis de Broglie (1892-1987)
Derivation of de Broglie’s Equation:Ewaves = Eparticles
(examples)
Objectives #1-3 Atomic Theory
*Work Function (Photoelectric Effect)Φ = hfo
Φ = work functionminimum energy required to remove
electron from surface of metalfo = threshold frequency
minimum frequency required to remove electrons from surface of metal
(examples)
Photoelectric Effect (Albert Einstein Nobel Prize 1921)
Niels Bohr (1892-1987)
Objectives #1-3 Atomic Theory
*Bohr’s Equation:E = -2.178 X 10-18 J (z2/n2) OR ∆E = -
2.178 X 10-18 J (z2) X (1/n2final –
1/n2initial)
used for: determining energy changes when electrons change energy levels
for hydrogen; z = 1(examples)
Johannes Rydberg (1854-1919)
Objectives #1-3 Atomic Theory
*Rydberg Equation:1/λ = 1/91 nm (1/nL
2 – 1/nH2)
*used for: determining wavelength of photons released change energy levels
(Examples)*relationships of answers: the greater
the energy difference, the smaller the wavelength
Erwin Schrodinger (1887-1961)
Objectives #4-5 The Quantum Numbers and Quantum States
*Review of Quantum Theory:1. Quantum NumbersA. Principle (n)*energy level of shell of electron*n = 1,2,3…..*(old system) n = K, L, M, ….*indicates the number of sublevels in
energy level
Illustration of Principle Quantum Number
Objectives #4-5 The Quantum Numbers and Quantum States
B. Orbital (l)*indicates orbital shape*l = 0, n-1*s, p, d, f
Illustration of Orbital Quantum Number / Orbital Shapes
Objectives #4-5 The Quantum Numbers and Quantum States
C. Magnetic (ml)
*indicates orientation of orbital in space
*ml = 0, +/-1 1
*the number of ml values indicate the number of orbitals within sublevel
Illustration of Magnetic Quantum Number
Objectives #4-5 The Quantum Numbers and Quantum States
D. Spin (ms)
*indicates spin of electron*+1/2 or -1/2*allows for up to 2 electrons per orbital*s 2 electrons p 6 electrons d 10 electrons f 14 electrons
Illustration of Spin Quantum Number
Objectives #4-5 The Quantum Numbers and Quantum States
*Quantum Number Sets for Electrons in Atoms:
Illustration of Quantum States
Objectives #4-5 The Quantum Numbers and Quantum States
(examples of quantum number states problems)
Objectives #7-9 Electron Configurations of Ions / Orbital Filling and Periodic Trends
*valence electrons and occasionally the electrons contained within the d sublevel are involved in chemical bonding
*atoms tend to lose or gain electrons in such a way to complete octets (s2p6) or to from similarly stable arrangements called pseudo noble-gas configurations
(examples)*Orbital Filling and Periodic Trends1. Ionization EnergyGroup 1Group 2Group 15Group 17Group 18
Trends in Ionization Energy
Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy
*ionic bonds involve the transfer of valence electrons from a metal to a nonmetal
*the tendency for a metal to lose electrons depends on its ionization energy and the tendency of a nonmetal to gain electrons depends on its electron affinity
*the loss of an electron requires a gain of energy and is therefore an endothermic process
example: Na + energy › Na+1 + e-
*the gain of an electron releases energy and is therefore an exothermic process
example: Cl + e- › Cl-1 + energy
Formation of Sodium Chloride
Formation of Crystal Lattice
Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy
*combinations of elements with low ionization energies and high electron affinities will cause an extremely exothermic reaction and generally be the most stable*example: Na(s) + Cl2(g) › NaCl(s) + energy
*the energy produced when the ionic bond forms is referred to as the lattice energy; this energy is also equal to the energy required to break apart the ionic bond
*chemical bonding not only involves a rearrangement of electrons but it also involves changes in energy
Illustration of Born-Haber Cycle
Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy
*the formation of an ionic compound; such as the following reaction:
Na(s) + 1/2Cl(2)(g) › NaCl(s) + ∆Hof =
-410.9 kJwhere ∆Ho
f refers to the standard heat of formation which is the energy change involved when a compound is formed from its elements, involves a series of energy changing steps known as the Born-Haber cycle
