CHEM 110 Exam 1 - CHEM 110 // SP2016 // EXAM 1 PDF

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CHEM 110 // SP2016 // EXAM 1...

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CHEM 110, Exam 1 13 January: 1. Intro to Atomic Structure: atoms are made of protons, neutrons, and electrons, masses all measured in a.m.u. (atomic mass unit) a. Component

Mass (a.m.u.)

Charge (a.u.)

proton (p+)

~1.0

+1

electron (e-)

5.486 x 10^-4

-1

neutron (n^0)

~1.0

0

b. Isotopes: same atomic #, different # of electrons. For a given element, the number of protons is fixed, but the number of neutrons can vary c. Atomic weight: also known as average atomic mass, can use all of the isotopes to find average atomic mass of an element i. Atomic mass of 12C = 12 a.m.u. 1. 12C has 6 p+ and 6 n^0 and each weigh 1 a.m.u

d. Moles: 1 mole of anything (atoms, molecules) = 6.02 x 10^23, used to connect the atomic scale (atoms/molecules) to the macro scale (gram quantities of atoms/molecules) i. Avogadro’s number is 6.02 x 10^23

ii.

If 18 g of water = 1 mol of water (atomic weights of 2 H and 1 O), then we can say 18 g/mol for water iii. Atomic weight of C is 12.011 a.m.u, meaning that 1 mole of carbon has a mass of 12.011 g - can express atomic weight of carbon as 12.011 g/mol e. Atoms: almost entirely empty space i. Atom diameter = 0.1-0.5 nm ~10^-10 m ii. Nuclear diameter = 10^-5 nm = 10^-14 m iii. Atom diameter = 10,000 x nuclear diameter

f.

Forms of Energy: two forms, kinetic and potential, measured in J = joules or cal = calories/kcal/kJ i. Kinetic energy is the energy of motion, incl. mechanical (1/2mv^2), electrical (moving charge), light (photons), sound (molecules moving uniformly), heat (molecules moving randomly) ii. Potential energy is stored energy, incl. mechanical (mass in a place where a force can act), chemical (bonds), nuclear (binding energy), electrostatic (interaction of charged particles) g. Electrostatic energy: form of potential energy that results from the interaction of charged particles, energy is positive with like charges, negative with unlike charges, use Coulomb’s Law i. Q1 and Q2 = charges of the particles ii. r = distance between the particles iii. k = constant (disregard)

h. Internal Energy: the total energy (E) associated with a system, the sum of all sources of kinetic and potential energy, the capacity to do work or transfer heat i. q = heat, w = work i. heat (q) 1. absorbing is + 2. releasing is ii. work (w)

1. work done on the system + 2. work done by the system -

j.

State function: a function whose value does not depend on the pathway used to get to the present state - only depends on the current state (composition, T, P): does not depend on past history. Still get from A to B, whether it takes 2 or 8 steps, energy will be the same i. Changes in state functions are path-independent, state functions are written as uppercase letters (E, H, P, V, T, S), q and w are not state functions, but ΔE = q + w is a state function

k. Enthalpy: heat transferred at constant pressure, exothermic and endothermic only deal with heat, work done on or by the system is completely irrelevant

l.

Spectroscopy: the study of light interacting with matter, necessary for learning about the structure of an atom, electromagnetic radiation is light i. Electromagnetic radiation behaves as a wave in most situations

^ know relative order of different rays and colors, know the visible spectrum

Blackbody radiation is inconsistent with the wave nature of light. When solids are heated, they emit different wavelengths of radiation. Max Planck found that light behaves are a particle and a wave - it’s transferred in specific “packets” or “quanta” of energy, also known as photons

^ v = frequency of the photon of light h = planck’s constant, 6.626 x 10^-34 J-s Energy increases, wavelength increases 20 January: 2. Different types of light/different spectrums a. Example

Spectrum type

wavelength

laser

monochromatic

one

lightbulb

continuous

all

hg vapor

discrete/line

few

i.

different gases have different light spectra, different metals emit different spectra as well b. Hydrogen - only 4 emission lines in the visible spectrum, good example bc only 1P/1E-

ii.

i. Observation of line spectra implies that atoms have discrete (quantized) energy levels 1. E- gaining/losing energy and moving between the energy levels within an atom

iii.

iv.

