Science 10
Chemistry
A key role of Chemistry is to tell us all about patterns in the properties and behaviors of matter. Patterns are an important aspect of nature because they allow us to know things about substances. If something is part of a pattern, we can predict certain things about it. Of course, to ensure our knowledge is correct, we can test our predictions, and Science does this all the time. Also, if something is part of a pattern, we know how to classify it. This is important because making predictions about things, testing our predictions and using the results to classify things are ways of learning about the world around us.
You can see patterns in properties everywhere. For example, some things are explosive while others are stable even if you heat or pound them. Certain materials dissolve in water but others are unaffected even if you leave them soaking in it. Glassy materials are hard but plastics are soft and easily scratched.
And, what about patterns in the behaviors of chemicals? We will look for these patterns as we investigate the behaviors of chemicals, sometimes when they are by themselves, sometimes when they are in groups. Here are just a few types of behaviors that can be studied in chemistry: a startled octopus changing the color of its skin, eggs turning white and solid as they cook, a candle “disappearing” as it burns, a battery getting back its energy as it recharges, a leaf making food for its tree. In fact, if we look at the world around us (and don’t forget about the world inside ourselves), the variety of behaviors we could investigate through the branch of science called Chemistry is endless!
The Chemistry section of Science 10 focuses on tiny particles called atoms and how they join together (bond) to form larger particles called molecules (or in some cases, how molecules break down into the simpler atoms). We will learn about the chemical and physical properties of atoms and of molecules and what happens to these particles during the behaviors mentioned above.
The Particle Theory of Matter This theory has five points:
1. all matter is made of extremely tiny particles (much, much smaller than cells!)
2. each pure substance is made of just one kind of particle
3. particles attract each other
4. particles are always in motion
5. the temperature of the particles controls how much they move about.
Classifying Matter Classifying things is a way of learning about them because it helps us to see
similarities and differences. A simple classification system of matter is based on the physical states of matter, i.e., solid (s), liquid (l), and gas(g).
Q. Create a classification chart based on the hair colors of your classmates. Look around and record the
different hair colors and the number of students that fit into each hair color category. Then, use the
data in your classification chart to create a bar graph. Label the x–axis with the colors of hair. Up the
y–axis, place the numbers indicating how many students belong in each hair color category. Use a
large enough range of values. Which hair color is the most common?
Q. Create a classification chart based on the states of matter. Look around and record the number of
objects that fit into each category (each state.) Then, use the data in your classification chart to create
a bar graph. Label the x–axis with the states of matter. Up the y– axis, place the numbers indicating
how many objects belonged in each state. Use a large enough range of values. Which state is the most
common?
The following table or chart is a more complicated but more accurate way of classifying matter. This type of chart would be more difficult for you to use, because, at this point, you may not understand all the terms. And, if you were given a sample of something to classify, how would you know where to place it in the chart? An important part of chemistry is performing tests on materials to determine where they belong in the chart. The tests reveal certain properties of the substances and the properties tell us where in the table the materials belong.
Matter
Mixtures Pure Substances
Heterogenous Homogenous Elements Compounds
Mechanical Solutions
Mixtures
Changes We will look at two types of changes – physical and chemical. Physical changes do not
change the basic nature of the thing we are investigating. For example, if we are testing the strength of the metal handle of a hockey stick, we can bend it until it breaks but we still have the same material at the end of the investigation. But, a chemical change creates new matter. To test the combustion temperature of a piece of flame retardant material for safety clothing, the material could be put in an oven. At some point, it would start to smoke and then to burn. After the test, we would have a small pile of ash – quite a different material than we started with.
Q. Look at the following processes. In each one, some change is taking place. Describe what the change
is. Then, label the changes as “physical” or “chemical.” Explain your choice of labels. Remember, the
key is if something new is being made.
a) you sharpen your pencil
b) you erase a pencil line
c) you bake a cake
d) you use battery power to play your portable CD player.
Another way of distinguishing chemical and physical changes is the ease of reversing them. Physical changes can be reversed easily but chemical changes can be reversed only with difficulty, sometimes only through a complicated process. A problem here is that students new to the field of Chemistry may not be able to judge what is “easy” and “hard” to do. At this point in your study of Chemistry, a certain change may seem hard to reverse but later, you would realize it is actually easy to reverse. As you gain experience with the subject, this type of decision will become much easier to make.
Q. Which of the following changes do you think could be easily reversed?
a) water to ice
b) building a Lego model
c) exploding a firecracker
d) growing a rose.
