pages 3 to 33 “quantum chemistry” target completion date: october 1

Pages 3 to 33“Quantum Chemistry”Target Completion Date: October 1

ChemistryUnit 0: Review Section

About Slide Iconsimportant information.

• You should either note or highlight items from this slide. Some items from this slide WILL be on tests!

Very Important Sample Problems• Always hand-copy important sample problems in your note book, and refer

back to them when doing assignments. Similar problems will be on tests!

Look at this! (usually charts, diagrams or tables)• You don’t need to copy or memorize this, but you must read and understand

the diagrams or explanations here. Concepts will be tested, but not the details.

Information only. Don’t copy!• This is usually background information to make a topic more interesting or to

fill in details, or to give examples of how to use a table. Not directly tested.

Review Stuff• Not part of the material you will be tested on, but you are expected to

remember this from grade 10. It may be indirectly tested.Supplementary stuff

• Not material covered this year.

R

Pages with a PINK background are

supplementary . Not material for a test!

http://icons.iconarchive.com/icons/iconarchive/red-orb-alphabet/72/Exclamation-mark-icon.png

Measurement and Conversion Basics• All sciences, including chemistry, depend on

observations. Measuring is an important part of observing.

• There are many important types of measurement in chemistry, but the four most important are...– Mass kilograms (kg) or grams (g)– Volume Litres (L) or millilitres (mL)– Temperature degrees Celsius (°C) or kelvins (K)*– Time seconds (s), minutes (min), hours (h)

* Degrees Fahrenheit (°F) are used in the United States, but are never used in chemistry. Although kelvins are similar in magnitude to degrees celsius, kelvins do not need the degree symbol (°)

Measuring Tools• You must know how to use the following to

measure volume:– Graduated cylinder– Pipette– Burette

Burette

ConversionsR

Kill Hector the Decathlete Until he’s Deceased

with Centipedes and Millipedes

• You must be able to do ALL standard metric conversions, especially:– Litres to millilitres, millilitres to litres– Grams to kilograms, kilograms to grams

Quick ConversionsPrefix means

mega (M) million

… 100000

… 10000

kilo (k) 1000

hecto (h) 100

deca (da) 10

… unit 1

deci (d) 0.1

centi (c ) 0.01

milli (m) 0.001

… 0.0001

… × 10-5

micro (μ) × 10-6

The table on the left gives the eight most commonly used prefixes in the metric system.

It also includes five rows that do not have prefixes.

The middle row is for the unit: metre, litre, gram, newton, or any other legal metric unit.

This table can be used to quickly convert from one metric amount to an equivalent. Make a copy of this table on the margin of the front cover of your notebook, and learn how to use it. Lets do an example. Let’s find how many centimetres there are in 2.524 kmConversion: 2.524 km ? cm

2 524 km 00 cm Add extra zeros if necessary

There are five steps in the table between “kilo” and “centi”, so we have to move the decimal five places to the right. If we were going up the table we would move left. Answer: 2524 km = 252 400 cm

http://icons.iconarchive.com/icons/iconarchive/red-orb-alphabet/72/Exclamation-mark-icon.png

Other conversions

• Later this year you will need these:– Temperature: degrees Celsius (℃) to kelvins (K)• Add 273 to degrees Celsius to get kelvins.• Subtract 273 from kelvins to get degrees Celsius.

– Pressure: kilopascals (kPa) to millimetres (mmHg)• Multiply kilopascals by 760 and divide by 101.3 to get

mmHg• Multiply mmHg by 101.3 and divide by 760 to get

kilopascals

Density• Density is the relationship between the volume of an object and

its mass. Density is an important characteristic property of matter.

• This is a review formula from last year:

R

Where: ρ = the density of the object, in g/cm3 or g/mLm = the mass of the object, in gV = the volume of the object, in cm3 or mL

ρ Vmor

ρw = 1 g/mL = 1 g/cm3 The density of water is 1 g/mL. This is not true of other substances. Objects with less density than water will float. Objects with greater density will sink.

Solving Problems• When solving Chemistry problems on a test or exam,

it is important not only to find the correct answer, but to justify it. While solving the problem you should:1. Show your data, the information you used to solve the

problem.2. Show your work, including the formulas you used and

the substitutions you made.3. Write an answer statement, a sentence that clearly

states your final answer.4. Include the correct units for your answer. Never just give

a number—you must specify what the number means!

Showing Your Solution

1. divide the area where you will write your solution into four sections.

2. In the first section, write your data and the items you need to find

3. In the second section, write the formula(s) you think you need to use.

4. In the third section, show your calculations

5. In the final section, write your answer in a complete sentence with the correct units.

On the final examination, you must not only be able to find solutions to problems, you must also justify your answers by showing what you did.

Suggested Solution MethodProblem: A block of material has a length of 12.0 cm, a width of 5.0 cm,

and a height of 2.0 cm. Its mass is 50.0 g. Find its density.Arrange your solution like this:

List all the information you find in the problem, complete with units, and the symbols.

Data:l = 12 cm.w = 5.0 cmh =2.0 cmm =50.0gV = ?

To Find:ρ (density)=?

Write down all the formulas you intend to use:

Formulas: V = lwh ρ =Show the substitutions you make, and enough of your calculations to justify your solution:

Calculations:V =12cm x 5cm x 2 cm

= 120 cm3𝜌 = 50 g / 120 cm3

= 0.417g/cm3

Answer: The density of the block is 0.417 g/cm3 (or 0.417 g/mL)

Always state your answer in a complete sentence, with appropriate units.

Problems on Conversions and Density

1. Convert the following:a) 125 mL to L d) 30 mL to L g) 75 mL to Lb) 450 g to kg e) 4500 mL to L h) 0.035L to mLc) 2.5 L to mL f) 1.35 kg to g i) 0.56L to mL

2. Find the density of a 4cm x 3cm x 2cm block that has a mass or 480 g. Justify your solution.

3. Find the width of a cube whose density is 5 g/cm3 and whose mass is 135 g. Justify your solution.

Also: Do the worksheets entitled “Density” and “Metric Conversions”

Overview: UncertaintyInherent Errors in Measuring Devices

In chemistry we often use instruments to measure quantities. Unfortunately, all instruments have some degree of inaccuracy or error. I prefer to call this error “uncertainty”, since it is not a mistake on the observer’s part, but an unavoidable inaccuracy that comes from the instrument. In this section we will see how to write measurements that show we recognize the limitations of our instruments.

Appendix 5Page 394

Absolute Uncertainty (AU)In math, numbers are considered pure, abstract things. In math, 2.00, 2.0 and 2 are considered the same, they all represent number 2. In science, numbers are considered to be measurements, and all measurements have some degree of uncertainty. They are seldom considered perfect!

In science, 2 mL, 2.0 mL and 2.00 mL are different!The difference is in the precision of the instrument used to measure.

All instruments that we use to make measurements have an inherent error or absolute uncertainty. On some instruments, the absolute uncertainty is marked, on other instruments we make the following assumption:

Assumption: The absolute uncertainty of a measurement is usually* one half of a measuring instrument’s smallest graduation.

*at university level, or when using high-quality equipment AU measurements may be expected to be one fifth of the measure between the smallest marking instead of one half!

App.5Page 394

Example of Uncertainty.

• At first glance, the two graduated cylinders here seem identical, but look closer.

• The first one has a measurement of 32.0 ± 0.5 mL

• The second one 32.5 ± 0.5 mL• It is NOT correct to say that

the first measurement is just 32 mL!

App.5Page 394

How to Record Absolute Uncertainty

• When you first look at a graduated cylinder, it appears to contain 32 mL of liquid.

• Looking closer, you see it is about halfway between 32 and 33 mL, so you record the .5

• If you judged it to be ⅓ of the way you could write 32.3. If it was below the 32 instead of just above it, you might record 31.5. If you still see it as exactly 32 even after a closer look, then record it as 32.0

• Then you write the absolute uncertainty, the allowable error of the instrument – usually* half the measure between the smallest markings.

• In this case, the smallest markings represent one millilitre, so half the measure ( 0.5 mL) is the uncertainty.

32 .5 ±0.5 mLDoubtful, but Significant digit Absolute Uncertainty

Plus-minus sign Unit

( ) A pair of parenthesis may be placed around the measurement.

Adding and Subtracting with Absolute Uncertainties

• Frequently we make two measurements and subtract them to find a difference (Δ). When we subtract numbers that have an uncertainty we must ADD the absolute uncertainty values!

• Eg. While doing a density experiment we add an object to a graduated cylinder. The reading of the cylinder changes from (20.5±0.5)mL to (24.0±0.5)mL. The volume difference (ΔV) is (3.5±1.0)mL

• When adding two numbers with uncertainties, we also ADD the uncertainty.AUT = ΣAU or AUT = AU1 +AU2+…

Relative Uncertainty• Sometimes it is useful to know how much uncertainty

we have compared to the original measurement. To do this we can calculate the relative uncertainty (RU).

• RU of a measurement equals the absolute uncertainty divided by the absolute value of the original measurement.

• The resulting decimal number is usually converted to a percentage (by multiplying it by 100)

Click here for details on UNCERTAINTY.

App.5Page 394

%100% measure

AURU

RU% = percentage relative uncertainty

http://www.wellesley.edu/Chemistry/Chem105manual/Appendices/uncertainty_analysis.html

Example of Relative Uncertainty

• The graduated cylinder has a reading of (32.5±0.5) mL (absolute)

• To find its relative uncertainty, divide:0.5 ÷ 32.5 = 0.01538461

• Round off to a reasonable number of decimal places and convert to a percent:

0.015 x 100 = 1.5%• Write it like this:

32.5 ml ±1.5%

The nice thing about relative uncertainties it that they show you how small your error actually is.

App.5Page 394

Parenthesis is NOT used for Relative Uncertainty

Multiplying and Dividing with Uncertainties

• When you multiply and divide measurements, you cannot use Absolute Uncertainties.

• Instead, we must add the Relative Uncertainties after we multiply or divide.

• Example: an object weighs (58.3±1.0)g and has a volume of (32.5±0.5)mL. Find its density.

• (58.3±1.0) g ÷ (32.5±0.5) mL =• find percents: 1.0 ÷ 58.3 x 100 =1.7%, 0.5 ÷ 32.5 x 100=1.5%

• 58.3 g ±1.7% ÷ 32.5 mL ±1.5% = 1.79g/mL ± 3.2%

• Answer: The density is 1.79g/mL ± 3.2%

Can’t do it this way!

Use RU

Not inTEXT!

Correct precision

• It is considered improper in science to imply that a measurement is more precise than it really is.

• If you have a graduated cylinder that is marked in 1 mL increments, you can record it to between the two smallest marks: eg. 32.0 ±0.5 mL or 32.5 ±0.5 mL are acceptable readings. • With the same graduated cylinder, it would be wrong to

write 32 ±0.5 mL or 32 ±0.5 mL or even 32.00 ±0.5 mL

• In science 32 mL, 32.0 mL and 32.00 mL have different meanings with respect to precision.

