September 29th - Graphing!

Today’s class was rather 'dense'. Not only did we go over density, but we also went over graphing! We began by quickly reviewing dimensional analysis.  

One interesting point that we examined was the amount of cubed centimetres in a cubed meter. This proved to be quite a number. To find this out, we execute the following steps:

• Examine each side of the large square
• Measure the dimensions (in this case it is 100 cm by 100 cm by 100cm)
• We multiply this, obtain 1.00 x 10⁶ cm³. Therefore, one million cubic centimetres fit in a cubic meter!

Another fact was that 1000L = 1 m³. Therefore, 1L = 10cm³.

Beginning the new lesson, we learned that density is the mass of an object divided by its volume. The equation is:

It's rather simple. Just remember to put those significant digits into action!

Here's an example we went over:

Determine the density of a statue that has a mass of 135kg and a volume of 65L.

d = 135kg / 65L = 2.1 kg / L

An important point to make here is that the limit of significant digits, in this example, is two. This is because 65L has only two significant digits. These digits are not exact. Because of this, in multiplication and division, they will limit the final answer’s certainty. Examples of exact digits include counting (ie. counting 4 washers) or 24 hours in one day. 

We can have fun with density. For example, we can challenge ourselves with experiments and projects:

Well, now that we veered slightly off topic, let’s move to graphing. There are five fundamental parts of a graph:

1.    Labelled axes (with units)

2.    Appropriate scale

3.    Title (ie. y vs. x)

4.    Data points

5.    Line of best fit

Note: To properly label your axis, remember that the independent variable is the factor that is changed or manipulated during the experiment, and goes on the x-axis. The dependent variable is the factor with a value that depends on the independent variable, and is plotted on the y-axis.

Now that we have a graph, with it we should be able to:

1.    Read the graph

2.    Find the slope (rise/run)

3.    Fine the area under the graph

Graphs are very useful. They give us a visual representation of our data. Graphing has many applications.

With a graph, we can multiply and divide the units to find the area or even create a new graph. We can split up the multiple straight line segments and use the formulae we know to find the area. We can split up a graph like so:

Using the formula for the area of a triangle or a rectangle, we can find our answer. With this new information, we can even plot a new graph! The new graph can also be a valuable source of information! Cool!

Posted by Andrew.

September 27th - All About Dimensional Analysis

This class, we learned all about dimensional analysis. Dimensional analysis is a problem-solving method that utilizes the fact that any number or expression can be multiplied by one without changing it's value. Just as we convert between currencies or units of length in our daily lives, it is necessary to convert between units in chemistry.  Dimensional analysis has 4 steps:

1. Identify the units you want to end up with
2. Find the conversion factors
3. Place the units in their appropriate places
4. Cancel out the units

For example, if we wanted to find out how many centimetres are in 6.00 inches, we would do the following:

We know that we want to find out the number of centimetres in 6.00 in. Therefore, the unit that we must end up with is cm. We must also find the conversion factor between the units:

With this information, we can solve the equation, like so:

As you can tell from the above equation, significant digits matter in dimensional analysis. Look at our previous posts for a review of significant digits.

In the next example, we try to figure out the number of seconds in 2 years:

Our answer must be in seconds. We know that there are 60 seconds in a minute, 60 minutes in an hour, 24 hours in a day, and 365 days in a year. Our next job is to place the units in their appropriate places, multiply/divide, cancel out the units, and give the final answer. Here's the full equation:

Dimensional analysis can only be used on units that are related. For example, it's impossible to convert a unit of time to a unit of mass. 

Here's a helpful video that explains dimensional analysis from the ground up:

Posted by Michael.

September 23rd - Powers of Ten, Sig Digs, and More!

On September 22, scientists at CERN announced that they had recorded particles moving faster than the speed of light. They say that if their claim is correct, they may have disproved Einstein's Theory of Special Relativity, which states that nothing can travel faster than light. In their experiment, they recorded neutrinos arriving 60 nanoseconds before light at a detector in Gran Sasso, Italy. The sixty nanosecond (10^-9) difference is a significant number, as we learned in the previous section. If confirmed true, this would force physicists to rethink much that they know about the universe.