Mechanical Mixtures Recall that the classification table above lists two types of mixtures –
mechanical mixtures (heterogenous) and solutions (homogenous.) the components of mechanical mixtures are loosely combined – you can see the components. (Sometimes a magnifying glass may be needed.) Ordinary mechanical mixtures are things like beach sand; you can separate a sample into the individual grains just by letting them slowly slide off the side of your hand. A suspension is another example of a mechanical mixture because, again, you can se the components. Think of muddy water. Let the suspension sit for a while and gravity will drag down the heavy particles of dirt, leaving a somewhat clear layer of water at the top. A colloid is a special type of suspension in which the particles are so small gravity alone can not pull them to the bottom of the container. Magnification is especially useful to see the components of this type of mixture. Also, shining light in through the side of the container allows us to see the very small particles because they scatter it just like tiny particles in the atmosphere scatter blue light. Which of these types of mixtures do you think is easier to “pick apart?”
Q. You mix small amounts of the following pairs of ingredients. Devise a way of separating each mixture
into its components. Later, we may test your ideas in lab.
a) salt and pepper
b) salt and sand
c) sugar and salt
d) salt and iron filings.
Q. Why are these mixtures called heterogenous?
Solutions This is the other branch of Mixtures from the classification table. A solution consists of
two components – the solvent that does the dissolving and the solute that gets dissolved. Often, the solvent makes up more of the solution than the solute. One property of a solution is clarity. Look through a salt water solution -- you “see” nothing. Alloys are solutions made by heating two or more metals, stirring them and letting the solution cool. We could say an alloy is a type of solid solution. Alloys have been important in the history of Technology and Industry. Solutions are more difficult to separate into their components.
Q. Which is the solvent and which is the solute in the following solutions?
a) the atmosphere and CO2
b) coffee and water
c) chocolate powder and milk
d) blood and O2 from our lungs
e) water and cola flavoring
Q. Research. Choose two of the following alloys and for each, create an original 100 word summary
including: approximate date of discovery, components, uses. Use the Internet or encyclopedias or
reference books on technology or invention. Choose from among: bronze, steel, Wood’s metal, the
metal in coins, yellow gold
Pure Substances There are two types of pure substances – elements and compounds. A sample
of an element contains just one type of atom while a sample of a compound contains at least two different types of atoms.
Elements Think of a handful of Legos. Each color is a type of atom. Grab a handful of red ones and you have a sample of an element because it contains only one kind of Lego, i.e., only one kind of atom. Should the Legos pieces all be separate or stuck together? It depends on the element. Oxygen occurs as pairs of atoms forming molecules (O2) but gold is just single atoms (Au). So, of the hundred or so elements listed in the Periodic Table of the Elements, some occur as separate atoms like in tin (Sn) or as molecules like in chlorine (Cl2). Hint: most of the metallic elements occur as separate atoms (some examples are Al, Li, and Cu) while non-metallic elements often occur as molecules (some examples are F2, I2, N2, P4, and S8).
Compounds A compound contains at least two different elements and so it must contain at least two types of atoms. And, if it must contain at least two types of atoms, it occurs only as molecules because the term molecule means a group of two or more atoms. Remember, there can be molecules of elements like hydrogen (H2) and molecules of compounds like sulfur dioxide (SO2) or sodium hydroxide (NaOH). So, there is no such thing as saying, “here is the atom of this compound.” We would have to say, “here is the molecule of this compound.” Could we say, “here is an atom of this molecule”? Sure, but we would have to mean the atom we’re talking about is just a part of the molecule and not the whole thing.
Back to the Legos. Because all compounds occur as molecules, they contain at least two types of atoms, and so we would start sticking together different colored Lego pieces. The simplest example of a “Lego compound” molecule would contain two pieces, each one a different color. The FAO Swartz toy store on 5th Avenue in New York contains enormous Lego animal sculptures six or seven feet tall with thousands and thousands of various colored pieces. One of these sculptures would be like the huge molecule of a super-complex compound. With over one hundred elements for nature to join in various combinations, the number of compounds is enormous – millions!
To progress in Chemistry, one must learn the symbols for atoms and molecules and the symbols for elements and compounds. Also, one must learn the names of common elements and of common compounds. The Periodic Table of the Elements will help us with this.
The Periodic Table Our universe contains over 100 hundred known chemical elements with a fascinating range of properties: color, shape, mass, reactivity, shininess, odor, state – the list goes on and on. Memory alone is not enough to organize and keep track of all this data so Chemists use the Periodic Table to help them.