Exercise on Uncertainty• Do the sheet “uncertainty”• The sheet will be corrected in class.• Procedures for an in-class exercise

• Make sure your first and last name are on the sheet.• Complete as much of the sheet as you can in the time

allotted. Use a pencil or dark colour pen.• When the time is up, follow the teacher’s instructions

regarding corrections. Correct with a red pen.• When the sheet has been corrected, put it into your

assignment folder or duotang. Keep it here until at least the end of the current term.

Overview: Significant FiguresKnowing how much to round an answer.

In the sciences, we have an particular way of determining how much precision we need in the observations and answers we record. The method of rounding is called significant digits or significant figures. There is a detailed section in the appendix to your textbook on pages 394 to 397. Unfortunately, a few of the details given there are, well… I won’t say wrong, let’s just call them “uncertain”.

Significant Figures(A.K.A. Significant Digits)

• In science, we use significant digits as a guide to how precise an observation is, and as a guide to how much we should round off the results we obtain by doing math with those observations.

App.5Page 395

Rules for Significant Figures Interpreting Significant Digits

1. Non-zero digits are ALWAYS significant2. Zeros between significant digits are ALWAYS

significant.3. Zeros at the beginning of a number are NEVER

significant.4. Zeros at the end of a number MAY be

significant, but only if trusted.5. Exponents, multiples, signs, absolute errors

etc. are NEVER significant.

App.5Page 395

Examples of Rule 1, 2 and 3

Rule 1. Non-zero digits are ALWAYS significant.1.234 has 4 significant digits145 has 3 significant digits19567.2 has 6 significant digitsRule 2. Zeros between significant digits ARE significant.1001 has 4 significant digits5007.4 has 5 significant digits20000.6 has 6 significant digitsRule 3. Zeros at the beginning are NEVER significant.007 has 1 significant digit0.0000005 has 1 significant digit0.025 has 2 significant digits

Explaining Rule 4Rule 4. Zeros at the end of a number MAY be significant.Your textbook says that they are ALWAYS significant, but this is contrary to what most textbooks say. If there is a decimal point, there is no problem. All textbooks agree, the zeros are ALL significant.

3.00000 has 6 significant digits5.10 has 3 significant digits10.00 has 4 significant digitsIf there is NO decimal, the situation is ambiguous, and a bit of a JUDGEMENT CALL. If you trust the source to be precise, then you count all the zeros at the end. If you have reason to believe the person was estimating, then you don’t count the zeros at the end.

5000 has 1 to 4 significant digits250 has 2 or 3 significant figures123 000 000 has 3 to 9 significant figures

In a test situation, assume the numbers are precise, unless something in the question states otherwise.

Estimated source Trusted precise source

Rule 5Rule 5: Exponents and their bases, perfect multiples, uncertainties (error values), signs etc. are NEVER significant.6.02 x 1023 has 3 significant digits504.1 mL x 3 has 4 significant digits5.3 ±0.5 mL has 2 significant digits– 5.432 x 10-5 has 4 significant digitsπ × 8.45 has 3 significant digitsIn each case, the blue part is significant, the green part is NOT significant.

Note: The term Significance in this usage is not the same as importance. A digit may be “insignificant” but still very important. The significant digits guide you to the correct way of rounding numbers to show precision. The insignificant digits may serve as “placeholders”, making sure the decimal point is in the right place. An important job indeed, but not one that adds to the precision of the answer.

Avoiding Ambiguity• We mentioned before that measurements ending in zeros with

no decimal were ambiguous. Their accuracy depends on how they were measured, and that doesn’t always show up in the number.

• For example, if you measure 200 mL in a cylinder with markings of 1 mL it will be more accurate than if you measured it in one with markings 10 mL apart, and much better than a beaker whose markings were 100 mL apart.

How can we show someone reading our lab notes the number of what our 200 mL really means?

One answer is Scientific Notation!

Not inTEXT!

Same Number, Different PrecisionNumber Precise to

200.000 6 significant digits



200 * Ambiguous, 1 to 3 SD*

2.00 x 102 3 significant digits

2.0 x 102 2 significant digits

2 x 102 1 significant digit

*This could represent one significant digit, or two significant digits, or three significant digits depending on how precise the measuring equipment was. If I am careless enough to write a number like this on a test, you should assume I mean 3 S.D., but you have my permission to point out my mistake!

Avoid using numbers like 200 mL. Instead write them in scientific notation.

2.0x102 mL means you measured it to the nearest 10 mL (2 S.D.)

2.00 x 102 mL means you measured it to the nearest 1 mL (3 S.D.)

2.0000 x102 mL means you measured it to the nearest 0.01 mL… a very fine level of accuracy indeed!

Another way of showing the difference is to include the absolute uncertainty!

Math with Significant Figures• Adding and Subtracting:

• All units must be the same (can’t add different units!)• Line up all the measurements at their decimal points.• Round off all numbers to match the shortest number of

decimals.• Add or subtract as normal.

5.34576 L

547.1 mL

5345.7 6 mL55.1 43mL

547.1 mL

55.1 43mL

Example: add the following measurements.

Decimals lined up

5345.8 mL

Round off

55.1 mL

5948.0 mL

This unit is not the same as the

others!(litres vs. millilitres)

The answer is 5948.0 mL. Note that the answer has 5 sig. digits, even though one of the measurements had only 4 sig. digits. This can happen with addition.

• Multiplication and Division:• Different units may be multiplied or divided if there is a

formula to justify it.• The main rule in multiplying and dividing is that you

cannot have an answer with more significant digits than your “weakest” measurement (the one with fewest significant digits)• After doing the math, round off your answer to match

the weakest measurement.

Math with Significant Figures

You are the weakest link.

Goodbye!Multiply 2.53 g/mL by 75.35 mL

2.53 x 75.35 = 190.6355

=191 g

Weakest measurement

only 3 S.D.

Answer has only 3 S.D.

About the unit: x mL = g

Justification:m=ρV

• Perfect numbers• Occasionally we consider a number to be perfect. For example, if you

are told to “double a quantity” the 2 you multiply by is considered perfect. It does not affect the significant digits of your answer, neither increasing or decreasing them. Mole ratios in stoichiometry are also considered perfect, as are universal constants like pi. Perfect numbers have no units.

• Other operations• Generally, use the same rule as for multiplying for square roots,

exponents etc. That is, your answer can have no more significant digits than your weakest measurement.

• Mixed operations• When doing mixed operations in science, you will usually do the

additions or subtractions first (there should be brackets around them), then the other operations.

Math with Significant Figures

Problems on Significant Figures1. How many significant digits are in each measurement:

a) 123.45 mL c) 007 spies e) 0.0023 m b) 4.500 x103 mL d) times 5 f) 4000 kg

2. A Coulter counter is a device which counts the blood cells in a sample as they pass through a beam of light. A laboratory technician records 20000 wbc in a blood sample. At a demonstration a reporter says there were 20000 protesters. Both numbers are the same, which one has more significant figures? Why?

3. Find the volume of a prism that measures 2.3 cm by 3.55 cm by 2.14159 cm.

4. Add the measurements: 2.500 kg, 354.2 g, 153.78 gAlso: Do the worksheet entitled “Significant Figures”

Periodic Classification

OverviewThe periodic table is a useful arrangement of the elements, into regions, families and periods that have important meanings. It is also a source of much additional information about the elements. With careful interpretation of the table, we can find the number of protons an atom has, the approximate number of neutrons, and the arrangement of electrons in the atom and in its ions.

Review1

Topic 1: Organization of MatterPage 4

H2O2 atoms of hydrogen

1 atom of oxygen

0.1.1

36

R

H

CO

N

SCl

O

O

H

H

H

Cl

NH3

CO2

SCl2H

Ne

One atom

several thousand atoms

Ne

DNA

• 0.1.1 Atoms and Molecules– All matter is composed of atoms.– The atoms that make up most matter are

assembled into molecules. • A molecule may contain one atom, or it may contain

several thousand atoms, or any number between.

– A molecule is represented by its formula• Water molecules, for example, are represented by the

formula H2O, shown below:• A few large molecules have abbreviations, not formulas,

like DeoxyriboNucleicAcid

Na Cl

• Chemical Formulas and Ions– Some matter is formed from ions instead of

normal atoms or molecules.• For the most part, we treat ions like regular atoms,

and ionic compounds like molecules but there are a few very technical differences.

– Ions are atoms or clusters of atoms that have become positively or negatively charged by losing or gaining one or more electrons.• Positive ions are called cations (ca+ions), • Negative ions are called anions (aNions)• Metals always form cations (+), non-metals usually

form anions (-)

Page 4

Notice the slightly

stronger wording with

respect to metals than nonmetals!

Na+ Cl–cation anion0.1.2

37

Differences between ionic and covalent compounds

Ionic Compounds Covalent (molecular) Compounds

Ionic bonds “give” or “take” electrons Covalent bonds “share” electrons

Ionic compounds don’t have distinct molecules. Clusters of ions are sometimes referred to as “formula units” rather than “molecules”.

Covalent compounds have distinct, strongly bonded molecules. This is why some people call covalent compounds “molecular” compounds.

Ionic compounds are solid at room temperature.

Covalent compounds may be solid, liquid or gas at room temperature.

Ionic compounds usually have a high melting point. That’s why they are solid.

Covalent solids usually have a low melting point.

Ionic solids are usually hard, but brittle Covalent solids are usually softer

Ionic compounds are usually more soluble in water, but less soluble in non-polar solvents like acetone.

Covalent solids are usually less soluble in water, but more likely to dissolve in non-polar solvents like acetone.

38

Sample IonsElement ions Ion names alternate names

Sodium Na Na+ Sodium ion

Calcium Ca Ca2+ Calcium ion

Aluminum Al Al3+ Aluminum ion

Tin Sn Sn4+,Sn2+ Tin(IV) ion, Tin(II) ion Stannous, Stannic

Copper Cu Cu2+, Cu+ Copper(II) ion, Copper(I) ion Cuprous, Cupric

Iron Fe Fe3+, Fe2+ Iron(III) ion, Iron(II) ions Ferrous, Ferric

Carbon C C2+, C4+, C4- Carbon(II), Carbon(IV), Carbide

Carbon can form both anions and cations as well as covalent bonds

Nitrogen N N3- Nitride ion,

Phosphorus

P P3- Phosphide ion,

Oxygen O O2- Oxide ion,

Sulphur S S2- Sulphide ion Sulfide ion

Fluorine F F1- Fluoride ion

Chlorine Cl Cl1- Chloride ionNotice that some elements can form more than one type of ion. Compounds of the same element can differ quite a bit, for example, red iron oxide (rust) has Fe3+ ions, black iron oxide (wustite) contains Fe2+ ions. Note also, that most negative ions have the name ending changed to –ide.

