Matter and Qualitative Analysis
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Homework: Do the Questions on Page 2 Questions 1 through 15
This system was developed to provide workers and students with complete and accurate information regarding hazardous products. All chemical products that are used in business, workplaces, industry or schools must contain standardized labels and must be accompanied by Material Safety Data Sheet (MSDS) in a location convenient to the worker or student. The MSDS sheet must provide clear and precise information about the product. Clear and standardized labeling of products is an important component of WHMIS. The labels must be present on the product's original container or a label must be created and added to any container the product is added to.
General safety rules: Always store materials in their proper designated areas. Always wear protective clothing, along with face and eye protection
House Hold Product Symbols (HHPS)
The Canadian Hazardous Products Act requires manufacturers of consumer products containing chemicals to include a symbol that specifies both the nature of the primary hazard and the degree of this hazard. In addition, any secondary hazards, first-aid treatment, storage, and disposal must be noted. These are the House Hold Product Symbols (HHPS):
These symbols are surrounded by one of two signs shapes. The number of sides indicates the degree of danger. The stop sign with 8 sides poses a greater danger than the yield sign with 3 sides.
• Homework: Read Appendix B2: Safety In the Laboratory
Pages 480 to 484
• Activity 1.1 Identifying a Mystery Powder
Observation - A statement that is based on what you can see, hear, taste, touch or swell.
Inference - A judgment or opinion that is based on observations and/or conclusions from testing
Empirical Knowledge - Knowledge coming directly from observations.
Theoretical Knowledge - Knowledge based on ideas that are created to explain observations.
Theory - An explanation of a large number of related observations.
Law - A statement that has been proven to be true over and over again.
Models - A representation of a theoretical concept.
• Homework: Page 12
Section 1.2 Questions 1 through 6 Answers
• Activity 1.2 The Burning Candle
• Extention Exercise 1.2 Building Scientific Knowledge
Models of the Atoms Read Page 13 through 15
Empedocles and Democritus
Empedocles in the fifth century BC, proposed that all matter was made up of the four basic particles of matter, Earth, Air, Fire and Water. Later on Democritus , doing a though experiment came up with the idea that all matter was made up of tiny indivisible particles. These particles were so small that they could not be seen with the naked eye and they could not be cut or subdivided into smaller bits. He called these particles "atomus".
Dalton's Atomic Theory
150 years ago John Dalton formulated the idea that the world around us is made up of large numbers of identical very small particles called molecules, and that the many types of different kinds of molecules are simply differently arranged groups of atoms. It is this idea of the molecule that lies at the heart of chemistry.
The understanding of the nature of matter which is called the atomic theory of matter, first postulated by John Dalton, is the basis of all modern chemistry. As stated by Dalton, the atomic theory of matter consists of five postulates:
• Each element is made up of its own type of atom
• Atoms of different elements have different properties
• Atoms of two or more elements can combine in constant ratios to form compounds
• Atoms cannot be created, destroyed, or subdivided in a chemical change
Sub Atomic Particles
In the 1880's advancements in technology like glass blowing, electricity and vacuums allowed for the creation of cathode ray tubes. J.J. Thomson, using cathode ray tubes theorized that there was something smaller than the atom and that there was a small, negatively charged particle inside the atom which were later called electrons. Thomson came up with a model which he called his Raisin Bun Model. The raisins were the negatively charged electrons and the bun was the positive charged remainder of the atom.
In 1911 Ernest Rutherford shot radioactive particles at gold foil. He show that the nucleus was very small compared with the rest of the atom and that all the positive charge was located in this dense central nucleus. The electrons were in orbit around the nucleus much like the planets orbiting the sun. Rutherford later named the positively charged particles protons.
In 1932, James Chadwick discovered a third particle that had no charge, which he called a neutron.
Neutrons are the same as a proton except that they have no charge.