Some sources believe that Einstein’s theory of relativity can be broken. It states that nothing can travel at the speed of light. Assuming that neutrinos are in fact faster, their findings would be very difficult to explain. How can it be possible to go from 5 m/s to 7m/s without passing through 6 m/s? We look forward to hearing from other science labs that are currently trying to replicate the experiment.

Now back to chemistry! We began by learning about the magnitude of powers of ten. Each time you move down a power of ten (ie. From 10¹⁰ to 10⁹) you lose 90% of whatever you are measuring. One notable power of ten is 10¹⁶ metres, which is a light-year This is followed by 10¹⁷ metres, which is ten light-years. Another interesting power is 10^-10, which is an angstrom.

This interesting video gives us a sense of how big or small these numbers really are:

The next topic we covered was significant digits. We were given a worksheet and were instructed to complete it. Some important ideas about significant digits are:

• Any non-zero number is significant (ie. 456 has 3 significant digits)
• Zeroes that aren’t place keepers are significant (ie. 304 has 3 significant digits)
• Estimated zeroes are significant (ie. 1.040 has 4 significant digits)

Note: The number 4000 has four significant digits. When we see a whole number with multiple zeroes, we should assume there is a decimal place after the last zero. This decimal place indicates that the zeroes are significant digits (they are part of the measurement). If the number was written in the form 4 x 10³, it would have one significant digit.

To determine the number of significant digits when adding or subtracting, follow these steps:

1. Add or subtract as you normally would
2. Determine the lowest decimal position of the original numbers
3. Round your answer to that decimal position

ie. 3.1 + 2.11 = 5.21. The lowest decimal position of the original numbers is in the tenths spot, so the answer is rounded to 5.2.

To determine the number of significant digits when multiplying or dividing:

1. Multiply or divide as you normally would
2. Find the number with the lowest number of significant digits of the original numbers
3. Round your answer, making sure that the answer has the same number of significant digits as the number from the second step

ie. 2 x 1.25 = 2.5. The lowest number of significant digits in the original numbers is 1, so we round the answer to 3.

We concluded the lesson with a few examples of scientific notation. To recall, a number written in proper scientific notation is between 1 and 9 (inclusive) and written to an according power.

What can this be?

 

A neutrino of course!

Posted by Andrew.

September 21st - Error, Uncertainty, and Prefixes

Today, as the title suggests, we learned about the types of error, the possible causes of error, the uncertainty principle, and SI prefixes.

The class started with a collection of the chemical balancing worksheet that was for homework. Next, we got into groups of 2 or 3 and were out handed a worksheet. We were instructed to use our textbooks and any other resources we could find to answer the questions on the worksheet. 

The first section of the worksheet asked us to list 13 common SI prefixes, their symbols, and their multipliers. The SI prefixes are:

The above diagram also lists one bonus SI prefix! The smallest of the listed divisions, 'atto-', can be symbolized by the lower-case letter 'a'. It's multiple is 10⁻¹⁸ (fun fact: 'atto' is derived from the Danish word atten, meaning "eighteen"). 

In Chemistry 11, the five prefixes that we will use the most are 'mega-' (M, 10⁶), 'kilo-' (m, 10³), 'centi-' (c, 10⁻²), 'milli-' (m, 10⁻³), and 'micro-' (μ, 10⁻⁶).

Our next task was to list at least one example of a place where each prefix is used. Here are a select few that came up with:

• 'femto-' is used in the term 'femtotechnology' (technology on the scale of a femtometre)
• 'tera-' and 'giga-' are used in the terms 'terabyte' and 'gigabyte' (used for digital information storage)
• 'kilo-' and 'milli-' are used in the terms 'kilogram' and 'milligram' (used to measure mass)
• 'nano-' and 'micro-' are used in the terms 'nanosecond' and 'microsecond' (used to measure time)

The next major topics of the day were accuracy, precision, and error. Accuracy refers to how close measurements are to the expected value. Precision refers to how close measurements are to each other. In the image below, you can see the difference between precise throws and accurate throws in a game of darts:

Now that we know what precision and accuracy are, we can learn about error. Error refers to the amount by which a sample differs from its expected value. Error and accuracy can be expressed through either absolute error or percent error. Absolute error refers to the difference between the measured value of a quantity and the actual value. It can be calculated with the following formula:

Percent error refers to the percentage given by dividing the absolute error by the accepted measurement, then multiplying the quotient by 100, like so:

It is important to note that error is inescapable in science. Measuring instruments are never completely free of flaws, so the measurement will always be off (the amount it is off depends on the specific instrument). Ambient conditions may change during an experiment. Human error could also occur in estimation. Estimation is always used in measurement (the amount of estimation depends on the specific instrument).

