General Chemistry I - CHEM 123

Learning Goals/Objectives:

 After class, reading, and practice, the students will be able to

1

8/28

R A2

Welcome, Distribute Lab Manual;
Online Resources;

Intro - What Chemists Do:  Synthesize, Purify, Analyze;

Atoms, Molecules, Mixtures, and Separations

Understanding Chemistry: Macroscopic, Nanoscopic, Mathematical, and Symbolic

• state what a chemists does
• differentiate between an element, compound and molecule
• differentiate between homogeneous (solutions) and heterogeneous mixtures and give examples
• describe different methods to separate mixtures
• state the difference between physical and chemical processes

2

9/1

M A4

-review syllabus; 

-review mixtures,  their complexities and their separations; chemical vs physical reactions

-  Understanding chemical and biological complexity: Macroscopic, Nanoscopic, Mathematical, and Symbolic

 -Problem Solving

- Dimensional Analysis

• state the differences between macroscopic, nanoscopic, symbolic, and mathematical descriptions of matter
• give nanoscopic and mathematical descriptions of a measurable macroscopic property of a gas
• describe general features of problem solving
• use dimensional analyses to solve problems involving unit conversion
• recognize when to use dimensional analyses to solve mathematical problems.

3

9/3

W A6

- Is Lake Sag filled with oil and sulfuric acid?

- Laws of Matter;

- Dalton's Atomic Theory;

- The Divisible Atom:  cathode rays/demo;
- elephants,  mosquitoes, and e/m

• state the laws of mass conservation, definite composition, multiple proportion and how they led to Dalton's atomic theory
• state the differential between empirical laws and scientific theories
• draw a sketch of a cathode ray tube
• identify dependent (varying and constant), independent variables in the cathode ray experiment,
• state what  variable(s) or derivative properies of an electron can be determined from a cathode ray tube

4

9/5

F A2

- Millikan Oil Drop, calculation of e;
- Plum-Pudding model;

- Rutherford, nuclear atom; planetary model;

- Atomic Masses; isotopes and  weighted
   averages;

- first ionization energy – shells;

- Lewis Dot structure

- intro to bonding

• state the origin/types of forces on a falling object in a gravitation field and a charged object in the presence of another charged object, using arrows to show the direction of the forces
• write expressions for the gravitation and
• draw a sketch of a Millikan oil drop apparatus, and Rutherford's gold foil experiment.
• state what macroscopic variable(s) or derivative property of an electron can be determined from the Millikan oil drop experiment, and from the Rutherford experiment
• describe difference between isotopes of an element, how their atomic masses are determined, and how the average atomic mass of an element is calculated
• define 1st ionization energy for a gas phase atom
• interpret trends in 1st ionization energies of elements as a function of the arrangement of electrons around atoms
• explain how electronic structures around atoms of a given element affects the likelihood that an element will gain or release an electron to form an ion and form ionic bonds, or share electrons and form covalent bonds
• draw Lewis dot structures of atoms and use them to predict simple covalent compounds of an element with H and ionic bonds with oppositely charged ions

5

9/9

T A4

- Lewis Dot structure

- ionic, covalent, and polar covalent bonding

- FON and electronegativity;
- Counting Electrons:  formal charge; partial charge; octet/duet rule;

- Lewis Structures 

• explain how electronic structures around atoms of a given element affects the likelihood that an element will gain or release an electron to form an ion and form ionic bonds, or share electrons and form covalent bonds
• draw Lewis dot structures of atoms and use them to predict simple covalent compounds of an element with H and ionic bonds with oppositely charged ions

6

9/11
R A6

Quiz:  Lewis Structures

- Lewis Structures; Salts, Acids, Acid reactivity and Lewis structures;

- pushing electrons: Reaction Mechanisms;

- acids and abases

- polyatomic or molecular ions; naming molecular ions 

• use Lewis structures to predict and draw reactions mechanisms (using curved arrows) for acid/base reactions
• define acids and based based on electron pushing mechanisms
• write the formula and draw the Lewis structures for common oxyacids
• draw the Lewis structures and give names for molecular ions derived from their parent oxyacid
• write formula and names for salts of metals and molecular ions from oxyacids

7

9/15

M A2

Qz: Reaction Mechanism - Electron Push

- Salts of molecular ions;

Chapt 3

- Chemical Equations; balancing chem equations;

- mole and molar mass

- simple stoichiometry:  gravimetric analyses - prep for Lab 3 (9/16)

• write and balance a chemical equations using symbolic and nanoscopic representations for reactants and products
• define and calculate molecular mass, mole, and molar mass for any chemical substance
• using symbolic and nanoscopic representations in chemical equations, solve mathematical stoichiometry problems using dimensional analyses

8

9/17 W A4

-  Examples typical types of stoich: - limiting/excess reagent, % composition, empirical formula, combustion analysis

• develop a problem solving strategy for solving stoichiometry problems involving limiting/excess reagent, combustion analyses, % composition, and empirical formula. 
• successfully solve stoichiometry problems involving limiting/excess reagent, combustion analyses, % composition, and empirical formula. 