Fundamental Sub Atomic Particles of the Atom
• 1.3 Self-Quiz
• 1.3 Extension Exercise: Case Study
Rutherford's Gold Foil Experiment
• 1.1 - 1.3 Self Quiz
of the Elements: A Chemist's Shorthand
For example: 1123Na or 2311Na The smaller number is always the Atomic number. The larger number is called the Mass number. The second version will be used in these pages, with the mass number on top and the atomic number on the bottom. The mass number is also never found on a periodic table as a nice whole number. It usually has some decimal places after it for reasons which we will cover shortly.
The atomic number tells the number of protons that are found in the nucleus. It also tells you the number of electrons that the element has in its outside shells. If the atomic number of nickel is 28 then every atom of nickel has 28 protons in its nucleus and 28 electrons outside the nucleus.
The mass number is a sum. It is the number of protons + the number of neutrons together in the nucleus. Why not the electrons? The electrons are so tiny that they basically have no mass so we ignore them.
You can calculate the number of neutrons in a nucleus very easily.
Number of neutrons = mass number - atomic number
In nickel, 5928Ni, the 28 is the atomic number and 59 is the mass number.
Therefore the number of protons = 28,
Isotopes and Atomic
Example. The fissionable isotope of uranium is U-235. The nonfissionable isotope U-238 makes up most of naturally occurring uranium. Since uranium has the atomic number 92, a nucleus of U-235 contains 92 protons and 143 neutrons (235-92) while a nucleus of U-238 contains 92 protons and 146 neutrons(238-92). When used in a chemical equation the symbols are 23592U and 23892U.
Although the atomic number of an atom is automatically understood in the element symbol, nuclear physicists and some chemists choose to write the atomic number deliberately in nuclear reactions. The atomic number can be written as a subscript preceding the element symbol, as 92U, but most regular chemists prefer to omit it when used in sentence form.
Look at a periodic table. The masses are not nice whole numbers! This is because there are many different isotopes of the same elements. For example there are two isotopes of chlorine; Cl-35 and Cl-37.
Stop here and do the exercise on Elemental Symbols, Fundamental Particles
The Bohr Model of the Atom
The problem with finding out that electrons where capable of existing only at certain energy levels was coming up with a model to explain these levels. In 1913 Neils Bohr (1885-1962), a Danish physicist, proposed a theoretical model for the hydrogen atom. He chose hydrogen because its atoms are the simplest, having only one electron about the nucleus, and because it produces the simplest spectrum with the fewest lines. In his model, Bohr imagined the electron to move around the nucleus following fixed paths, or orbits, much as a planet moves around the sun.
Homework: Section 1.3 Questions Page 15
Homework: Section 1.6 Questions Page 22
comes in a large range of frequencies and wavelengths. The range is referred
to as "the electromagnetic spectrum". To give a few examples:
in the range of 104 to 1012 Hz the electromagnetic
spectrum has a portion called radio waves. From 1012 to 1014
is the range called infrared. From 1017 to 1019
we have what are called the X-rays. Microwaves are part of the radio
wave portion in the shorter wavelength portion of the radio waves.
Infrared radiation consists of the range of frequencies that can make molecules of most substances vibrate internally. An increase in internal vibration is measured by an increase in temperature.
104 105 106 107 108 109 1010 1011 1012 1013 1014 1015 1016 1017 1018 1019
| | | | | | | | | | | | | | | |
| Radio waves | Infrared | | Ultraviolet | X-rays |
| Long | Short | TV | Microwaves | ^ ^
| waves | waves | | _ | | |
Visible Light Gamma rays
As the wavelength gets smaller the frequency gets larger. The larger the frequency the more energy light can carry. If you look at the chart above you should see that Infrared waves have less energy than X-rays.
