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Chemistry Review Section. Pages 3 to 33 “Quantum Chemistry” Target Completion Date: October 1. Pages with a PINK background are supplementary . Not material for a test!. About Slide Icons. Very Important Points

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pages 3 to 33 quantum chemistry target completion date october 1

Chemistry

Review Section

Pages 3 to 33

“Quantum Chemistry”

Target Completion Date: October 1

about slide icons

Pages with a PINK background are supplementary . Not material for a test!

About Slide Icons

Very Important Points

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

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 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.

R

conversions

R

Conversions
  • You must be able to do ALL standard metric conversions, especially:
    • Litres to millilitres, millilitres to litres
    • Grams to kilograms, kilograms to grams
  • Other conversions you will learn during the course of the year:
    • Temperature: degrees celcius (℃) to kelvins (K)
    • Pressure: kilopascals (kPa) to millimetres (mmHg)
quick conversions
Quick Conversions

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 km

Conversion: 2.524 km  ? cm

00 cm Add extra zeros if necessary

2 524 km

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

density

R

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:

ρ V

m

or

Where: ρ= the density of the object, in g/cm3 or g/mL

m = the mass of the object, in g

V = the volume of the object, in cm3 or mL

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.

ρw= 1 g/mL = 1 g/cm3

solving problems
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:
    • Show your data, the information you used to solve the problem.
    • Show your work, including the formulas you used and the substitutions you made.
    • Write an answer statement, a sentence that clearly states your final answer.
    • Include the correct units for your answer. Never just give a number—you must specify what the number means!
suggested solution method
Suggested Solution Method

Problem: 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 cm

h =2.0 cm

m =50g

V = ?

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:

V =12cm x 5cm x 2 cm

= 120 cm3

𝜌= 50 g / 120 cm3

= 0.46 g/cm3

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

Answer: The density of the block is 0.46 g/cm3 (or 0.46 g/mL)

problems on conversions and density
Problems on Conversions and Density
  • Convert the following:
    • 125 mL to L d) 30 mL to L g) 75 mL to L
    • 450 g to kg e) 4500 mL to L h) 0.035L to mL
    • 2.5 L to mL f) 1.35 kg to g i) 0.56L to mL
  • Find the density of a 4cm x 3cm x 2cm block that has a mass or 480 g. Justify your solution.
  • 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 significant figures knowing how much to round an answer
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”.

uncertainty
Uncertainty

In math, numbers are considered pure, abstract things. In math, 2.00, 2.0 and 2 are considered the same, they all represent perfect number 2.

In science, numbers are considered to be measurements, and all measurements have some degree of uncertainty. They are almost never considered perfect!

The absolute uncertainty of a measurement is usually ½ of a measuring instruments smallest gradation. If a graduated cylinder is marked in millilitres, then each measurement taken with that cylinder has a ±0.5 mL uncertainty.

In science, 2 mL, 2.0 mL and 2.00 mL are different!

correct precision
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. 20.0 ±0.5 mL or 25.5±0.5 mL are acceptable readings.
      • With the same graduated cylinder, it would be wrong to write 20 ±0.5 mL or 25 ±0.5 mL or even 20.00 ±0.5 mL
  • In science 20 mL, 20.0 mL and 20.00 mL have different meanings with respect to precision.
rules for significant figures interpreting significant digits
Rules for Significant Figures Interpreting Significant Digits
  • All non-zero digits are ALWAYS significant
  • Zeros between significant digits are ALWAYS significant.
  • Zeros at the beginning of a number are NEVER significant.
  • Zeros at the end of a number MAY be significant.
  • Exponents, multiples, signs, absolute errors etc. are NEVER significant.
examples of rule 1 2 and 3
Examples of Rule 1, 2 and 3

Rule 1. Non-zero digits are ALWAYS significant.

1.234 has 4 significant digits

145 has 3 significant digits

19567.2 has 6 significant digits

Rule 2. Zeros between significant digits ARE significant.

