SCH4U - Unit 1 - Version C.pdf

7/25/2019 SCH4U - Unit 1 - Version C.pdf

1/64

SCH4UGrade 12

University Chemistry

Version C


2/64

SCH4U Chemistry Introduction

Copyright 2008, Durham Continuing Education Page 2 of 64

Introduction

Welcome to the Grade 12 University Chemistry Course, SCH4U. This full-credit courseis part of the Ontario Secondary School curriculum. Prerequisite Grade 11 UniversityChemistry.

This course enables students to deepen their understanding of chemistry through thestudy of organic chemistry, energy changes and rates of reaction, chemical systemsand equilibrium, electrochemistry, and atomic and molecular structure. Students willfurther develop problem solving and laboratory skills as they investigate chemicalprocesses, at the same time refining their ability to communicate scientific information.Emphasis will be placed on the importance of chemistry in daily life, and on evaluatingthe impact of chemical technology on the environment.

How to Work Through This Course

Each of the units is made up of four lessons. Each lesson has a series of assignmentsto be completed. In this course you must complete ALLassignments. Be sure to readthrough all the material presented in each lesson before trying to complete theassignments.

Important Symbols

Questions with this symbol are Key Questions. They give you anopportunity to show your understanding of the course content. Ensure that

you complete these thoroughly as they will be evaluated.

Questions with this symbol are Support Questions. They do not need tobesubmitted to the marker, but they will help you understand the coursematerial more fully. Answers for support questions are included at the end ofeach unit. Refer to these for suggestions of how to properly structure theanswers to questions.

Remember, you must complete the KEY QUESTIONS successfully in order toachieve the credit in this course. Remember to write the unit number, lessonnumber and key question number on all assignments. Make sure that yourassignments are submitted in the proper order.


3/64



What You Must Do To Get a Credit

In order to be granted a credit in this course, you must

Successfully complete the Key Questions for each unit and submit them for

evaluation within the required time frame. This course is made up of 5 units.

Complete the mid-term examafter Unit 3.

Complete and pass a final examination.

After you submit lessons for evaluation, begin work on your next lesson(s) rightaway! Do not wait until you receive your evaluated assignments from the marker.

Your Final Mark

Each Unit has 4 lessons each worth 2% (10% per Unit x 4 Units) 40% Midterm Test 30%

Final Examination 30%

Materials

This course is self-contained and does not require a textbook. You will require linedpaper, graph paper, a ruler, a scientific calculator and a writing utensil.

Expectations

The overall expectations you will cover in each unit are listed on the first page of eachunit.

Term


4/64



Table of Contents

Unit 1 Structure and Properties

Lesson 1 Atomic TheoriesLesson 2 Quantum MechanicsLesson 3 Chemical BondingLesson 4 Intermolecular Forces

Unit 2 Organic Chemistry

Lesson 5 HydrocarbonsLesson 6 Functional GroupsLesson 7 Types of Organic Reactions

Lesson 8 Polymers

Unit 3 Rates of Reactions

Lesson 9 ThermochemistryLesson 10 Enthalpies of ReactionsLesson 11 Energy OptionsLesson 12 Chemical Kenetics

Unit 4 Electrochemistry

Lesson 13 Oxidation and Reduction ReactionsLesson 14 The Activity Series of MetalsLesson 15 Galvanic CellsLesson 16 Electrolytic Cells

Unit 5 Chemical Systems and Equilibrium

Lesson 17 Introducing EquilibriumLesson 18 The Equilibrium ConstantLesson 19 Acid and Bases EquilibriumLesson 20 Solubility Equilibriums


5/64

SCH4UGrade 12


Lesson 1 Atomic Properties


6/64

SCH4U Chemistry Unit 1 Lesson 1


Unit 1: Structure and Properties

Have you ever seen a picture of the first computer? Figure 1.1 below depicts what thefirst computer looked like. It may seem odd and funny to look back at such pictures, butmost technologies are constant works in progress. Computers now come in tiny

devices such as phones and laptops.

The development of a modern day atomic model is similar in nature to that of computertechnologies. There have been many modifications along the way to current modernatomic theory. In this unit you will learn about these models and modification. You willalso learn more about the molecular structures, and chemical bonding.

Figure 1.1: The First Computer

Overall Expectations

demonstrate an understanding of quantum mechanical theory, and explain howtypes of chemical bonding account for the properties of ionic, molecular, covalentnetwork, and metallic substances;

investigate and compare the properties of solids and liquids, and use bondingtheory to predict the shape of simple molecules;

describe products and technologies whose development has depended onunderstanding molecular structure, and technologies that have advanced theknowledge of atomic and molecular theory.


7/64



Lesson 1: Atomic Theories

The atomis the smallest unit of an element that retains the chemical properties of thatelement. In this lesson, you will learn about the various theories that led up to thecurrent model of the atom.

What You Will Learn

After completing this lesson, you will;

explain the experimental observations and inferences made by Rutherford and Bohrin developing the planetary model of the hydrogen atom;

describe some applications of principles relating to atomic and molecular structure inanalytical chemistry and medical diagnosis (e.g., infrared spectroscopy, X-raycrystallography, nuclear medicine, medical applications of spectroscopy);

describe advances in Canadian research on atomic and molecular theory

The Development of the Atomic Model

Matteris anything that has mass and takes up space. This means everything aroundyou (including yourself!) is matter. Matter is made up of tiny particles called atoms.The term atom was first coined by the Greek philosopher Democritus, who proposedthat the atom was the smallest particle that could not be subdivided.As experimentation and the scientific method gained importance, the model of the atombegan to evolve. Lets take a look at the scientists involved in the development of theatomic model.

