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1.1 Organic Chemistry The study of carbon-containing molecules and their reactions What happens to a molecule during a reaction? A collision Bonds break/form.

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Presentation on theme: "1.1 Organic Chemistry The study of carbon-containing molecules and their reactions What happens to a molecule during a reaction? A collision Bonds break/form."— Presentation transcript:

1 1.1 Organic Chemistry The study of carbon-containing molecules and their reactions What happens to a molecule during a reaction? A collision Bonds break/form WHAT IS A BOND? The BIG question: WHY do reactions occur? We will need at least 2 semesters of your time to answer this question FOCUS ON THE ELECTRONS Copyright 2012 John Wiley & Sons, Inc.

2 1.1 Organic Chemistry Why do we distinguish between organic and inorganic compounds? Why are organic compounds important? Copyright 2012 John Wiley & Sons, Inc.

3 1.2 Structural Theory In the mid 1800s, it was first suggested that substances are defined by a specific arrangement of atoms. Why is a compound’s formula NOT adequate to define it? What term do we use to describe different substances with the same formula? Copyright 2012 John Wiley & Sons, Inc.

4 1.2 Structural Theory Atoms that are most commonly bonded to carbon include N, O, H, and halides (F, Cl, Br, I). With some exceptions, each element generally forms a specific number of bonds with other atoms Practice with SkillBuilder 1.1 Copyright 2012 John Wiley & Sons, Inc.

5 1.3 Covalent Bonding A covalent bond is a PAIR of electrons shared between two atoms. For example… Copyright 2012 John Wiley & Sons, Inc.

6 1.3 Covalent Bonding How do potential energy and stability relate?
What forces keep the bond at the optimal length? Copyright 2012 John Wiley & Sons, Inc.

7 1.3 Atomic Structure A review from General Chemistry
Protons (+1) and neutrons (neutral) reside in the nucleus Electrons (-1) reside outside the nucleus. WHERE? Some electrons are close to the nucleus and others are far away, WHY? Look at carbon for example. Which electrons are the valence electrons? Why are valence electrons important? Copyright 2012 John Wiley & Sons, Inc.

8 1.3 Counting Valence Electrons
You can always calculate the number of valence electron by analyzing the e- configuration. Look at phosphorus. Or, for Group A elements only, just look at the Group number (Roman Numeral) on the periodic table Practice with SkillBuilder 1.2 Copyright 2012 John Wiley & Sons, Inc.

9 1.3 Simple Lewis Structures
For simple Lewis Structures… Draw the individual atoms using dots to represent the valence electrons. Put the atoms together so they share PAIRS of electrons to make complete octets. WHAT is an octet? Take NH3, for example… Practice with SkillBuilder 1.3 Copyright 2012 John Wiley & Sons, Inc.

10 1.3 Simple Lewis Structures
For simple Lewis Structures… Draw the individual atoms using dots to represent the valence electrons. Put the atoms together so they share PAIRS of electrons to make complete octets. WHAT is an octet? Try drawing the structure for C2H2 Copyright 2012 John Wiley & Sons, Inc.

11 1.4 Formal Charge What term do we use to describe atoms with an unbalanced or FORMAL charge? How does formal charge affect the stability of an atom? Atoms in molecules (sharing electrons) can also have unbalanced charge, which must be analyzed, because it affects stability To calculate FORMAL charge for an atom, compare the number of valence electrons that should be associated with the atom to the number of valence electrons that are actually associated with an atom Copyright 2012 John Wiley & Sons, Inc.

12 1.4 Formal Charge Consider the formal charge example below. Calculate the formal charge on each atom. or Carbon should have 4 valence e-s, because it is in group IVA on the periodic table. Carbon actually has 8 valence e-s. It needs 8 for its octet, but only 4 count towards its charge. WHY? The 4 it actually has balance out the 4 it should have, so it does not have formal charge. Its neutral. Copyright 2012 John Wiley & Sons, Inc.

13 1.4 Formal Charge Analyze the formal charge of the oxygen atom. or
Oxygen should have 6 valence e-s, because it is in group VIA on the periodic table. It actually has 8 valence e-s. It needs 8 for its octet, but only 7 count towards its charge. WHY? If it actually has 7, but it should only have 6, what is its formal charge? Practice with SkillBuilder 1.4 Copyright 2012 John Wiley & Sons, Inc.