*these steps are as follows:
Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy
1. Sublimation or Vaporization of nongaseous reaction components:
Here: Na(s) › Na(g) 108 kJ which represents the energy of sublimation or vaporization (an endothermic process)
2. Breaking the bonds of any gaseous components:Here 1/2Cl2(g) › Cl(g) 122 kJ which represents the
dissociation energy (an endothermic process)(now that all reactants are gaseous, ions must be
formed)
Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy
3. Formation of the positive ion:Here: Na(g) › Na+1
(g) + e- 496 kJ which represents the ionization energy (an endothermic process)
4. Formation of the negative ion:Here: Cl(g) + e- › Cl(g)
-1 -349 kJ which represents the electron affinity affinity (an exothermic process)
5. Formation of the ionic compound by combining the two ions formed together:
Here: Na+1(g) + Cl(g)
-1 › NaCl(s) -788 kJ
which represents the lattice energy (an exothermic process)*the overall energy change, ∆Ho
f, is equal to the sum of all these changes:
∆Hof = ∆Ho
fNa + ∆HofCl + IENa -EACl - ∆Hlattice
Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy
*Relationship of lattice energy and ionic chargeConsider the following lattice energy data from the above
example problems:NaCl 788 kJLiF 1030 kJMgCl2 2326 kJ
KCl 701 kJ**strength of ionic bonds:KCl ‹ NaCl ‹ LiF ‹ MgCl2**size of ions:LiF ‹ MgCl2 ‹ NaCl ‹ KCl
**charge of ions:Na +1, Cl -1 Li +1, F -1 K +1, Cl -1 Mg +2, Cl-1
Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy
**formula: Eel. = KQ1Q2/d where “K” is a constant of electrical charge, where “Q1” and “Q2” are the charges of the ions involved, where “d” distance separating the ions
As the magnitude of the charges in an ionic compound increases, the lattice energy increases (affects lattice energy the most)
As the size of the ions involved decrease, the lattice energy increases
(examples)
Objectives #13-14 Phase Changes and Phase Diagrams
Objectives #13-14 Phase Changes and Phase Diagrams
*Important Parts of a Heating/Cooling Curve(see curve in lecture guide)A. Specific heat of solid added (endothermic)B. Specific heat of liquid added (endothermic)C. Specific heat of gas added (exothermic)D. Melting (heat of fusion added) (endothermic)E. Freezing (heat of solidification released)
(exothermic)F. Boiling (heat of vaporization added) (endothermic)G. Condensing (heat of condensation released)
(exothermic)
Objectives #13-14 Phase Changes and Phase Diagrams
*where the graph is increasing or decreasing, specific heat is being added or subtracted (which results in the temperature changing)
*where the graph is not changing, a phase change is occurring and these is no change in the temperature of the substance
*key equations:to change temperature: Q =mc∆tto change phase:∆H = moles of material X molar heat of phase
change(examples)
Objectives #13-14 Phase Changes and Phase Diagrams
Objectives #13-14 Phase Changes and Phase Diagrams
*Interpreting Phase Diagrams*a phase diagram allows one to determine the phase that a
substance is in at a given temperature and pressure*the phase diagram only shows one substance in its various
phases*a typical phase diagram:(see diagram in lecture guide)*the boundaries between different phase regions represent areas
of equilibrium in which the two phase changes are occurring at the same rate; for example at the liquid – gas boundary, molecules of gaseous vapor are moving into the liquid phase while molecules of liquid are moving into the gaseous phase
*if a point on the diagram does not fall on any line, only one phase is present
Objectives #13-14 Phase Changes and Phase Diagrams
*the following lines on the graph represent phase change boundaries:Line Segment Phase Change
Boundary
A-B Liquid – gas (vaporization ↔ condensation
A-C Solid – gas (sublimation ↔ deposition)
A-D Solid – liquid (melting ↔ freezing)
Objectives #13-14 Phase Changes and Phase Diagrams
Critical Point the endpoint of the vapor-pressure curve; beyond this point of critical temperature and critical pressure, the liquid and gas phases can not be distinguished from each other
Normal Boiling Point the location on the vapor-pressure curve where the vapor pressure is 1 atm