In 1829 the German chemist Dobereiner suggested there might be a natural sequence in which to organize elements. His ideas lead him to create groups of threes called triads. By 1854, a Harvard professor, J. Cooke, had reorganized the elements into six groups, some of which contained the triads. In 1862, the French geology professor A. de Chancourtois reshaped Cooke’s groups into a pyramid pattern. Then in 1865 the English chemist J. Newlands created a seven column table in which he positioned the elements so that an element had properties similar to the element eight behind it or eight ahead of it. This idea of similar properties occurring each eighth element was called his law of octaves. Finally, during 1868–1871, two researchers, Mendeleev (Russian) and Meyer (German) independently created the modern arrangement and look of the Periodic table.
The modern version of the Periodic table is somewhat rectangular and is organized into columns (groups or families) and rows (Periods or Series) and each element has its own spot in the Table. Where an element is located in the Table gives us clues about its properties and how these properties will compare with those of the surrounding elements. To start learning about the Periodic Table, we must become familiar with some of the more common elements and their properties.
Q. Look at your Periodic Table and answer the following questions about it.
1. In a very general way, the Periodic Table resembles what geometric shape?
2. a) What are the proper names for a column? b) for a row?
3. a) So, the Periodic table has how many Periods? b) how many groups?
4. There are how many chemical elements listed in the Periodic Table?
5. What are the two main sections of the table?
6. There are how many elements in each section?
7. For each of the following elements, write the symbols and atomic numbers:
a) hydrogen b) oxygen c) iron d) copper e) tin f) carbon g) lead h) iodine i) sulphur
j) helium k) chlorine l) nitrogen m) mercury n) nickel o) neon
8. a) Generally, how does the atomic number change as you move to each new element down a group?
b) How about as you move to each new element across a row?
9. a) In which section of the Periodic Table do we find most of the gaseous elements?
b) How about the solid elements?
10. Write the names and symbols of the chemicals with the following atomic numbers:
1, 3, 4, 10, 17, 23, 34, 45, 52, 60 and 90.
11. Name, give the symbol and atomic number for any ten metals.
12. Now, the same for five non-metals.
Q. Research. Create an original 150 summary on the alchemists. These secretive researchers lead the
way for the development of chemistry (although chemistry did not develop directly from alchemy.)
Use the Internet or encyclopedias or reference books on technology or invention.
Atomic Theory Recall the Legos. If you were asked to make a sample of some element, say oxygen for example, how many colors would you snap together? The same or different colors? What if you were asked to make a sample of water? How many colors and how many pieces of each color would you use? John Dalton’s Atomic Theory helps us answer these questions. Once you know the answers to these types of questions – memory and practice. Note: before Dalton published his ideas, the investigations of four famous researchers, Davy and Faraday (Br.) and Lavoisier and Proust (Fr.), suggested atoms formed molecules by joining together (bonding) in definite ratios. This is how Science works; each new researcher uses the ideas of their predecessors.
Nomenclature Since there are millions of chemicals, formal rules for creating an official name for each chemical, a name accepted around the world, is a must. Otherwise, we would be faced with many common names for some of the same chemicals. This is a problem for gardeners – popular plants have one official botanical name (that few people know) and many, confusing common names. Nomenclature involves two skills which are the opposite of each other. One is looking at the name of a chemical and writing its formula; the other is looking at the formula and writing the name. We will learn how to name and write formulas for some chemicals belonging to a few large, common classes of chemicals.
Warning: nomenclature is a straightforward procedure but it becomes a zoo if you do not do make the simple preparations listed below.
Ionic Compounds Compounds of this type contain metallic and non-metallic ions.
So, a) know where to find the metallic and the non-metallic ions in the Periodic Table (PT).
b) know the charges of the metallic and non-metallic ions. The charges are listed in the PT.
c) know the metallic ions with more than one size of charge. Just use the PT.
d) know the polyatomic ions and their charges from the PT.
Some ionic compounds with simple ions:
sodium + chlorine = Na + + Cl – = NaCl = sodium chloride. The metallic ion name is the same as its element name but the non-metallic ion name changes a bit from its element name. See the “ide”. See that the + and – charges balance. The + charge from the Na balances the – charge of the Cl.
calcium + chlorine = Ca 2+ + Cl – = CaCl2 = calcium chloride. See how the Ca’s charge goes with the Cl to make the subscript of 2. This makes the charges balance: a 2+ charge = a pair of – charges.
aluminum + oxygen = Al 3+ + O 2– = Al2O3 = aluminum oxide. See how the Al’s charge gives the O a subscript of 3 and the O’s charge gives the Al a subscript of 2. Again, this makes the charges balance. Do you agree that 2 x 3+ makes 6+ (for the Al) and this does equal 3 x 2– which makes 6 – (for the O)?
silver + sulfur = Ag + + S 2– = Ag2S = silver sulfide. The sulfur’s charge gives the silver a subscript of 2. And, a pair of + charges = a 2– charge.
tin (II) + fluorine = Sn 2+ + F – = SnF2 = tin (II) fluoride. What’s the (II) about? Tin and some other metallic elements have a variety of charges. The value in the brackets tells the size of the charge.
tin (IV) + fluorine = Sn 4+ + F – = SnF4 = tin (IV) fluoride.