Making a Model: Key Scientists

JOHN DALTON (1809)Dalton was an English schoolteacher came up with his atomic theory based on manyyears of experimentation by many scientists.

Daltons Atomic Theory1. All matter is composed of tiny particles called atoms2. Atoms can be neither subdivided nor changed into one another3. Atoms cannot be created or destroyed4. All atoms of one element are the same in shape, size, mass and all other properties

5. All atoms of one element differ in these properties from atoms of all other elements6. Chemical change is the union or separation of atoms7. Atoms combine in small whole-number ratios such as 1:1. 1:2, 2:3, etc.


8/64



Figure 1.1: Daltons Bill iard ball model

Daltons model did not account for the subatomic particles found in an atom (protons,neutrons, and electrons).

J.J. THOMPSON (1897)

studied the deflection of cathode rays by electric and magnetic fields

results suggested that the atom was not the smallest unit of matter; there weresubatomic particles within the atom.

J.J. Thompson proposed that the atom is a sphere of uniform positive electricity inwhich negative electrons were embedded like raisins in plum pudding (or chocolatechips in a cookie)

Figure 1.2 Thompsons raisin bun model for the atom

RUTHERFORD (1909)

Radium gives off three different types of radiation (alpha and beta particles, andgamma rays)

Alpha particlesare Helium nuclei (2 protons, 2 neutrons)

Using these principles, Rutherford performed his famous gold foil experiment

The Gold Foil Experiment

Rutherford hypothesized that If Thompsons model is true, then the high speedpositively charged alpha particles should pass through the gold foil without beingdeflected. Although most alpha particles passed through the gold foil, some weredeflected, and some even reflected back towards the source. Since opposite chargesattract, this means that there must be a positive charge present in the centre or nucleusof the atom. From these observations, Rutherford formulated his own nuclear model ofthe atom.


9/64



Rutherfords nuclear model of the atom (1911) the mass and the positive charge in the gold atoms is concentrated in a very small

region most of the atom is empty space this was dubbed the beehive model of the atom

Figure 1.3 Rutherfords beehive model

BOHR (1913) explained why the hydrogen atom does not collapse

explained the line spectrum of hydrogen

predicted undiscovered lines in the ultraviolet region of the Hydrogen spectrum

Figure 1.4 The Line Spectrum for Hydrogen (a line spectrum is created whenelectricity is applied to elemental gas. Light is then passed through the sample and aprism. The pattern revealed is called a line spectrum. Each element has its ownunique spectrum)


10/64



Bohrs theory states:

There are specific allowed energy levels (n) in which an electron can move. The energy of an electron in each level is quantized. The larger the nvalue, the more energy an electron possesses. Each energy level corresponds to an orbit, a circular path in which the electron can

move around the nucleus. An electron can travel in one of the allowed orbits without loss of energy An electron can jump from one allowed orbit to another. The jump cannot be

gradual it must occur all at once. Only certain energies can be absorbed or emitted as the electron changes orbits.

Two principles of Quantum Mechanics and some of Bohrs terminology can be used todevelop a straightforward view of the electron structure of the first 20 elements.

1. Electrons exist in energy levels in atoms. The number of the energy level, n, iscalled the principle quantum number.

2. Each energy level can hold up to 2n2electrons.

The 1stenergy level can hold 2The 2ndenergy level can hold 8The 3

rdcan hold 18

Energy Level Population = electrons in ground state

Na) 2e-)8e-)1e- P)2e-)8e-)5e-

CHADWICK (1932) discovered the neutron

Subatomic Particles

Atoms can be broken down into smaller subatomic particles: The table belowsummarizes some key information about the three subatomic particles (protons,neutrons and electrons).

Table 1.1: Subatomic particles

Subatomic particle Charge Location in atom Relative mass

Proton Positive (+1) nucleus 1

Neutron Neutral (0) nucleus 1

Electron Negative (-1) Orbits nucleus 1/1800


11/64



The number of protons in the nucleus is called the atomic number. This numberdetermines the identity of an atom. Atoms are electrically neutral; therefore the numberof protons in an atom must equal the number of electrons in an atom. For example,oxygen has an atomic number of 8 and has 8 protons in its nucleus and 8 electronsorbiting around the nucleus. Atoms also have a number called the mass number. The

mass number is the number of protons plus the number of neutrons an atom has. Letslook at the element oxygen for an example. Oxygen has a mass number of 16. Thenumber of neutrons = mass number atomic number or 16-8 =8. Thus oxygen has 8protons, 8 electrons and 8 neutrons.

Elements are organized on a chart called the periodic table. You should have receiveda periodic table at the start of this course. The elements are organized into verticalcolumns called groups, and horizontal rows called periods.

Figure 1.5: The Periodic Table of Elements

Lets examine how you can retrieve information about subatomic particles from theperiodic table.

Example 1: Determine the number of protons, neutrons, and electrons a potassiumelement has.


12/64



Solution 1:

The atomic number of potassium is 19, and since the atomic number is equal to thenumber of protons and electrons, a Potassium atom has 19 protons and 19 electrons.The number of neutrons is 40-19 or 21 neutrons.

Support Questions

(Reminder: these questions are not to be submitted but reinforce the material taughtand are strongly recommended DO NOT write in this book).

1. Reproduce this chart in your notes and your periodic table to complete the missinginformation:

ElementAtomicNumber

MassNumber

Number ofProtons

Number ofElectrons

Number ofNeutrons

Hydrogen

2

74

5

Carbon

7

16

9

10

Sodium

12

27

1415

Sulphur

17

40

19

20

19

K

39.10

Atomic number

Mass number


13/64



Structure of the Atom: Bohr-Rutherford Diagrams

Electrons move around the nucleus in circular paths called shellslike planets around theSun.