Normal Melting Point the location on the solid-liquid curve where the melting (freezing) point is at 1 atm
Triple Point the point where all 3 curves intersect and all three phases are in equilibrium
Objectives #13-14 Phase Changes and Phase Diagrams
*some general relationships and observations to note:
*if the solid-liquid line curves to the right with increasing pressure, then the melting point is also increasing (this is the norm)
*if the solid-liquid line curves to the left with increasing pressure, then the melting point is decreasing (this is not the norm; water follows this pattern)
Illustration of Phase Diagram for Water
Objectives #15-17 Structure of Solids, Properties and Applications
*solids come in two general types: crystalline or amorphous*crystalline solids contain particles arranged in a well-defined
pattern called a crystal lattice with flat faces and definite angles; examples include NaCl or diamonds
*amorphous solids lack any well defined structure; examples include wax, rubber, or glass
*the crystal lattice of a crystalline solid, which is a three dimensional array showing the location of individual particles, is actually made up of many repeating individual parts called the unit cell; for example the repeating pattern on wall paper
*the simplest common type of unit cell is the cubic unit cell where all sides are equal in length and consist of all 90o angles
Objectives #15-17 Structure of Solids, Properties and Applications
*the three types of cubic unit cells are:primitive or simple cubic – lattice points
only occur at the corners of unit cellbody-centered cubic – lattice points
occur at the corners and in the centerface-centered cubic – lattice points
occur at the corners and faces
Types of Cubic Unit Cells
Objectives #15-17 Structure of Solids, Properties and Applications
*except for the atom in the center of the body-centered cubic unit cell, all of the atoms located at the lattice points are actually shared to various degrees by other unit cells; in order to determine the net number of atoms in a unit cell and thus its chemical formula, one must know the fraction of an atom that occurs in each position of the unit cell as follows:
Objectives #15-17 Structure of Solids, Properties and Applications
Position in Unit Cell Fraction in Unit Cell
Center 1
Face ½
Edge ¼
Corner 1/8
Objectives #15-17 Structure of Solids, Properties and Applications
(examples)*X-Ray Crystallography*the layers of atoms in the crystal lattice acts as an
effective diffraction grating that can be used to scatter a beam of x-rays
*the diffraction, or scattering, of the x-rays produces a characteristic pattern of light and dark areas on a x-ray detector
*by examining the areas of light and dark and measuring the angles of deflection, the original crystal structure of the material can be deduced
*this analytical technique has been used to determine the structure of DNA and other molecular crystals
X-Ray Crystallography
Example of Body Centered Unit Cell in Cesium Chloride
Example of Face Centered Unit Cell in Sodium Chloride
Objectives #15-17 Structure of Solids, Properties and Applications
*Types of Bonding in Solids and their Influence on Properties
Type of Solid
Form of Unit Particle
Forces Between Particles
Properties Examples
Molecular Atoms or molecules
London, DD, HB
Fairly soft, low melting points, poor conductors
Gases, sugars, dry ice
Covalent-Network
Atoms connected in a network of covalent bonds
Covalent bonds
Very hard, high melting points, poor conduction
Diamond, quartz
Ionic Positive and negative ions
Electrostatic attractions
Hard and brittle, high melting point, poor conductor
Salts such as NaCl
Objectives #15-17 Structure of Solids, Properties and Applications
*Types of Bonding in Solids and their Influence on Properties
Metallic Atoms Metallic bonds
Soft to very hard, low to high melting point, malleable and ductile, excellent conductor
Metallic elements such as copper
Dry Ice – Example of Molecular Solid
Diamond – Example of Covalent Network Solid
Sodium Chloride – Example of Ionic Solid
Copper – Example of Metallic Solid
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