2. The lines are the result of electronic transitions; electrons moving from one allowed energy state to another H atom has 1+ and 1-, Bohr proposed that the electron is in one of several possible orbits around the proton 1. Bohr used the line spectrum to figure out the energy of an electron in each energy level 2. Energy of an electron increases with increasing distance from the nucleus a. spacing between the energy levels gets exponentially smaller and smaller as the distance from the nucleus increases i. n=1 → n=2 > n=5 → n=6 1. gap between energy levels 1 and 2 is much larger than between 5 and 6

Bohr Model a. Putting electron into orbital: forming an attractive interaction, energy will be negative

b. Reverse the process: try to remove the electron, energy will be positive c. Large transition = shorter wavelength d. Small transition = longer wavelength 3. dslfjslk

22 January: 4. Matter also has wave-like properties - all matter has a wavelength based on its mass and velocity (and Planck’s constant)

a. for larger objects (baseballs, cells), the wavelength is too small to observe, but the wavelength for electrons is profound b. Heisenberg Uncertainty Principle - it is not possible to simultaneously know both the position and velocity (momentum, mv), of a particle with complete certainty i. derives from wavelike nature of matter, important when dealing with subatomic matter

ii. all electrons have velocity, can’t specify their exact location iii. contradicts Bohr’s model, electrons do not exist in perfect orbitals c. Orbitals - Schrodinger found that orbitals are regions of space that describe the most probable locations of electrons i. described electrons as clouds of electron density, which is described as an orbital ii. Schrodinger equation (not needed) is used to predict orbital shape and size iii. orbitals tell where the electron is and the energy of the electron 1. specifies the probability of finding an electron in a given region of space (orbitals have shapes) 2. specifies the energy of the electron 3. characterized by quantum numbers, which give info about size, shape, orientation, and energy level 4. only certain orbital shapes exist for each number a. Principal quantum # (n) i. energy level or shell, must be an integer (1 → infinity), gives info about size ii. gives info about energy 1. larger n = farther from nucleus, higher (less negative) energy b. Azimuthal quantum l i. integer ranging from 0 to n - 1 ii. gives info about shape iii. based on the energy level that it is located in 1. n = 1, l = 0 2. n = 2, l = 0 or +1 3. n = 3, l = 0, 1, or 2

iv.

s = sphere, p = 1 on top, 1 on bottom, d = 4, clover leaf c. Magnetic quantum number (ml) i. determines the orientation of the orbital, ranges from -l to +l (integers) ii. l = 0, ml = 0 iii. l = 1, ml = -1, 0, +1 iv. l = 2, ml = -2, -1, 0, +1, +2

5. shells, subshells, + orbitals a. n^2 = number of energy states in a given shell b. all orbitals in a given subshell are degenerate i. all have the same energy 5. Structure determines function - electron configurations are the electronic structure of the atom, properties of elements are determined by size (n), shape of orbitals (l), and atomic number. 4 main elemental properties

a. atomic size - atomic radius is how the size of the atom is measured (from the nucleus to the edge of the electron cloud), size increases going down the column - size increases because of the increasing number of shells being populated by electrons i. right to left, atomic radius increases

ii. iii.

up to down, atomic radius increases; more shells being filled with electrons more protons = smaller radius

b. ionization energy - the energy needed to remove an electron from an atom, valence electrons are more easily removed. Ionization energy decreases going down a column

ii. iii.

1. exceptions to ^: filled & half-filled subshells are relatively stable and are difficult to add/remove electrons from (ex. Mg, Ar, P) IE is always positive (endothermic process) as it requires energy to remove an electron from an atom ex. Mg → Mg+ = 1E1, Mg → Mg2+ = 1E2 1. first and second ionization energies

iv.

IE3 is much higher because removing an electron from the octet is very, very difficult v. metals have IE below 1,000 KJ/mol, nonmetals have higher IE’s c. electron affinity - energy released/absorbed when an electron is added to an atom or ion in the gas phase, typically negative. trends in electron affinity are not obvious i. exothermic 1. halogens - one added electron completes the octet 2. group I metals - added electron results in filled S2 subshell ii. endothermic 1. group 2A metals (Be, Mg) - already have a filled NS2 subshell 2. noble gases (group 8A) - already have complete octet (filled NP6) d. reactivity - reactivity of metals is related to the ionization energy, reactivity increases as ionization energy decreases...