Electrons are spinning so fast in their orbits that they seem to form a solid shell aroundthe nucleus. Electrons cannot exist between these orbits, but can move up or downfrom one orbit to another. Electrons are more stable when they are at lower energy,closer to the nucleus. Each orbit has a maximum number of electrons that it can hold.

The number of electrons found in the orbits of the first twenty elements:

1storbit (Kshell) holds 2 electrons

2nd

orbit (Lshell) holds 8 electrons

3rdorbit (Mshell) holds 8electrons

4thorbit (Nshell) holds 2electron

Drawing Bohr-Rutherford Diagrams

To draw Bohr-Rutherford diagrams, use the following steps:

1. Using the periodic table, determine the number of protons, neutrons, and electronsfor the element.

2. Draw a circle to represent the nucleus of the atom. The number of protons and

neutrons are written inside this circle.3. Electrons are drawn in circular orbits around the nucleus. Remember that lower

orbits will fill up first!


14/64



Example 2: Carbon

Draw the Bohr diagram for an atom of the element Carbon:

Solution 2:

Following the steps outlined above:

1. Since carbon has a mass number of 12 and an atomic number of 6, it has 6 protons,6 electrons, and 6 neutrons.

2. Draw the nucleus and write the 6 protons (6p+) and 6 neutrons (6n0) inside.

6p+

6n0

3. Then add your electron shells. Remember the K shell holds 2 electrons, and the Lshell holds the remaining 4 electrons. The shells total 6 electrons.

6p+

6n0


15/64



Example 3: Nitrogen

Draw a Bohr Diagram for an atom of the element nitrogen

Solution 3:

14N

7

7p+

7n0

Often this can be written in a short-hand manner as follows:

N)2e-)5e-

This method will be used more frequently in this course.

Support Questions

2. Draw the Bohr diagrams for the first twenty elements on the periodic table (i.e.elements with atomic number 1-20). State any patterns you may observe based onthe locations of the elements on the periodic table.


16/64



Key Question #1

1. Summarize, using labelled diagrams in chart form, the evolution of atomic theory

from Dalton to the Rutherford model. (10 marks)

2. Write a short biography for oneof the following Canadians that made advances inatomic and molecular theory. (15 marks)

a. Harriet Brooksb. R.J. Leroyc. Richard Bader

Include the following;

i) A brief history of their life (birth death, family, where they lived, etc),ii) their research, and,iii) the significance/relevance of their research to the development of atomic and

molecular theory.


17/64

SCH4UGrade 12


Lesson 2 Quantum Mechanics


18/64



Lesson 2: Quantum Mechanics

The development of the Bohr model led to further study of the atomic model based ofwave properties of electrons, a branch of chemistry known as Quantum Mechanics. Itcan be a very complex area of chemistry, so this lesson will cover the fundamental

concepts.

What You Will Learn


describe the quantum mechanical model of the atom (e.g., orbitals, electronprobability density) and the contributions of individuals to this model (e.g., those ofPlanck, de Broglie, Einstein, Heisenberg, and Schrdinger);

list characteristics of the s, p, d, and f blocks of elements, and explain therelationship between position of elements in the periodic table, their properties, and

their electron configurations; write electron configurations for elements in the periodic table, using the Pauli

exclusion principle and Hunds rule;

Schrodinger (1924) postulated that sometimes electrons behave as particles, andsometimes like waves. Because of this we cannotmeasure both the position andvelocity of an electron at the same time. This exclusion is referred to as the PauliExclusion Principle. What this really means is that we cannot determine themomentum (velocity) and position (electron address) of an electron at the same time.The best we can do is to calculate the PROBABILITY of an electron being in a certainplace at a certain time. There are calculations that can be done to describe the region in

space where the electron is MOST LIKELY to be found, however we will not focus onthe mathematical in this course,

Regions where electrons are most likely to be found are called orbitals.

For every value of n,there are ntypes of orbitals and n2actual orbitals

1stenergy level has 1 type of orbital and (1

2) orbital.

2ndenergy level has 2 types of orbital and (22) 4 orbitals3rdenergy level has 3 types of orbital and (32) 9 orbitals

Each electron has a set of four numbers, called quantum numbersthat specify itcompletely; no two electrons in the same atom can have the same four. As mentionedabove, quantum numbers are determined by solving a mathematical wave equation,which we will not do in this course rather we will focus on what each quantum numberrepresents.


19/64



The Principle Quantum Number, n

The letter nwas used by Bohr to identify the orbits and energies. In other words, ntellsyou which of the "main" energy levels the electrons are in.

identifies the energyof an orbital values: 1,2,3 to infinity (4)

You can use the analogy of uneven steps on a staircase to describe the energy levels.If an electron falls from a higher energy level such as n=2 to n=1, the differencebetween the two is released as a photon of light.

Figure 2.1 Energy Levels (Source: Nelson Chemistry 12)

The Secondary Quantum Number, lThe secondary quantum number, lwas introduced by Arnold Somerfield in 1915 toexplain the line splitting of the hydrogen spectrum. The secondary quantum numberlidentifies the electron subshells or sublevels that part of a main energy level.

identifies the shapeof the orbital

values: 0 to (n-1) i.e: 0, 1, 2, 3, n-1s p d f


20/64



The number of sublevels equals the value of the principle quantum number. Forexample if n=3, then there are three sublevels, l= 0, 1, 2

Figure 2.2 Energy Sublevels (Source: Nelson Chemist ry 12)


21/64



Support Questions

3. What is the difference in the electron orbits proposed by Bohr and those of

Somerfield?

