Chem 1103 Exam 3

Previous Chapter Info

Prefixes (conversion)

* mega (M)- 10^6
* kilo (k)- 10^3
* deka (da)- 10
* deci (d)- 10^-1
* centi (c)- 10^-2
* milli (m)- 10^-3
* micro (the weird u symbol)- 10^-6
* nano (n)- 10^-9
* pico (p)- 10^-12

Acetate

C₂H₃O₂⁻

Carbonate

CO₃²⁻

Hydrogen Carbonate (aka bicarbonate)

HCO₃⁻

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Hydroxide

OH⁻

Nitrate

NO₃⁻

Nitrite

NO₂⁻

Chromate

CrO₄²⁻

Dichromate

Cr₂O₇²⁻

Phosphate

PO₄³⁻

Hydrogen Phosphate

HPO₄²⁻

Dihydrogen Phosphate

H2PO4-

Ammonium

NH₄⁺

Hydronium

H3O+

Hypochlorite

ClO⁻

Chlorite

ClO₂⁻

Chlorate

ClO₃⁻

Perchlorate

ClO₄⁻

Permanganate

MnO₄⁻

Sulfate

SO₄²⁻

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Sulfite

SO₃²⁻

hydrogen sulfite (aka bisulfite)

HSO₃⁻

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hydrogen sulfate (aka bisulfate)

HSO₄⁻

Peroxide

O₂²⁻

\

Cyanide

CN⁻

Chapter 7-11

Visible light

is a type of __**electromagnetic radiation**__

Wave properties of electromagnetic radiation

* **frequency** (ν, *nu*)
* cycles per second (1 / s)
* **wavelength** (λ, *lambda*)
* the distance a wave travels in one cycle; the distance between adjacent wave peaks
* **amplitude**
* the height of a wave crest or depth of a trough

**Speed of light**

3\.00x10^8 m/s

amplitude and wavelength

* no relationship between the two

Frequency and wavelength

* __**inverse relationship**__ between the frequency of a wave and its wavelength
* For waves traveling at the same speed, the shorter the wavelength, the more frequently they pass

frequency/ wavelength formula

**v =  c/λ**

energy and wavelength

Inversely proportional

electromagnetic spectrum (from Low wavelength/ high energy to high wavelength/low energy)

* Gamma Rays
* X-Rays
* Ultra-violet
* Visible range
* infrared
* microwaves
* radio

Visible Range

* 400-750 nm

color

* “White” light is a mixture of __**ALL**__ the colors of visible light
* wavelength of colors decreasing order: ROY G BIV
* color= when object absorbs some of wavelengths of white light but reflects others

Refraction

* When a light wave passes from one medium into another, the speed of the wave changes
* Particles of matter do not undergo refraction

Dispersion

* White light separates into its component colors when it passes through a prism
* Each incoming wave is refracted at a slightly different angle

Interference

* interaction between waves

Constructive interference

* waves interact so they add to make a larger wave
* IN phase

Destructive interference

* The waves interact so they cancel each other
* OUT of phase

Diffraction

* NOT refraction
* When traveling waves encounter an obstacle or opening in a barrier, they “*move*” through or around it
* Particles do not diffract
* either go thru slit or dont

Blackbody radiation

* energy radiated by any object or system that absorbs all incident radiation
* Black Body Radiation illustrates that **temperature** is related to **energy**

Quantum Theory

* color/ intensity of emitted light changes as the temperature changes
* **COLOR** is related to *n*  and λ
* THUS **energy** has to be related to frequency and wavelength somehow
* made by Max Plank
* determined that a hot, glowing object could emit (or absorb) only ***certain*** quantities of energy

Energy and Frequency Formula

* E=*n*h*v*
* *E = energy of the radiation*
* *n* = quantum number; a positive integer (1, 2, 3…)
* v = frequency
* h= Planks Constant (6.626x10^-34)

Planks Constant

6/626x10^-34

The Quantum Theory of Energy

* Any object can emit or absorb ONLY __**certain**__ quantities of energy
* energy is quantized
* occurs in fixed quantities rather than continuous


* Each fixed quantity of energy is called a quantum
* atom changes energy “state” by emitting or absorbing one or more quanta of energy

Energy Changes

* Δ*E =* Δ*n*h*v*
* *E = energy of the radiation*
* *n* = quantum number; a positive integer (1, 2, 3…)
* v = frequency
* h= Planks Constant (6/626x10^-34)