9

9/19

F A6

-  Examples typical types of stoich: - limiting/excess reagent, % composition, empirical formula, combustion analysis

• explain the similarities between the extensive macroscopic properties of density and solution concentration
• define the following types of solutions concentrations:  % (g/V) and molarity (M)
• calculate then number of moles of a substance from density and solution concentration information
• solve problems involving how to make solutions of different molarities from solid reagents and from concentrated stock solutions.

10

9/23

T B2

• solve solution stoichiometry problems given volume and/or molarities of reactant/products in aqueous solution

11

9/26  F B4

12

9/30

T B6

-Review Test; 

- Lb 4: % Acetic Acid Solutions - are they the same?

- What's in the Beaker; Electrolytes (Lab 2 Data)

- Precipitation Reactions; Molecular, Ionic, and Net Ionic Equations; Solubility Rules (See pg 112; also OLSG) 

• after reviewing test 1, start to develop new strategies for problem solving and test taking
• use mathematical analyses to compare the likelihood that calculated macroscopic observable values for two different samples are the same or different;
• define and give examples of strong and weak electrolytes, and soluble and insoluble substances in water
• better correlate and connect class and lab experiences in developing critical thinking and analysis skills
• write the formulas, draw Lewis structures, and draw nanoscopic representation of the actual chemical species found in water after the addition of ionic salts and molecule solids, liquids and gases
• write the formulas for soluble and insoluble chemical species present after the addition of two aqueous solutions
• state the simple solubility rules
• using solubility rules, write molecular, ionic, and net ionic equations for precipitation reactions.

13

10/6

M B2

- Overview:  Precipitation, Acid/Base and Redox reactions; - - Molecular, Ionic, and Net Ionic Equations; Solubility Rules (See pg 112; also OLSG)
- Acid/Base Rx.

- Strong and Weak Acids and bases; conjugate acids/bases;

-Redox reactions: oxidizing and reducing agents

14

10/8 W B4

-Acid/Base:  strong Acid/bases, Weak acid bases

- ranking acids and bases in reactions: 

-Start Redox reactions

• given paired reactions with a common acid or base, rank order all acids and bases in the reactions based on your knowledge of common strong acids/bases and weak acids and bases
• given structural similar acids, predict trends in the strength of the acid by analyzing the Lewis structure of the conjugate bases
• state the relatives meanings of high or low reactivity, energy and compared to strong and weak bases.

15

10/10 F B6

- Redox reactions;

- Strength oxid agent/reducing agent (EN, Core Charge)

- predicting direction of acids/base and redox reaction from charts

- oxidation # (done in Lab 10/9:  analysis of Cu ammonia complex; 

- balance redox eq. - acid.  (You will only need to know the 1/2 rx method for balancing Redox equations, and only in acidic solutions)

• predict reactions products of acid/base and redox reactions from tables showing the relative strength of acids/bases and oxidizing/reducing agents
• use the 1/2 reaction method to balance redox reactions in acidic conditions.

16

10/14 T B2

-Name that reaction:  
 

Intro to Quantum Chemistry: the Dalton, the Plum Pudding (Thompson), The Planetary Atom (Rutherford) - review;

Forces, acceleration, oscillating charges and fields;

The instability of the planetary atom
 

• develop a conceptual framework to identify and predict the products of precipitation, acid/base and redox reactions
• state how a net force can change the motion of a fixed object or one moving at constant velocity
• state why the planetary model of atoms violates classical physical views of an orbiting electron as an oscillating charge

17

10/16 R B4

Quiz:  OLSG Web Links for 10/16

- properties of waves:  reflection, refraction, diffraction, interference (constructive/destructive), superposition (adding);

- light as a wave

- Force, Energy, and work;

• define force, energy, and work and state how they are interrelated
• define and differentiate potential and kinetic energy
• state the properties of waves
• write formulas and explain the variables in the formulas for:
  • force as a function of mass
  • gravitational and electromagnetic force
  • speed of a wave
  • momentum of a particle
  • gravitational and electromagnetic energy