4. Recreate and complete this table in your own notes.

Primary energylevel

Principle quantumnumber

Possiblesecondaryquantum numbers

Number ofsublevels perprimary level

1

2

3

4

5

5. Write a general rule that can be used to predict all possible values from lowest tohighest, of the secondary quantum number for any value of the principal quantumnumber.

The Magnetic Quantum Number, ml

Recall from lesson one that a line spectrum is created by the passing of electricitythrough a gaseous element sample contained within a gas discharged tube. Light ispassed through the sample and then a prism to obtain the spectrum of the element.

If a gas discharge tube is placed near a strong magnet, some single lines split into newlines that were not initially present. This occurrence was discovered by Pieter Zeemanis 1897 and is called the Zeeman Effect. For example he observed a single linetransformed into three lines when the magnet was applied. This effect was explainedusing another quantum number, the magnetic quantum number, ml. The magneticquantum number explains that orbits could exist at various angles.

identifies direction of orientation(x, y, z) to an external magnetic field

values: -l0+ l i.e. l = 2 p orbital split into 3 planes (-1, 0, +1)


22/64



Figure 2.3: Spectra lines can be split in the presence of a magnetic field(Source: Nelson Chemistry 12)

Table 2.1: Values for Magnetic Quantum NumberValue of l 0 to n-1 Values of ml -lto +l

0 0

1 -1, 0, +1

2 -2, -1, 0, +1, +2

3 -3,-2,-1,0,+1,+2,+3


23/64



TheSpin Quantum Number ms

Atoms have their own magnetism which is unique to the magnetism of a group ofatoms. This magnetism is referred to as paramagnetism. It was suggested that eachelectron spins on its own axis. It can spin in either a clockwise or counterclockwise

direction.

identifies the spin directionof an electron

only 2 values: + (8) and - (9)

Table 2.2 Summary of Quantum NumbersPrincipleQuantumNumber, n

SecondaryQuantumNumber, l

MagneticQuantumNumber, ml

Spin QuantumNumber, ms

1 0 0 +1/2, -1/2

2 0

1

-1,0,+1 +1/2, -1/2

3 012

-2,-1,0,+1,+2 +1/2, -1/2

Electron Orbitals: Modernizing the Atom

Early Bohr atomic model theory were based on the idea of an electron travelling insome kind of path or orbit, however modern day theory is that of an electron orbital.Recall that an orbital is a volume of space where an electron is most likely to be found.

The first two quantum numbers, nand ldescribe electrons that have different energiesunder normal circumstances in multi-electron atoms. The other two quantum numbersapply to abnormal conditions in which a magnetic field is applied.

We will focus on the energy shells (n), or the primary quantum number, and subshells(l), or the second quantum number. We will now communicate the values of l withletters to denote the orbitals rather than numbers

Table 2.3 Values and Letters for the Secondary Quantum Number

Value of l 0 1 2 3

Letter of

designation

s p d f

Name ofdesignation

sharp principal diffuse fundamental

It is common for chemists to use the number for the main energy level and letter for theenergy sublevel. For example: a 1s orbital, a 2p orbital and a 3d orbital.


24/64



Table 2.4 Number of Orbitals for each Energy Sublevel

Value of l Sublevel symbol Number of orbitals

0 s 1

1 p 3

2 d 5

3 f 7

Energy Level Diagrams

Modern day atomic theory infers that electrons in an atom have different energies. Theatomic spectra indicate the energy sublevels, defined by a quantum number. An energylevel diagram shows the relative energies of the electrons in various orbitals.

Figure 2.4: Energy Level Diagram. The circles represent orbitals which contain twoelectrons. (Source: Nelson Chemistry 12)

Each electron orbital is represented by a circle and can contain two electrons. Theenergy of the electrons increases with increasing principle quantum number, n. For agiven value of n, the sublevels increase in energy, in order, s


25/64



Rules for Completing an Energy Level Diagram electron orbitals are filled from lowest energy up (Aufbau Principle) electrons go into the orbital corresponding to the lowest energy level available. an orbital can hold up to 2 electrons

electrons in the same sublevel will not pair until all the orbitals in the sublevelhave at least 1 electron (Hunds Rule)

Figure 2.5: Energy-level diagrams for lithium, carbon and fluor ine. (Source:Nelson Chemistry 12)

Aufbau Diagram for Filling Orbitals

Figure 2. Memory Tip for Filling Orbitals (SourceNelson Chemistry 12)

To use this diagram to help with drawing energy leveldiagrams, start at the bottom left side and add electronsin the order shown by the diagonal arrows. Work yourway up to the upper right-hand corner.


26/64



Example 1: Draw the energy level diagram for an oxygen atom

Solution 1: Oxygen has an atomic number of 8, so we must place these 8 electrons inenergy levels.

The first two electrons are placed in the 1s energy level

1s

The next two are placed in the 2s energy level

2s

The next three electrons are placed singly in each of the 2p orbitals

2p

Finally the last electron is paired up to fill one of 2p energy sublevels

2p

Now draw your energy diagram.

1s

2s

2p

Oxygen (O)


27/64



Example 2: Draw the energy level diagram for nitrogen anion.

Solution 2:

Nitrogen (N) 1s

2

2s

2

2p

3

Nitrogen anion (N

-3

) 1s

2

2s

2

2p

6

Use the same procedure except remember that a nitrogen anion will have 7 + 3additional electrons.

