Chem - Final Exam

calories to joules

1 cal. = 4.184 j.

Kelvin

Measures average kinetic energy

Chemical Bonds and Energy

Forming = Release Energy

Breaking = Consume Energy

System vs. Surrounding

System = singled out portion for energy change

Surroundings = everything else not included in system

First Law of Thermodynamics

Energy can be converted b/w forms. NEVER created NOR destroyed

Open vs. Closed vs. Isolated System

O: matter & energy CAN be exchanged w/ surroundings

C: energy can be exchanged w/ surroundings. NOT matter

I: NEITHER matter NOR energy can be exchanged w/ surroundings

E

Internal Energy

sum of all kinetic and potential energy of the components of the system

Change in Energy formula

del_E = q + w

change in energy = heat + work (into the system)

Endo vs. Exo thermic

Endo = heat absorbed into system

Exo = heat released from the system

State Function

a property of a system determined solely by the specific system conditions

only consider start and end states, NOT path taken to reach these states

H + FORMULA

Enthalpy: total heat content of a system

H = E + PV =

enthalpy = internal energy + pressure x volume

Pressure-Volume Work FORMULA

W = - P x del_V

dec in V → inc. in W

del_H FORMULA

del_H = del_E + P x del_V (w/ constant pressure)

therefore: del_H = q_p

→ change in enthalpy = heat gained / lost

del_E FORMULA

del_E = n x C_v x del_T

change in energy = mol. of solute x spec. heat cap. x change in temp.

q FORMULA

q = m x c x del_T = n x del_H = -C_cal x del_T

heat transfer = mass x spec. heat cap. x change in temp.

heat transfer = mol. of substance x enthalpy of system

heat transfer = -1 x calorimeter heat cap. x change in temp.

+q = heat into system. -q = heat out of system

Hess’ Law + FORMULA

del_H.rxn = sum(del_H.steps)

if a reaction can be carried out in a series of steps, the enthalpy change of the reaction is the sum of the steps’ enthalpy changes

del_H.rxn = sum(del_H.prod) - sum(del_H.reac.)

therefore, a reaction’s change in enthalpy is equivalent to the sum of enthalpy of the products formed subtracted by the sum of enthalpy of the reactants decomposed

AKA del_H.rxn = del_H.formed - del_H.decomposed

del_H.f

enthalpy of formation

del_H associated w. substance change

Enthalpy Standard State Conditions

P: 1 atm.

T: 25 C or 298 K

del_H^deg

standard enthalpy change

enthalpy of reactants and products in standard states

del_H.f^deg

standard enthalpy of formation

change in enthalpy for the reaction that forms 1 mol. of compound from its elements w/ all its elements in its standard states

if an element is in its natural form, del_H.f^deg = 0

del_H.f^deg FORMULA

elements (standard state) → compound (1 mol. standard state) , del_H.rxn = del_H.f^deg

Bond Enthalpy

del_H of breaking a particular bond in 1 mol. of a gas

always positive b/c breaking bonds takes energy

ONLY an ESTIMATION when del_H.f^deg not available

del_H.rxn for Bond Enthalpy FORMULA

del_H.rxn = sum(del_H.broken) - sum(del_H.formed)

if del_H.rxn > 0 , more bonds broken

if del_H.rxn < 0, mroe bonds formed

Specific Heat Capacity Units

C.s = J / g x ^deg.C or J / g x ^deg.K

Wavelength and Frequency of Electromagnetic Radiation

lambda x nu = c

wavelength x frequency = speed of light (3E8)

All Electromagnetic Radiation

By inc. lambda or dec. nu

Gamma, X-ray, UV, visible light, Infrared, Microwaves, Radio

Hertz

Hz. unit of frequency as units/second or s^-1

Energy of Photon

E = h x nu

h: Planck’s constant = 6.676E-34 J x s

nu: frequency of radiation

according to quantum theory, E only exists as integer multiples of h bc energy released in discrete or quantum chunks of energy

Mono vs Poly chromatic

M: radiation composed of a single wavelength

P: radiation composed of several emag. of diff wavelengths

Continuous vs. Line Spectrum

spectrum: polychromatic radiation separated into its several component wavelengths

C: spectrum w/ radiation over all wavelengths

L: spectrum w/ radiation of specific wavelengths

Rydberg Equation

Calculates all spectral lines

(lambda)^-1 = R.H( 1/n.1² - 1/n.2²)

n.2 > n.1

Bohr’s Model Postulates of H atoms

1. Only orbits of certain radii, corresponding to different energies, are permitted for electrons.

2. Electrons in permitted orbit are in an “allowed” energy state, keeping them radiated and at constant energy

3. Energy emitted/absorbed by electrons causes changes between energy states

   → in quantities of E = h x nu (quants)