To acknowledge error, a plus-minus sign (±) is used when giving a measurement. For example, on some graduated cylinders, the smallest division is 1mL (please note that the measurement of water in a graduated cylinder is always taken from the bottom of the mensicus, which is the curved surface of a liquid in a container). An appropriate error measurement for this particular graduated cylinder would be ±0.5mm. To calculate the error measurement, you simply divide the smallest division of the instrument by 2.

The last thing we learned about was the Heisenberg uncertainty principle. This principle states that no matter how hard you try, experiments will always be affected by those who perform them. In some way, shape, or form, you are changing the experiment. For example, if you add light for better observation, you can influence the behaviour of quanta (quantum particles). If you want to know the velocity of a quark, you have to measure it. In order to measure it, you have to add energy. For this reason, scientists can never be certain that their answer is completely accurate. Here's an in-depth look at the principle:

 Posted by Michael.

September 19th - The Classification of Matter

We kicked off today’s class by reviewing chemical equations and balancing. During this time we encountered an interesting chemical reaction:

            MgSO₄^.6H₂O --> MgSO₄ + 6H₂O

This is called a hydrated compound or a hydrate. This is a compound that contains water molecules in their solid structure. The vertical ^. represents a weak bond between the water and the rest of the compound and is written in the vertical center of the line. The coefficient before water indicates the number of water molecules attached.

Shortly after, we moved on to the classification of matter. Nearly all matter can be classified as either a homogeneous substance or a heterogeneous substance. 

Homogeneous substances consist of only one visible component. Examples include:

• distilled water
• oxygen
• graphite
• air
• salt water

No matter how close we look at a homogeneous substance, there will be only one visible component.

On the other hand, heterogeneous substances contain more than one visible component. Examples include: 

• granite
• sand 
• cloudy water 
• blood

This flow chart can give you a better understanding of the classification of matter:

As shown in the diagram, there are two types of pure substances: elements and compounds. Elements are substances that cannot be broken down into simpler substances by chemical reactions. Examples include oxygen, iron, and magnesium. Compounds, on the other hand, are substances that are made up of two or more elements and can be changed into elements (or other compounds) by chemical reactions. This includes water and sucrose. In general, acids and bases are pure substances.

Determining whether something is an element or compound is not an easy task. The distinction is usually only visible at the atomic level. However, there are certain methods that we can use to find out the composition of a substance. 

In class, we witnessed an electric current pass through a bowl with a few pickles in it. When the electric current passed through the pickles, we saw sparks. When the current flowed through the water, electrolysis took place. Electrolysis split the compound, water, into its constituent elements.

Afterwards, we discussed what we knew about solutions. We learned that a solution is a homogeneous mixture of two or more substances. Solutions usually involve liquids, but don’t have to. We also learned about solvents and solutes. In a solution, the solvent is the component present in the great amount that does the dissolving. The solute is the component present in the smaller amount.

Our next topic was mixtures. Many mixtures are easy to identify, but some can be easily confused as pure substances. In heterogeneous mixtures, the different parts are clearly visible. In homogeneous mixtures, they are not (such as salt water, air, and brass).

There are many ways to separate mixtures, and they include:

• manually (by hand)
• filtration (heterogeneous mixture only)
• distillation
• crystallization
• chromatography 

All of the above methods are physical changes. 

Posted by Andrew.

September 15 - Physical Changes vs. Chemical Changes

Today we learned all about physical and chemical changes!

The lesson started with a quick review of the homework, which included identifying types of reactions, balancing chemical equations, and utilizing phase symbols. You can read all about these fun topics in our previous blog entry from September 13th (scroll down)! We also completed a little more balancing practice. 