Electron Configurations

Electron Configurations provide the same information as energy level diagrams, but in amore concise format. An electron configuration depicts the type of electron inincreasing energy level. For example the electron configuration for an oxygen atom is1s22s22p2.

Example 3: Write the electron configuration for the following elements:a) sodium atomb) fluorine ion

Solution 3:

Sodium Atom1s22s22p63s1

Fluorine ion1s22s22p6


28/64



Shorthand Form of Electron Configuration

The shorthand form abbreviates writing the electron configuration by using the nearestNobel Gas in its written format. For example:

Cl: 1s2

2s2

2p6

3s2

3p5

becomes [Ne]3s2

3p5

Sn:1s22s22p63s24s23d104p65s24d105p2 becomes [Kr]5s24d105p2

Support Questions

6. Draw an energy level diagram for the followinga) Carbon

b) Cl-1

7. Write the complete ground state electron configurations for the following:a) lithiumb) oxygenc) calciumd) titaniume) rubidiumf) leadg) erbium

8. Write the abbreviated (shorthand method) ground state electron configurations forthe following:a) heliumb) nitrogenc) chlorined) irone) zincf) bariumg) polonium


29/64



Key Question #2

1. Calculate the maximum number of electrons with principal quantum number;

(8 marks)a) 1b) 2c) 3d) 4

2. Show the permissible values of land mfor; (8 marks)a) n=1b) n=2c) n=3

3. Draw energy level diagrams for beryllium, magnesium and calcium ions. What is thesimilarity in these diagrams? (10 marks 3 each for diagram 1 for statement)

4. The sodium ion and the neon atom are isoelectronic (have the same electronconfiguration). (8 marks 4 marks each)a) Write the electron configuration for the sodium ion and neon atom.b) Describe and explain the similarities and differences in the properties of these

two chemical entities.

5. One important application of Quantum Mechanics is laser technology. Construct aninformation pamphlet including;

uses of laser technology

how it applies quantum mechanics

advantage/disadvantages of laser technology

NOTE: Dont forget to include the references for your information (Wikipedia orsearch engine references are not acceptable look for professional sites or journalarticles) (25 marks 20 for content, references and design, 5 for grammar/spelling)


30/64

SCH4UGrade 12


Lesson 3 Chemical Bonding


31/64



Lesson 3: Chemical Bonding

Chemical compounds are formed by the joining of two or more atoms. You will reviewsome of the basics of chemical bonding that you learned in Grade 11 Chemistry;however you will also extend your knowledge of chemical bonding using Quantum

Mechanics.

What You Will Learn


predict molecular shape for simple molecules and ions, using the VSEPR model;

use appropriate scientific vocabulary to communicate ideas related to structure andbonding (e.g., orbital, absorption spectrum, quantum, photon, dipole);

predict the polarity of various substances, using molecular shape and theelectronegativity values of the elements of the substances

Ionic and Covalent Bonding: The Octet Rule

Atomsform bonds to become more chemically stable. The most chemically stableelements on the periodic table are the noble gases. We know this because they areextremely unreactive and tend not to form compounds.

According to the octet rule, atoms bond in order to achieve the same electronconfiguration as a noble gas. This rule is called the octet rule because all the noblegases (except helium) have eight valence electrons.

Generally, when an atom tends to gain or lose electrons, an atom will have the sameelectron configuration (arrangement of electrons) as a noble gas. Atoms that haveidentical electron configurations are said to be isoelectronic.

An atom that has lost an electron to become stable is referred to as an ion. If an atomloses electrons, it becomes positively charged and is referred to as a cation.


32/64



Lewis Structures

Another way to draw atoms and ions is to use Lewis structures. Lewis structures depictonly the valence electrons an element has. The figure below shows the Lewisstructures for some common elements.

To draw Lewis structures, use the following steps:

1. Write the atomic symbol of the element. This will represent the atomic nucleus. It isconsidered to have 4 sides.

2. Place the valence electrons, one per side, and then pair up if necessary.

Example 1: Chlorine - Draw the Lewis structure for a chlorine atom

Solution 1:

Cl

x x

x x

xxx

The table below shows the Lewis structures for the first twenty elements:

Figure 3.1: Lewis Structures for the first twenty elements

Notice that the group number (vertical column) indicates how many valence electronsan atom has. Thus, lithium is in group oneand has one valence electron, whilefluorine is in group sevenand has seven valence electrons.


33/64



Ions can also be represented by Lewis symbols. The Lewis symbol is enclosed by abracket as with the Bohr diagram:

Example 2: Lewis Chlorine Ion - Draw the Lewis structure for a chlorine ion

Solution 2:

Cl

x x

x x

xx

xx

-1

Lewis Structures and Quantum Mechanics

There is unity between the atomic and bonding theories you learned in Grade 11

chemistry. The octet rule comes from the maximum of two electrons in the s orbital andthe six electrons in the p orbitals.

Figure 3.2 - Correlation between Atomic and Bonding Theories(Source: Nelson Chemistry 12)


34/64



Support Questions

9. Recreate and complete the following table in your own notes;

Element AtomicSymbol

ElectronConfiguration

Number ofvalence

electrons

Valence

Oxygen

Chlorine

Sodium

Phosphorus

10. Identify the elements and write the Lewis structure for the following electronconfigurations:

a) 1s22s22p4 b) 1s22s22p63s23p3c) [Ar]4s23d104p5 d) [Kr]5s1

Electronegativity

Electronegativity is a measure of an atoms ability to attract electrons in a chemicalbond. It is a periodic property.