Energy Formulas

E = h*v*  =  hc /λ

* E= hc/ λ
* E= hv
* V= c/ λ

Threshold Freq. in Wave Model & Real World

Wave model:

* intensity is responsible for observed E and e- will break off when it has absorbed enough light of any color

Real world

* the e- only breaks free when it is hit w certain color of light (certain v), regardless of brightness

Time Lag in Wave Model & Real World

Wave Model

* if the light is dim, less E is absorbed, so the e- should have to spend more time absorbing before it can break free

Real World

* current begins to flow immediatley when it is hit w appropriate color of light, again, regardless of brightness

Photon Theory

Threshold Frequency:

* intensity represents the number of photons, not the E.
* E is related to v, so an e- must absorb a photon of a certain minimum color to break free

Time Lag

* photon either has enough energy to free e- in one hit or it doesnt; the e- cannot store energy until it has enough

Line spectrum

* series of fine lines at specific frequencies separated by “black spaces”
* Each atom of a particular element has its own unique line spectra (aka emission spectra)

Bohr’s Model of the Hydrogen Atom

* made of ORBITS not orbitals!


1. The H atom has only certain energy levels, stationary states


1. The higher the energy level, the farther the orbit is from the nucleus
2. The atom does not radiate energy while in one of its stationary states
3. The atom changes to another stationary state __**ONLY**__ by absorbing or emitting a photon


1. The energy of the photon (hn) equals the difference between the energies of the two energy states

quantum numbers and electron orbit

* n (quantum) positive integer that reps radius of e orbit
* lower the n value, the smaller the radius of the orbit, and the lower the energy level
* When the electron is in an orbit closer to the nucleus (lower n), more energy is required to move it out of that orbit than when it is in an orbit farther from the nucleus (higher *n*)

Ground state

* When the electron is in the first orbit (n=1), closest to nucleus, H atom is in its lowest (1st) energy level

excited state

* if electron in any orbit further from nucleus, atom in excited state
* second orbit (n=2) = first excited state, third orbit (n=3) = second excited state etc etc

Absorption & Bohr Model

If a H atom absorbs a photon whose energy equals the difference between lower and higher energy levels, the electron moves to the outer (higher energy) orbit

Emission & Bohr model

If a H atom in a higher energy level (electron in a farther orbit) returns to a lower energy level (electron in a closer orbit), the atom emits a photon whose energy equals the difference between the two levels

Quantum staircase

The energy difference between two consecutive orbits decreases as *n* increases

* absorption & emission = inversely related

Rydberg’s equation: NRG transition problem (constant provided)

* used to to solve for the wavelength of a spectral line or energy-level transitions

Limitations of Bohr’s Model

* ONLY works for Hydrogen
* fails completely when you introduce more than one electron to the system
* MAJOR FLAW: assumes *electrons move in fixed, defined orbits*

Emission Spectrum

* Occurs when atoms in an excited state *emit* photons as they return to a lower energy state
* Some elements produce an intense spectral line that is evidence of their presence


* **Flame tests –** performed by placing a granule of an ionic compound or a drop of its solution in a flame

Absorption Spectrum

* “opposite” of an emission spectra
* Produced when atoms ***absorb*** photons of certain wavelengths and become excited
* Sodium’s absorption spectrum shows dark lines at the same wavelengths as the yellow-orange lines in sodium's emission spectrum

**Theory of Relativity**

**matter and energy are alternate forms of the same entity**

de Broglie Wavelength equation

* an equation for the wavelength of any particle of mass m moving at speed u (substituted for c)
* Matter behaves as though it moves in waves
* An object’s wavelength is **inversely** proportional to it’s mass and speed

Heisenberg’s Uncertainty Principle

* impossible to know, simultaneously, the position ***and*** momentum of an particle

locations of electrons

* dont know exact position of an electron, but can determine where it *probably* might be
* Solving the wave function gives the ***probability density***, a measure of the ***probability*** of finding an electron of a particular energy in a particular region of the atom

Quantum numbers and Atomic Orbitals

* atomic orbital specified by 4 quantum numbers:
* Principal quantum number
* Angular momentum quantum number
* Magnetic quantum number
* spin quantum number

