Honors Chem

1. Honors Chem What we should know, chapters 1-5

2. Chapter 1: Intro to Chemistry • Chemistry: the study of matter and its interactions, typically considered on the atomic level (physics handles the bigger items) • Scientific method: not a linear process but a cyclical one, where reevaluation and review is part and parcel all the time, as revision is always occurring. We must remain aware that we are constantly learning new things about everything that we persist in studying. • Lab safety: something that is in your heads…not something you need to look up…this carries over into everything you do in life.

3. Chapter 2: Matter and Change • Properties • Intensive: inherent to the material, not how much there is • Extensive: dependent upon the amount you happen to be dealing with • Physical: a property which doesn’t change readily • Melting/boiling point, color, luster, xl form, state, malleable, ductile, etc • Chemical: only 2: reactivity, flammability/combustibility • Definitely different before and after, in chemical composition

4. Chapter 2: matter and change, continued • States or phases of matter • Three “normal” ones: solid, liquid, gas • Solid: definite shape, definite volume • Liquid: indefinite shape, definite volume • Gas: indefinite shape, indefinite volume (fills what is available) • Two “extreme” ones: high temp: plasma; low temp (0K): Bose-Einstein condensate • Phase changes • Exothermic: going to lower energy state • Freezing, condensing, depositing • Endothermic: going to higher energy state • Melting, vaporizing, sublimating

5. Chapter 2: matter and change, continued • Mixtures • Homogeneous: unable to distinguish parts • Heterogeneous: able to distinguish parts • Some of either of these may be solutions, involving liquids or gases, like air • Separating • Filtration • Distillation • Other means: magnets, sieves, etc

6. Chapter 2: matter and change, continued • Elements and compounds • Elements: unique combinations of protons, neutrons, electrons • Atomic number: number of protons, electrons (if neutral); unique for each element • Mass number: number of protons and neutrons for a specific isotope of an element • Atomic mass: weighted average of mass numbers for naturally occurring isotopes of a given element • Compound: unique combination of atoms of different elements, forming an electronically neutral molecule, with its own unique properties • Can be broken down chemically into the elements it is formed from

7. Chapter 2: matter and change, continued • Chemical reactions • Signs: color change, temperature change, evolution of gases, precipitation • Reactants form new products • Law of Conservation of Mass/matter: matter is neither created nor destroyed in a chemical reaction; all must be accounted for

8. Chapter 3: Scientific Measurement • Problem-solving • Must know what you are looking for…sounds simple, but often overlooked • Write equations down before you touch a calculator • NO NAKED NUMBERS…units are your safety net. • Measurements • SI units • Significant figures and scientific notation: know how precise your measurements are • Percent error: how good are your numbers?

9. Chapter 4: atomic structure • Models of the atom • Democritus: (400BC)“atomos,”, good idea, shot down by Aristotle • Dalton: (c.1800) • All elements are composed of tiny particles called atoms • All atoms of an element are the same; atoms of different elements are different • Atoms of different elements form compounds with simple whole-number ratios • Atoms of elements can chemically react but don’t turn into new elements, just new compounds of the original elements • Segue: size of atoms: 0.5-2 e-10 meters…pretty small.

10. Chapter 4: atomic structure, continued • Subatomic particle discoveries • JJ Thomson (c. 1904): discovered electrons; working with cathode ray tubes and charged plates, realized negative particles being deflected by negative plates but attracted to positive plates • Millikan (c. 1916): worked out mass of electron, utilizing charge • Electrons weigh approximately 1/1840 of a proton • Goldstein (c. 1886): discovered protons • Chadwick (c. 1932): discovered neutrons • Atomic Nucleus: Rutherford and the gold foil experiment • Atoms are largely empty space: think nucleus as pea in football field • Nucleus has positive charge; we know due to all protons there

11. Chapter 4: atomic structure, continued • Review: atomic #, mass #, isotopes, atomic mass (wt’d. avg. of naturally occurring isotopes) • Periodic table • Periods/rows: properties repeat “periodically,” as you go to next higher level • Groups/columns/families: similar properties