Predicting Bond Type Using Electronegativi tyYou can use the difference between electronegativities of two atoms to determine if the

bond formed between the two atoms is ionic or covalent, or polar covalent. Thesymbol EN stands for the difference in electronegativity between two values

3.3 1.7 0.5 0

MOSTLY IONIC POLAR COVALENT COVALENT


35/64



Table 3.1 Electronegativities of Various Elements

Element Electronegativity

H 2.1

Metals Li 1.0

Be 1.5

Na 0.9Mg 1.2

K 0.8

Ca 1.0

Nonmetals C 2.5

N 3.0

O 3.5

F 4.0

P 2.1

S 2.5

Cl 3.0

Example 3: Determine if the elements below would form ionic, covalent or polarcovalent bonds:

Solution 3:

Substance EN Element 1 EN Element 2 EN Ionic orCovalent?

KF0.8 4.0 3.2 ionic

O2 3.5 3.5 0 covalent

HCl2.1 3.0 0.9 Polar covalent

Bonds within Molecules: Intramolecular Bonds

When an ion loses or gains an electron, it is forming an ion. An atom always loses orgains an electron in conjunction with another atom forming a chemical bond. There arethree main types of intramolecular bondswe will explore in this lesson, ionic, covalentand polar covalent.

Ionic Bonding

electrons are transferred from one atom to another

usually occurs between metals and non-metals.


36/64



Example 4: Potassium and Fluoride - Using Lewis structures draw the formation of abond between potassium and fluorine.

Solution 4:

Step 1: Write out the correct number of valence electrons for each atom.

K + F

x x

x x

xxx

Step 2: Analyze the valence electrons. Potassium will lose its one valence electron toobtain a full octet and fluorine will gain one.

K F

x x

x x

xxx

Potassium loses its one valence electron, becoming a cation with a charge of +1, andfluorine gains one valence electron, becoming an anion with a charge of -1.

KF

x x

x x

xxx

+1 -1

Notice how the ions are written with square brackets and the overall charge is indicatedin the upper right hand corner.

When naming ionic compound, we name the cation first, then the anion. The ending ofthe anion is replaced with ide.

Thus, the name of this compound is potassium fluoride. The positive potassiumcation (+) is attracted to the negative fluorine anion (-). This is what forms the ionicbond.


37/64



Example 5: Magnesium and Nitrogen - Using Lewis structures draw the formation of abond between magnesium and nitrogen.

Solution 5:

Again, first draw the Lewis structures

Mg + N xx

x

x

x

In this case, magnesium will lose two electrons and nitrogen will gain. However, sincethis does not balance we need three magnesium atoms and two nitrogen atoms tomake this balance. The transfer of the electrons from magnesium to nitrogen is shownbelow.

MgN x

x

x

x

x

Mg N xx

x

x

x

Mg

We can then summarize this with brackets. The upper right hand corner indicates thecharge of the ion and the lower right hand corner indicates the number of ions.


38/64



This compound is named magnesium nitride.

Support Questions

11. Recreate and complete the following table. Use Lewis structures to depict atoms.

Bondformation

Name ofcompound

Chemicalformula

Anion Cation

a) lithiumand fluorine

b) calciumandphosphorus

Covalent Bonding

Covalent bondsform when two or more non-metals shareone or more pairs ofelectrons. As a result of forming covalent bonds through sharing electrons, the atomsend up with a stable electron arrangement in their outer orbit similar to that of a noblegas.


39/64



Example 6: Using Bohr diagrams, draw the formation of Chlorine gas (Cl2)

Solution 6:

Chlorine gas is a molecule that consists of two chlorine atoms held together with a

covalent bond.

Each chlorine atom has 7 electrons in its outer orbit and needs to gain 1 electron tobecome stable.

Two chlorine atoms share a pair of electrons to form a covalent bond. Each chlorineatom now has 8 electrons in its outer orbit (which forms a stable octet).

Example 7: Draw the formation of chlorine gas using Lewis structures

Solution 7:

Other examples of covalently bonded molecules include:

Methane (CH4) Water (H2O) Ammonia (NH3)


40/64



Drawing Lewis Structures for Polyatomic Ions

When an ion is composed of more that one atom, it is termed a polyatomic ion.Polyatomic ionsare groups of atoms that tend to stay together and carry an overallionic charge.

Table 3.2 Common Polyatomic Ions and Their Ionic Charges

Name of polyatomic ion Ion formula Ionic charge

nitrate NO3

1

hydroxide OH 1

bicarbonate HCO3 1

chlorate ClO3 1

carbonate CO3 2

sulfate SO4 2

phosphate PO4 3

Drawing Lewis Structures for polyatomic Ions can be tricky. There are some general rulesfor drawing these structures. We will use the nitrate ion (NO3

-1) as our example.

Example 8: Draw the Lewis structure for nitrate, NO3-1

Solution 8:

Step 1: Arrange atoms symmetrically around the central atom (usually listed first, andsingular although not usually hydrogen or oxygen). In this case, nitrogen is our centralatom.

Step 2: Total valence electrons for all atoms. Add one electron for each -1 charge, subtractfor each +1 charge. For nitrate there are 5 valence electrons for nitrogen + 1 electron for -1charge and 18 electrons for the oxygen atoms. This totals 24 electrons.


41/64



Step 3:Place a bonding pair of electrons between the central atom and each ofThe surrounding atoms

Step 4:Complete the octets of the surrounding atoms

Step 5:If the central atom does not have an octet, move lone pairs from the surroundingatoms to form double or triple bonds.