Principal quantum number

* n
* positive whole # (1, 2, 3…)
* Indicates the relative *distance from the nucleus (tells u how far u r from nucleus)*
* Specifies the energy level
* orbital energy (size)

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Angular momentum quantum number

* (*l*)
* integer from 0 to ***n*** **– 1**
* shape of the orbital
* S,P,D,F

Magnetic quantum number

* (*ml*)
* integer from –*l* to +*l*
* Describes the 3D orientation of the orbital in the space around the nucleus (what orientation L is in n state)

Spin Quantum Number

* Ms
* +1/2 or -1/2
* direction of e- spin

how to get orbitals from quantum numbers

Energy Levels

* The levels (given by n) are divided into sublevels (or subshells), given by the *l* value
* *l* = 0 is an *s* sublevel
* *l* = 1 is an *p* sublevel
* *l* = 2 is an *d* sublevel
* *l* = 3 is an *f* sublevel

S Orbital shape

* spherical shape w/ nucleus in center
* has only ONE *ml* value

P Orbitals shape

* have two regions (lobes) of high probability of finding an electron, one on either side of the nucleus

D Orbitals shape

Pauli Exclusion Principle

* each orbital may contain a maximum of 2 electrons, which must have opposite spins

aufbau principle

* electrons are always placed in the lowest energy sublevel available

**Hund’s rule**

* when orbitals of equal energy are available, the lowest energy electron configuration has the **maximum number of unpaired electrons** with parallel spins

S orbital electrons

* (*l* = 0)
* max number of e—s = 2
* *ml* = 0, so there is only one atomic orbital

P orbital electrons

* (*l* = 1)
* max number of e—s = 6
* *ml* = -1, 0, +1 → three atomic orbitals

d orbital electrons

* (*l* = 2)
* max number of e—s = 10
* *ml* = -2, -1, 0, +1, +2 → five atomic orbitals

f orbital electrons

* (*l* = 3)
* max number of e—s = 14
* *ml* = -3, -2, -1, 0, +1, +2, +3 → seven atomic orbitals

Nuclear Charge (*Z*)

* A higher nuclear charge (more protons) increases nucleus-electron attractions, lowering the sublevel energy and stabilizes the atom (lower E = good!)

Shielding

* each electron “feels” presence of others so each electron __**shields**__ the others from the nuclear charge (charge of the nucleus)
* Essentially, each e— is blocking some of the nucleus’s attraction from other nearby e—

**effective nuclear charge (*****Z***eff)

* “full” nuclear charge is reduced to an __**effective nuclear charge (*****Z***____eff)__, the nuclear charge an electron __*actually*__ experiences

Penetration

* increases nuclear attraction and decreases shielding
* The better an outer electron is at penetrating through the electron cloud of inner electrons, the more attraction it will have for the nucleus

stability of sublevels

s < p < d < f

e configuration

Half-filled exceptions!

* **Cr** (Z=24) → \[Ar\] 4s2 3d4 → __[Ar] 4s13d5__
* Mo → __5s1 4d5__
* Cu (Z=29) → \[Ar\] 4s1 3d10
* Ag→ 5s1 d10
* Au→ 6s1 4f15 5d10

writing e config→ types of electrons

1. full is based off atomic #
2. condensed is valence electrons
3. inner electrons are the electrons which get replaced by a noble gas

atomic size

\*transition metalls increase down but dont really change ACROSS

ionization energy

* energy required to remove e

exceptions to ionization energy trend

* nitrogen: 1s2 2s2 2p3
* stable half-filled structure
* Taking an e— from N would make it **less stable** (higher energy)
* Oxygen: 1s2 2s2 2p4
* one e— beyond stable
* Taking an e— from oxygen would make the atom **more stable** (lower energy)
* It is easier to remove an e— from O (creating stability) than it is to remove one from N (destroying stability)
* True for Be/B, N/O, Mg/Al, P/S, Ca/Ga, As/Se

Successive Ionization Energies

* For a given element, IE1, IE2, and so on, increase because each electron is pulled away from a species with a higher positive charge
* This increase includes an enormous jump __***after***__ the __**last valence electron**__ has been removed because *much* more energy is needed to remove an inner (core) electron

ID an element from its IEs

* identify the largest increase in IE
* occurs after last Ve is removed
* that increase identifies Ve
* period X (will be given) element which has that # valence electron