Meta

l Io

ns

(+)

Non-M

eta

l Io

ns

(-)

+ Ca

tions

– An

ions

39

H

H

O OO

OBig Fat Ions

• Polyatomic ions are ions that are composed of a cluster of atoms, instead of a single atom.

• For example, the nitrate ion (NO3–) looks like

this:• But it acts like a single, negatively charged

particle in reactions.• Polyatomic ions are sometimes called radicals.• They are not the same as molecules.

O

O

ON-

SO

O

PO

O

N HH+

2–

3–Cl O

–

O

O

ON- Na+ + NO3- NaNO3

Na+

Common Polyatomic Ions (see p.422)Formula Name (ionic charge) Formula Name (ionic charge)

PO4 3- Phosphate ion (3-) NO3 - Nitrate (1-)

PO3 3- Phosphite ion (3-) NO2- Nitrite (1-)

SO4 2- Sulphate ion (2-) ClO4 - Perchlorate (1-)

SO3 2- Sulphite ion (2-) ClO3 - Chlorate (1-)

CO3 2- Carbonate (2-) ClO2 - Chlorite (1-)

CrO4 2- Chromate (2-) ClO - Hypochlorite (1-)

C2O4 2- Oxalate (2-) MnO4 - Permanganate (1-)

SiO3 2- Silicate (2-) H2PO4 - Dihydrogen phosphate (1-)

HPO4 2- Hydrogen phosphate (2-) . H2PO3- Dihydrogen phosphite (1-)

HPO3 2- Hydrogen phosphite (2-) HSO4- Hydrogen sulphate (AKA: bisulphate) (1-)

Cr2O7 2- Dichromate (2-) HSO3 - Hydrogen sulphite (AKA: busulphite) (1-)

C2H3O2 - Acetate (AKA: ethanoate) (1-) HCO3 - Hydrogen carbonate (AKA: bicarbonate) (1-)

OH- Hydroxide (1-) NH4 + Ammonium (1+)

CN - cyanide (1-) H3O+ Hydronium (also written as H+) (1+)

3-

2-

1-

1-

41This information is important when naming ternary ionic compounds. Click to skip ahead to Ionic Naming Rules

Representation of Atoms

OverviewSince the time of classical Greece, humans have tried to represent what matter was made of. Because the particles of matter are too small to see, we have used models to represent our concepts of atoms and molecules.

Review2

Representation of Atoms• Early Representations– Democritus (c.450 BCE)

• first suggested that matter was made of particles.

– John Dalton (1800) • represented the atoms as spheres (like microscopic

bowling balls)

– J.J. Thomson represented the atom as a “plum pudding” of positive charge with negative charged electrons scattered inside “like rasins”

– You studied the historic importance of these models last year, so you will not be tested on them this year. We will concentrate on the three most widely used representations on the slides that follow.

H

C

O

N

+-

- -

-

-

0.2.0

P

S

Cl

43

Dalton modelsOriginal

and Modern

R

Rutherford-Bohr Model– Rutherford discovered that the atom has a

dense nucleus containing positively charged protons.

– Negatively charged electrons move around this nucleus in paths that resemble an orbit.

– Later, Bohr calculated that there were different orbital energy levels or “shells” that could hold different numbers of electrons.

– A pattern of “Bohr numbers” corresponds to the formula 2n2 where n is a whole number.• Bohr numbers: 2,8, 18, 32, 50…

Page 5

Early Rutherford model

Revised Bohr model

0.2.1Page 5

44

R

The Simplified Atomic Model– The simplified atomic model that we often use

today adds neutrons (discovered by James Chadwick after the Bohr-Rutherford model had been proposed) to the protons in the nucleus.

– We often draw this in a simplified way, showing the nucleus as a full circle, and the electron “shells” as half-circles.

Page 5

11p+

12n0Na2e- 8e- 1e-

Z=11, configuration: 2,8,1

Nucleus: If asked for a complete simplified model, give the #protons and #neutrons (if known) in the nucleus. Otherwise, just draw a full circle.

The Atomic Number, Z, is the number of protons in the element. The configuration is the arrangement of the electrons in the shells

Symbol: The symbol of the element Electrons: 2 in first shell, 8 in 2nd 1 in 3rd

0.2.2Page 6

45

WARNING

• Be careful how you draw them!• The diagram must show the nucleus!

Unacceptable!

Nucleus is not shown. Nucleus is confused with 1st shell

Nucleus shown as solid circle.Labelled with element symbol beside.

Nucleus shown as full circle.Labelled with #protons and neutrons.

Unacceptable!

ACCEPTABLE ACCEPTABLE

The Sub-atomic ParticlesParticle Symbol Charge Actual Mass

(g) Rounded mass

(amu)Location in atom

Proton p+ 1+ 1.672x10 -24 ≈1 u (1679/1680) Nucleus

Neutron n0 0 1.674x10 -24 ≈1 u (1680/1680) Nucleus

Electron e- 1- 9.109x10 -28 ≈0 u (1/1680) Shells

nucleons

Page 6

47

Lewis Model: (AKA Lewis electron dot notation)– Lewis notation is a way of drawing a representation

of the valence electrons of an atom– When sketching an atom, write the symbol, and then

arrange dots around it to represent its valence electrons.

– Example: N has 5 valence electrons N– The “odd” or unpaired electrons are available for the

purpose of bonding. – Because there are 3 electrons available for bonding,

we say nitrogen has a valence of 3.– When bonding, atoms gain, lose or share electrons in

order to get a total of 8* electrons around each atom.

1

2

3

4

5 2 paired electrons

3 “odd” unpaired electrons

0.2.3Page 7

48

R

The preferred way of drawing Lewis diagrams of the first ten elements is shown below:

However, the dots may be moved around to show different arrangements. All of the drawings of Beryllium shown below might be correct in some circumstances.

Sometimes showing the bonding between atoms requires clever sharing of dots, as in the drawing of a nitrogen molecule (N2) shown here:

49

Sometimes electrons are removed from one atom to others in order to get 8

The Modern Model(Optional Enrichment)

• The Modern Model of the Atom– Of course, the Rutherford-Bohr model and the

Simplified Model do not perfectly represent what happens inside the atom. No model can!

– A more complete model, The Modern or Electron-Cloud model exists, but is more complicated and extremely difficult to draw.

– The Modern Model more accurately explains the relationship between the atom and the periodic table, and allows you to produce simplified models of elements in the transition area of the periodic table.

50


• The 2-8-8 vs. 2-8-18 problem.– You have probably been taught how to draw

Simplified Models for the first 20 elements – If so, you have noticed that for the elements

potassium and calcium, the third shell only holds 8 electrons—but Bohr said it should hold up to 18!

– The models you have been taught can’t explain why, but the modern model includes a concept called “orbitals” or subshells, and a filling pattern called the “aufbau diagram” that explain this .

51


• You are not required to learn the Aufbau diagram or the modern electron cloud model, but if time permits, I will show you how it works near the end of the review section. In the meantime:

52

You must know that the third shell CAN hold up to 18 electrons, but often doesn’t.And you must learn how the periodic table can be used to figure out the electron arrangement of many elements past the first 20.

But that is part of the next lesson…

Atomic Model Exercises

1. Draw Simplified Models of the first 20 elements.2. Draw Lewis Models of the first 20 elements.3. Convert the following:

a) 125 cm to m d) 320 mL to L g) 750 mL to Lb) 280 g to kg e) 45000 mm to km h) 0.0035km to cmc) 4.63 L to mL f) 5.52 kg to g i) 0.45L to mL

Periodic Classification

OverviewThe periodic table is a useful arrangement of the elements, into regions, families and periods that have important meanings. It is also a source of much additional information about the elements. With careful interpretation of the table, we can find the number of protons an atom has, the approximate number of neutrons, and the arrangement of electrons in the atom and in its ions.

Review3

Information in your Periodic Table

O8 2- Ionic charge

Atomic Radius

Density (g/L gas)

(g/mL solid/liquid)

Oxygen

15.999

Name

Atomic weight (amu)

0.65

1.43

3.44

1314

-218.3-182.9

Atomic number (Z)

Electronegativity

Ionization EnergyMelting Point (°C)

Boiling Point (°C)

(or g/mol)

Symbol

The symbol is a 1 or 2 letter abbreviation of the element’s name, or sometimes its Latin or German name. The first letter is always uppercase. If there is a second letter it MUST be written in lowercase. (eg. For sodium, Na is correct, na or NA are absolutely unacceptable!)

The number of protons

The English name of the element

Also the molar mass in g/mol

Electronegativity is a rating of how well the atom attracts electrons, on a scale from 0 to 4

Ionization Energy is how much energy it takes to remove an electron (kj/mol)

55

In-line Notation of Element Information

• An alternative to the periodic table is in-line notation of elements and isotopes. Note that the arrangement of information in this notation system is not the same as the arrangement in most periodic tables.

• Examples of inline notation:

• In-line notation is designed to be more compact, but less complete presentation of the information in a full periodic table.

Inline Isotope, Ion and Molecule Notation

C14

6 2

Mass Number (AKA. Isotope Number)

Represents the number of nucleons

in THIS particular atom

Atomic number “Z” represents the

number of protons in this atom

Valence Number(used in bonding)

4+4 Ionic Charge(used for ions)

Oxidation Number(used in electrochemistry)

Number of atoms in a molecule

–

8 neutrons

4+Total Nucleons

Minus Protons (s) Phase Marker(solid, liquid, gas

aqueous)

Looking at the Examples Again• Examples of inline notation:

means an isotope of Lithium with 3 protons and 4 neutrons means a lithium ion with a charge of 1+ means that the normal valence of lithium is 1 (forms 1 bond)

means that in a particular compound, lithium has oxidation #+1 means a compound that contains 2 atoms of lithium and 1 atom of oxygen

Mass #(total nucleons)

Protons Neutrons Electrons Charge Oxidation or Valence

Avg. mass

7 3 4 3 Charge= 0 6.99

≈7 3 ≈ 4 2 Charge= 1+ 6.99

14 6 8 6 Charge = 0 12.012C4+ 12 6 6 2 Charge= 4+ 12.0

C -4 ≈ 12 6 ≈ 6 ≈10 Oxid. #= -4 12.018O2 18 8 10 8 Valence= 2 16.0

C146

FAQ• What is the difference between mass number and

atomic mass (AKA atomic weight) ?– Mass number is the actual number of nucleons in a

particular atom (or isotope). It is mostly used in nuclear chemistry to distinguish isotopes of the same element (eg Uranium-238 238U vs. Uranium-235: 235U). It is always a whole number.

– Atomic mass is the average mass of all the atoms of an element, as it is recorded in periodic tables. This average was found by taking the average mass of many samples using a mass spectrometer. It is normally a decimal number.

FAQ• What is the difference between valence, charge and

oxidation number ?– Valence is the number of bonds an atom is likely to form.

Specifically, it is the number of “odd” electrons in an atom available for forming bonds.