Step 6:Draw the Lewis structure and enclose polyatomic ions within square bracketsshowing the ion charge.


42/64



Support Questions

12. Draw Lewis Structures for each of the following molecules or polyatomic ions.

a) ClO4-1 b) CN-1 c)HCO3

-2

Valence Bond Theory

The valence bond theoryexplains why and how electrons are shared between atoms.The VB theory postulates how individual atoms, each with its own orbitals and electronscome together and form covalent bonds in a molecule.

A bond between two atoms is formed when a pair of electrons is shared by two

overlapping orbitals,

For example, in a hydrogen molecule, the two 1s orbitals from each H atoms overlapand share electrons.

H H H2(g)

1s1 1s1 1s2

In water, there is overlap of the 1s orbital of hydrogen and the 2 p orbital of oxygen.

Figure 3.3: Atoms overlapping orbitals(Source: Nelson Chemistry 12)


43/64



According to Valence Bond Theory:

A half-filled orbital in one atom can overlap with another half-filled orbital of asecond atom to form a new, bonding orbital

The new bonding orbital contains two electrons of opposite spin

When atom bond, they arrange in space for maximum overlap

Hybrid Orbitals

You may have noticed in the figures above that when the orbital bond and overlap theychange shape. They form new orbitals called hybrid orbitals. Hybridizationis theprocess of combining two or more atomic orbitals to create new orbitals, called hybridsthat will fulfill the geometric demands of the system.

When an atom hybridizes, it will restructure its original set of s and p atomic orbitals intoa new set of hybrid orbitals. The process of hybridization is driven by the needs of theatoms to produce specific geometric patterns.

Example 9: What are the bonding orbitals of the BF3molecule?

Solution 9:

Boron has a valence of +3 and Fluorine is -1.

Write the electron configuration of boron. Focus on the valence electron orbitals. In thiscase the 2s has two electrons, and p orbital has one.

B: 1s22s22p1

2s2 2px1 2py 2pz

In this case, one of the 2s2 electrons can be promoted to an empty p orbital

2s1 2px1 2py 2pz


44/64



By promoting an s electron to an empty p orbital, there are one s orbitals and two porbitals available for hybridization. A total of three identical sp2(i.e one s, two porbitals)

Therefore the hybridization is sp2

More forms of hybridization are summarized in the table following.

Table 3.3: Forms of hybridization


45/64



Double and Triple Bonds

Two kinds of orbital overlap are possible.

a) The end-to-end overlap of s orbitals, p orbitals, hybrid orbitals, or some pair of

these orbitals. These form sigma bonds(

)b) Side by sideoverlap forms pi () bond

A double bond contains a sigma and a pi bond

Consider the molecule C2H4

In this case we only have partial hybridization. We have three new sp2 orbital and onenormal 2pz orbital.


46/64



This is shown in figure 3.4 following:

Figure 3.4: Partial hybrid ization(Source: Nelson Chemistry 12)

Partial hybridization can also be used to explain how triple bonds formConsider ethyne gas C2H2

In this case, two sp hybrid orbitals are formed for each carbon, and two unhybridizedorbitals are formed.


47/64



Figure 3.5 Formation of a Triple Bond(Source: Nelson Chemistry 12)

Support Questions

13. What atomic orbital or orbitals are available for bonding for each of the followingatoms?a) H

b) Fc) Sd) Br

14. Provide ground state and promoted state electron configurations for each of thefollowing atoms and indicate the type of hybridization involved when each atomforms a compound;a) carbon in CH4b) boron in BH3c) beryllium in BeH2

15. When are bonds formed?

16. Provide an explanation for bonding in each of the following;a) C2Cl4b) C2F2


48/64



The Valence Shell Electron Pair Repulsion (VSEPR) model:

One of the more important properties of any molecule is its shape. It is very important toknow the shape of a molecule if one is to understand its reactions. It is also desirable tohave a simple method to predictthe geometries of compounds.

The underlying assumptions made by the VSEPR method are the following.

Atoms in a molecule are bound together by electron pairs. These are calledbonding pairs. More than one set of bonding pairs of electrons may bind any twoatoms together (multiple bonding).

Some atoms in a molecule may also possess pairs of electrons not involved inbonding. These are called lone pairs.

The bonding pairs and lone pairs around any particular atom in a molecule adoptpositions in which their mutual interactions are minimized. The logic here is simple.Electron pairs are negatively charged and will getas far apart f rom each other as

possible. Lone pairs occupy more space than bonding electron pairs. Double bonds occupy more space than single bonds.

Using VESPR Theory

Example 10: What is the shape of the BeH2?

Solution 10:

Step 1: Draw the Lewis Structure

Be HH

If we consider the central atom, Be, it has 2 bonding pairs and no lone pairs. It has thegeneral formula AX2and is considered linear.


49/64



Table 3.4 VESPR predicts Molecular Shapes

Example Predict the shape of a sulphate ion, SO42-

The central atom sulphur, has 6 valence electrons plus the two additional electronstotalling eight. This gives 4 bonding pairs and 0 lone pairs. Thus the general formula is

AX4, and the molecule is tetrahedral.

Support Questions

17. Use Lewis Structures and VESPR theory to predict the shapes of the followingmolecules:a) CO2b) HCNc) BF3d) SiCl4e) CH4f) OCl2


50/64



Polar Covalent Bonds

A polar covalent intramolecular bond is formed when there is unequal sharing betweenvalence electrons resulting in dipoles (ends) that are slightly positive or slightly negative

Example 11: Draw the polar bond formed in a molecule of carbon tetrachloride (CCl4)between the elements carbon and chlorine

Solution 11:

Referring to table 3.3 above, chlorine has an electronegativity of 3.0 and carbon 2.5.Thus the difference, En is 0.5, and the intramolecular bond is considered polarcovalent. Since the chlorine has a higher electronegativity, the electrons are pulledcloser to the chlorines atomic nucleus. This makes the chlorine ends or dipoles slightlynegative (-). Conversely, the electrons are farther away from carbons nucleus, makingit slightly positive (+).