Energy states of H atom

E = (-h x c x R.H) (1/n²) = (-2.18E-18) (1/n²)

as n → inf. , radius to e- → inf. e- separates & therfore E = 0

lower energy = more electrons = more stable atom

del_E = (-h x c x R.H) (1/n.f² - 1/n.i²)

n.f > n.i = Photon absorbed = del_E > 0

n.f < n.i = Photon emitted = del_E < 0

ground vs. excited states

G: lowest energy state of e- AKA n = 1.

bottom rung of the energy ladder

E: when an e- is in a higher energy state AKA n > 1

DeBroglie

matter waves: describes wave characterstics of a moving particle

electrons move about nuclei like waves, so they must have lambda

lambda = h / mv

frequncy = plank’s (6.626E-34) / momentum (mass x velocity)

Heisenberg

Uncertainty Principle

inherent uncertainty in precision of particle’s simultaneous position & momentum bc of extremely small mass

del_x x del_mv >= h / 4 x pi

since del_x.e- = 1E-9, del_mv cannot be measured since e- too small

since usually m»»», del_x unnoticable

Schrodinger

Proposed eq. where e- act as a wave and a particle

AKA wave or quantum mechanics

treat e- as a wave like a plucked guitar sting

n = 1: fundemental, n=2: 1st overtone, n=3: 2nd overtone etc.

Describes e- as wave functions (psi)

(psi)² describes e- location of energy states = e- density

Orbitals

allowed energy states of e- described by wave fxns (psi²)

exist as electron clouds of probabilistic position

n

principal quantum number. n > 0 (int)

associated w/ orbital size

inc. n = larger orbital size = looser found to nucleus = more energy

l

angular momentum quantum number. range (int): 0 - (n-1)

defines shape of orbital. 0 = s, 1 = p, 2 = d, 3 = f

inc. l = inc. orbital energy
→ ns < ns < np < nf for each level n

m.l

magnetic quantum number: range (int): -l - l

describes orientation/number of orbitals in space

Electron Shells

orbitals w/ same value n

1. n shell consists of exactly n shells

2. each l subshell consists of (2l+1) orbitals
   → each orbital corresponds to diff. n value
   → all s = 1, all p = 3, all d = 5 etc.

3. Total # of orbitals = n²

e- in ground state when in lowest orbitals. if else, exited state

Subshell

set of orbitals w/ same value n & l

s orbitals

lowest energy orbital. spherically symmetrical. l = 0, m.l = 0

e- density of a given distance is the same regardless of direction

n inc. = e- density more spread = likely farther from nucleus

p orbitals

“dumbell-shaped”. l = 1, m.l = -1,0,1 so 3 orbitals

density concentrated on either side of nucleus
→ separated by nucleus to make 2 lobes

each oriented on axis: p.x, p.y, and p.z

inc. n = inc. in size = longer lobes

d orbitals

“4-leaf clover + dumbbell doughnut”.

l = 2, m.l = -2,-1,0,1,2 so 5 orbitals

3 orbitals on planes, 1 along y=x line, 1 doughnut dumbbell

→ d.xy, d.xz, d.yx, d.x²-y², dz

f orbitals

l = 3, m.l = -3,-2,-1,0,1,2,3 so 7 orbitals

weird shapes + only worry about spd for bonding

degenerate

all orbitals of a given subshell (l) have same energy

electron spin

intrinsic property where e- behaves like a sphere spinning on an axis

m.s

spin magnetic quantum number. values: -1/2, +1/2

indicates direction of spin and there fore direction of magnetic field

each orbital can holds max 2 e- where each has opp. spin

e- configuration

how e- are distributed among various orbitals

orbitals filled in order of increasing energy

Notation: ns^(#e-) for each subsequent orbital

orbital diagram

boxes per orbital per subshell w/ 2 opposite arrows max (for m.s)

Hund’s Rule

when filling degenerate (same l) orbitals, lowest energy is attained when #e- having same spin is maximized

therefore, fill each orbital w/ one spin first then pair

condensed e- configuration

describes e- distribution w. reference to closes noble gas

1. Bracket noble gas prior to element = core electrons

2. List e- config. of remaining n. = valence electrons

result from each period on periodic table = inc. n

ex. Na: [Ne]3s^1

Transition Metal

4th row periodic. 10 elements form Se to Z

after 4s filled, then 3d fill in TM, then fill 4p

Lanthanide vs. Actinide Elements

L: rare earth metals. 14 elements filling 4f

A: all radioactive. 14 elements filling 5f

Mendeleev

Developed the earliest form of the Periodic table

organized chemicals by similar physical & chemical properties periodically by increasing atomic weight

PT Trend: Metallic

Increases from Right to Left and Top to Bottom

Z.eff

Effective nuclear charge

The partially screened nuclear charge by electrons

Approximate: Z.eff = Z - S

Z = actual nuclear charge = # protons

S = Screening constant = # non-valence electrons

PT Trend: Z.eff

Increases from Left to Right on all Periods

→ more protons w/ constant core e- = higher Z.eff

Slight Increase from Top to Bottom

→ broader e- cloads = less e- shielding = slight higher Z.eff

Z.eff Chart

Orbitals: n, n-1, n-2, n-3

s & p e-: 0.35, 0.85, 1, 1

d & f e-: 0.35, 1, 1, 1

Calculate S by #e- - 1 (b/c on the outermost e-)

Each e- value determined by n

Use normal Z.eff equation.