  • has 4 significant digits

5007.4 has 5 significant digits

20000.6 has 6 significant digits

Rule 3. Zeros at the beginning are NEVER significant.

007 has 1 significant digit

0.0000005 has 1 significant digit

0.025 has 2 significant digits

explaining rule 4
Explaining Rule 4

Rule 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 digits

5.10 has 3 significant digits

10.00 has 4 significant digits

If 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 any of the zeros at the end.

5000 has 1 or 4 significant digits

250 has 2 or 3 significant figures

123 000 000 has 3 or 9 significant figures

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

Trusted precise source

Estimated source

rule 5
Rule 5

Rule 5: Exponents and their bases, perfect multiples, uncertainties (error values), signs etc. are NEVER significant.

6.02x 1023 has 3 significant digits

504.1 mL x 3has 4 significant digits

5.3±0.5 mL has 2 significant digits

–5.432x 10-5has 4 significant digits

In each case, the blue part is significant, the greenpart 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, but not one that adds to the precision of the answer.

same number different precision
Same Number, Different Precision

Try to avoid the “ambiguous” situation in your answers. If an answer ends in zero, or worse, in several zeros, indicate whether it should be interpreted as “exactly” or “approximately”.

Better still, convert it to scientific notation, and leave only the zeros you know are accurate.

Eg. If your answer is 2500 mL, but you only measured to the nearest 10mL, then write 2.50 x 103 mL. That way every one will know its accurate to 3 significant figures

*Your textbook says to call this 3 significant figures. Traditional measurement would call it 1 significant figure. Written this way it is ambiguous.

Avoid writing answers that end with zeros and no decimal!

math with significant figures
Math with Significant Figures
  • The basic rule for math is that you do not improve significance by multiplying or dividing numbers:

53.81 m x 2.43 m = 131 m2NOT130.7583 m2 !!!

Why? Because the least precise measurement had 3 significant digits, so our answer should not have more than 3 significant digits!

The technique for addition and subtraction is slightly different (see p.396 ) but the concept is the same. You cannot make your result better than your measurements!

topic 1 organization of matter

Page 4

R

Topic 1: Organization of Matter

0.1.1

O

O

C

H

  • 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:

CO2

H

N

NH3

H

One atom

Ne

Ne

Cl

Cl

S

several thousand atoms

O

DNA

2 atoms of hydrogen

1 atom of oxygen

SCl2

18

H2O

H

H

slide19

Page 4

Cl–

Cl

Na

Na+

cation

anion

0.1.2

  • Chemical Formulas and Ions
    • Some matter is formed from ions instead of normal atoms or molecules.
      • For the most part, we treat ions the same way as regular atoms, 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 almost always form cations (+), non-metals may form anions (-)

Notice the slightly stronger wording with respect to metals than nonmetals!

19

sample ions
Sample Ions

Metal Ions (+)

+ Cations

– Anions

Non-Metal Ions (-)

Notice 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.

21

big fat ions polyatomic ions

2–

O

Big Fat Ions(Polyatomic Ions)

+

H

O

S

O

N

H

O

H

O

O

3–

P

H

O

O

Cl

O

  • Polyatomicions 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

N-

O

O

O

Na+

Na+ + NO3- NaNO3

N-

O

O

common polyatomic ions p 422
Common Polyatomic Ions (p.422)

3-

2-

1-

1-

23

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

representation of atoms

R

Representation of Atoms

0.2.0

H

  • Early Representations
    • Democritus (c.450 BC) 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.

C

N

O

P

S

Cl

Dalton models

Original and Modern

+

-

-

-

-

-

24

slide25

R

Page 5

0.2.1

1. 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.

Early Rutherford model

Revised Bohr model

25

slide26

Page 5

0.2.2

2. 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.

Symbol: The symbol of the element

Electrons: 2 in first shell, 8 in 2nd 1 in 3rd

11p+

12n0

Na

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

2e- 8e- 1e-

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

Z=11, configuration: 2,8,1

26

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

Unacceptable!