14 1.5 Polar Covalent Bonds Covalent bonds are electrons pairs that exist in an orbital shared between two atoms. What do you think that orbital looks like? Just like an atomic orbital, the electrons could be anywhere within that orbital region. What factors determine which atom in the bond will attract the shared electrons more? Copyright 2012 John Wiley & Sons, Inc.

15 1.5 Polar Covalent Bonds Covalent bonds are either polar or nonpolar
Nonpolar Covalent –bonded atoms share electrons evenly Polar Covalent – One of the atoms attracts electrons more than the other Electronegativty - how strongly an atom attracts shared electrons Copyright 2012 John Wiley & Sons, Inc.

16 1.5 Polar Covalent Bonds Electrons tend to shift away from lower electronegativity atoms to higher electronegativity atoms. The greater the difference in electronegativity, the more polar the bond. Copyright 2012 John Wiley & Sons, Inc.

17 1.5 Polar Covalent Bonds Can a bond have both covalent and ionic character? Practice with SkillBuilder 1.5 Copyright 2012 John Wiley & Sons, Inc.

18 1.6 Atomic Orbitals General Chemistry review
In the 1920s, Quantum Mechanics was established as a theory to explain the wave properties of electrons The solution to wave equations for electrons provides us with visual pictures called orbitals Copyright 2012 John Wiley & Sons, Inc.

19 1.6 Atomic Orbitals General Chemistry review
The type or orbital be identified by its shape An orbital is a region where there is a calculated 90% probability of finding an electron. The remaining 10% probability tapers off as you move away from the nucleus Copyright 2012 John Wiley & Sons, Inc.

20 1.6 Atomic Orbitals Electrons behave as both particles and waves. How can they be BOTH? Maybe the theory is not yet complete The theory does match experimental data, and it has predictive capability. Like a wave on a lake, an electron’s wavefunction can be (+), (-), or ZERO. Copyright 2012 John Wiley & Sons, Inc.

21 1.6 Atomic Orbitals Because they are generated mathematically from wavefunctions, orbital regions can also be (-), (+), or ZERO The sign of the wave function has nothing to do with electrical charge. In this p-orbital, there is a nodal plane. The sign of the wavefunction will be important when we look at orbital overlapping in bonds. Copyright 2012 John Wiley & Sons, Inc.

22 1.6 Atomic Orbitals Electrons are most stable (lowest in energy) if they are in the 1s orbital? The 1s orbital is full once there are two electrons in it. Why can’t it fit more? The 2s orbital is filled next. The 2s orbital has a node. WHERE? Copyright 2012 John Wiley & Sons, Inc.

23 1.6 Atomic Orbitals Once the 2s is full, electrons fill into the three degenerate 2p orbitals Where are the nodes in each of the 2p orbitals? Copyright 2012 John Wiley & Sons, Inc.

24 1.6 Atomic Orbitals Common elements and their electron configurations
Practice with SkillBuilder 1.6 Copyright 2012 John Wiley & Sons, Inc.

25 1.6 Atomic Orbitals What are the rules that govern our placement of electrons ? Copyright 2012 John Wiley & Sons, Inc.

26 1.7 Valence Bond Theory A bond occurs when atomic orbitals overlap. Overlapping orbitals is like overlapping waves Only constructive interference results in a bond Copyright 2012 John Wiley & Sons, Inc.

27 1.7 Valence Bond Theory The bond for a H2 molecule results from constructive interference Where do the bonded electrons spend most of their time? Copyright 2012 John Wiley & Sons, Inc.

28 1.8 Molecular Orbital Theory
Atomic orbital wavefunctions overlap to form MOs that extend over the entire molecule. MOs are a more complete analysis of bonds, because they include both constructive and destructive interference. The number of MOs created must be equal to the number of AOs that were used. H2 MOs Copyright 2012 John Wiley & Sons, Inc.

29 1.8 Molecular Orbital Theory
Why is the antibonding orbital higher in energy? When the AOs overlap, why do the electrons go into the bonding MO rather than the antibonding MO? Copyright 2012 John Wiley & Sons, Inc.

