Acids & Bases- Oct. 8 Class

In today's class we learned the characteristics of acids and bases.

Acids
- Solid, Liquid, or gas at SATP (25 degrees celcius, 100 kPa)
- Form conducting aqueous solutions
- Turn blue litmus red
- Dissolve in water to produce H+
- Taste sour


Bases
- Turn red litmus blue
- Slippery
- Non-conductive
- Dissolve in water to produce OH-




Naming Acids
- Acids are aqueous (dissolve in water)
- Hydrogen compounds are acids

• HCL(aq) --> Hydrochloric
• AcidH2SO4(aq) --> Sulfuric Acid
  

- Hydrogen appears first in the formula unless it is part of a polyatomic group

• CH3COOH(aq) --> Acetic Acid

- Classical rules use the suffix "ic" and/or the prefix "hydro"

• Eg. Hydrochloric Acid

- IUPAC system uses the aqueous hydrogen compound

• Eg. HCl(aq) --> Aqueous Hydrogen Chloride

Naming Bases
- For now, all bases will be aqueous solutions of ionic hydroxides

• NaOH
• Ba(OH)2
  

- Use the cation name followed by hydroxide

• Sodium Hydroxide
  
• Barium Hydroxide
  

Examples
- HI(aq) --> Hydrochloric Acid
- H3PO4(aq) --> Phosphoric Acid
- H3PO4(aq) --> Phosphorous Acid
- HNO3(aq) --> Nitric Acid
- HNO2(aq) --> Nitrous Acid
- Mg(OH)2 --> Magnesium Hyrdroxide
- HBr(aq) --> Hydrobromic Acid
- HOOCCOOH(aq) -- Oxalic Acid


How to make your own PH indicator between an acid and base!

October 6th 2009

Today in chemistry class Mr. Doktor showed us ...

Hydrates

- Some compounds can form lattices that bond to water molecules
Example: Copper Sulfate
- These crystals contain water inside them which can be released by heating
- Without water the compound is often preceded by “anhydrous”

Naming Hydrates

1. Write the name of the chemical formula.
2. Add a prefix indicating the number of water molecules.
3. Add hydrate after the prefix

Molecular Compound

- Composed of two or more non-metals.
- Low melting point and boiling point.
- Share (not exchange) electrons
- Usually end in –gen or –ine
Example: Hydrogen, Oxygen…
- 7 molecules are diatomic
They are… Hydrogen, Nitrogen, Oxygen, Fluorine, Cholrine, Bromine and Iodine
Two molecules are polyatomics.

Chemical Nomenclature - Oct 2 Class

In the beginning of class, Mr. Doktor showed us an example of electrolysis. He used pickles, water, and a battery charger. He then touched the pickle with the two wires, and you could see smoke come out and the pickle being electricuted. He then touched the water with the wires and smoke came out as well. You could see the hydrogen and oxygen as the smoke and you could smell this funky burning smell. You should not try this experiment at home due to dangerous results such as being electricuted or starting a fire.



Here is a video of an example of electrolysis.

http://www.youtube.com/watch?v=Or22ktW8btc&feature=related
This is a quick video that shows how hydrogen and oxygen atoms blow up a balloon. In my assumption, i believe that the balloon blew up due to evaporating water.



After that experiment in class, Mr. Doktor gave us examples and questions about naming compounds. Here are the examples done in class:
Name:
1. PbS2 - Lead (IV) Sulphide
2. MgO - Magnesium Oxide
3. CuCl2 - Copper (II) Chloride
4. Cr2O3 - Chromium (III) Oxide

Chemical Nomenclature
- Naming chemical compounds has been a very difficult task and different systems have been used through the centuries.
- Today the most common system is IUPAC for most chemicals.
- Ions
* Binary Ionic
- Polyatomic Ions
- Molecular Compounds
- Acids



Chemical Formulas
Be aware of the differences between ion and compound formulas.
Zn²+ - ion charge
BaCl2 - number of ions (Subscript)

Naming Ions
- For metals, use the name of the element and add ion
Eg. Al³+ = Aluminium Ion
- For non-metals, remove the original ending and add -ide
Eg. F- = Fluorine -> Fluoride

Polyatomic Ions have special names.

Binary Ionic
Steps:
1. Write the formula for the cation first and then the formula for the anion. (Cation = positive because "cat" = "pussytive" & Anion = Negative)
2. Criss cross charges moving the numbers below.
3. Reduce ion numbers to lowest common multiples, omit 1 and omit charges.
Eg.


Example
- Write the chemical formulas of:
Aluminum Fluoride = AlF3
Sodium Oxide = Na2O
Iron (III) Sulphide = Fe2S3

Multivalent Ions
- Some elements can form more than one Ion
Eg. Iron -> Fe³+ or Fe²+ / Copper -> Cu²+ or Cu¹+
- The more common ion is the top one of the P.T.
- IUPAC uses roman numerals in the parenthesis to show the charge.
- Classical (ie old) systems uses latin names of elements and the suffixes -ic (larger charge) and -ous (smaller charge)
FeO -> Ferrous Oxide
Fe2O3 -> Ferric Oxide

Other Classical Names
- Ferr = Iron
- Cupp = Copper
- Mercur = Mercury
- Stann = Tin
- Aunn = Gold
- Plumb = Lead
Chemical Formula: Sodium Nitrate -> NaNO3 (Na+/NO3-)
Barium Phosphate -> Ba3(PO4)2

Classification of Matter - September 30 Class

In this class, we learned about Homogeneous and Heterogeneous substances.. We distinguished that tap water is a heterogeneous substance because there are other chemicals that are put into the tap water when running through the drains, sinks, etc. Mr. Doktor also passed around a beaker filled with aluminum pieces and salt, a heterogeneous mixture that didn't react together. We also learned that there are certain ways of separating mixtures, which are mostly physical changes.

