answer the prompt belowwrite a lab report about experiment 12 in microbiology :Identification of enterobacteriaceae following the instruction provided

                                                                              Name __________________________________

Date ____________________________              Section_________________________________

INTRODUCTION

The atom is the smallest unit of matter. The structure of an atom determines whether or not it is reactive with other atoms, and what types of bonds will form between atoms. Our current understanding of the atom is based on two models – the quantum mechanical model and the Bohr model, both of which were developed in the early twentieth century.  These models explain the role of electrons and the reactivity of an atom, and how the physical and chemical properties of the elements result from the number of electrons each element possesses.  The quantum mechanical model helped scientists to formulate the modern-day periodic table, and further explains chemical bonding. 

The quantum mechanical model would not have been possible without the analysis of the characteristics of light.  Light does not have mass and therefore cannot be matter.  It is a form of electromagnetic radiation, energy travelling at a speed of 3.0 x 108 m/s.  In addition to its speed, light has three characteristics that define it:  the wavelength, (l), the frequency, (n), and the amplitude, (A). The wavelength is the distance between two wave crests. It is commonly measured in nanometers (nm), although it can be measured in any unit of length. Frequency is the number of waves that pass through a point in a given amount of time, and is typically measured in Hertz (Hz), or s-1. (1 Hertz is one wave per second, or 1/s.)  The amplitude is the intensity of the light, represented by the distance from the center of a wave to the crest.

Figure 13.1 – The Relationship between Frequency, Wavelength, and Speed in a Wave

Chemistry by OpenStax is licensed under Creative Commons Attribution License v4.0

Figure 13.2 – Summary of the Relationship between Wavelength, Energy, and Frequency

Chemistry by OpenStax is licensed under Creative Commons Attribution License v4.0

Wavelength and frequency are inversely proportional – as one increases, the other decreases and vice versa.  Light is part of the electromagnetic spectrum, shown in Figure 13.1 below. The entire spectrum includes not only visible light, but also X-rays, microwaves, radio waves, infrared radiation, ultraviolet radiation, and gamma rays.  As the frequency of the radiation increases, so too does the energy of the light. For example, radio waves have long wavelengths (~10 meters), while gamma rays have very short wavelengths (~0.01 nm, or 1×10-11 meters). Radio waves are benign because they carry low amounts of energy while gamma rays are very high energy and present health risks to those exposed to large amounts of gamma radiation. For visible light, the color of the light is determined by its wavelength. Red light has a wavelength around 700 nm while blue light has a wavelength around 450 nm. Therefore, red light carries less energy than blue light.

Figure 13.3 – The Electromagnetic Spectrum

 CCO 1.0

When an atom is exposed to energy, it can absorb that energy, and an electron in its valence shell is promoted from a lower energy state to a higher energy state. At a later time, the electron will ‘relax’ back to a lower energy state, and when it does it gives off the excess energy as light.  The wavelength of the light corresponds to the energy that is emitted. Light emitted by electrons of an element can be separated into its constituent wavelengths, producing an emission spectrum.  The emission spectrum for atomic hydrogen is shown in Figure 13.2 below.  This spectrum is not continuous, and consists of bright spots at specific wavelengths for each element.

 Figure 13.4 – The Emission Spectrum for Atomic Hydrogen

 CCO 1.0

Using the emission above as an example, there are four distinct wavelengths for hydrogen at 410 nm, 434 nm, 486 nm, and 656 nm. As atoms of each element are different, the amounts of energy that are emitted are different, so the emission spectrum is a different for each element. In this laboratory experiment, we will use PhET simulations to explore the various models that describe the behavior of electrons in an atom.

PROCEDURE/DATA COLLECTION:

Part 1 – Phet Simulation

1. Obtain a laptop from the laptop cart in the lab. Download and open the Molecules and Light simulation:         https://phet.colorado.edu/en/simulation/legacy/hydrogen-atom

2. While the program is initializing, record the following information about a hydrogen atom:

            # of protons:

            # of electrons:

            Electron configuration:

            Outer shell:

3.  Using your textbook or the internet, define a photon. Is the photon a particle or a wave?  Is it both?  If so, how is that possible?

4.  What characteristic of light determines the color of a photon?

5. In the simulation, turn the light beam “on”.

6.  Making sure the “experiment” is highlighted in white, observe what is happening while photons encounter a hydrogen atom (the box with a question mark).  Write your observations below:

7.  When studying the structure of an atom, scientists witnessed something similar to what you are witnessing now.  They then deduced how the atom must be organized.  What do you think is making the photons deflect?  Is every color deflected?

8. Change the Light control from “White” to “Monochromatic”.  Are photons still being deflected?

9.  Click the “show spectrometer” box. 

10.  Change the wavelength of the photons by moving the handle up and down the spectrum.  Record your observations in the chart below:

Color
Observation

UV (97 nm)
   

Purple (410 nm)
   

Green (550 nm)
   

11. Click the “Show absorption wavelengths” box.

12.  What is the spectrometer box keeping track of? Is this consistent with your answer in #7? If not,   

        theorize about what might be happening as the photons encounter the hydrogen atom.

Part 2 – Electron Configurations

Part A – Counting Electrons for a Neutral Element

The first step in writing the electron configuration of a specific element is to determine the total number of electrons that the element has. The total number of electrons in a neutral element is equal to the total number of protons. Recall that the number of protons for any element is equal to its atomic number, and can be found on the periodic table.

For each of the following neutral elements, determine the total number of electrons.