• Homework Section 1.4 Questions Page 18 Answers
• 1.4 Extension Exercise - The Electromagnetic Spectrum
• 1.5 Activity - Identifying Gases Using the Line Spectrum
• 1.6 Extension Exercise Bohr-Rutherford Diagrams
• 1.6 Extension Exercise The Bohr model of the Hydrogen Atom
• 1.7 Activity Flame Tests
• Self Quiz 1.4 - 1.7
Formation of Ionic Compounds
Ionic bonds are formed when a metallic ion (a cation) with a positive charge is attracted to a non-metal ion (an anion) with a negative charge. This attraction between positive and negative ions forms an ionic compound. The bond that holds them together is an ionic bond. This bond is electrostatic in nature.
Ion compounds can be shown forming either of two ways:
1. Showing the ions already formed and showing the resulting compound OR
2. Starting with the neutral atoms and showing the resulting ions being created.
Lets start with number 1.
We've looked at the formation of ions. From the periodic table we get the following table of ions:
M stands for any metallic cation, X stands for any non-metallic anion.
M+ M2+ M3+ X3- X2- X-
Group 1 2 13 15 16 17
Li+ Be2+ Al3+ N3- O2- F-
Na+ Mg2+ P3- S2- Cl-
K+ Ca2+ As3- Se2- Br-
Rb+ Sr2+ Te2- I-
Rules for Writing Formulas for Ionic Compounds
1. The positive ion (cation) is given first in the formula. This is a chemistry custom.
2. The subscripts in the formula must produce an electrically neutral formula.
3. The subscripts should be the smallest whole numbers possible.
Write formulas for
Example: Al3+ Cl-1 Rb2+ P3- Li+ S2-
• Extension Exercise 1.11 The Formation of Ionic Compounds
So far we have looked at atoms that form ions and then are held together by electrostatic attraction between the positive and negative charge. These are collectively refered to as ionic compounds. However some molecules do not play the give and take electron game, and simply share the electrons in what will be called a covalent bond. Quite often the electrons are not shared equally and result in polar compounds. In a few rare examples the electrons are shared evenly and we find that these are non-polar covalent compounds.
A useful tool in understanding covalent bonds are Lewis symbols, named after G.N. Lewis (1875-1946). To draw a Lewis symbol for an element write the element's symbol surrounded by a number of dots. The number of dots represent the number of electrons in the valence shell in the outer shell of the element. For example, Li and H both have 1 electron in their outer shells, 2s and 1s respectfully. They would be written as Li . and H .
Each element in Group IA have the same Lewis symbol, H . , Li . , K . , Rb . , Cs .
The Lewis symbols for the eight A-group elements of period 2** are:
Group 1 2
13 14 15
16 17 18
Na . + . Cl : -------> Na+ + [ : Cl : ]-
The valence shell of the Na ion is emptied and so no dots remain. The outer shell of the chlorine, which has seven electrons, gains the sodium's electron to give a total of eight. The brackets are drawn around the choride ion to show that the electrons represented by the dots are its exclusive property.
Homework Section 1.11 Questions Page 35
Noble gases have very stable valence shells. When ions form, atoms tend to lose or gain electrons until a noble gas configuration is achieved. This noble gas configuration is also responsible for the number of electrons an atom can share and the number of covalent bonds that will form.
Hydrogen can obtain a stable full shell with only 2 electrons. It is the exception rather than the rule. H : H The Lewis structure indicates that both atoms have access to the electrons in the bond.
The Octet Rule
The valence shells of all the noble gases, except helium, all contain eight electrons. There is the tendency for many atoms to acheive this noble gas configuration of 8 electrons. This is the basis of the octet rule. When atoms react, they tend to achieve an outer shell having eight electrons. The octet rule can be used to explain the number of covalent bonds an atom forms. This number normally equals the number of electrons the atom needs to have a total of eight electrons.
In a chlorine atom, which has seven electrons in its valence shell, we only need one electron to complete the octet. Chlorine can of course gain total control of an electron and become a chloride ion. This is what happens when it becomes an ionic compound. When chlorine combines with another nonmetal, the transfer of electrons is not complete. In compounds like HCl and Cl2 the chlorine gets the one electron it needs by sharing.