Shortly after, we started a discussion about the ways in which matter can change. Despite what the title of this post may lead you to think, there are actually more that two types of changes. In fact, there are three categories that nearly all changes can be broken down into:

• Physical changes
• Chemical changes
• Nuclear changes

Physical changes, the first of the three types, involves only a change in state or shape. For example, crumpling paper changes its shape but does not change its chemical nature. Boiling water and smashing cars are two other examples. This fine gentleman also demonstrates a physical change:

Contrasting physical changes are chemical changes. In a chemical change, new substances are often formed. Depending on the specific reaction, certain chemical properties may change, such as the conductivity or acidity. Here's the chemical equation for the decomposition of nitrogen triiodide:

2 NI₃ (s) → N₂ (g) + 3 I₂ (g) + energy

As you've probably guessed, this exothermic reaction is a chemical change. Be careful, though! Some changes can be deceiving. For example, dissolving table salt in water involves the breaking of chemical bonds but is often described as being a physical change. If you remove the water through evaporation, you end up with your original substance; sodium chloride! Some examples of chemical change include the burning of wood, the digestion of food, and the formation of rust on iron. In class, Mr. Doktor demonstrated the chemical change that occurs between lead (II) nitrate and potassium iodide. Two transparent solutions reacted to form a precipitate (lead (II) iodide) and a salt (potassium nitrate). Below you can find the balanced chemical equation and a video of the reaction.

Pb(NO₃)₂ (aq) + 2KI (aq) → PbI₂ (s) + 2KNO₃ (aq)

Often times, a change in state (like one from solid to gas) is confused as a chemical change. However, the chemicals involved will not change. When a change/transition between states occurs, it can be called a phase change. This graph, showing the relationship between heat and temperature, demonstrates the peculiar phase changes of H₂O:

 Here's another nifty diagram to help you learn about phase changes:

Posted by Michael.

September 13th - Nothing But Balancing!

Today we learned all about word equations and balancing chemical equations.

We started class today with a review of the homework questions about lab safety. These questions reviewed:

• the basic safety rules to be followed in a laboratory (such as following the directions of the teacher and disposing of chemicals properly)
• the importance of tying back hair and rolling up sleeves (so that they do not droop into an open flame or come into contact with chemicals)
• examples of how basic chemistry safety can also be applied to other aspects of life, such as changing a tire or baking (for example, it is important to follow directions while changing a tire or baking so that you do not get into an accident or cut yourself in the kitchen)

 As you can see in the photo above, nobody is too cool to practice safe behaviour in the lab.

Shortly after going over the questions, we learned all about word equations and balancing chemical equations. For example, you may be reading your textbook when you encounter a sentence such as this:

Methane gas reacts with oxygen gas to produce carbon dioxide gas and water.

This is known as a word equation, the simplest form of an equation. It gives us the names of the reactants (the chemicals that react together) and the product(s) (the chemicals that are produced). It could also be written as:

methane + oxygen --> carbon dioxide + water

We learned that chemical reactions can also be represented by sentences known as chemical equations (or formula equations). Here's an example of a chemical equation:

CH₄ + O₂ --> CO₂ + H₂O

The letters that you see represent the atoms, molecules or compounds in the reaction. For example, on the left side of the equation, you can see that both methane and oxygen are present. You will also notice a plus sign (+) and an arrow ( --> ). The plus sign means 'reacts with', and the arrow means 'to produce'. The direction that the arrow is pointing indicates the direction of the reaction. The subscripts indicate the number of each element in the molecule, atom, or compound. If a molecule or compound is in brackets with a subscript on the outside, then the subscript applies to every element in the brackets. A chemical equation can be read the same way as the word equation.

It is also important to understand the products in the reaction. This is a combustion reaction, where a fuel reacts with oxygen to produce carbon dioxide and water. In this specific equation, the atoms of carbon and hydrogen break apart and end up in different products. The old bonds break apart and are replaced with new bonds, resulting in completely new molecules. However, you may be wondering why there are fewer atoms of hydrogen in the products than in the reactants. This, my friends, is where the law of conservation of matter comes into play. This law states that all matter has mass, which cannot be created or destroyed in a chemical reaction. In other words, those hydrogen atoms have to go somewhere. But where? To solve this problem, we need to perform the next step, called balancing.

