CHEM 111 - Experiment 2 (both parts!)

what are some limitations of using the Beer-Lambert law and calibration graphs?

- the line is not always straight, so you have to use a line of best fit! - assumes that the radiation reaching the sample is of a *single wavelength*—that is, that the radiation is purely monochromatic - there is stray radiation

White light

- the total electromagnetic spectrum - ranges from about 400 nm - 800 nm - made up of the colors of the rainbow

Percent Error calculation

% error = [ experimental value - accepted value ] / accepted value ] x 100

how would we find the *concentration* of an unknown solution?

(C) = A/(b x Ɛ) remember... A = [ log(I^0/I) ] there is a *linear relationship between concentration and absorbance* A = Slope x concentration + y intercept plot the numbers given in desmos! m = _______ b = ________ plug them into our equation! solve and round to three sig figs vids for help: https://www.khanacademy.org/science/ap-chemistry-beta/x2eef969c74e0d802:intermolecular-forces-and-properties/x2eef969c74e0d802:beer-lambert-law/v/spectrophotometry-example

Ok here's some math using the Beer Lambert Equation! Just a plug and chug no worries here! A solution of blue dye has a *concentration of 0.081 M* and a *molar absorptivity coefficient ε = 5.81 M-1 cm1* at *590 nm*. If the *path length of the cell is 1 cm*, *what is the absorbance of the dye* solution at 590 nm measured on a spectrophotometer?

*A = ebc* Therefore: A = 5.81 M-1 x 1 cm x 0.081 M *A = 0.471 M-1 cm*

How to determine frequencies and energies associated with wavelengths?? YAY MATH

*E = hv* E : energy (J) h : 6.626 x 10^-34 (Js) (Planks constant) v : frequency ( Hz or sec^-1 or [1/sec] ) use this equation to find *energy*! c = λv c : speed of light (3.0 x 10^8 m/sec) λ : wavelength (m) v : frequency ( Hz or sec^-1 or [1/sec] ) use this equation to find *frequency*!

Differences between a *line* and *continuous spectrum* and be able to identify *examples of each*

*line spectrum* : produced by excited atoms in the gas phase and and contain only certain frequencies, all other frequencies are absent. *ex)* every chemical on the periodic table has a unique line spectrum! *continuous spectrum* : produced by the sun and heated solids. The emitted radiation contains all frequencies within a region of the electromagnetic spectrum. *ex)* a rainbow and light from a lightbulb

How would we identify a maximum wavelength from a visible spectrum?

1) find the wavelength 2) Find the transmittance percentage 3) use the equation A = 2 - log (%T) 4) this will give you your absorption! 5) plot the wavelengths with their absorption to find the maximum one in an absorption spectrum !

What is a line spectrum? Give an example of a line spectrum

A line spectrum is one that *shows only certain colors* (or specific wavelengths) of light. *The visible light spectrum is a great example of this!* It is a very narrow range of electromagnetic radiation detectable by human eyes, has wavelengths of *400—750 nm*, and the different wavelengths are perceived as different colors of the rainbow.

What is a "blank" sample? What is the benefit to using a blank sample?

A spectrometer needs to be calibrated against a blank sample/solution so that measurements taken can use the blank solution's absorbance as a a zero reference.

Beer's Law (variables and how each of the variables effects the observed absorbance!)

A=ebc A : absorbance [ log(I^0/I) ] e : molar absorptivity b : path length c : concentration relates the amount of material in a sample to the absorption of light passing through it If you get stuck or need that visual again: https://www.youtube.com/watch?v=zuUvQN8KXOk

What is an *absorption spectrum* and how does it differ from an *emission spectrum*?

An *absorption spectrum* is a plot of absorbance (amount of light absorbed) as a function of wavelength of incident light. The absorption of photons of the appropriate energy promotes electrons from a lower orbit to a higher one. An *emission spectrum* differs in that it is a *spectrum produced by excited atoms* (when they jump down). *absorption* -> colors absorbed and photons absorbed *emission* -> colors *emitted, given off*, and the energy released when jumping back down from the excited state

How can we use Beer's Law to determine the concentration of an unknown solution?

Beer's law states that if a *monochromatic radiation* is allowed to fall on a solution, then the *amount of light absorbed or transmitted* is an *exponential function* of the *concentration* of the *absorbing substance* and of the *length of the path of light* through the sample.

Planck's Law

E=hv E - the photon energy in Joules v - frequency of the radiation (Hz or s-1) h - Plank's constant (6.63 x 10^-34 Js)

mmMm math! but you go this luv. What is the *wavelength* (in nm) of a photon that has *frequency = 4.40 x 1014 Hz*?

Well... Wavelength = Speed of light / frequency Wavelength = 299,792,458 m/s (/) 4.40 x 1014 Hz = 682

c = λv

c - speed of light (3 x 10^8 ms^-1) λ - wavelength of the radiation (usually in nm) v - frequency

How would we use the *Bohr equation* to calculate the *wavelengths* expected for the *Balmer series* of the *hydrogen atom*? Everything you need to know!! :D

n = discrete energy states in a hydrogen atom (1,2,3, etc) n = 1 is the lowest energy level - when a hydrogen atom jumps down from an energy level it releases a photon with a *wavelength* THAT's what we're trying to calculate! *example time* (n = 4 -> n = 2) Use *Bohr's Equation* (look at picture) and plug in the numbers to solve! the 2.178 is called Rydberg's constant! If you get stuck: https://www.youtube.com/watch?v=mXxsT1ut35Q

Electromagnetic radiation from lowest to highest energy:

radio waves, infrared, visible light, UV, Xray, gamma rays

Define spectroscopy

the study of how how *photons* of light *interact with matter*

Why is λmax is chosen for analysis with the Beer-Lambert Law??

this absorption spectrum displays *how much light*, through a band of frequencies, is able to *pass through a sample* to reach the detector where there are *peaks*, it means that *not as much of the light of that wavelength is getting to the detector* this means it's being *absorbed* by the sample! if we keep collecting this data throughout a chemical reaction, it can show us *how rapidly the absorbance changes*, and thus, how the *concentration* of a particular substance is *changing* over time

colors of the spectrum from *highest to lowest frequency* :

violet, indigo, blue, green, yellow, orange, red

A dye solution strongly absorbs violet light. What color does the solution appear?

yellow!

*Energy* and *frequency* are *directly proportional* to each other!

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