– Charge is the actual electric charge an atom will get when it forms an IONIC bond. Some elements have more than one possible charge. When representing charge, the sign is written after the number. Eg 2+, 3‒ etc.

– Oxidation number is the charge assigned to an atom inside a compound, determined by a set of arbitrary rules. This is used mainly in electrochemistry to determine if and oxidation/reduction reaction can occur. Oxidation signs are written before the numbers. Eg. -2, +3 etc.

The Periodic Tablewith Primary Regions shaded

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18I Solid Liqui

dGas VIII

1 H II Synthetic III IV V VI VII He

2 Li Be metal Metal-oid

Non-metal B C N O F Ne

3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

Al Si P S Cl Ar

4 K Ca Sc Ti V Cr Mn

Fe Co Ni Cu Zn Ga Ge As Se Br Kr

5 Rb Sr Y Zr Nb Mo

Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

6 Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

7 Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu

7 Ac Th Pa U Np Pu Am

Cm

Bk Cf Es Fm

Md

No Lr

↑ The properties and region associations of these 10 elements are hypothetical ↑

62The heavy “staircase” line was the traditional separation between metals & non-metals but we now know it is not a sharp division.

The Periodic Tablewith Families Shaded

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18I III IV V VI VII VIII

1 H II Transition Elements He

2 Li Be B C N O F Ne

3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo




6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu


Cm

Bk Cf Es Fm

Md

No Lr

↑ The properties and family associations of most elements in period 7 are hypothetical↑

IA:

Alka

li M

etal

s

IIA

: Al

kalin

e Ea

rths

VIIIA

: N

oble

Gas

es

VIIA

: H

alog

ens

VI: O

xyge

n Fa

mily

V: N

itrog

en F

amily

IVA:

Car

bon

Fam

ily

IIIA:

Bor

on F

amily

Lanthanides

Actinides

IB:

Coi

n M

etal

s

Iron Triad

632nd Transition or “Rare Earth” Elements

The Periodic Tableand Valence Electrons (electrons in outermost shell)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18I III IV V VI VII VIII

1 H II Transition Elements He


3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo




6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu


Cm

Bk Cf Es Fm

Md

No Lr

↑ The properties and family associations of these synthetic elements are hypothetical ↑

ON

E (

I)

TW

O (

II)

EIG

HT

(VIII

)

SE

VEN

(V

II)

SIX

(V

I)

FIV

E (

V)

FOU

R (

IV)

TH

REE

(III)

64

TH

REE

(III)

SE

VEN

(V

II)

SIX

(V

I)

FIV

E (

V)

FOU

R (

IV)

TWO

ON

E

If the square is the same colour as the arrow above, it obeys its family with respect to valence. If the square is rainbow shaded, it is polyvalent, and not obeying its family rules. If the square is partly shaded, then it obeys its family rules most of the time.

The Periodic Table with Periods shaded

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18I VIII

1 H II III IV V VI VII He


3 Na Mg

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo




6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu


Cm

Bk Cf Es Fm

Md

No Lr

↑ The properties and family associations of these 10 elements are hypothetical ↑

3rd Period = 3 shells

5th Period = 5 shells





1st Period = 1 shells

2nd Period = 2 shells


65The periods of the table show how many shells of electrons an element normally has.

How to Use the Periodic Table to Find the Electron Arrangement of an Atom

Eg. Find the electron arrangement of Iodine (I)

Iodine is at the intersection of Period 5 and Family VII. Its number is 53. It has a total of five shells, 7 electrons in the outermost shell, and will have 53p+, and normally 53 e-. From this we can USUALLY figure out the electron arrangement.

53p+

2 7e-8 18 Total 53, So far: 35, left: 1818

1 H II A

Transition Elements III A

IVA

VA

VIA

VIIIA

He


3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo



SE

VEN

(V

II)

5th Period = 5 shells 53

Five shells

Note: The inner shells usually contain Bohr numbers

Periodic Table Exercises

• Write the name and symbol of each of the first 20 elements. (bragging rights if you can do it without looking!)

The Naming of Compounds

OverviewIn this section you must learn to name compounds based on their formulas, and also to find their formulas based on their names. This is a bit trickier than you might think, since different naming methods are used for different types of compounds.

You will need to distinguish between ionic, covalent and organic compounds, and also learn some common polyatomic ions.

Review4

Some elements with FORMULAS as well as SYMBOLSElement (systematic name)

Formula Common name Other names

dioxygen O2 Oxygen gas Diatomic oxygen

Trioxygen O3 Ozone Triatomic oxygen

dihydrogen H2 Hydrogen gas Diatomic hydrogen

dinitrogen N2 Nitrogen gas Diatomic nitrogen

dibromine Br2 Bromine Diatomic bromine

diiodine I2 Iodine Diatomic iodine

dichlorine Cl2 Chlorine gas Diatomic chlorine

difluorine F2 Fluorine gas Diatomic fluorine

tetraphosporus P4 White phosphorus Yellow phosphorus

amorphous phosphorus Pn Red phosphorus

hexasulphur S6 Orange sulphur hexathiane

octasulphur S8 Yellow sulphur brimstone

carbon (diamond) CcF8 or C8n diamond

amorphous carbon Cn or C Soot, charcoal Carbon black

carbon (graphite) CmP6 or C6n Graphite Pencil “lead”

buckminster fullerine C60 Fullerine C-60 “Buckyballs”

Naming Compounds

• There are four sets of rules for naming compounds:– The binary ionic rules:

• For compounds containing only two elements, joined by an ionic bond (usually a metal and non-metal).

– Ternary ionic rules: • For compounds containing 3 or more elements, including a

polyatomic ion.

– The covalent rules:• For two elements joined by covalent bonds (usually two non-

metals)

– The organic rules:• Used for compounds that contain carbon atoms bonded to

each other covalently.70

Compound Types

Ionic

Binary Ionic

Ternary Ionic

Covalent Organic

Is it a common exception? See “exceptions”Does it end with COOH? Organic (acid)Does it end with CH2OH? Organic (alcohol)Does it end with OH? Ternary Ionic (base)Does it end with a radical? Ternary Ionic Does it start with H? Binary Ionic (acid)Does it begin with a metal? Binary Ionic (salt)Does it begin with NH4? Ternary Ionic (ammonium)Does it contain >1 carbons? OrganicDid you answer NO to all? Covalent

ExceptionsH2O, H2O2, CH4, NH3Covalent/Organic

The Exceptions• H2O is usually called WATER (common name)

• Other acceptable names are dihydrogen monoxide, dihydrogen oxide, hydrogen monoxide, hydrogen hydroxide and oxidane.

• H2O2 is usually called hydrogen peroxide (ionic rule)• Other names are dihydrogen dioxide, dioxidane, and oxidanyl.

• CH4 is usually called methane (organic rule)• Pronounced either Meth-ane (Amer.) or me-thane (Brit.)• Other names include: carbon tetrahydride, tetrahydriocarbon.

• NH3 is usually called ammonia (organic functional rule)• Other names can include azane, hydrogen nitride, trihydrogen

nitride and nitrogen trihydride,

This list is by no means exclusive. Many other exceptions exist, but they are unlikely to be encountered in a high school chemistry course.

The Binary Ionic Rules

– First name the element on the left side of the compound’s formula.

– Then name the element on the right hand side of the compound’s formula, but change the suffix to “ide”

• For example:NaCl sodium chloride BaCl2 barium chloride

CaO calcium oxide K2S potassium sulphide

Al2O3 aluminum oxide Ca2C calcium carbide

Na+ Cl-

Ca2+ O2-

Ba2+Cl- Cl-

Al AlO2-

O2-

O2-

S2- K+K+

Ca2+ Ca2+C4-

1+ 1–

3+3+

73

We don’t need prefixes because ionic compounds ALWAYS follow the crossover rule

Common Non-metal Ion Names

Element (symbol) Negative Ion (charge) Element (formula) Negative Ion (charge)

Boron (B) Boride (B5-) Phosphorus (P4)

Phosphide (P3-)

Carbon (C) Carbide (C4-) Sulphur (S8) Sulphide (S2-)

Silicon (Si) Silicide (Si4-) Fluorine (F2) Fluoride (F–)

Arsenic (As) Arsenide (As3-) Hydrogen (H2) Hydride (H–)

Selenium (Se)

Selenide (Se2-) Chlorine (Cl2) Chloride (Cl–)

Other monatomic negative ions occur rarely. If you encounter one, use the atomic name, with the last syllable altered to ide as sounds best. Eg. Antinide or Polonide

Bromine (Br2) Bromide (Br–)

Iodine (I2) Iodide (I–)

Nitrogen (N2) Nitride (N3-)

Oxygen (O2) Oxide (O2-)

74

Ionic Rules No No!• When naming an ionic compound (and that

includes most compounds that contain a metal)

YOU SHOULD NOT USE A PREFIX!• Do NOT say: calcium difluoride for CaF2

• It’s Wrong. The correct name is just calcium fluoride.

• Do NOT say: dialuminum trioxide for Al2O3 • It’s Wrong. The correct name is aluminum oxide.

There are, or rather there USED to be, a few exceptions to this. Chromium dioxide was an accepted name for CrO2, and is still used occasionally. Now the name chromium(IV)oxide is preferred for the compound, since it obeys the ionic rules.

Ionic Rules:Dealing with Polyvalent Metals• Some metal elements have more than one possible

valence. Copper, for example, can have a charge of 1+ or 2+, depending on which compound it is in (eg. CuCl or CuCl2). Since we don’t use prefixes in naming ionic compounds, we shouldn’t use copper dichloride. We need a new rule!

• If a metal is polyvalent, we include its current valence in roman numerals inside parenthesis within an ionic compound name, for example: – CuCl = Copper (I) chloride (not copper monochloride)– CuCl2 = Copper (II) chloride (not copper dichloride!)

This copper ion must have a charge

of 2+

This copper ion has

a charge of 1+

76

Polyvalent ElementsThe elements with flashing circles have more than one positive valence.