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.

ACCEPTABLE

ACCEPTABLE

slide29

R

0.2.3

3. 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.
  • When bonding, atoms gain, lose or share electrons in order to get a total of 8 electrons around each atom.

2 paired electrons

1

5

3 “odd” unpaired electrons

4

2

3

29

slide30

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

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

30

the modern model optional enrichment
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.

31

the modern model optional enrichment1
The Modern Model(Optional Enrichment)
  • 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 “aufbaudiagram” that explain this .

32

the modern model optional enrichment2
The Modern Model(Optional Enrichment)
  • 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:
    • 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…

33

atomic model exercises
Atomic Model Exercises
  • Draw Simplified Models of the first 20 elements.
  • Draw Lewis Models of the first 20 elements.
  • Convert the following:
    • 125 cm to m d) 320 mL to L g) 750 mL to L
    • 280 g to kg e) 45000 mm to km h) 0.0035km to cm
    • 4.63 L to mL f) 5.52 kg to g i) 0.45L to mL
periodic classification
Periodic Classification

Overview

The 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.

in line notation of element information
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.
in line notation for a carbon 14 atom carbon 14 is an isotope or alternate form of normal carbon
In-Line Notation for a Carbon-14 atom(carbon-14 is an isotope, or alternate form of normal carbon)

An average carbon atom weighs 12.01 amu according to the periodic table. But no atom of carbon has that exact weight. For every thousand atoms that weigh exactly 12 amu, a few weigh more. This one weighs 14 amu

Valence (4)

or

Ionic Charge (4–)

or

Oxidation # (–4)

Isotope or Mass number. Represents the number of nucleons in a particular atom

C

14

4–

Valence is the number of bonds the atom is likely to form. Ionic charge is the most likely charge an ion will have.

Atomic number “Z” represents the number of protons in this atom

2

6

Subtracting the Mass # and the Atomic # “Z” gives the number of neutrons in the atom

Number of atoms in a molecule, such as C2H4

(6p+, 8n0, 10e-)

Configuration of this atom

8

information in your periodic table
Information in your Periodic Table
  • 8

Atomic number (Z)

The number of protons

2-

Ionic charge

3.44

0.65

Atomic Radius

Electronegativity

O

1314

1.43

Density (g/L gas)

(g/mL solid/liquid)

Ionization Energy

-218.3

-182.9

Melting Point (°C)

Symbol

Boiling Point (°C)

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

Name

Oxygen

The English name of the element

15.999

Atomic weight (amu)

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

(or g/mol)

Also the molar mass in g/mol

The symbol is a 1 or 2 letter abbreviation of the element’s name, or sometimes its Latin 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!)

38

the periodic table with regions shaded
The Periodic Tablewith Regions shaded

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

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

39

the periodic table with families shaded
The Periodic Tablewith Families Shaded

VIIIA: Noble Gases

IIA: Alkaline Earths

IA: Alkali Metals

VIIA: Halogens

VI: Oxygen Family

IVA: Carbon Family

IIIA: Boron Family

V: Nitrogen Family

IB: Coin Metals

Iron Triad

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

Lanthanides

Actinides

40

the periodic table and valence electrons electrons in outermost shell
The Periodic Tableand Valence Electrons (electrons in outermost shell)

THREE (III)

EIGHT (VIII)

SIX (VI)

ONE (I)

TWO (II)

FOUR (IV)

FIVE (V)

SEVEN (VII)

FIVE (V)

SEVEN (VII)

THREE (III)

FOUR (IV)

SIX (VI)

TWO

ONE

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

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.

41

the periodic table with periods shaded
The Periodic Table with Periods shaded

1st Period = 1 shells

2nd Period = 2 shells

3rd Period = 3 shells

4th Period = 4 shells

5th Period = 5 shells

6th Period = 6 shells

7th Period = 7 shells

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

6th Period = 6 shells

7th Period = 7 shells

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

42

use the periodic table to find the electron arrangement of an atom
Use the Periodic Table to Find the Electron Arrangement of an Atom

Eg. Find the electron arrangement of Iodine (I)

SEVEN (VII)

5th Period = 5 shells

53

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.