Notes:
Classification of Matter

• Understanding matter begins with how we name it. We can divide matter into two parts: Homogeneous substances and heterogeneous substances
  
• Homogeneous: consists of only one visible component
  - distilled water, oxygen, graphite
  
  
  
• Heterogeneous: contain more than one visible component
  - chocolate chip cookie, granite

Pure Substances:

• There are 2 types of Pure Substances:
  Element: substances that cannot be broken down into simple substances by chemical reactions
  - oxygen, iron, magnesium
• Compound: substances thar are made up of 2 or more elements & can change into elements (or other compounds by chemical reactions)
  - water, sugar

Telling the difference

• It is often very difficult to know if something is an element or a compound
  - The differences are only "visible" on the atomic level
• One method is to connect the substance to an electric current. This technique. called electrolysis can split the compound apart into its constituent compounds

Solution

• A solution is a homogeneous mixture of 2 or more substances
  - Solutions usually involve but don't have to (fog, steel)
• The compound present in greater mount is the solvent
  - Water is the most ommon solvent
  - The symbol (aq) is used when something is dissolved in water
• The compound present in smaller amount is the solute
  - In salt water, salt is the solvent

Mixtures

• Many mixtures are easy to identify (chocolate chip cookies) but others are easily confused as pure substances
• In heterogeneous mixtures.. the different parts are clearly visible (granite, sand)
• In homogenous mixtures.. the different parts are NOT visible (salt, water, air, brass)

Separating Mixtures

• There are many methods to separate mixtures, depending on how the type of mixture
  - By Hand (Heterogenous mixture only)
  - Filtration (Heterogeneous mixture only)
  - Distillation
  - Crystallization
  - Chromatography
• All of these are physical changes

Matter and its Changes- Sept 28 Class

In Monday's lesson we learned the States of Matter, Changes in Matter, Physical Changes, Phase Changes, Chemical Changes, and the Conservation of Matter. Also, we got to watch a youtube video of mythbusters as an example to figure out the chemical and physical changes that occurred.


What is Matter?
-Anything that has mass and occupies space
-Matter can exist in many different states, the most common are:

• Solid, Liquid, Gas
  
• Plasma, Aqueous (dissolves in water), Amorphous
  
  

-Solids: holds one shape and has a definite volume (strong bonds)
-Liquids: can change shape, but has a definite volume (weak bonds)
-Gas: can change shape & volume (no bonds)
-Aqueous: something dissolved in water
-Plasma --> ionization


Changes in Matter
- Matter can undergo many changes
- Nearly all changes can be broken down into 3 categories:

• Physical Changes
  
• Chemical Changes
  
• Nuclear Changes
  

Physical Changes
- Involves changing shape or state of matter (crushing, tearing, etc.)
- No new substances are formed (Eg. Boiling water, cutting wood,
smashing cars)




Phase Changes
- Changing from a solid to a gas can often be confused as a chemical change
- Chemicals remain the same


Chemical Changes
- New substances are formed
- Properties of the matter change

• Conductivity, acidity, color, etc.
• Eg. Iron rusting, burning wood, digested food

Physical change of water into hydrogen peroxide


Conservation of Matter
- In Physical and Chemical Changes, matter is neither created nor destroyed.

This is the link of the video we watched of mythbusters:
http://www.youtube.com/watch?v=MWJU6sbf8Ng

Chapter 1 Review- September 21st Class

Today in Class we finished doing our NaCl labs and write-ups. We also had time to review for the upcoming chapter one test.

Below are some helpful study tools to study the various things in chapter one that we will be tested on.

Sodium Chloride Lab - Sept 17 Class

In the beginning of class, we reviewed Dimensional Analysis since we didn't get a chance to go over it thoroughly in the previous class. Here are some examples of unit conversions:


If you have trouble converting milligrams into kilograms, then you could always break it down into grams first. The 2 kilograms and the other 2 grams cancel each other out, leaving you with 120.000,000mg (120,000 x 1000).


This problem includes dealing with measurements and time as well. We must remember to incorporate both information. The Litres and the seconds each have another pair to cancel each other out, leaving you with 435 000mL/min (7.25 x 60 x 1000).


Once again we have broken down kg into g because it is easier to convert g into Mg. This problem also includes time that needs conversion as well. The kilograms, seconds, and grams all cancel out, and the end result is 626.4Mg/h (174 x 3600 x 1000/1,000,000).


After we did a little bit of review, we started our very first lab called Sodium Chloride.


















Purpose
To experiment the maximum amount of salt you can dissolve in 10mL, 20mL, 30mL, and 40mL.

Materials
- Sodium Chloride
- 150mL of Distilled Water
- 100mL graduate cylinder
- 50mL beaker
- 100mL beaker
- Weight paper
- Electronic Scale
- Lab coat
- Safety Glasses

Procedure
1. Gather all the materials an put them on the lab bench.
2. Measure 10mL of distilled water using the graduated cylinder. Transfer this water into a 50mL beaker.
3. Weigh 50 g of Sodium Chloride.
4. Add Sodium Chloride to the water until it stop dissolving and the first salt crystals begin appearing on the bottom of the beaker.
5. Measure the mass of salt remaining. Record the difference in salt as the amount added to the 10mL of water.
6. Repeat steps 2-5 for 20mL, 30mL, and 40mL of water.
7. Record all your data in the table below.
8. Create a graph of Mass of Salt vs Volume of Water.

Observations

The results can and may vary.





http://www.youtube.com/watch?v=Kxbc8nuv_0k

This video is similar to the lab we did today. The only differences were the scales, how they measured the ingredients, and they shook the salt and water together instead of stirring it with a glass rod.