30 1.8 Molecular Orbital Theory
Imagine a He2 molecule. How would its MOs compare to those for H2? In general, if a molecule has all of it MOs occupied, will be stable or unstable? How would the energy of the He2 compare to 2 He? Why does Helium exist in its atomic form rather than in molecular form? Copyright 2012 John Wiley & Sons, Inc.

31 1.8 Molecular Orbital Theory
Consider TWO of the many MOs that exist for CH3Br There are many areas of atomic orbital overlap Notice how the MOs extend over the entire molecule Each picture below represents ONE orbital. Copyright 2012 John Wiley & Sons, Inc.

32 1.8 Molecular Orbital Theory
How many electrons can fit into the areas represented? In the ground state, electrons occupy some MOs and not others, WHY? Depending on the circumstances, we will use both MO and valence bond theory to explain phenomena Copyright 2012 John Wiley & Sons, Inc.

33 Study Guide for sections 1.1-1.8
DAY 1, Terms to know: Sections Isomers, valence electrons, octet rule, lone pair, ion, formal charge, subshell, orbital, electronegativity, polarity, node, wavefunction sign, wavesign, electron configuration, orbital diagram, atomic orbital, molecular orbital, bonding, antibonding DAY 1, Specific outcomes and skills that may be tested on exam 1: Sections Be able to predict the number of bonds each relevant atom typically has. Be able to explain how the energy changes as covalent bonds get stretch or compressed and WHY and why there is an optimal length for each covalent bond between atoms. Be able to determine the number of protons, neutron, and electrons in an atom, and how many electrons are valence electrons. Be able to draw Lewis structures for molecules including molecules with heteroatoms and with single, double, and/or triple bonds with the most reasonable location for all bonding pairs of electrons and all lone pairs. Be able to calculate formal charge for any atom given how many bonds and how many lone pairs it has. Given an atom with a formal charge shown, know how to determine how many bonding and lone pair electrons are around it. Be able to count electrons around an atom to determine whether an atom has a complete octet and also count electrons around an atom to determine the formal charge of an atom. Be able to explain the difference between an energy level, subshell, and orbital. Be able to determine the relative electronegativity of atoms based on periodic trends. Be able to determine the relative polarity of individual bonds. Be able to explain how the shapes of the orbitals vary for s and p orbitals. Be able to describe how electrons have wave properties and how the wavesign squared gives the shapes of the orbitals and location of greatest electron density. Be able to explain how the wavesign is different from the charge. Be able to identify and/or describe where the nodes are for the orbitals in the first and second energy levels as well as the 3s. Be able to write electron configurations and orbital diagrams Be able to describe how both constructive and destructive interference work for wavefunctions and how they yield bonds and antibonds with specific shapes including both pi and sigma bonds. Be able to explain how atomic orbitals overlap and that the number of molecular orbitals that result is equal to the number of atomic orbitals overlapping. Be able to explain how electrons are more stable in the bonding MO than in the antibonding and two reasons why that relate to electronic repulsions and attractions. Be able to explain how each bonding MO is lower in energy than the nonbonding by the same quantity that anitbonding MO is greater in energy than nonbonding. Be able to draw the bonding and antibonding energies for small molecules involving two atoms bonding and how having electrons in the bonding and antibonding orbitals relates to the octet rule and how entropy plays into the likelihood of bonds forming. Be able to explain that molecular orbitals may have many nodes and assymetrical shapes, but 1 orbital can still only hold a maximum of 2 electrons.

34 Practice Problems for sections 1.1-1.8
Complete these problems outside of class until you are confident you have learned the SKILLS in this section outlined on the study guide and we will review some of them next class period

35 Prep for day 2 Must Watch videos: Other helpful videos:
(Lewis structures, crash course chemistry) (formal charge) (valence bond theory) (molecular orbital theory) (sp3 hybridization) (atomic orbital hybridization) (VSEPR, Tyler DeWitt) (Dipole moments) Other helpful videos: (Tyler DeWitt: ionic vs. molecular) (Lewis structures) (polar vs. nonpolar, crash course chemistry) (UC-Berkeley lessons10-11) (UC-Berkeley lessons 13-14) (MIT lectures 11-12) (UC-Irvine lectures 9-10) (UC-Irvine lectures 1-4) Read sections

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