What if the metallic element has just one size of charge? Just omit the Roman numerals in brackets after its name. So, the name for CaI2 is just calcium iodide. (Insta-review: Why does the iodine have a subscript of 2? It takes the charge from the calcium ion as its subscript. Look: calcium + iodine = Ca 2+ + I – = CaI2)
Here are a few more examples of ionic compounds but these contain polyatomic ions:
First, what is a polyatomic ion? A polyatomic ion looks like a molecule with a + or – charge. Remember, a real molecule is a complete thing and has a neutral charge. A polyatomic ion is just the + or – part of a larger molecule.
magnesium + hydroxide = Mg 2+ + OH – = Mg(OH)2 = magnesium hydroxide. Look at the hydroxide ion, OH – ; it’s a polyatomic ion because it contains more than one element, O and H is this case. See how the magnesium’s charge gives the OH a subscript of 2 (we’ve seen this charge crisscrossing many times before) but see how the brackets surround the OH – and its subscript of 2 is written outside the last bracket. Do the charges still balance? Sure, because the 2+ charge of the one Mg 2+ equals the two – charges, one – charge from each OH – .
lead (II) + chlorate = Pb 2+ + ClO3 1– = Pb (ClO3)2 See that the lead’s charge gives the polyatomic chlorate ion a subscript of 2 and the chlorate’s charge gives the lead a subscript of 1 (which we do not bother to write). And remember, if more than one of the polyatomic ion is needed, brackets are required. See that the lead needed two chlorates so brackets are placed around the formula for the chlorate and its subscript is written outside the last bracket. Look carefully: without the brackets we might think there is a subscript of 32 written after the O!! (Insta-review: Why do we say the chlorate ion is a polyatomic ion? Look at its formula: ClO3 1–. How many elements in it? Two elements, Cl and O. How many Cl’s? One. How many O’s? Three. How many of the whole chlorate polyatomic ion? Two: look at the subscript written outside the last bracket.)
cobalt (III) + silicate = Co 3+ + SiO3 2– = Co2 (SiO3)3 Again, just crisscross the charges to get the correct subscripts and be sure to include the brackets around the three silicate polyatomic ions needed by the cobalt. Look carefully: without the brackets around the SiO3, we might think there is a subscript of 33 written after the O!!
copper (I) + sulphite = Cu 1+ + SO3 2– = Cu2 SO3 Why no brackets here? Because only one of the polyatomic sulphite ion is needed. (Insta-review: Where did the copper’s subscript of two come from? From the sulphite’s charge. And, do the charges balance? Sure: the two + charges, one from each copper ion, balance the 2 – charges from the single sulphite polyatomic ion.)
Covalent Compounds This other large group of compounds contains just nonmetallic elements.
remember, ionic compounds contained ions. With the ionic compounds we just looked at, we could use the + and – charges to help us get the proper subscripts to use in the formula. Here we use prefixes in the compound name to tell the sizes of the subscripts. Lets see how this works in the following examples.
So, a) know the nonmetallic elements and their subscripts
b) know the meaning of various prefixes (this will come with practice).
carbon + oxygen = C + O2 = CO = carbon monoxide. The “mon” means one. See also that the name of the last element changes a bit to end in “ide”. (Remember we saw this with the ionic compounds.)
carbon + oxygen = C + O2 = CO2 = carbon dioxide. The “di” means two. But, two compounds of the same two elements? Don’t be surprised. Nature sometimes creates a variety of similar chemicals. The prefixes in their names allow us to indicate which one we mean.
sulfur + oxygen = S8 + O2 = SO2 = sulfur dioxide. Again, “di” means two and see the slight change at the end of the last element’s name.
nitrogen + hydrogen = N2 + H2 = NH3 = the name should be nitrogen trihydride but, for historical reasons, the common name of ammonia is used. We will encounter a few chemicals that have common names.
Acids This group of compounds is easy to recognize because the first element in the formula is
hydrogen. Acids are an old group of chemicals known for hundreds of years. We will be looking at only a few common examples and memory is important here because many of the names are affected by the element(s) attached to the hydrogen.