Five shells

53p+

2

8

18

18

7e-

Total 53, So far: 35, left: 18

periodic table exercises
Periodic Table Exercises
  • Write the name and symbol of each of the first 20 elements. (bragging rights if you can do it without looking!)
naming compounds
Naming Compounds
  • There are four sets of rules for naming compounds:
    • The binary ionicrules:
      • For compounds containing only two elements, joined by an ionic bond.
    • The ternary ionicrules:
      • 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.

45

the binary ionic rules
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

1+

1–

Cl-

Ba2+

Cl-

Na+

Cl-

K+

K+

S2-

Ca2+

O2-

Ca2+

C4-

Ca2+

O2-

Al

O2-

Al

3+

3+

O2-

46

ionic rules no no
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 acceptable 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. Monosodium glutamate is an organic compound that does not follow the rules.

dealing with polyvalent metals
Dealing with Polyvalent Metals
  • Some metal elements have more than one possible valence. Copper, for example, can have a valence of 1+ or 2+, depending on what 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 has a charge of 1+

This copper ion must have a charge of 2+

49

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

C

Cr

Mn

Fe

Co

Ni

Ti

Cu

Pd

Sn

Sb

Ru

Pt

Po

Tl

Bi

Au

Hg

Pb

Sm

Eu

U

50

examples of ionic compounds with polyvalent elements
Examples of Ionic Compounds with Polyvalent Elements

*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∙Fe2O3)

52

the ternary ionic rules
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:
    • NaNO3sodium nitrate CaCO3calcium carbonate
    • K2SO4potassium sulphate Ba(CN)2barium cyanide
    • Al2(CrO4)3aluminum chromate NH4Cl  ammonium chloride

Polyatomic ions: See Table 8.10 on p. 422

53

covalent rules
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  dinitrogentetrahydride
    • PF3 phosphorus trifluoride** P2O5 diphosphoruspentoxide
    • CO2  carbon dioxide** CO  carbon monoxide**

* The last “o” in mono or the “a” in tetra, penta, or hexais 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.

54

electronegativity how much an atom attracts electrons
Electronegativity(how much an atom attracts electrons)

One of the uses of electronegativity is to decide which element goes first in a formula or name. Usually the element with the lowest electronegativity goes first. Therefore it is called carbon dioxide (CO2), NOT dioxygen carbide (O2C).

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.

do use prefixes with covalent compounds
DO use prefixes with covalent compounds

*commonly called hydrogen peroxide.

simplification of covalent names
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
      • Many chemists simply say “hydrogen sulphide” instead of “dihydrogen sulphide” for the compound H2S. Since H2S is the only common sulphide of hydrogen, 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.
      • Water can be called hydrogen oxide, but it is perfectly acceptable to use “dihydrogen oxide” or even “dihydrogen monoxide”

57

finding formulas from compound names
Finding Formulasfrom Compound Names
  • For covalent compounds, the name usually tells you the formula:
      • For example: dinitrogenpentoxide = 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 O2crossoverNa2O

58

the crossover rule and simple ionic compounds
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

59

the crossover rule and covalent compounds
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 dichloride

2) 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.

60

the crossover rule simplifying ionic compounds
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.