18

10/20 M B6

- Problem with planetary model:  discrete energies, radii, instability;  fields

- light as a particle:  Photoelectric effect

- light interactions with matter:  absorption/emission spectra gases;

-

• state two characteristics of how waves carry energy
• interpret a graph of E vs freq for a photoelectron and state experimental observations concerning photelectrons that can NOT be explained by assuming light is a wave
• explain how experimental observations concerning photelectrons can be explained by assuming light is a particle (photon)
• qualitatively describe atomic absorption and emission spectra

19

10/22 W C2

-  Empirical Equation to predict line spectra: Rydberg Eq;
-   Theoretical Equation to predict line spectra:  Bohr Equation/Bohr Atom

- electron as a wave:  deBroglie, Bohr, Schrodinger, and Heisenberg

• explain how discrete (vs continuous) atomic emission and absorption line spectra are consistent with a planetary model of the atom when the orbital energies of the electrons are quantized;
• explain why the planetary model of the atom is incompatible with classical physics
• explain how Bohr used standing waves to model electrons around H atoms to create a "stable atom"
• interpret the deBroglie equation and from the Schrodinger Equation, the wave function () and W²
• explain how quantum numbers n, l, ml, and ms are derived and their relationships to s, p, d, and f orbitals
• write box and shorthand notion for the electron configuration of atoms

20

10/24 F C4

 - review wavelike properties of atoms;

- Orbitals; quantum numbers, and electron Configurations; 

• explain how quantum numbers n, l, ml, and ms are derived and their relationships to s, p, d, and f orbitals
• write box and shorthand notion for the electron configuration of atoms

21

10/28 T C6

-  UV/Vis absorption/emission; ESR, NMR;

review potential and kinetic energies of electrons

- review Columbs law

- Zeff;

-  Periodic Trends in atomic and ionic radii; 

• define Zeff
• Using Zeff values and a knowledge of electron configurations, predict trends in atom and ion size, ionization energy, electron affinity,

22

10/30

R C2

- periodic trends in ionization energy, electron affinity

- ionic bonds; Lattice Energy, Born-Haber cycle;

• calculate, given energies for a series of individual reactions that would add up to the formation of a salt like NaCl from its component elements, Na(s) and Cl₂ (g), the energy change on salt formation,  and the energy change (lattice energy) required to break up the salt into separate ions in the gas phase;
• Using Coloumbs law, predict for similar salts the relative size of their lattice energies

23

11/3

F C4

24

11/5 W C6

• convert Lewis structures of of organic molecules that contain all atoms and lone pairs into line drawings;  
• convert line drawings of organic molecules into full Lewis structure with the correct number of bonds, atoms and lone pairs;
• draw the Lewis structures for molecules with expanded octets, deficient octets, free radicals
• draw Lewis structure with different resonance structures and determine which is most stable (lowest in energy) and hence most likely to exist

25

11/7 F C2

- Lewis Structures 2;

- Bond Energy, Quant Mechanical Explanation covalent
    bond;

- VSEPR; (2008: on 11/5)

• draw Lewis structures of triatomic molecules with different atom connectivities and determine which is most stable (lowest in energy) and hence most likely to exist
• predict the geometrical arrangement of electron clouds emanating from a central atom using the VSEPR model
• using dotted lines/wedges and bolded line/wedges to depict the 3D geometry of electron clouds emanating from the central atom in which the clouds are arranged in a tetrahedral, trigonal (triangular) bipyramidal, and octahedral geometries.
• predict the geometric shapes and bond angles for atoms bonded to a central atom from the geometry of the electron clouds emanating from the central atom

26

11/11 T C4

- VSEPR, Polarity, Dipoles;

• define a dipole mathematically and determine from the arrangement of electron clouds around a central atom if the atoms has a dipole
• draw a  +--> symbol on a Lewis structure of a molecule indicating the direction of the dipole.

27

11/13 R C6

- Review VSEPR geometry

- Qz: VSEPR

- Valence Bond Theory - overlap of atomic orbitals

- Hybridization

• Using box diagrams for electronic configurations and drawn shapes for atomic orbitals, show how single and double bonds can form between two atoms by overlap of atomic orbitals in simple diatomic molecules;
• draw orbital pictures to show the different between sigma and pi bonds
• using box diagrams for electronic configurations and drawn shapes for orbitals to show how equivalent hybrid atomic orbitals consistent with the observed and VSEPR-predicted geometry can be generated through combination of the same number of s and p orbitals
• differentiate between sp³, sp² and sp hybridized orbitals and how they can be used to form sigma bonds of the correct geometry with atomic orbitals of H or hybridized orbitals of other atoms.