Non vs. Bonding Atomic Radius

NB: Shortest distance separating 2 nuclei i.e. atomic radii x2

B: ½ dist. b/w 2 bonded atoms nucleus. referred to for atomic size
→ hard to find for Noble Gases b/c don’t like to bond

nb > b

PT Trend: Atomic Radii

Increases from Top to Bottom

→ higher n = greater outer shell e- = bigger size

Decrease from Left to Right

→ higher Z.eff = more e- near nucleus = smaller size

Cation vs Anion Ionic Radii

Cations smaller than parent

→ + charge = less e- = decreased e- cloud = smaller

Anions larger than parent

→ - charge = more e- = increased e- cloud = larger

PT Trend: Ionic Radii

Ions w/ Same Charge: Increase from Top to Bottom

→ increased n = increased outsershell e- = larger size

isoelectronic series

ions grouped by same # e-

list elements by increase atomic # & therefore increased Z

→ bc # e- is constant, ionic radius dec. as Z inc.

higher Z + more e- towards nucleus = smaller size

ionization energy

min. energy to remove an e- from ground state

→ from n=1 to n=inf. is a removed e- and ionized atom

greater ionization energy = higher diff. of e- removal

I.1 < I.2 < I.3

largest increases w/ smaller orbital levels b/c closer to nucleus

iregularities result from orbital occupancy

→ easier to lose e- back down to a full shell

First vs. Second Ionization Energy

1: E required to remove first e- from neutral atom

2: E required to remove the second e- from ionized atom

1 < 2

PT Trend: First Ionization Energies

Increases from Left to Right

→ closer to full e- shell = want to lose e- less = higher I

Decreases from Top to Bottom

→ more core e- = more e- screening = easier to lose e- = lower I

s & p block elements have larger range of I than TMs

→ I inc. slowly b/w TMs and even less b/w f-block metals

PT Trend: Ionization Energy and Atomic Size

smaller atom = higher I

→ w/ inc. Z.eff and dec. atomic size = e- closer to nuclues = higher enrgy to remove e- = higher ionization e-

L2R: inc. Z.eff + dec. atomic size = inc. I

T2B: Inc. atomic size + rel. constant Z.eff = dec. I

e- configuration of ions

1. e- always removed from largest “n” orbital first

2. if more than 1e- occupy largest “n”, move to largest “l” orbital

3. atoms gaining e- place them in the lowest “n” & “l” orbital first

electron affinity

change in energy when atom gains e- and becomes an anion

greater attraction b/w nucleus & e- = EA more negative

I vs. EA energy

I: del_E > 0 b/c energy is put in to remove e-

EA: del_E < 0 b/c energy released when e- attached

PT Trend: EA

Not as clear as Z.eff, atomic size, or I

More negative from Left to Right

→ closer to filled orbital = more need for e- = more energy released

• Noble gases &. filled orbitals elements EA > 0 or slight neg.

→ need energy to add e- to filled shells

Generally Constant from Top to Bottom

→ dec. Z and dec. e- repulsion cancel out

metallic character

extent to which an element exhibits metal properties

Metals CHARACTERISTICS

shiny luster, heat & electric conductivity, malleable, ductile, solid @ room temp, and hi melting point

low I → form cations easily via removing s or p valence e-

compounds containing M tend to be ionic

metal oxide = base. M oxide + water = M hydroxide (base)

metal oxide + acid → salt + water

Nonmetals CHARACTERISTICS

varied state @ room temp, no luster or conductivity, and low melting point. range from super hard to super soft

form diatomic molecules

very neg. EA + rel. large atomic size → like to gain e- to form anions
→ taken from metals to from ionic compounds

compounds w/ ONLY NM tend to be molecular

NM oxide = acid. NM oxide + water → acid

NM oxide + base → salt + water

Metalloids CHARACTERISTICS

has some properties of M and NM. some w/, some w/o

best used in electrical semiconductors b/c intermediate conductivity b/w Ms and NMs

Group 1A CHARACTERISTICS

Alkali Metals

soft, metallic solids w/ silvery luster and high conductivity

low densities and low melting points
→ T2B: inc. density and atomic rad = dec. melting point & I

exist in nature only as compounds. Na and K in highest abundance

react vigorously w/ H2O (to make H2) and O2 (reg, peroxide, and hyperoxide)

emits color when placed in a flame

Group 2A CHARACTERISTICS

Alkaline Earth Metals

regular M properties. denser & higher melt. than AM
→ T2B: inc. density and atomic rad = dec. melting point & I