61

reverse crossover rule for finding the valence of uncertain ions
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? FeO
    • They are both iron oxide, but which iron oxide (there are several types!)
    • Fe O Fe O

2 2

2

3

1

1

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

Fe has a valence of 3, so the name of the compound is:

Iron(III)oxide

Fe’s proper valence here is 2

Iron(II)oxide

the organic rules not studied this year
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 (…tone), organic acids, 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.

practice
Practice
  • Page 12, Question #9
  • Practice sheets:
      • Naming ionic compounds
      • Naming covalent compounds
      • Naming mixed compounds

64

the mole concept and the enumeration of matter
The Mole Conceptand the Enumeration of Matter

0.5.1

  • 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

65

molar mass
Molar Mass

0.5.2

  • Molar massis 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

66

diatomic and polyatomic elements
Diatomic and Polyatomic Elements
  • Diatomic elements: There are seven elements whose molecules normally contain two atoms: I2, H2, N2, Br2, O2, Cl2and 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 and phosphorus, 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)

67

How to Remember the Diatomic Elements: IHave No Bright Or Clever Friends

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

m

Actual mass

(g)

# moles =

Molar mass =

n

M

Molar mass

(g/mol)

# moles

(mol)

68

Actual mass = # moles x molar mass

practice1
Practice
  • Page 14, #12, 13, 14
  • Practice sheet:
      • Moles and Molar mass

69

physical changes
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.

70

phase change

0.6.1

Phase Change
  • 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 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.

71

slide72

Gas

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.

Vaporization

Exothermic

Process

Liquid Condensation

Endothermic

Process

Solid Condensation

Sublimation

Terminology associated with

Change of Phase

Liquid

Solid

Liquid

Melting (fusion)

Freezing (solidification)

72

phase markers
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)

74

dissolution and solubility

0.6.2

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.

75

dissolving physical change
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)

76

dissolution of ammonia gas in water an extreme case of solubility at 25 c
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!

50g of NH3(g)

+

100

150

g

g

72.058 litres NH3(g)

+

mL

101

mL

100

55 g of NH3

5 g

+

+

g

100

150

g

77

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!

slide78

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.

78

solubility curves graphs of solubility vs temperature see page 16
Solubility Curves(Graphs of Solubility vs. Temperature. See page 16)
  • Notice how most of the solidsbecome 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

79

concentration and dilution
Concentration and Dilution

0.6.3

  • 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.

80

molar concentration molarity
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/L

CHCl= 2.0 M

The correct unit for molar concentration is mol/L, although this is sometimes abbreviated with a capital M for molarity

81

molar concentration formula
Molar Concentration Formula

Molar Concentration =

n

# moles

(mol)

C =

V =

V

C

Concentration

(mol/L)

Volume

(L)

n = CV

82

dilution
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 dilution

C2 is the concentration after dilution

V2 is the volume after dilution

83

assignments
Assignments
  • Read pages 15 to 18
  • Do page 16
      • Questions 15 to 17
  • Do page 19:
      • Questions 18 to 23

84

electrolytes
Electrolytes

0.6.4

  • 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.

85

electrolyte characteristics
Electrolyte Characteristics

pH Scale

The 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.

86

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

assignments1
Assignments
  • Read page 19
  • Do page 20
      • Questions 24-27
      • Question 28

87

0 7 chemical changes
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  Products

Reactants on the Left side of equation

Products on the Right side of equation

88

slide89

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:

Chemical equation

Reactants

Change to

Product

4Fe(s)+ 3O2(g)2Fe2O3(s)

Coefficients

4:3:2

(used for balancing)

Phases

(s) Solid (l) liquid (g) gas (aq) dissolved in water

Indexes* 2,2,3

Number of atoms in the molecules

89

*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.

conservation of mass
Conservation of Mass

0.7.1

  • 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.

O

O

H

O

H

O

H

H

H

H

H

H

2H2

+ O2

 2H2O

mreactants= mproducts

90

balancing equations
Balancing Equations

0.7.2

p. 22

  • 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.

91

stoichiometry
Stoichiometry

0.7.3

p. 23

  • 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.

92

steps for stoichiometry
Steps for Stoichiometry
  • Balance the equation, or verify that the equation you have been given is properly balanced.
  • Use the coefficients to find the mole ratios
  • 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.
  • Write an x under the mole ratio of the substance you are looking for. Ignore the other substances for now
  • Change the : to =; Solve for x by cross multiplying.
  • The result is the answer in moles.
    • If you need an answer in grams, convert using the mole formula (with the proper molar mass!)