28

11/17 M D2

- Homework Qz:  O hybridization in H₂CO

- hybridization and VSEPR:  sp, sp², sp³,  sp³d, sp³d²

-  Polarity and Dipole Moment

- Molecular Orbital Theory

- Bond Order

- Spartan Electron Densities:CH₄ to HF, H₂ to LiH,  H₃O⁺ and formal charges

-  Bond Energies and Trends:  H-H and H-F, N₂ and H₂

-  Avg bond energies: H₂, F₂, I₂, HF, and HI:  electronegatiivty

• explain the difference between an atomic orbital and a molecular orbital
• using a box diagram show how two equivalent atomic orbitals of equivalent energy can be combined to form two molecular orbitals, one a bonding molecular orbital and one an antibonding orbital.
• using a wave function m for an atomic orbital on an atom A and m for an atomic orbital on an atom B, show how they can mathematically combine to produce a bonding and an antibonding molecular orbital
• place electrons in molecular orbitals for diatomic molecules, given a box energy diagram showing the relative energy of the bonding, antibonding, and nonbonding orbitals.
• from a filled (with electrons) molecular orbital energy diagram, predict the bond order between the two interacting atoms
• draw pictures illustration electron density derived from Spartan calculations for simple molecules studied in the Spartan lab
• describe the difference in electron density of a water molecule calculated using Spartan compared to the tradition Lewis structure.
• Define bond energy
• Using Lewis structures, predict trends in bond energies between diatomic molecules

29

11/19 W D4

11/20

R D5

in lab

- Energy, Energy Conservation, Internal Energy,

- State Functions.

-  PV Work, ΔE = q + w.  Signs of q, W, E;

-  Enthalpy, ΔH; First Law

-  ΔH and bond dissociation energies or bond enthalpies

• state the differences between force, energy and work
• state the 1st Law of Thermodynamics in words and mathematically
• state the differences between the system, surroudings, and the universe
• state two ways to transfer energy
• define work
• write an equation for the energy changes of a system, ΔE as a function of q and w.
• write an equation for the energy changes of a system, ΔE as a function of q and w. when W involves volume changes of a gas at constant external pressure
•  predict the sign of ΔE, q, and W for a variety of circumstances
• Write an equation that defines the change in enthalpy mathematically
• define enthalpy verbally and mathematically
• calculate ΔH for a chemical reaction from bond dissociation energies (or bond enthalpies
• draw energies diagrams and use them to calculate ΔH for a chemical reaction from bond dissociation energies (or bond enthalpies)

30

11/21 F D6

-Standard State

-ΔH and Standard enthalpy formation

- Hess's Law,

• define standard enthalpy of formation
• calculate ΔH for a chemical reaction from standard enthalpies of formation
• combine thermochemical equations to  calculate DH values for net reactions

31

11/25 T D2

- Heat Capacity

- Calculation ΔH rx's from: calorimetry:
- Coffee Cup Calorimetry.

• define heat capacity, specific heat (c), and molar heat capacity
• solve quantitative and qualitative problems given specific heat values  and q=mcΔT

32

12/2

T D4

• define pressure
• explain how a closed end column containing Hg or water when its open end is immersed in an open resevoir of either Hg or water, respectively, can be used to determine atmospheric pressure
• draw graph and write mathematical equations for the graphs that show how V (dependent variable) depends on P (constant n, T), T (constant P, n) and n (constant P, T)
• combine the mathematical equations above to give the ideal gas law
• explain the different between empirical laws and theories
• list the assumptions of the Kinetic Molecular Theory (KMT).
• Describe how the meaning of temperature can be mathematically derived by combining the empirical ideal gas law and the theoretical equation for P from KMT

33

12/4 R D6

34

12/8

M D2

• Solve typical problems based on the ideal gas law and KMT.
• Differential between effusion and diffusion

35

12/10 W D4

Problem Solving - MCQ

Topic 3 - Concept Inventory Multiple Choice Quest

36

12/12 F D6

Course Survey

Problem Solving - MCQ

Topic 3 - More Review Multiple Choice Questions

• Measurement Quiz
• Naming Ionic Compounds Quiz
• Precipitation Reactions Quiz
• Acid/Base Quiz
• Oxidation/Reduction Quiz
• Copper Experiment Quiz
• Isoelectronic Structures Quiz
• Molecular Geometry 1 Quiz
• Molecular Geometry 2 Quiz

Equations to know and love

Final Review Concepts

Final Exam:  TBA (Dec 16-18)

Note:  You will not be allowed to bring a graphing calculator into the final exam.  A regular calculator that can do scientific notation will be adequate.