I1 still low but higher than AMs → less reactive than AM
→ larger AEM more reactive. Mg = slow. Ca and below = reactive

Mg and Ca most common in nature as ionic compounds

Hydrogen Group CHARACTERISTICS

very high I1 compared to Ms

reacts w/ NM shows tendency to hold onto e-
→ forms molecular compounds w/ NM rather than ionics like A/EM

forms H+ in presence of water w/ Ms to form acids

gains e- from metal w/ low I1 like AM

Group 6A CHARACTERISTICS

Oxygen Group

O: colorless gas. Exists as O2 (more stable) or O3 (ozone)

tends to attract electrons AKA “oxidizes” elements

forming NM oxides = exothermic (releases heat)

S: rare & toxic in high doses. Exists as S.8 rings

Te: even more complexe structure with Te-Te chains

Thermal Stability: Decreases from Top to Bottom

Group 7A CHARACTERISTICS

Halogens

all NMs w/ melt & boil increasing w/ atomic #

all exist as diatomic molecules

EA very negative → strips substances of e- → very reactive
→ T2B: Decrease in Reactivity

F: Very Very Reactive → Very Exothermic

Cl: Used in electrolysis: elec. current turns. Cl- into Cl2

Reacts w/ most metals + H+ to form ionic compounds

Group 8A CHARACTERISTICS

Noble Gases

NM gases @ room temp.

all exist as monoatomic
→ completely filled s & p orbitals = unreactive

group w/ largest I1

Lewis e- Dot Structure

simple diagram to depict the valence e- of an atom for bonding

consists of elemental symbol and max. 8 dots for each valence e-

fill Top, Bottom, Left, Right then double up the dots

groups will have the same dot structure among all elements

Octet Rule

atoms gain/lose/share e- to reach 8 valence e-

mostly applies to atoms w/ s & p shell valence e-

Elements Guaranteed: C, N, O, F

LS: Ionic Bonding

1. move e’- dot from 1 structure to another indicated w/ arrow

2. bracket the structure and indicate charge of ion

Ionic Structures CHARACTERISTICS

brittle and crystalline w/ high melt. cleavable

results from rigid well defined arrangements from e-static forces

lattice energy

amount of energy released when 2 gaseous ions combine to form an ionic solid. highly exothermic

allows ion bonds to be release a lot of energy despite e- transfer using energy

Q.ion inc. = E.lattice inc. AKA more charge = more energy released

Min./Max/ Ionic Charge

± 3

I rapidly increased w/ each successive ion

covalent bond

chemical bond where atoms share a pair of electrons

LS: drawn as e- dot pair between elemental symbols
→ also represented for a single line for each bond

Covalent Bond Length and Energy

more e- shared = shorter bond = more energy = more stable

bond polarity

measures how un/equally the e- in a covalent bond are shared

nonpolar = equally shared, polar = 1 atom has greater e- attraction

→ ionic = e- attraction so high that e- shifts into the other atom

electronegativity

ability of an atom in a molecule to attract e- to itself
→ very neg. EA and high I = e- attracted = high e-neg.

use difference of e-neg. to determine polarity

diff. = 0: NP. 0 < diff. < 2: P. diff >= 2: ionic
→ greater the diff. = increased polarity

as e- shifts towards 1 atom, e- density shift towards e-neg. atom

PT Trend: electronegativity

Increases Left to Right

→ follows EA more negative and I increasing

Decreases Top to Bottom

→ decreases w/ atomic size

Polar molecule

molecule where the centers of + and - charge do NOT coincide

LS: represented with delta + and delta - over each atom
→ also represented w/ the + arrow from + to - charge (for e- density)

more e-neg. = delta -

dipole vs dipole moment

D: when 2 electrical charges of equal magnitude and opposite sign are separated

DM: measurement of magnitude of dipole

mew = Qr. higher mew = higher polarity

→ NP: mew = 0.

Polarity: Ionic vs. Covalent Bonds

Both technically share e-. Compounds either I or C dominant.

C: molecular = low melt/boil and nonelectrolyte. NM + NM

I: ionic = brittle, lattice structure, and strong electrolyte. M + NM

Formal Charge

atom charge if each e- pair is shared equally

purely for BOOKKEEPING valence e- of each LS

1. All unshared e- assigned to atom they are found in

2. All shared e- count as half # for each atom.

3. FC = valence e- of parent - all assigned e-

How to find dominant Lewis Structure?

1. Dominant structure generally have formal charges closes to zero, especially around central atom

2. Dominant structure tend to have - charges lie on more e-neg. atoms