12. Chapter 5: Electrons in Atoms • Timeline of atomic models: additions to Dalton, et al, who we already have listed • 1913: Nils Bohr: concept of energy levels • 1923: DeBroglie: electrons have wave properties; realization that they act like both particles and waves • 1926: Schrodinger: mathematical equations lead to electron cloud model (delocalized locations for electrons). Also called quantum mechanical model

13. Chapter 5: Electrons in Atoms • Atomic orbitals • S orbital: all energy levels, capacity of 2 electrons • P orbitals: all but 1st energy level, 3x2 electron capacity (6) • D orbitals: all but 1st and 2nd levels, 5x2 electron capacity (10) • F orbitals: all but 1st, 2nd, 3rd levels, 7x2 electron capacity (14) • G orbitals: start at 5th level, after 8s2, 9x2 capacity (18), none there yet

14. Chapter 5: Electrons in Atoms • Electron Configurations • Methods to show where the electrons are arranged in the orbitals around a given element’s atoms • Three rules for completing electron configurations: • Aufbau principle: electrons always occupy the orbital of lowest energy first • Pauli Exclusion principle: an orbital may hold at most two electrons and they must have opposite spin properties; convention has us writing the “up” arrow first • Hund’s rule: electrons fill at a rate of one per equal orbital at a given level first before the second electrons fill these same orbitals (no orbitals with two electrons while another orbital of the same level/type has none) • Shorthand writing: can utilize [noble gas] of next lower energy level to represent the filled orbitals to that point…need to know how to read, at least

15. Chapter 5: Electrons in Atoms • Exceptions to general rules of electron configurations • Cr: 1s2 2s2 2p6 3s2 3p6 4s1 3d5 • Cu: 1s2 2s2 2p6 3s2 3p6 4s13d10 • These occur due to the knowledge that having the d orbitals balanced creates a lower energy scenario than having the 4s orbital filled and the d orbitals unfilled…yes, it is very subtle, but measurable to those who research this

16. Chapter 5: Electrons in Atoms • Electrons and light • Timeout…we haven’t talked about electromagnetic energy at all yet • EM waves: energy transferring without a medium • Amplitude: height of wave from rest position, this is a measure of energy • Wavelength: how long (distance) between the same spot on repeating waves; λ (lambda) • Period: how long (time) between the same spot on repeating waves; ν (nu) • Frequency: inverse of period: number of wavelengths/time • Important equation: c=λν This tells us that speed = wavelength x frequency. Since the speed of em waves is constant (c=2.998e8m/s), this allows one to calculate the missing information • Different wavelengths/frequencies have very different energies!!! • Visible light is a very narrow band in this spectrum of waves!

17. Chapter 5: Electrons in Atoms • Atomic spectra: when atoms absorb energy, their electrons get excited and jump to higher energy levels • If we measure here, it is an absorption spectrum • If we measure when the electrons “get rid” of this energy, it is an emission spectrum (think of the lines you saw thru the spectroscopes at the gas tubes) • Different lines: different transitions • From higher levels to level 3: Paschen series, typically in infrared range (heat to us) • From higher levels to level 2: Balmer series, visible light/lines to us • From higher levels to level 1: Lyman series, UV range • Each of these represent larger amounts of energy being released

18. Chapter 5: Electrons in Atoms • Quantum Mechanics: 1905: Albert Einstein describing light as packets (quanta) of energy…not bad for a patent clerk who had major issues in school with his math classes, right? • This started a whole range of discussion/argument in the world of physics at this time: is light wave-like or particle-like? DeBroglie weighed in in the 1920’s, proving both sides of this argument…ended up with a Nobel prize for these efforts…ended up discussion that all objects have some degree of wavelike motion…usually too small to be detected (read discussion on page 145), so we follow classic mechanics in our descriptions of motion (yep, I know, physics)

19. Chapter 5: Electrons in Atoms • Final notes from this chapter • Heisenberg Uncertainty principle: it is impossible to know both the velocity and position of a particle at the same time…this means that we really cannot predict the exact location, but rather the general area, as our effort to locate the object tends to then change its location (I always remember it that the experimenter becomes part of the experiment) http://abyss.uoregon.edu/~js/21st_century_science/lectures/lec14.html Remember the flame test lab…similar to emission spectra from gas tubes, with a different source of the energy that the electrons were absorbing.

20. Keep your perspective:

21. And more…