The Lewis Dot Structure for Acetic Acid (CH3COOH)

Created by MakeTheBrainHappy.

Above is the Lewis Dot Structure for Acetic Acid (CH3COOH). You could alternatively also draw the structure by including two dots for every bond. As you can see every single element has a filled valence shell with the two oxygen's each containing two lone pairs of electrons, the only instance of this phenomena within the Lewis Structure. In a sense this is a modified structure of methane (CH4) with the replacement of one hydrogen with the replacement of a carboxylic group (-COOH). As we will see the properties we observe within the Lewis structure have a significant impact on the properties of acetic acid.

The Greek Philosopher Theophrastus. Source

What is the history of acetic acid?

Due to the importance of different alcohols such as beer and wine in early civilizations, vinegar became one of the earliest chemical substances that was familiar to ancient peoples. Vinegar is formed by natural fermentation processes and contains approximately 5% acetic acid. One of the earliest mentions of acetic acid was by the Greek Philosopher Theophrastus who explained how to form different pigments, including those for white and green colors, with vinegar as an important constituent ingredient. 

Is Acetic Acid (CH3COOH) Polar or Nonpolar?

Acetic Acid (CH3COOH) is a polar molecule due to the presence of the functional group COOH, a carboxylic acid. It is also an acid in solution, releasing a small number of protons into solution which form H3O+. Due to its polar qualities it is found as a liquid at standard temperature and pressure in relation to its rather light molar mass. It has a melting point is between 16˚C to 17˚C while the boiling point is between 118˚C and 119˚C. These are again elevated due to the polarity of acetic acid. 

How is Acetic Acid (CH3COOH) utilized in the real world?

Acetic acid is useful due to its properties as a polar solvent and building block for other molecules, containing both a methyl (CH3) and COOH functional group. Nearly one third of produced acetic acid is utilized in order to produce "Elmer's glue" material. Inks, paints and coatings are also created via a reaction involving acetic acid. It is frequently utilized as a polar solvent in lab research settings and thereby has been involved in certain medical practices. As mentioned above acetic acid is also present in vinegar which has a variety of household uses; however, the acetic acid is diluted in water to a greater degree than in a research lab.  

Downloading FASTQ Files Quickly utilizing IBM’s Aspera Connect on a LINUX or MAC Machine

Please note that this guide is based on this BioStars thread and this GitHub repository. 

Procedure:

1. Download the latest version of Aspera Connect from the IBM “featured client software” section (you may need to install a browser extension as well)

2. Download the file “ena-fast-download.py” from wwood’s GitHub repository

3. Open the script in your favorite python editor and scroll down to the bottom. Add the specific file path to the ascp field (addition highlighted below)

aspera_commands = []

    for url in ftp_urls:

        quiet_args = ''

        if args.quiet:

            quiet_args = ' -Q'

        cmd = "/Users/USER/Applications/Aspera\ Connect.app/Contents/Resources/ascp{} -T -l 300m -P33001 {} -i {} era-fasp@fasp.sra.ebi.ac.uk:{} {}".format(

            quiet_args,

            args.ascp_args,

            ssh_key_file,

            url.replace('ftp.sra.ebi.ac.uk',''), output_directory)

        logging.info("Running command: {}".format(cmd))

        subprocess.check_call(cmd,shell=True)

logging.info("All done.") 

4. Save the file. Move the file to the directory where you want FASTQ files to be deposited into. Run the following command in terminal for each accession number (example highlighted):

./ena-fast-download.py ERR1739691 --ssh_key osx 

Note: Sequential commands can be done through a looped text file call or with the ; operator (ex. ./ena-fast-download.py key1 --ssh_key osx ; /ena-fast-download.py key2 --ssh_key osx)

Note 2: For this method it is recommended that you utilize a bash terminal.

The Lewis Dot Structure for KCl

Created by MakeTheBrainHappy.