Practice:

Mg       12 total electrons
Fe        26 total electrons
Example #1:

Ag        ___ total electrons
Cl         ___ total electrons
Ca        ___ total electrons
Part B – Writing Electron Configurations

Within an atom, electrons reside in orbitals that are located at specific, fixed distances away from the nucleus. These orbitals are specified by both a number and a letter. The number, called the principle quantum number (n = 1, 2, 3, …), specifies the principle shell. The principle shells that are closest to the nucleus (smaller n value) are at a lower energy than those further from the nucleus (larger n value). Within each principle shell, there are subshells (specified by the letters s, p, d and f) that indicate specific areas where the probability of finding an electron is greatest.

Writing electron configurations can be thought of as ‘filling’ the electron cloud, beginning with the lowest energy orbital first. The order in which the orbitals should be filled can be determined using the diagram below. Starting at 1s, the red arrows and dotted lines follow the order of filling. For example, as the diagram demonstrates, the 3d subshell begins filling only after the 4s has been filled. Orbitals and subshells can also be partially filled if the element does not have enough electrons to completely fill it.

The maximum amount of electrons you can place in a specific subshell depends on the shape of the subshell (represented by the letters s, p, d, or f). Once a subshell is full (contains the maximum number of electrons), the next subshell then begins filling. The maximum number of electrons that can be placed in a specific subshell is shown below:

Subshell
Maximum Number of Electrons

s
2

p
6

d
10

f
14

The number of electrons that are contained within a specific subshell are represented as a numerical superscript. For example, the electron configuration for lithium (Li), which has 3 total electrons, is 1s22s1. This shows that 2 electrons exist in the 1s subshell and 1 electron exists in the 2s subshell.

Write the complete electron configuration for each of the following elements:

Practice:

a) Si     1s22s22p63s23p2                              b)   Co    1s22s22p63s23p64s23d7

Example #2:

Na        _____________________                  b)    Mn                        _____________________
Electron configurations can also be written in the shorthand noble gas form, which is especially useful for larger atoms that have many electrons. If we take a look at the electron configuration for sodium (Na), for example, it contains the electron configuration of the noble gas neon (Ne), plus one additional electron. We can put this noble gas in brackets and then write the rest of the electron configuration next to it, as shown below:

Write the shorthand noble gas electron configuration for each of the following elements:

Practice:

Br        [Ar] 4s23d104p5                b)            Ir          [Xe] 6s24f145d7

Example #3:

Te        _____________________      b)   Ba  _____________________
       c)  Po       _____________________

Part C – Writing Electron Configurations for Ions

Ions form when a neutral atom gains or loses electrons. The charge of an ion is given by:

Metal atoms will lose electrons, and the number of electrons lost is equal to the positive charge. For example, a Li+ ion started with 3 electrons (as determined from the atomic number), and lost one electron (because it has a +1 charge), so has 2 remaining electrons. The electron configuration for Li+ is 1s2.

Write the electron configuration for each of the following ions:

Practice:

            a)  Ca2+                 1s22s22p63s23p6                  b)  N3-      1s22s22p6

Example #4:

Br –       _____________________      b) Pb2+ _____________________

 

 

 

POST-LABORATORY QUESTIONS:

1.  Ultraviolet light has sufficient energy to cause damage to DNA and the skin.  Compare the energy of UV light to those of IR and Visible. Explain why UV light is more dangerous than the other two lights.

2.  An electron in a hydrogen atom goes into two transitions: the electron falls (a) from the n = 3 level to the n = 2 level, (b) from the n = 3 level to the n = 1 level.  Compare the radiation emitted by the two transitions in each term. Circle the correct choice.

Term
Transitions

Frequency
The frequency in the transition (n = 3  →  n = 2) is higher or lower   than the frequency in the transition (n = 3  →  n = 1)

Energy per photon
The energy per photon in the transition (n = 3  →  n = 2) is larger or smaller   than the Energy of a photon in the transition (n = 3  →  n = 1)

Wavelength
The wavelength in the transition (n = 3  →  n = 2) is longer or shorter   than the wavelength in the transition (n = 3  →  n = 1)

3.  Arrange a set of wavelengths for light in order of increasing frequency: λ (250 nm), λ (300 nm), and λ

     (350 nm). Explain your arrangement of the light in frequency.

4. Give the electron configuration for each of the following atoms & ions.  For a-d give the full configuration (do not abbreviate) and for the others use noble gas shorthand format.

a) Kr

b)  Mg2+

c) C

d) F–

e) Fe2+

f) I–

g) P

h) Ge

5. For each of the atoms or ions in question 4, give the number of valence and unpaired electrons based on the electron configurations.

a) # of valence electrons –

    # of unpaired electrons –

b) # of valence electrons –

    # of unpaired electrons –

c) # of valence electrons –

    # of unpaired electrons –

d) # of valence electrons –

    # of unpaired electrons –

e) # of valence electrons –

    # of unpaired electrons –

f) # of valence electrons –

    # of unpaired electrons –

g) # of valence electrons –

    # of unpaired electrons –

h) # of valence electrons –

    # of unpaired electrons –

PRE-LABORATORY ASSIGNMENT

1. Define the following terminologies:

            a) Electromagnetic radiation (light)

            b) The speed of light

            c) The frequency of light

            d) The wavelength of light

 2. Fill in the blank in the Table.  For the given criteria rank the following types of light from the smallest to the largest in order: radio waves, gamma rays, microwaves, red light, and blue light. 

 
smallest
small
middle
large
largest

Wavelength (λ)
gamma rays
 
 
 
 

Frequency (ν)
radio waves
 
 
 
 

Energy per photon
 
   
 
 
 

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