H . + . Cl : ------> H : Cl : OR H - Cl
Further examples: Cl2 CH4 NH3 H2O
These are examples of structural formulas that show how the atoms in a molecule are attached to each other.
When atoms of Group 15 or Group 16 react with each other or themselves there are not enough electrons to share to form a single covalent bond that will satisfy the octet rule. However there is no reason that we have to restrict the bonding in the compound to a single bond. Elements like N tend to form triple bonds.
: N = N :
Notice that in a Lewis diagram the three shared electrons are placed between the two atoms. We count all of these electrons as though they belong to both of the atoms. Each nitrogen therefore has an octet. The triple bond is usually represented by three dashes, so the bonding N2 molecule is normally shown as
: N = N :
Double bonds are also possible. For example the molecule CO2 contains two sets of double bonds.
: O :: C :: O :.. ..
• Homework: Section 1.12 Questions Page 39
Polar Bonds and Electronegativity
When two identical atoms form a covalent bond they have equal pull on the covalent electron pair. The electrons used for bonding are therefore shared equally and no one atom has more of the electron pair than the other. However when a molecule like HCl forms, the chlorine atom has a much stronger attraction for the single covalent bond and the electron pair ends up spending most of its time closer to the chlorine atom than it does the hydrogen atom. Since the negative charges of the electrons spend most of their time in the vicinity of the chlorine, the chlorine atom takes on a partial negative charge. This charge is more than what the neutral chlorine atom had but not as strong as the charge acquired by a chloride ion that takes the electron and keeps it. The negative charge gained by the chlorine atom is balanced out by a partial positive charge gained by the hydrogen as the hydrogen's electron is pulled out of place.
The charges on each atom are less than the full +1 and -1 found in ions. For this reason they are called partial charges. A bond that carries partial positive and negative charges on opposite ends is called a polar bond, or a polar covalent bond. The term polar comes from the idea that the opposite charges are at opposite poles of the bond. Because there are two poles of charge involved, the bond is said to be a dipole.
The polar bond in HCl causes the molecule to act as if the entire molecule had opposite charges on it. For this reason the HCl is a polar molecule. The extent of the polarity in the dipole can be calculated very simply.
Paulings Table of Electronegativities
Just how polar a a polar bond is can be calculated using the Table of Electronegativities above.
Electronegativity is defined as the amount of attraction a nucleus has for an electron. Metals have low electronegativities because the shed electrons readily. Electronegativities are high in atoms that like to gain electrons in order to fill out their shells. The trends in the periodic table are that electronegativities increase as you go up a group and from left to right across a row. Take any two electronegativites and find their difference. The one with the high electronegativity will be negative compared to the other. If the difference is greater than 1.7 then the bond formed will be ionic. If the difference is zero then the bond will be non-polar covalent. Any bond formed with a difference between 0.1 and 1.6 is considered polar covalent. A polar covalent bond with a difference is 1.6 would be very polar compared to one with a difference of 0.1
Percent Ionic Character of a Single Chemical Bond
Polar and Non-Polar Covalent Bonds
• Homework: Read Page 41
Polar and Non-Polar Covalent Molecules
• Homework: Read Page 42-44
• Questions Section 1.12 Page 45 Answers
• Investigation 1.13 Classifying Solids Using Physical Properties
• Self Quiz 1.11 - 1.13
The vast number of chemical reactions can be classified in any number of ways. Under one scheme they can be categorized either as oxidation-reduction (electron transfer) reactions or non-oxidation-reduction reactions. Another completely different but common classification scheme recognizes four major reaction types:
(1) combination or synthesis reactions
(2) decomposition reactions
(3) substitution or single replacement reactions
(4) metathesis or double displacement reactions
The Four Major Types of Reactions
Name General Reaction Pattern
Combination or Synthesis A + B ----> AB
Decomposition AB ----> A + B
Substitution or Single Replacement A + BC ----> B + AC
Metathesis or Double Displacement AB + CD ----> AD + CB
Combination or Synthesis Reactions Two or more reactants unite to form a single product.