93

example
Example

Problem: 8 grams of hydrogen are burned with oxygen to make water. How much oxygen was used?

  • Balance the equation, or verify that the equation you have been given is properly balanced.
  • Use the coefficients to find the mole ratios
  • 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.
  • Write an x under the mole ratio of the substance you are looking for. Ignore others
  • Change : to =. Solve for x by cross multiplying.
  • 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 8g hydrogen. To convert it to moles we must divide by the molar mass of hydrogen, 2.0; That gives us 4 moles of hydrogen. Write this under the corresponding mole ratio 2 : 1 : 2

4

Step 4: write an x2 : 1 : 2

4 x

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

4 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.

94

assignments2
Assignments
  • Read pages 21-24
  • Do Question 30 on page 23
  • Do Questions 31 and 32 on page 24

95

0 8 examples of chemical reactions
0.8 Examples of Chemical Reactions
  • There are many types of chemical reaction. Among the most important types 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

96

acid base reactions
Acid – Base Reactions

0.8.1

  • 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.

The words Neutralization and

Titration

are also associated with this process

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

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

97

synthesis decomposition precipitation
Synthesis, Decomposition, Precipitation

0.8.2

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

98

endothermic and exothermic
Endothermic and Exothermic

0.8.3

  • 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

99

oxidation and combustion
Oxidation and Combustion

0.8.4

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

100

photosynthesis and respiration
Photosynthesis and Respiration

0.8.5

p. 27

  • 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

101

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.

assignments3
Assignments
  • Page 25 #33
  • Page 26 #35-36
  • Page 27 #37-38
chemical bonds

0.9.0

Chemical Bonds

Salt, ionic crystal lattice

Sugar, covalent molecule

  • 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 three 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.

Not Studied

Studied

Studied

103

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

ionic bonding
Ionic Bonding

0.9.1

p. 28

  • 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.

Sodium has an “extra” electron in its outer shell

Cl–

Chlorine “needs” another electron in its outer shell

Cl

Na

Na+

cation

anion

X.= 3.16

X= 0.93

ΔX = 2.23

A crystal lattice structure with alternating ions

Cl

Cl–

Alternating particles do not overlap.

Na+

N

Cl

104

Cl

A covalent molecule

A sodium chloride formula unit

electronegativities supplemental
Electronegativities(Supplemental)
  • 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.
covalent bonding
Covalent Bonding

0.9.2

p. 29

  • 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.

O

O

C

Shared electrons in overlapping shells

106

illustrating covalent bonds
Illustrating Covalent Bonds

With 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.

Cl

N

Cl

 Another way to illustrate covalent bonds is with overlapping circles

107

Cl

energy

R

Energy

0.10.0

p. 30

  • Energy is the ability to do work or make a change.
      • There are many types of energy, a few of which are listed in the table below:

108

* The word “potential” is often inserted to indicate that these associated with potential energy.

kinetic energy

0.10.1

p. 30

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 energy

m= mass of the object

v= velocity of the object

109

0 10 2 potential energy

0.10.2

p. 30

0.10.2 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)

110

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R

  • 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 joules

m = mass of the object in kilograms

g = gravitational acceleration (9.8 m/s2 on Earth)

h = height of the object above a reference point (such as the ground)

0.10.2

p. 30

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

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conservation of energy

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Conservation of Energy

0.10.3

p. 30

  • 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 Energy

Ep = Potential Energy

Ek = Kinetic Energy

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0 10 4 thermal energy temperature

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0.10.4 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 joules

m = mass of the substance heated in grams (usually the water in a calorimeter)

c = specific heat capacity of the substance heated, in j/g∙°C

ΔT = the change in temperature in °C

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0 11 fluids

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0.11 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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end of module 1
END of MODULE 1
  • Prepare for the module 1 test.
    • Read up to page 33 in your text book
    • Prepare study notes.

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