1+ 2+ 3+ 4+ 5+ 6+ 7+ 4- 3- 2- 1- 0I VIII

1 H II III IV V VI VII He

2 Li Be Non-metal B C N O F Ne

3 Na Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

Al Si P S Cl Ar



5 Rb Sr Y Zr Nb Mo



7 Fr Ra U n c e r t a i n

6 La Ce Pr Nd Pm

Sm

Eu Gd

Tb Dy Ho Er Tm

Yb Lu


Cm

Bk Cf Es Fm

Md

No Lr

FeCr

Sn

Pb

Sm Eu

Cu

Au

Mn Co

Hg

Ru Pd

Pt

Ti Ni

U

C

Tl

Sb

Bi Po

77

The Common Polyvalent IonsFormula Charge Stock Name (new name) Classical Name (old name)

Cu+ 1+ Copper (I) ion Cuprous ionCu2+ 2+ Copper (II) ion Cupric ionFe2+ 2+ Iron (II) ion Ferrous ionFe3+ 3+ Iron (III) ion Ferric ionSn2+ 2+ Tin (II) ion Stannous ionSn4+ 4+ Tin (IV) ion Stannic ionPb2+ 2+ Lead (II) ion Plumbous ionPb4+ 4+ Lead (IV) ion Plumbic ionMn2+ 2+ Manganese (II) ion Manganous ionMn3+ 3+ Manganese (III) ion Manganic ionCr2+ 2+ Chromium (II) ion Chromous ionCr3+ 3+ Chromium (III) ion Chromic ion

Hg+, Hg22+ 1+ Mercury (I) ion Mercurous ion

Hg2+ 2+ Mercury (II) ion Mercuric ion78

Examples of Ionic Compounds with Polyvalent Elements

Formula Common Name

Stock (new) Name Classical (old) Name or Incorrect name

ions

FeO Wustite Iron(II)oxide Ferrous oxide Fe2+, O2-

Fe2O3 rust Iron(III)oxide Ferric oxide Fe3+, O2-

Fe3O4 Iron(II,III)oxide Ferosso Ferric oxide* Fe2+, Fe3+, O2-

Cu2O Cuprite (red) Copper(I)oxide Cuprous oxide Cu+, O2-

CuO (Black green) Copper(II)oxide Cupric oxide Cu2+, O2-

CrO Chrome black Chromium(II)oxide Chromous oxide Cr2+, O2-

Cr2O3 Chrome green Chromium(III)oxide

Chromic oxide Cr3+, O2-

CrO2 Crolyn Chromium(IV)oxide

Chromium dioxide Cr4+, O2-

CrO3 Chromic acid Chromium(VI)oxide

Chromium trioxide Cr6+, O2-

PbCl2 cotunnite Lead(II)chloride Plumbous chloride Pb2+, Cl-

PbO2 platternite Lead(IV)oxide Plumbic oxide Pb4+, O2-*Ferrosso ferric oxide is a unique combination of Iron(II)oxide and Iron(III)oxide together in a crystalline ionic structure Its formula can also be given as (FeO Fe∙ 2O3) 79

The Ternary Ionic Rules

– First name the metallic element (or ammonium ion) on the left of the formula.

– Then name the polyatomic ion on the right side of the formula. • If the compound is an ammonium salt, then name the

non-metal ion, changing it to end in “ide”

• Examples:– NaNO3sodium nitrate CaCO3calcium carbonate

– K2SO4potassium sulphate Ba(CN)2barium cyanide

– Al2(CrO4)3aluminum chromate NH4Cl ammonium chloride

Polyatomic ions: See Table 8.10 on p. 422 or click here

80

Covalent Rules– Name the less electronegative element on the left.– Name the more electronegative element on the

right, changing its suffix to “ide”– Add prefixes to each element to indicate the

number of atoms in the formula:• Mono*=1, di=2, tri=3, tetra*=4, penta*=5, hexa*=6

• Examples:– CCl4 carbon tetrachloride** N2H4 dinitrogen tetrahydride

– PF3 phosphorus trifluoride** P2O5 diphosphorus pentoxide

– CO2 carbon dioxide** CO carbon monoxide**

* The last “o” in mono or the “a” in tetra, penta, or hexa is usually dropped before “oxide” to sound better. (eg. “Carbon monoxide”, not “carbon monooxide”)** The “mono” prefix is usually dropped from the first element of the compound, except when that would cause confusion between two similar compounds. 81

Another use of electronegativity is to find how ionic or covalent a bond is. Click the arrow above to skip to the section on bonds

Simplified System(how much an atom attracts electrons)

I Solid VIII

1 2.2H

II III IV V VI VII 0.0He

2 1.0Li

1.6Be

0.70.9

1.01.4

1.51.9

2.02.4

2.52.9

3.03.1

3.23.5

3.64.0

2.0B

2.6C

3.0N

3.4O

4.0F

0.0

Ne

3 0.9Na

1.3Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

1.6Al

1.9Si

2.2P

2.6S

3.2Cl

0.0Ar

4 0.8K

1.0Ca

1.4Sc

1.5Ti

1.6V

1.7Cr

1.6Mn

1.8Fe

1.9Co

1.9Ni

1.9Cu

1.7Zn

1.8

Ga

2.0

Ge

2.2As

2.6Se

3.0Br

0.0Kr

5 0.8Rb

0.9Sr

1.2Y

1.3Zr

1.6Nb

2.2Mo

2.1Tc

2.2Ru

2.3Rh

2.2Pd

1.9Ag

1.7Cd

1.8In

2.0Sn

2.0

Sb

2.1Te

2.7I

0.0Xe

6 0.8Cs

0.9Ba

1.3Hf

1.5Ta

1.7W

1.9Re

2.2

Os

2.2Ir

2.2Pt

2.4Au

1.9Hg

1.8Tl

1.8Pb

1.9Bi

2.0Po

2.2At

0.0Rn

7 0.7Fr

0.9Ra

Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

There are a few exceptions, like CH4 and NH3, where the more electronegative elements are written first. These formulas have been used for years, and are based on organic chemistry concepts, so it’s unlikely we will change them.

Bonds

I Solid V VIII

1 II III IV 2.2H

VI VII 0.0He

2 1.0Li

1.6Be

2.0B

2.6C

3.0N

3.4O

4.0F

0.0

Ne

3 0.9Na

1.3Mg

III B

IV B

V B

VI B

VII B

VIIIB

IB

IIB

1.6Al

1.9Si

2.2P

2.6S

3.2Cl

0.0Ar

4 0.8K

1.0Ca

1.4Sc

1.5Ti

1.6V

1.7Cr

1.6Mn

1.8Fe

1.9Co

1.9Ni

1.9Cu

1.7Zn

1.8

Ga

2.0

Ge

2.2As

2.6Se

3.0Br

0.0Kr

5 0.8Rb

0.9Sr

1.2Y

1.3Zr

1.6Nb

2.2Mo

2.1Tc

2.2Ru

2.3Rh

2.2Pd

1.9Ag

1.7Cd

1.8In

2.0Sn

2.0

Sb

2.1Te

2.7I

0.0Xe

6 0.8Cs

0.9Ba

1.3Hf

1.5Ta

1.7W

1.9Re

2.2

Os

2.2Ir

2.2Pt

2.4Au

1.9Hg

1.8Tl

1.8Pb

1.9Bi

2.0Po

2.2At

0.0Rn

7 0.7Fr

0.9Ra

Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo

DO use prefixes with covalent compounds

# atoms Covalent prefix Examples

1 Mono…, mon… carbon monoxide (CO), mononitrogen monoxide(NO)

2 Di… carbon dioxide (CO2), dihydrogen dioxide* (H2O2)

3 Tri… nitrogen trichloride (NCl3)

4 Tetra…, tetr… carbon tetrachloride (CCl4), tetramethyl lead ((CH3)4Pb)

5 Penta…, pent… diphophorus pentoxide (P2O5), nitrogen pentafluoride (NF5)

6 Hexa…, hex… sulphur hexafluoride (SF6)

7 Hepta…, hept… bromine heptafluoride (BrF7), heptose (C7H14O7)

8 Octo…, oct… diphosporus octafluoride (P2F8) , octane (C8H18)

9 Nona…, non… nonane (C9H20)

10 Deca…, dec… Decane (C10H22)

*commonly called hydrogen peroxide.

Simplification of Covalent Names• IUPAC (The International Union of Physicists and Chemists) which

oversees naming conventions, allows some simplifications to the systematic names of covalent compounds.– The “mono” prefix may be dropped from an element, unless doing so

could result in confusion.• We usually say “carbon dioxide” rather than “monocarbon dioxide”• However, we always say “carbon monoxide” for CO, since there are two common

oxides of carbon (CO2 and CO)

– A prefix may be dropped from a formula if there is no ambiguity in the formula, and if the formula obeys the crossover rule.• Many chemists simply say “hydrogen sulphide” instead of “dihydrogen sulphide”

for the compound H2S. Since H2S is the only common sulphide of hydrogen and obeys the crossover rule, this doesn’t cause confusion.

– Knowing when simplification is allowed is a matter of experience. Until you become familiar with the conventions, it is safer to use all the prefixes. It’s not wrong to include them all (in covalent compounds).• Water can be called “dihydrogen monoxide”, but it is acceptable to use

“hydrogen oxide” since water obeys crossover. 84

Finding Formulasfrom Compound Names

• For covalent compounds, the name usually tells you the formula:

• For example: dinitrogen pentoxide = N2O5

• However:• If the name has been simplified by dropping a prefix you may have

to use the crossover rule, discussed later.• For example: “sulphur fluoride” has had a prefix dropped, so

S(valence=2) F(valence=1) crossover SF2

• “Sulphur fluoride” is the short name for the compound more accurately called sulphur difluoride.

• For ionic compounds, the name never tells you the formula.

• You always use the crossover rule to find the formula.• Example: Sodium oxide is Na1 and O2crossoverNa2O

85

The Crossover Ruleand simple ionic compounds

• The crossover rule is used to find the formula of a compound when the name has no prefixes (ie. all ionic compounds and some covalent compounds that have had a prefix removed)

• Example 1: What is the formula of aluminum sulphide?• Aluminum sulphide : Al S• Ions: Al3+ S2- • Valences (remove signs): Al3 S2

• Cross over: Al2S3

• The formula of aluminum sulphide is Al2S3

86

87

The Crossover Ruleand covalent compounds

• The crossover rule can also be used for covalent compounds if prefixes have been dropped from a name. When a covalent compound’s name has no prefixes at all, check it with the crossover rule.

• Example 1: What is the formula of “sulphur chloride”?• Sulphur chloride: S Cl• Oxidation numbers: S2- Cl- • Valences (remove signs): S2 Cl1

• Cross over: S1Cl2 or SCl2

• The formula of “sulphur chloride” is Al2S3

Notes: 1) The compound “sulphur chloride” should properly be called sulphur dichloride2) The prefixes trump the crossover rule. If any prefixes were used in the name, then they take precedence over whatever formula the crossover rule would give you.

The Crossover Rulesimplifying ionic compounds

• Ionic compounds can often be simplified• Example 1: What is the formula of the compound

made from Barium ions (Ba2+) and Carbide ions (C4-)?• Ions: Ba2+ C4- • Remove the signs Ba2 C4

• Cross over: Ba4C2

• Cancel (divide both by 2) Ba2C

• The formula of barium carbide is Ba2C

Note: Do not simplify covalent compounds by cancellation. Covalent compound formulas must reflect the compound names that include prefixes.

88

Reverse Crossover Rulefor finding the valence of uncertain ions

• Sometimes we can use the crossover rule in reverse to find the valence or ionic charge of an ion we are not certain of, such as an ion of polyvalent metal.

• For example, what is the name of Fe2O3? of FeO?– They are both iron oxide, but which iron oxide (there are

several types!)