S + O2 ---------> SO2
2 S + 3 O2 --------->
iron oxygen iron (II) oxide
Decomposition Reactions A single reactant is decomposed or broken down into two or more
CaCO3 ----------> CaO
2 H2O -----------> 2 H2
2 KClO3 -----------> 2 KCl +
Zn + 2 HCl ---------->
H2 + ZnCl2
Cu + 2 AgNO3 ----------->
2 Ag + Cu(NO3)2
H2 + 2 AgNO3
-----------> 2 Ag + 2 HNO3
2 Na + 2 H2O -----------> 2
NaOH + H2
HCl + NaOH ----------->
NaCl + HOH
BaCl2 + 2 AgNO3 ---------->
2 AgCl + Ba(NO3)2
CaCO3 + 2 HCl ----------->
CaCl2 + H2CO3
Balancing Chemical Reactions
There are a great number of chemical reactions. In order to be a chemical reaction at least one new substance must be produced. Recall that there are definite signs that a new substance gets produced
a) a change in colour
b) the formation of a gas
c) the formation of a precipitate
d) the release or absorption of energy (heat).
You may get only one of these, or a combination of them.
Chemical reactions are typically written one of three ways. There are word equations, skeleton equations and balanced equations.
Word equations are used to describe the reaction in sentence form or in a literally form. As an example we will use the results from a lab that you have already done. Magnesium burns in air to produce a white powder. The white powder has been experimentally determined to be a compound of magnesium and oxygen. The product is magnesium oxide. The word equation would be:
magnesium + oxygen -------> magnesium oxide
The "+" means "reacts with", and the arrow means "to produce". The word equation can therefore be read as follows: "Magnesium reacts with oxygen to produce magnesium oxide." Magnesium and oxygen are reactants and the magnesium oxide is the lone product.
A word equation gives limited information. It identifies only the reactant and products. It does not give their formulae, nor does it tell you the masses of reactants needed or the mass of product produced.
However at this point in time, you should be able to make up chemical formula based upon the name.
Do the following worksheet on English into Chemical Equations
Elements in Skeleton Equations
There are 109 elements. A few need to be treated in a special way because of how they bond with each other. You never find elemental oxygen by itself. Elemental oxygen is always O2. Oxygen is one of the diatomic elements.
The other diatomic elements are: H2, F2, Cl2, Br2, I2, O2, and N2.
Please note that these are all gases. When substituing these into skeleton equations make sure that you use the correct formulas. All other elements can be treated as if they were monoatomic. (Act like lone atoms.)
Skeleton equations are simply the bare bones of a chemical equation. The chemical formulae are substituted into the word equation. The skeleton equation for the reaction above is:
Mg + O2
The formulas are written first. Each formula should be checked
at this time to make sure that they are correct. If they are not correct
then the equation probably will not balance later.
A skeleton equation again provides limited information. It tells you what the chemical formulas are for the reactants and products, but again it tells you nothing about how much. If you observe the skeleton equation above carefully, you'll see that there is 1 Mg atom on the left and 1 Mg atom on the right. At least that part is balanced. There is however, 2 O atoms on the left and 1 O atom on the right. If even one atom does not match up evenly on both sides then the equation is unbalanced.
A chemical equation must be balanced. The number of atoms of each type must be the same on both sides of the equation. You must never change the subscripts inside a formula. After all, these have been discovered experimentally and cannot be changed at the whim of a chemistry student or teacher. The only way to balance an equation is to place numbers, called coefficients, in front of whole formulas. A coefficient applies to the entire molecule that follows it.