The law of conservation of mass tells us that there has to be an equal number of specific atoms on both sides of the equation. When this law is observed, it is considered balanced. We begin by counting the number of atoms on each side of the equation. Here is the chemical formula from above:

CH₄ + O₂  --> CO₂ + H₂O

This tells us:

That we have one carbon atom (reactant).

That we have four hydrogen atoms (reactants).

That we have two oxygen atoms (reactants).

That we have one carbon atom (product).

That we have two hydrogen atoms (products).

That we have three oxygen atoms (products).

This equation is unbalanced, so we must use coefficients to balance it. For example, we multiply 
the number of hydrogen atoms by 2 on the right side to make the numbers even. This way, we 
can follow the law of conservation of matter. If you do it correctly, it should end up looking like this:

CH₄ + 2 O₂  --> CO₂ + 2 H₂O

Here is an awesome video of the reaction you have just balanced:

This skill is very important, as it will allow us to better understand the rest of the curriculum.

But we're not done yet! One of the important, but subtle, parts of balancing equations is the inclusion of the phase symbol. Phase symbols are subscripts that indicate the state of matter of the chemical. The four states of matter that we study in Chemistry 11 are solids (s), liquids (l), gases (g), and aqueous (aq) solutions (dissolved in water). To complete the chemical equation, we must add the phase symbols:

CH₄ (g) + 2 O₂ (g) --> CO₂ (g) + 2 H₂O (l)

In the above reaction, none of the reactants or products are solutions. In chemistry, a solution is something dissolved in water (the aqueous state). It is important that chemical and word equations indicate the state of matter in the form of phase symbols or words.

But wait, there's more! Here's something fun to keep you interested:

We also learned about diatomic and polyatomic molecules. Diatomic ('di' meaning two, and 'atomic' referring to an atom) molecules are molecules that exist in pairs of two. The diatomic molecules are:

H₂, N₂, O₂, F₂, Cl₂, Br₂, and I₂

_{Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, and Iodine}

Polyatomic molecules also have more than one of the same element attached together. The important ones to know for Chemistry 11 are:

P_4, S₈

Tetraphosphorus, Octasulfur

Be careful, though! It is often easy to forget about these exceptions when writing or balancing equations!

Posted by Michael.

September 9th - Safety, Styrofoam and Scoopulas

Today is the day of our first official blog entry!

We started today’s class by going to a lab station and identifying lab equipment. We managed to identify everything from scoopulas to evaporating dishes, as well as distinguished between Erlenmeyer and volumetric flasks. However, a triangular utensil was among the equipment, and we were unsure of what it could be used for. Later on, we learned it was a pipe stem triangle. These triangles are made from galvanized wire and porcelain pipe stemming, and can be used to hold a container over a Bunsen burner.

Shortly after, we made up a list of the ten most important safety rules in the lab. A few of the important rules our class brainstormed include:

• Do not leave your Bunsen burner unattended.
• Always use proper safety gear (lab coat, goggles, gloves, etc.)
• Always follow the teachers directions and procedures.
• Always add acid to water (instead of water to an acid) to prevent sudden overheating.
• Do not engage in horseplay.

Nearing the end of our class, we witnessed an astonishing chemical reaction. We observed how a Styrofoam cup disappeared in a container filled with acetone. The reason for this reaction is because the Styrofoam cup and acetone are both non-polar. After a little internet research, we learned that something that is non-polar will dissolve something else that is non-polar. We're looking forward to learning about this kind of reaction in class.

Here's a neat video of the reaction:

To conclude the class we received a WHMIS (Workplace Hazardous Materials Information System) worksheet with safety symbols. These symbols can be found on products that could be harmful if improperly used, if stored improperly, if they come into contact with flesh, or pose any other type of hazard. The WHMIS safety symbols include:

Posted by Andrew.

Welcome to our blog!

Welcome to thescienceofmatter.blogspot.com! This blog belongs to Andrew and Michael, and we hope you'll learn a thing or two from this page. Now, here's an awesome chemistry video to get you in the mood:

Posted by Michael.