– Fe O Fe O2 3 1 1

Fe has a valence of 3, so the name of the compound is:Iron(III)oxide

There’s a problem here! Oxygen hardly ever has a valence of 1. Let’s double both valences.

2 2

Fe’s proper valence here is 2Iron(II)oxide

The Organic Rules(not studied this year)

A system of names for organic compound exists that is based on the number of carbon atoms they have (as a prefix), and the type of compound they are (as a suffix): alkane (…ane), alkene (…ene) alcohol (…ol), aldehyde (…hyde), ketone (…one), organic acids, etc.# carbons Prefixes examples

1 Methyl, Formyl Methane, methanol, formaldehyde, formic acid

2 Ethyl, Acetyl Ethane, ethanol, acetaldehyde, acetone, acetic acid

3 Propyl, Propane, propanol, propanoic acid

4 Butyl Butane, butanol, butanoic acid

5 Pentyl Pentane, pentanol, pentanoic acid After this the prefixes resemble those for inorganic compounds, 6=hex, 7=hept, 8=oct, etc.

As you may notice, the common names of some chemicals come from the organic system, such as methane, the common name of carbon tetrahydride (CH4) . For more information on organic nomenclature, see the wikipedia article.

http://en.wikipedia.org/wiki/IUPAC_nomenclature_of_organic_chemistry

Practice

• Page 12, Question #9• Practice sheets:

• Naming ionic compounds• Naming covalent compounds• Naming mixed compounds

91

Enumeration of Matter

OverviewIn this section we will review the concept of the mole and the mole formula.

We will also see how the mole concept applies to molecules and to diatomic and polyatomic elements.

This is a required concept for stoichiometry, which we will cover later.

Review5

The Mole Conceptand the Enumeration of Matter

• The Mole: The mole is a unit used to count atoms, ions, molecules, and other fundamental particles.

• A mole corresponds to Avogadro’s Number of particles: 6.02 x 1023 particles.

NA =6.02 x 1023 = 602 000 000 000 000 000 000 000= six hundred and two sextillion= the particles in a mole.

0.5.1

93

Molar Mass• Molar mass is the mass of one mole of atoms

or molecules.• The symbol for molar mass is M (not MM!)

• For elements, molar mass corresponds to the atomic mass found in the periodic table, but expressed in grams/mol rather than amu. For example, the molar mass of carbon, M(C )= 12.011 g/mol, (frequently rounded to 12.0 g/mol)

• For compounds, M is the sum of the masses of all the atoms in the molecule or all the ions in the formula. For example, the molar mass of carbon dioxide molecules is:

M(CO2) =44.009 g/mol, (frequently rounded to 44.0 g/mol)

• that is: 2M(C)+2M(O) or 12.001 +2(15.999) g/mol

0.5.2

94

Alternate Method of Calculating Molar Mass(very useful if you don’t have a scientific calculator)

• Find the molar mass of Na2CO3:

Na = 2 atoms x 23.0 amu = 46.0

C = 1 atom x 12.0 amu = 12.0O = 3 atoms x 16.0 amu = 48.0

106.0 g/mol

Na2 O3C

+

Multiply the number of each atom (as given by

the formula) , by the atomic mass of the atom (as given in periodic table)

Then add the total masses together

Change the unit from amu g/mol

Diatomic and Polyatomic Elements• Diatomic elements: There are seven elements

whose molecules normally contain two atoms: I2, H2, N2, Br2, O2, Cl2 and F2.

• If finding the molar mass of these elements, remember to double the mass of one atom.

• M (I2) = 253.808 g/mol (not 126.904 g/mol!)

• Polyatomic elements: a few elements, such as sulphur, phosphorus, and sometimes carbon occur in larger molecules (eg. S8 or P4)

• If a formula like this has been used in a balanced equation, remember to multiply the atomic mass by the appropriate amount (eg. M(S8)=256.52 g/mol)

96How to Remember the Diatomic Elements: I Have No Bright Or Clever Friends

The Mole FormulaThe mole formula is used to convert from grams to moles and vice-versa

n

m

M

Actual mass(g)

# moles(mol)

Molar mass(g/mol)

𝑚=𝑛𝑀

𝑛=𝑚𝑀 𝑀=

𝑚𝑛

Actual mass = # moles x molar mass

# moles = Molar mass =

97

Practice

• Page 14, #12, 13, 14• Practice sheet:

• Moles and Molar mass

98

Physical Changes

OverviewWe will look at types of physical change that are important in chemistry, including Change of form (deformation)Change of phase (change of state)Change of mixture (dissolution)

Review6

Physical Changes• A physical change occurs when a substance undergoes

a modification in its appearance or form, but does not alter its nature or characteristic properties.

• In a physical change the molecules or ionic formula of the substance do not change.

• There are 3 main categories of physical change• Change of form, caused by crushing, cutting, grinding, bending,

denting, etc.• Change of phase or state, caused by melting, boiling, freezing,

evaporation, condensation, sublimation, etc.• Change of mixture, caused by dissolving (dissolution without

reaction), blending, stirring together dry ingredients, mixing paints, etc.

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Change of Form(AKA. Deformation)

• Only one type of deformation is of much importance in chemistry, and that is when we grind up reactants into fine powder so that they will react more quickly.

• We will go into this in greater detail later in the course, when we study the rates of reaction• For now, let’s just say that ground up materials usually

react faster.• Remember, a powdered solid is still a solid

Phase Change(AKA. Change of State)

• As a pure substance is heated, its particles move faster. It changes from a solid state to a liquid state and then to a gaseous state. Your textbook refers to this as “phase change”

• Change of “phase” is a physical change, since the particles of the pure substance do not (usually) change.

Picky note: What your textbook calls “phase change” should more properly be called “change of state”. Although “phase” and “state” are frequently used as synonyms, the word phase has a broader meaning in chemistry. There are three main states of pure matter (solid, liquid, and gas) , “phase” includes these three, but may also apply to many other possible phases of matter– including aqueous (a solid dissolved in water), gel (a jelly-like colloidal mixture) etc. In addition, phase can refer to a boundary between two similar phases that don’t mix, for example, a liquid mixture could have an oily phase and a watery phase that contact each other but do not mix.

0.6.1

102

LiquidSolid

Gas

Melting (fusion)

Freezing (solidification)

Vaporization

Liquid

Liquid CondensationSubl

imati

on

Solid

Con

dens

ation

(Dep

ositi

on)

Terminology associated with

Change of Phaseor State

Rapid vaporization is called “boiling”,

Slow vaporization is “evaporation”

Sublimation occurs when a material

“evaporates” from a solid straight to a gas, like dry ice or iodine.

ExothermicProcess

EndothermicProcess

103

Comparison of the States of MatterSolid Liquid Gas

Shape Definite Variable Variable

Volume Definite Definite Variable

Compressibility Incompressible Incompressible Compressible

Fluidity Not Fluid Fluid (flows) Fluid (flows)

Particle separation Close together Close together Far apart

Motion of particles1 VIBRATION only ROTATION2 and vibration

TRANSLATION, Rotation, and vibration

104

Notes: 1. The most characteristic (ie.abundant) motion of the phase is underlined.2. Rotation implies some tumbling motion, so molecules can move around a little.

Phase Markers

• During the course of the year, you will often notice small letters in parenthesis added formulas in equation. These “phase markers” are inserted whenever it is important to know what state or phase the reactants or products are.

• The most important phase markers are:• (s) = solid: the substance is a solid or a powder• (l) = liquid: the substance is a pure liquid• (g) = gaseous: the substance is a gas• (aq) = aqueous: the substance is dissolved in water

Eg: NaCl(s) H2O(l) NH3(g) NaCl(aq)

105

Dissolution and Solubility

• In dissolution, one or more solutes are mixed into a solvent to create a solution.

• During dissolution:• The mass of the substances does not change.• The total volume is usually slightly less than the sum of

the volumes of the components (since some particles pass into the spaces between other particles)• When the solvent cannot dissolve any more of the

solute, the solution is saturated.

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Dissolving = Physical Change• Remember that dissolution is normally considered

a physical change, not a chemical one. The material mixes with the solvent, but is not significantly altered by it

• In a few cases a material will react with the solvent, rather than just dissolve. For example, trying to dissolve sodium in water, or baking soda in vinegar will produce a reaction. In this case a chemical change has occurred as well.

eg: Na(s) + H2O(l) NaOH(aq) + H2(g)

• Ionic compounds may “dissociate” while dissolving, that is, their ions may separate by some distance. While this may seem like a chemical change, it is not a permanent condition, and is considered to be a physical change.

eg: NaCl(aq) Na+(aq) + Cl-

(aq) (dissociation of salt)107

Dissolution of Ammonia Gas in Water (an extreme case of solubility at 25°C)

• 100g of water + 50g of ammonia 150g of ammonia solution

• 100 mL of water + 72058 mL of ammonia 101 mL of NH3(aq) solution

• If you try to dissolve more than 50g of ammonia in 100 mL of water, you won’t be able to. There will be leftover ammonia!

100

72.058 litres NH3(g)

100

150

50g of NH3(g)150

101

Ammonia is a great example, because water can absorb what seems like a huge amount of ammonia gas before it becomes saturated. Mass-wise, its actually half the weight of the water, but volume-wise its over 720 times greater!

55 g of NH31005 g

+

+

+ +

gg

g g

mL mL

108

• Solubility indicates the maximum amount of solute that can dissolve in a given volume of solvent at a given temperature.

• Solubility is usually expressed as grams of solute per 100 mL of solvent (g/100mL).

• A substance’s solubility can vary with temperature:

• Solubility of solids usually increases with temperature• Solubility of gases usually decreases with temperature• Solubility of gases can also be affected by pressure.

109

Solubility Curves(Graphs of Solubility vs. Temperature. See page 16)

• Notice how most of the solids become more soluble at higher temperatures – KNO3, for example, starts at a

mere 10 g/100 mL at 0°C, but goes right off the top of the chart by 70°C

• Notice that most of the gases become less soluble at high temperatures – NH3 goes from 90 g/100mL at 0°C

to less than 10 g/100 mL at 100°C

110

Concentration and Dilution• Concentration is the ratio of dissolved solute to total

amount of solution.• General formula for concentration is:

• But concentration can be expressed in many different units, including:

• g/L (grams per Litre) g/mL (grams per millilitre)• % (by volume) % (by mass)• ppm (parts per million) mol/L (molar concentration)

• Molar concentration is the most important.

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111

Molar Concentration(molarity)

• The molar concentration is the number of moles of solute that is dissolved in one mole of the solution.

• Molar concentration can be represented by the letter C, or by square brackets [] or occasionally by a capital M used as a unit (molarity). Any of the following notations could represent a 2.0 mol/L solution of hydrochloric acid:

CHCl = 2.0 mol/L

[HCl] = 2.0 mol/LCHCl = 2.0 M

The correct unit for molar concentration is mol/L, although this is sometimes

abbreviated with a capital M for molarity

112

Molar Concentration Formula

n

C VConcentration

(mol/L)

# moles(mol)

Volume(L)

C =

Molar Concentration =

V =

n = CV113

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Dilution

• Dilution is a physical change that lowers the concentration of a solution by adding more solvent.