A way of indicating the strength of an acidic solution is to give its pH value. Acids have pH values from ranging from 1 (strong) to 6.9 (weak). A neutral solution like freshly distilled water has a pH of 7.
HCl = hydrochloric acid
HNO3 = nitric acid
H2SO4 = sulphuric acid
H3PO4 = phosphoric acid
Bases Compounds belonging to this other old group of chemicals have properties the opposite of acids.
They are easily recognized because all their formulas finish with the hydroxide (OH – ) polyatomic ion. We could say that bases are just a special class of ionic compound. Also, their pH values run from 7.1 (weak) up to 14 (strong).
The nomenclature of bases is easy if you know the metallic elements and their charges and remember the idea of criss-crossing charges to get the correct subscripts so that the + and – charges balance.
sodium + hydroxide = Na + + OH – = NaOH = sodium hydroxide.
aluminum + hydroxide = Al 3+ + OH – = Al(OH)3 = aluminum hydroxide. See how we criss-cross the charges to get the proper subscripts so the + and – charges will balance? And, since the OH – is a polyatomic ion, we include the brackets. Why no brackets in the NaOH example above? Only one OH – is needed.
calcium + hydroxide = Ca 2+ + OH – = Ca(OH)2 = calcium hydroxide.
iron + hydroxide = Fe 3+ + OH – = Fe(OH)3 = iron (III) hydroxide. What’s the (III) about? Remember that some metallic elements have a variety of charges. The value in the brackets tells the size of the charge. Look at this: iron + hydroxide = Fe 2+ + OH – = Fe(OH)2 = iron (II) hydroxide. What if the metallic element has just one size of charge? Just omit the brackets after its name. So, the name for Mg(OH)2 is just magnesium hydroxide, not magnesium (II) hydroxide. (Insta-review: Why does the hydroxide have a subscript of 2? It comes from the 2+ charge of the magnesium. Look: magnesium + hydroxide = Mg 2+ + OH – = Mg(OH)2.)
Chemical Reactions Chemical equations are a type of sentence describing chemical changes (chemical reactions) during which atoms or molecules break apart or reform into new materials. The reaction, when it is balanced, shows you how many of each type of atom or molecule is involved in the process. It is one thing to know what types of atoms or molecules might be involved in some type of chemical change; it is another thing, an important thing, to be able to say how many are involved. Balancing equations requires patience, attention and practice.
This is the general look of an equation: reactant(s) ➔product(s). The reactants (reagents) are the atoms or molecules present at the start of the reaction process and the products are what is created at the end of the process. The ➔ represents the reaction process, i.e, the various steps and conditions that change the reactants into the products. One or more conditions are needed to start a reaction, depending on the particular situation. Some common requirements to start reactions are: subjecting the reactants to a specific temperature range, sending electricity through them, pressurizing the reactants, exposing them to particular light levels, or, mixing the reagents in certain quantities.
Remember, the goal of balancing is simple: to show how many of each element is involved. A balanced equation will have the same number of each element before and after the reaction.
a) write the correct formula for each reactant and for each product – correct symbols and subscripts.
b) place coefficients in front of each reactant and product so the totals of the elements are equal on both
sides.
Some examples of Composition reactions: element + element ➔ compound
carbon + oxygen ➔ carbon monoxide 2 C + O2 ➔ 2 CO
calcium + chlorine ➔ calcium chloride Ca + Cl2 ➔ CaCl2 (self-balances)
sulfur + oxygen ➔ sulfur dioxide S8 + 8 O2 ➔ 8 SO2
Some examples of Decomposition reactions: compound ➔ element + element
silver sulfide ➔ silver + sulfur 8 Ag2S ➔ 16 Ag + S8
ammonia ➔ nitrogen + hydrogen 2 NH3 ➔ N2 + 3 H2
tin (IV) fluoride ➔ tin + fluorine SnF4 ➔ Sn + 2 F2
Some examples of Single Replacement reactions: element + compound ➔ element + compound
chlorine + sodium iodide ➔ iodine + sodium chloride Cl2 + 2 NaI ➔ I2 + 2 NaCl
aluminum + iron (III) oxide ➔ iron + aluminum oxide 2 Al + Fe2O3 ➔ 2 Fe + Al2O3
Some examples of Double Replacement reactions: compound + compound ➔ compound + compound
sulfuric acid + calcium phosphate ➔ hydrogen phosphate + calcium sulfate
3 H2SO4 + Ca3(PO4)2 ➔ 2 H3PO4 + 3 CaSO4
aluminum iodide + lead (II) chromate ➔ aluminum chromate + lead (II) iodide
2 AlI3 + 3 PbCrO4 ➔ Al2(CrO4)3 + 3 PbI2