Please find above the Lewis Dot Structure for KCl (Potassium Chloride). As per usual you could replace the one bond with two electrons. In the case for KCl the electronegativity difference between potassium and chloride is so strong (.82 vs. 3.16, respectively) that the bond is considered ionic. The electrons aren't really considered "shared"; rather, the valence electrons are nearly completely coopted by the chloride. This property as illustrated by the Lewis Dot Structure gives KCl many of its properties as will be explored in the following paragraphs.

KCl Dissolving in Water. Source

How does KCl (potassium chloride) act as a solute?

As a result of its ionic character as shown in the Lewis Dot Structure it has a great partial positive and partial negative charge on either end. Therefore it dissolves very well in polar solvents such as water (shown above). Water molecules essentially surround the different individual ions and thereby disassociate the salt into solution. However since you need many water molecules per ion there can be a point where you saturate the solution. This means that no more salt can be disassociated unless you add more of the polar solvent.

KCl Molecule. Created with MolView.

Is KCl (potassium chloride) polar or nonpolar?

As you may have guessed from the above paragraph, KCl is a polar molecule due to the great electronegativity difference discussed before. This incredibly large difference means that the compound is a solid at standard pressure and temperature with a melting point of 770˚C and a boiling point of 1420˚C. The act of disassociation as mentioned above is not equivalent to melting the pure compound into a liquid as we are discussing here. These properties are also presupposed in the Lewis Dot Structure for KCl.

KCl Salt. Source

How is KCl (potassium chloride) utilized in the real world?

There are a wide variety of uses for potassium chloride. There are different fertilizers which are composed of KCl for usage in agriculture. Due to the necessity of potassium in biological systems the salt form shown above is often consumed by those who are potassium-deficient. In regular food it can be utilized as a table salt (NaCl) substitute in order to lower the actual Na+ salt content in someones diet. Lastly it is found in many different niche industrial uses as a "ice" salt or in the creation of potassium metal for example. In conclusion, the different properties that allow for these utilizations are exemplified by the Lewis Dot Structure for KCl.

The Lewis Dot Structure for CS2

CS2 Lewis Structure. Created by MakeTheBrainHappy

This is the Lewis Dot Structure for Carbon Disulfide (CS2). As we've discussed before you could replace each bond with two electrons which would represent well how through the sharing of electrons each atom achieves a full octet. The combination of four electrons for the bonds and four electrons in lone pairs equals eight for sulfur while bonds supply eight electrons for the central carbon atom. This is a common feature for carbon as it is unlikely for this element in particular to have lone pair electrons.

The lewis dot structure for CS2 also predicts many of the properties of the molecule. Due to the presence of large sulfide atoms over for comparison oxygen atoms in CO2 the molecule has a greater capacity for temporary london dispersion forces. As a result there can be more induced dipoles which increases the solubility of CS2. The compound is slightly soluble within water and much more soluble in less polar compounds such as alcohol or even nonpolar compounds like benzene.

CS2 Liquid State. Source

How is CS2 utilized in the real world?

Due to the stronger induced dipoles CS2 also has a greater melting and boiling point. This means that the compound is a liquid at standard temperature and pressure. This concept is discussed further in our article on the polarity of CS2 which you can learn more about in the next section. There are only a limited number of uses for CS2 due to its pretty toxic properties. One of these uses is as a solvent for other nonpolar compounds. These include many elements like phosphorus and sulfur and also larger compounds like fats, resins and rubber. There also other uses for different chemical reactions involved in metallurgy and rubber production. One of the most popular reactions involving CS2 produces carbon tetrachloride (CCl4).

Is CS2 polar or nonpolar?

If you are interested in learning more about this subject, feel free to read this article on whether CS2 is polar or nonpolar. The article will also discuss some of the ways in which CS2 differs from CO2. In addition it will also go into the underlying reasons behind the determination of CS2s polarity.

Volcanic eruption. Source

How can CS2 be formed naturally?

CS2 (carbon disulfide) can be formed naturally by volcanic eruptions or in marshes. The way this reaction occurs is actually very compatible with how you can see the molecule within the Lewis Dot Structure. Elemental carbon and sulfide combine at very high temperatures to form CS2. However, it should be noted that there is only a very small amount of the compound formed in this process.