Blog #8

This past week we focused on describing and balancing chemical reactions, and also worked on rearranging atoms. For balancing chemical reactions, we had to put coefficients in front of the reactants and the products to have equal amounts of each element for both of them. We noticed when doing this that after the reaction occurred no mass was added or subtracted in the product. These coefficients tell us how much of that substance you have. The subscripts tell how many atoms you have of the substance like in NH3, there are 1 Nitrogen and 3 Hydrogens. Balancing some equations can be tricky when there are an even and odd match up of an element. An example would be the reaction 2CH4+3O2=2CO+4H2O. How I got the coefficients was by first putting a 2 in front of the H2O to have 4 hydrogens and 2 oxygens. On the product side I had 3 oxygens in total while the reactant side had 2. How you fix that is by putting a 3/2 in front of the O2 (reactant) and multiplying everything else by two, which is how I came to my final answer. Below is a whiteboard of some chemical reactions that have been balanced out, and particle drawings are shows as well for before and after the reaction.

We also described chemical reactions from doing many mini labs testing different things. We did combination reactions where as an example we took a strip of magnesium and ignited it with fire to see a bright flame which is shown below.

We also described chemical reactions from doing many mini labs testing different things. We did combination reactions where as an example we took a strip of magnesium and ignited it with fire to see a bright flame which is shown below.

After we added P-indicator, it showed that there was a presence of a hydroxide ion from the reaction. We then did decomposition reactions where for example solid ammonium carbonate was being heated, and started shaking rapidly because it was starting to form into a gas.

Then we had our single replacement reactions where an example of one was placing a small piece of wire copper in a well plate and adding a few drops of .1M AgNO3 to it, which kept increasingly forming little crystals after a few minutes. This reaction is single replacement because an oxidation reduction chemical reaction occurred where an element moved out of the solutions compound and into the copper forming the crystals.

After this we had double replacement reactions where for example we added a drop of .1M AgNO3 to a transparency sheet and then added .1M NiCl2 to it as well creating this product that was foamy looking, and a mix of both of their colors. It was called double reaction because the cations and anions switched between both reactions to form a new product.

Lastly, we did combustion reactions. We put 10 drops of isopropyl alcohol in a small dish and lit it on fire with a butane lighter. We took a cold watch glass and held it above the fire to see condensation form on the bottom of it. Oxygen was reacted with hydrocarbon to produce energy and light, which was why it was in the combustion category.

In all, I feel that my overall understanding of balancing chemical reactions is a bit shaky, and I usually need help with the harder problems. I need help with knowing exactly what steps to take when I’m given a problem so I don’t start confusing myself, which is what usually happens. I would give myself an overall rating of a 5 or 6 out of 10, but hopefully that will change in the next few days as finals approach!

Blog #7

This past week we focused on naming ionic and molecular compounds, and understanding their differences. With naming ionic compounds, we identified elements that were cations such as: Hydrogen, Lithium, and Magnesium, which are positively charged ions. Also we identified anions such as: Iodine, Chlorine, and Nitrogen, which are negatively charged ions. When looking at ionic compounds, we realized that the metal of the compound is always listed before the non-metal. Also, in an ionic compound both elements have to zero out to become balanced or neutral. For example with the compound Aluminum Oxide, there needs to be two aluminums, and three oxygens since aluminum has a 3+ charge while oxygen has a 2- charge. In this previous example, aluminum oxide is the made up of the regular element name for the metal, but for the non-metal an ide is added on as an ending so instead of saying Aluminum Oxygen, it would be Aluminum Oxide. Other examples would be K3N which is Potassium Nitride, or MgF2, which is Magnesium Fluoride. In some cases certain metals from multiple cations. The metal Copper can have 1 or 2 cations. The way to specify it is by putting the corresponding roman numeral behind the metal in the compound name giving the charge of the element. An example of this would be Cu2O, which is Copper (I) Oxide. Another metal that has this is lead. There can be 2 or 4 cations that correspond with it. Ionic compounds have ionic bonds, and can conduct electricity since its bonds are weaker and allow electrons to pass through.

In contrast to ionic compounds, molecular compounds are only made of non-metals, and have covalent bonds. They do not conduct electricity since there are no metals present, and their bonds are much stronger than ionic bonds, meaning that electrons can’t slip through. With these type of compounds, prefixes and suffixes are added to the element names depending on how many of each element there are. An example would be with the molecular formula BCl3 the name would be Boron trichloride, since there are 3 chlorides and 1 boron. The prefix mono is only used when there is one of the second element, such as ClF. With that compound it would be named Chlorine monofluoride. There are different rules with the spelling and use of certain prefixes and suffixes such as use tetra when having 4 of an element when there isn’t a vowel that would come after, and use tetr if a vowel comes after it. An example of this would be N2O4, which is Dinitrogen tetroxide since an o was the first letter after the prefix. As you can see these compounds are not only different from ionic compounds in that they are made of non-metals, but they also don’t zero eachother out to become neutral. The only time elements need to have equal charge is with ionic compounds.

With naming ionic compounds, we also drew out their ions and formula unit. Below is a few white boarded examples of some different compounds.

How we get the number of ions is simply by counting how many ions of each element is needed to equal each other in charge. An example would be in CaBr2 there would be 3 ions needed, one ion of Ca and 2 of Br.

Overall I have a good understanding of the main differences between Ionic compounds and Molecular compounds and how each work with classifying them. I would give myself an overall rating of a 9 because from last week to this week I really locked the concept of how to name these compounds and how to make sure that the ionic compounds have a balanced charge. One question I still have is what are Polyatomic ions, and how they compare to molecular and ionic compounds.