• The dilution formula is:

C1V1 = C2V2

Where: C1 is the concentration before dilution,V1 is the volume before dilutionC2 is the concentration after dilutionV2 is the volume after dilution

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Assignments

• Read pages 15 to 18• Do page 16

• Questions 15 to 17

• Do page 19:• Questions 18 to 23

115

Electrolytes

OverviewOur bodies are full of electrolytes: salts, acids and bases dissolved in our blood and tissues that are very important for the function of muscles and nerves.Most people think that water conducts electricity, but in fact, pure water is a poor conductor. It is the presence of dissolved electrolytes that gives the water in lakes, rivers and us the ability to conduct electricity.

Electrolytes

• Electrolytes are substances which, when dissolved in water, allow the solution to conduct electricity.

• Electrolytes are usually ionic compounds.• Electrolytes “dissociate” into positive and

negative ions when they dissolve.• There are three main types of electrolytes:

Acids, Bases, and Salts.• Most solid electrolytes do not conduct

electricity until they are dissolved.

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Electrolyte CharacteristicsAcids Bases Salts

Ions: Release H+ ions Release OH-ions Metal and non-metal ions

neutralization Neutralize bases Neutralize acids Products of neutralization

pH pH is less than 7 pH is greater than 7 pH variable, close to 7*

Litmus Turn litmus paper red Turn litmus paper blue Don’t change litmus*

Phenolphthalein Stays clear Turns red/purple Stays clear*

Formula H + non-metal Metal + OH Metal + non-metal

Dissociation eg: HCl(g) H+(aq) + Cl-

(aq) NaOH(s)Na+(aq)+OH-

(aq) NaCl(s)Na+(aq) + Cl-

(aq)

pH ScaleThe pH (positive Hydrogen potential) scale is used to measure the relative acidity or alkalinity of a solution. It is in theory open-ended, but in practice runs from 0 to 14.

Strong Acids Weak Acids Neutral Weak Base Strong Base

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

118* Some salts are slightly acidic (aluminum salts) or slightly basic (carbonates)

Assignments

• Read page 19• Do page 20

• Questions 24-27 • Question 28

119

Chemical Change & Stoichiometry

OverviewChemical changes are the root of chemistry, and the chemical equation is the fundamental tool we use to understand chemical changes.

Once you understand the chemical equation, you can use the techniques of stoichiometry to find the amounts of materials needed or used in chemical changes.

Review7

0.7 Chemical Changes

• Chemical changes occur when substances (reactants) react to form new substances (products).

• The products differ from the reactants:• They have different characteristic properties.• They have different molecular or ionic arrangements.

Reactants ProductsReactants on the

Left side of equation

Products on the Right side of

equation

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• Indications that a chemical change has taken place include:

• Release of a gas (effervescence)• Significant change in colour• Formation of a precipitate (solid from two solutions)• Change of energy in the form of heat, light or explosion.

• Parts of a chemical equation:

4 Fe (s) + 3 O2 (g) 2 Fe2O3 (s)

Reactants Product

Chemical equation

Coefficients4:3:2

(used for balancing)

Change to

Phases(s) Solid (l) liquid (g) gas

(aq) dissolved in water

Indexes* 2,2,3

Number of atoms in the molecules

*Yes, I am fully aware that the dictionary says that the correct plural of index is indices, but for clarity I am using the term the text uses. 122

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O OHH

HH H

Conservation of Mass• During a chemical reaction, mass is neither

lost nor gained• The total mass of all the reactants is equal to

the total mass of the products.• This is because no atoms are created or destroyed

during the reactions. The atoms are just rearranged.

• The balancing of chemical equations is based on the law of conservation of mass.

H OH O

2H2 + O2 2H2OH

m reactants = m products

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123

Balancing Equations• Balancing means adding coefficients in front of the

formulas of an equation so that it will conform to the law of conservation of mass

• A word equation names the reactants and products• A skeleton equation is an unbalanced equation• A balanced equation respects conservation of mass.

• Rules for balancing equations:• Only coefficients may be added or changed. The indexes in

formulas must not be changed.• You do not need to write the coefficient 1. It is understood.• Balanced equations should be reduced to the lowest terms.• When an equation is properly balanced, the total number of

atoms of each element on the left and right sides will be equal.

0.7.2p. 22

124

Stoichiometry

• Stoichiometry is the study of the relationships between the amounts of substances (reactants and products) that take part in a chemical reaction.

• Stoichiometry can be used to:• Calculate the amount of reactants need for a reaction• Calculate the expected amount of product from a

reaction.

0.7.3p. 23

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Steps for Stoichiometry1. Balance the equation, or verify that the equation you have

been given is properly balanced.2. Use the coefficients to find the mole ratios3. Write the amount in moles of the known reactant under

the corresponding mole ratio number.• If the amount is given in grams, convert it to moles using the

mole formula.

4. Write an x under the mole ratio of the substance you are looking for. Ignore the other substances for now

5. Change the : to =; Solve for x by cross multiplying.6. The result is the answer in moles.• If you need an answer in grams, convert using the mole formula

(with the proper molar mass!) 126

Example: moles to moles

H2 + O2 H2O

Mole Ratio Row: 2 1 2

Actual Mass Row: m= 8.0 g

Problem: 8.0 grams of hydrogen are burned with oxygen to make water. How much oxygen was used?

Molar Mass Row: M=2.0 g/mol

Moles row: n =4.0 mol

2 2

n= 2.0 mol

M=18.0 g/mol

m=36.0 g

molxmol

1

0.4

2

Example: gram to gram

H2 + O2 H2O

Mole Ratio Row: 2 1 2

Actual Mass Row: m= 8.0 g


Molar Mass Row: M=2.0 g/mol

Moles row: n =4.0 mol

2 2

n= 2.0 mol

M=18.0 g/mol

m=36.0 g

molxmol

1

0.4

2

Example1. Balance the equation, or verify

that the equation you have been given is properly balanced.

2. Use the coefficients to find the mole ratios

3. Write the amount in moles of the known reactant under the corresponding mole ratio number.• If the amount is given in grams,

convert it to moles using the mole formula.

4. Write an x under the mole ratio of the substance you are looking for. Ignore others

5. Change : to =. Solve for x by cross multiplying.

6. The result is the answer in moles.• If you need an answer in grams,

convert using the mole formula (with the proper molar mass!)


Step 1: H2 + O2 H2O (skeleton)

2H2 + O2 2H2O (balanced)

Step 2: mole ratios 2 : 1 : 2

Step 3: known reactant is 8.0 g hydrogen. To convert it to moles we must divide by the molar mass of hydrogen, 2.0; That gives us 4.0 moles of hydrogen. Write this under the corresponding mole ratio 2 : 1 : 2

4.0

Step 4: write an x under 2_ : 1 : 24.0 x

Step 5: cross multiply 2 = 1 so x = 2.0 mol 4.0 x

Step 6: to get the answer in grams, multiply the 2 mol by the molar mass of oxygen (32 g/mol) to give us the answer 64 g of oxygen is used.

129

Limiting and Excess Reagents• Sometimes the gram amounts of both reactants are

given to you in a stoichiometry problem. • It is possible that the given amounts could produce

different amounts of product. • You must do the stoichiometry twice... Once for each

reactant. The actual amount of product will be the lower of the two results.

• The reactant that produced the lower result is called the limiting reagent. The one that produced the greater result is called the excess reagent. Not all of the excess reagent will be used up.

Example of Limiting and Excess• 10 grams of hydrogen are burned with 64 grams of oxygen.

Which is the limiting reagent? How much of the excess reagent will be left over?

• Step 1. Balance equation: 2 H2 + O2 2 H2O• Step 2. Molar ratios• Step 3. Known reactant 1• Step 4. Ignore reactant 2• Step 5. Cross multiply• Step 6. Find answer• Repeat for reactant 2

2 : 1 : 210 g = 5 mol x

molx

mol

mol

mol 2

5

2

Answer 1 = 5 mol water or 90 g

X64 g = 2 mol x

molx

mol

mol

mol 2

2

1

Answer 2 = 4 mol water or 72 g

Answer 2 is the correct one.Oxygen is the limiting reactantThere will be an excess of one mole of hydrogen which will weigh 2 grams.

Assignments

• Read pages 21-24• Do Question 30 on page 23• Do Questions 31 and 32 on page 24

132

Examples of Chemical Reactions

OverviewNeutralization, Synthesis, Decomposition, Precipitation, Oxidation, Combustion.

These are all terms associated with specific chemical reactions.

Review8

0.8 Examples of Chemical Reactions

• There are many types of chemical reaction. Among the most important categories are:

• Acid-base reactions• Synthesis, Decomposition and Precipitation Reactions• Endothermic and Exothermic Reactions• Oxidation and Combustion• Photosynthesis and Respiration

• These are just a few of the types. Some your textbook does not mention include:

• Single Replacement • Double Replacement

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NOTE: A particular reaction can belong to several categories simultaneously. The burning of hydrogen is a synthesis, an exothermic reaction and an oxidation reaction.

Acid – Base Reactions• When an acid and a base are mixed:

• The H+ ions from the acid join the OH– ions from the base to make H2O, that is water.• The other ions, usually a metal and a non-metal ion,

join to form a salt whose nature depends on the reagents.• If the original solutions contained equal amounts of H+

and OH-, then the mixed solution will be neutral.• If there was a surplus of H+ or OH- ions, then the

resulting solution will be slightly acid or slightly basic.

In general: ACID(aq) + BASE(aq) WATER(l) + A SALT(aq)

Example: HNO3(aq) + KOH(aq) H2O(l) + KNO3(aq)

The words Neutralization

and Titration are also

associated with this process

0.8.1

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Synthesis, Decomposition, Precipitation

• Synthesis is when two or more reactants combine to form a single product.

• Eg. 2Na(s) + Cl2(g) 2 NaCl (s)

• Decomposition is when a single reactant breaks into two or more products.

• Eg. 2 H2O(l) 2H2(g) + O2(g)

• Precipitation is when a solid powder is formed by the mixing of two solutions.

• Eg

0.8.2

136

Endothermic and Exothermic

• Endothermic reactions are chemical reactions that absorb energy. Endothermic reactions usually make their immediate surroundings cooler.

• Reactants + Energy Products• Exothermic reactions release heat. They often

make their surroundings warmer.• Reactants Products + Energy

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Oxidation and Combustion0.8.4

• Oxidation is process where a substance combines with an oxidizer (usually O2, but O3, F2, Cl2, N2O and other substances work as oxidizers too).

• Your textbook incorrectly states that an oxide is always formed, but sometimes chlorides or fluorides are formed by oxidation.