Start with a skeleton equation: Mg + O2 -------> MgO
The oxygen atoms did not match so place a 2 in front of the MgO so that there are now two O on the right.
Mg + O2 --------> 2 MgONow the oxygen atoms balance but the magnesium atoms no longer match. We have placed a 2 in front of the MgO which means we have 2 Mg atoms and 2 O atoms in total. If we place a 2 in front of the Mg on the left side then we get a balanced equation.
2 Mg + O2 --------> 2 MgOIf you check the number of atoms on each side of the equation you'll see that it is balanced. The method used above is called balancing by inspection. It is used on the simpler types of equations.
Homework: Section 1.14 Questions Page 53
Do the following worksheet on Balancing Chemical Equations
Do the following worksheet on Translating English into Chemical Equations
Solubility Rules to Predict Precipitation Formation
Solubility is a measure of how a substance dissolves in water at a given temperature and pressure. A substance that does not dissolve well in water is called insoluble . Chalk is a substance that has a low solubility. Substances like sodium chloride, NaCl, that dissolve well in water are called soluble.
Look at the Solubility Table until you can understand how to use it.
Predicting The Formation of A Precipitate
Follow these steps to correctly determine if a product is an insoluble precipitate.
Determine the products (if any) when a solution of sodium sulfate is mixed with a solution of lead(II) nitrate. If a reaction occurs, summarize the reaction as a balanced chemical equation.
Determine the products (if any) when a solution of sodium chloride is mixed with a solution of silver nitrate. If a reaction occurs, summarize the reaction as a balanced chemical equation.
Read Section 1.15 Pages 54 through 57
Practice Exercise Page 57
and Net Ionic Equations
ionic equations or total ionic equations - an equation that shows all the soluble compounds written in their dissolved form and all the insoluble compounds written normally as compounds.
net ionic equations - shows only the ions that are actually involved in the reaction. All other ions (the spectators) are left out.
How to Create Net ionic Equations
Following these steps the same why each time will create a net ionic equation
Step 1: Identify the type of reaction and decide on the possible products.
Step 2: Write out the word equation for the reaction
Step 3: Look up the solubility of the products using the solubility table.
Step 4: Indicate the states of the reactants and products, either (aq) for aqueous or (ppt) for precipitate.
Step 5: Write the correct formulas for the reaction.
Step 6: Balance the equation
Step 7: Write out the total ionic equation leaving any precipitates in compound or undissolved form
Step 8: Cancel out any ions that are the same on both sides (the spectator ions)
Step 9: Write out the equation showing only what is left. This is your net ionic equation.
Example 1 - Write the total ionic equation and net ionic equation between barium sulphide and sodium sulfate.
Example 2 - Write the total ionic equation and net ionic equation between sodium chloride and lead(II) nitrate.
• Homework: Practice Question Page 60 Answer
• Investigation 1.15 Precipitation Reactions in Solution
• Homework: Section 1.15 Questions Page 62 Answer
Qualitative Chemical Analysis
Chemists use the information in the solubility table to create a diagnostic tool called the Qualitative analysis. Chemists can use the solubility rules to determine if certain ions are present in a solution by conducting double displacement reactions. They use solution of ions that will cause the ions they are looking for to precipitate out.
If you suspect that a solution may have acetate ions in it, then you'd add silver ions in the form of silver nitrate to it. If there are acetate ions in the solution they will precipitate out because silver acetate is insoluble. If you add the silver nitrate and nothing happens then you can infer that the acetate ions are not present.
precipitate - The heavier solid material that will fall to the bottom of the test tube
supernate - The clear fluid layer sitting on top of the precipitate.
• Activity 1.17 Determining the Presence of Ions in Solution
• Self Quiz 1.14 - 1.17
• Summary For Unit Evaluation
• Homework: Unit Review Questions Page 71
Questions 1 through 30 Answers
• Unit 1 Self Quiz