• Slow oxidation takes time to happen.• Eg. The rusting of iron: 4 Fe + 3 O2 2 Fe2O3

• Combustion is rapid oxidation that produces heat and flames.

• Eg. Combustion of gasoline: 2 C8H18 + 25 O2 16 CO2 + 18H2O

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Photosynthesis and Respiration

• Life on Earth depends on two related chemical processes:

• Photosynthesis is the chemical reaction in which organisms, such as plants, transform radiant energy from sunlight into stored chemical energy.

6 CO2 + 6 H2O + energy C6H12O6 + 6 O2

• Respiration is the process by which organisms release stored chemical energy in sugars and other organic compounds in living cells.

C6H12O6 + 6 O2 6 CO2 + 6 H2O + energy

0.8.5p. 27

Error in textbook: On p. 27, respiration is referred to as a “combustion” reaction.What the textbook means, of course, is that is an “oxidation” reaction. 139

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Single and Double

• Replacement (or displacement) occurs when atoms of an element detach from one compound and attach to a new compound.

• Single replacement pattern A+BC AB+C example: Fe + CuSO4 Cu + FeSO4

• All simple metal with acid reactions are single replacements

• Double replacement, AB + CD AD + CB example: AgNO3(aq) + HCl(aq) AgCl(s)↓ + HNO3(aq)

• Many precipitation reactions are double replacements.

DisplacementReplacement

Fe Cu

Assignments

• Page 25 #33• Page 26 #35-36• Page 27 #37-38

Chemical Bonds

OverviewElectrical forces hold matter together. These forces come from the attraction between negative electrons and positive protons in atoms.

We call these forces chemical bonds, and there are several types of them.

Ultimately, it is chemical bonds that prevent matter and objects from spontaneously disintegrating .

Review9

Chemical Bonds• Chemical bonds are the forces that bind atoms together

into larger structures, such as molecules or crystal lattices.• Chemical bonds are the result of exchange or sharing of

electrons between two atoms, which causes the formation of a compound or diatomic or polyatomic element.

• There are many types of chemical bond. The four most important are:

• Metallic: metal to metal, found in alloys• Ionic: metal to non-metal, found in salts• Covalent: non-metal to non-metal, found in molecules.• Weak bonds: Van der Waals force, Hydrogen bonds

Not Studied

Studied

Studied

Sugar, covalent moleculeSalt, ionic crystal lattice

Your textbook has little about metallic bonds, but since we don’t study alloys in detail, this is not a problem.

0.9.0

143

Not Studied

Metallic Bonds and Weak Bonds(optional)

• A metallic bond consists of trillions of positive metal ions sharing a vast pool of negative electrons.

• Atoms of several different types of metal can share the same pool of electrons. That’s why metals can form an infinite number of alloys instead of specific compounds.

• If a piece of metal is dented or deformed, the disrupted pool of electrons instantly reforms bonds.

• This accounts for the great malleability and ductility of metals.

• The pool of electrons allows easy passage of other electrons through the material

• That is why metals are such good conductors

• Weak bonds:• Hydrogen Bonds: weak forces between polar-covalent molecules. These

account for some crystal structures, like ice.• Van der Waals Forces: weak forces between particles of solids

Ionic Bonding• An ionic bond forms when electrons are exchanged

between two atoms.

• This type of bond forms when one of the elements has a much higher electronegativity (X) than the other. This usually happens between a metal atom and a non-metal atom.

• Ionic bonds are between negative and positive ions• Ionic bonds do not form strong, distinct molecules. In most

ionic solids, the ions form a crystal lattice of alternating positive and negative particles. Some chemists prefer the term “formula units” to “molecules” when talking about ionic compounds.

Alternating particles do not overlap.

A crystal lattice structure with

alternating ions

0.9.1p. 28

NaNa+ Cl–cation anionCl

Sodium has an “extra” electron in its outer shell

145

Na+ Cl–

A sodium chloride formula unit

N ClCl

Cl A covalent molecule

Chlorine “needs” another electron in its outer shell

X.= 3.16X.= 0.93 ΔX = 2.23

Electro-negativityand Bond Type

• The electronegativity (X )(Greek letter chi or curly x) of an element can be found from the periodic table in front of your textbook.

• It indicates how much an element attracts electrons.• The greater the electronegativity difference between two

elements, the more likely they will form an ionic bond.• No bond is 100% ionic or 100% covalent, but we treat them that

way for simplicity.• The character of a bond is based on several things, in addition to

electronegativity, so the chart below is an approximation.

Δ X Character of bond Name of Bond type

1.7 to 3.9 Over 50% ionic Ionic

0.4 to 1.7 10% to 50% ionic Polar-Covalent

0.0 to 0.4 Less than 10% ionic Covalent

Chart of Electro-negativity

Covalent Bonding

• A covalent bond forms when electrons are shared between two atoms.

• This type of bond forms when two elements have similar electronegativity. This usually happens between two identical atoms, or between two non-metal atoms.• Covalent bonds can be single (sharing one pair of

electrons), double (sharing two pairs) or triple (sharing three pairs)• Covalent compounds form true, strong molecules.

They are sometimes referred to as molecular compounds.

C OO

Shared electrons in overlapping shells

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Illustrating Covalent BondsWith Rutherford-Bohr models:

With Lewis electron dot diagrams:

In either case, we draw the atoms to show a stable number of electrons (usually 8) in the outer shell of each atom involved in the covalent bond.

N ClCl

Cl Another way to illustrate covalent bonds is with overlapping circles

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Electrons shared between Carbon and one Hydrogen atom.

Electrons shared between Carbon and another Hydrogen atom.

AnotherAnother

Or you can just circle an electron pair to show they are shared.

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Assignments

• Page 25 #33• Page 26 #35-36• Page 27 #37-38

Energy

OverviewChemical reactions involve energy. Some reactions absorb energy, others release it.

There are dozens of specific types of energy, but we can group them all into two main categories: Kinetic energy, and Potential Energy.

Energy• Energy is the ability to do work or make a change. • Energy is classed into Kinetic and Potential

• There are many sub-types of energy, a few examples of which are listed in the table below:

Form of Energy Associated with Example

Mechanical Kinetic Energy An object’s movement Car driving along a road

Thermal Kinetic Energy Agitation of particles Boiling water

Radiant Kinetic Energy Electromagnetic waves Light, microwaves, radio waves

Gravitational Potential Energy Object’s position above ground Water behind a dam

Elastic Potential Energy Compressed/stretched materials A spring that has been stretched

Electric Potential Energy Force between electric charges Charged particles in a storm cloud

Nuclear Potential Energy Stored in the nucleus of atoms Uranium in a reactor

Chemical Potential Energy Stored in the bonds of molecules Energy in gasoline or glucose

* Potential energy can rapidly change into kinetic, for example, electrical potential energy can create electricity, a form of kinetic energy. Nuclear energy can create thermal energy.

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Units of Energy• In chemistry, energy is measured in joules

• if there are more than a thousand joules then we can measure them in kilojoules. (1 kJ = 1000 J)

• A joule is the amount of energy required to accomplish any of the following:

• Moving a force of one newton a distance of one metre.(N m)∙• Passing a 1 A current through a 1Ω resistor for 1 s (A2∙Ω s)∙• Using 1 Watt of power for 1 second (W s)∙• Moving a coulomb of electrons through 1 volt (C V)∙• Keep a human body alive for 1/100th of a second.

• In other fields energy can be measured in units such as thermal calories (cal), food Calories (kcal or Cal), ergs, British Thermal Units (BTU), Kilowatt-hours (kW h), Tonnes of TNT ∙(or kilo or megatonnes).

Energy Comparisons• Energy units smaller than a joule

• 1 joule = 6.24 x 10 18 electron Volts (eV)• 1 joule = 10 000 000 ergs (erg)

• Energy units greater than a joule• 4.184x10 9 joules = 1 ton of TNT• 3 600 000 joules = 1 kilowatt-hour (kW h)∙• 4184 joules = 1 kilocalorie (kcal)*• 1055 joules = 1 British Thermal Unit (BTU)

• A joule is:• The energy required to lift a small apple (102g) 1 metre• The energy a 1 watt LED uses every second.

* A kilocalorie is also known as a food calorie. It is the same unit identified as a calorie (Cal) on food packages. A thermal calorie (cal) is 1/1000 of a food calorie, or only 4.18 joules.

Kinetic Energy

• Kinetic energy is the energy associated with the movement of an object, or with the movement of its particles (molecules).

• Kinetic energy depends on the mass of the object and the velocity of its motion.

Where: Ek= kinetic energym= mass of the objectv= velocity of the object

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Potential Energy

• Potential Energy is energy stored in a body that can be transformed into another form of energy.

• Potential energy is sometimes referred to as “hidden energy”, since it is difficult to observe and measure.

• There are several types of potential energy, including:

• Gravitational Potential Energy (important in physics)• Chemical Potential Energy (important in chemistry)

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• Gravitational Potential Energy is the product of an object’s mass, its height above the ground, and the gravitational acceleration.

• Chemical Potential Energy (Enthalpy) is associated with the energy in the bonds between the particles of a material.

Where: Ep = Gravitational Potential Energy in joulesm = mass of the object in kilogramsg = gravitational acceleration (9.8 m/s2 on Earth)h = height of the object above a reference point (such as the ground)

We will devote a section later in the course to calculating enthalpy.

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Conservation of Energy• The law of conservation of energy states that

energy cannot be created or destroyed in chemical reactions, but it can be changed from one form to another.

• Potential energy can change to kinetic and vice versa

• Mechanical Energy is the total energy of an isolated system.

Where: Em = total Mechanical EnergyEp = Potential EnergyEk = Kinetic Energy

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Thermal Energy & Temperature

• Thermal energy or “heat” is a form of energy possessed by a substance due to the agitation of its particles. It depends on:

• The mass of the substance• The temperature of the substance• The specific heat capacity of the substance

Where: Q = amount of heat energy in joulesm = mass of the substance heated in grams (usually the water in a calorimeter)c = specific heat capacity of the substance heated, in j/g∙°CΔT = the change in temperature in °C 158

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Fluids• Compressible & Incompressible Fluids

• Substances that flow, like liquids and gases, are fluids• Gases are compressible fluids• Liquids are incompressible fluids

• Pressure• Pressure is the force exerted on a surface.• The standard unit of pressure is the kilopascal (kPa)• Formula for pressure: Pressure = Force divided by Area.

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Standard Atmospheric Pressure:

Ps = 101.3 kPa

LiquidGas

END of MODULE 1

• Prepare for the module 1 test.– Make sure that you have read and understood all

pages up to page 33 in your text book.– Make sure you have downloaded and read the

study notes called “Chemistry Unit 0” from my web site: chem534.wikispaces.com

– Prepare your own study notes by highlighting all the sections you think are important, and then recopying the highlights into your notebook.

– practice conversions and significant figures

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pages 3 to 33 “quantum chemistry” target completion date: october 1

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