Molecular Polarity

Covalent bonds can be polar or nonpolar. However the polarity of a molecule as awhole is dependent on bond polarity andmolecular shape. Symmetrical molecules produce nonpolar molecules (whether the bonds are

polar or not). Asymmetrical molecules produce nonpolar molecules if the bonds are all

nonpolar

Example 12:Predict the Polarity of the ammonia, NH3 molecule including yourreasoning.

Solution 12:

First, draw the Lewis structure


51/64



Next, check the electronegativities. Nitrogen has an electronegativity of 3 and hydrogenof 2.1. This difference is 0.9, making the hydrogen end positive and the nitrogen endnegative.

Now we must also consider symmetry. Because ammonia is asymmetrical and

contains polar bonds, the molecule is polar.

Support Questions

18. Recreate and complete the table below. If the molecule is covalent, indicate if it ispolar covalent or not.

Substance EN Element 1 EN Element 2

EN Ionic orCovalent?

NaCl

Cl2

HF

19. Predict the polarity of following molecules.a) BF3b) OF2c) CI2d) PCl3


52/64



Key Question #3

1. Recreate and complete the following table; (24 marks)

Molecule orCompoundion

Lewisstructure

Name ofShape

BondPolarity

MolecularPolarity

HClCH4CH3ClCO2H2ONH3

2. Identify the types of hybrid orbitals found in molecules of the following substances;(8 marks 2 marks each)a) CCl4(l)b) BH3(g)c) BeI2(s)d) SiH4(g)

3. The polarity of a molecule is determined by bond polarity and molecular shape.

a) Compare the polarity of the bonds N-Cl and C-Cl. (2 marks)b) Predict whether the molecules, NH3(l)and CCl4(l)are polar or non-polar. Explain

your predictions. (4 marks)

4. Indicate the number of sigma and pi bonds in each of the following molecules:(8 marks 2 marks each)

a) H2Ob) C2H2c) C2H4d) C2H6


53/64

SCH4UGrade 12


Lesson 4 Intermolecular Forces


54/64



Lesson 4: Intermolecular Forces

In this lesson, we will examine the other types of bonding that have not yet beencovered in the course. We will first examine the forces that exist between molecules,and how such forces can lead to the formation of large aggregates such as crystals.

What you will learn


explain how the properties of a solid or liquid (e.g., hardness, electrical conductivity,surface tension) depend on the nature of the particles present and the types offorces between them (e.g., covalent bonds, Van der Waals forces, dipole forces, andmetallic bonds)

predict the type of solid (ionic, molecular, covalent network, or metallic) formed by asubstance, and describe its properties;

conduct experiments to observe and analyse the physical properties of differentsubstances, and to determine the type of bonding present

describe some specialized new materials that have been created on the basis of thefindings of research on the structure of matter, chemical bonding, and otherproperties of matter (e.g., bulletproof fabric, superconductors, superglue);

Bonds between Molecules: Intermolecular Bonds

Intermolecular bonds are the chemical bonds betweenmolecules. These bondsdetermine the physical state of molecular substances. These bonds are broken as a

substance undergoes a change of state. Intermolecular forces are much weaker thancovalent bonds.

There are generally three types of intermolecular bonds: London forces, dipole-dipoleforces, and hydrogen bonds. These intermolecular forces are collectively called Vander Waals forces. These are summarized in table 4.1 below.

Table 4.1- Vander Waals forces

Force Description Example

London forces Hold covalent molecules together. Very weakforces of attraction. Momentary dipoles are

created by the electrons contained within thecompound, which are constantly in motion

All moleculesMethane gas

(CH4)

Dipole-Dipole Hold polar covalent molecules together. Theseforces are stronger than London forces.

HydrogenChloride(HCl)

Hydrogenbonding

Formed between the electropositive Hydrogendipole and an electronegative dipole ofOxygen, Chlorine, or Fluorine.

Pure distilledwater


55/64



Using Vander Waals forces to predict Boiling Points

As the number of electrons in the molecule increases, the boiling point increases aswell. There are some general guidelines for predicting boiling points;

The more polar a molecule is, the higher the boiling point The greater the number of electrons, the stronger the London force and

therefore, higher boiling point Isoelectronic molecules have the same London force

Support Questions

20. Determine the type of Vanderwaals forces that would occur between the followingmolecules;

a) water, H2Ob) butane C4H10c) hydrogen chloride, HCl

21. Which of the following pure substances has a stronger dipole-dipole force than theother? Discuss your reasons for your conclusion.

a) hydrogen chloride or hydrogen fluorideb) chloromethane or iodomethanec) nitrogen tribromide or ammonia

d) water or hydrogen sulfide

The Structure and Properties of Solids

Although solids have many commonalities, the physical properties of solids vary greatlyin such physical properties such as hardness, melting point, mechanical characteristicsand conductivity.

The structure and properties of solids are related to the forces between the particles.There are four different categories of solids, summarized in Table 4.2 following.

Table 4.2: Categories of SolidsClass of Substance Elements combined Examples

Ionic Metal + nonmetal NaCl(s),CaCO3(s)

Metallic Metal (s) Cu(s), CuZn3(s)

Molecular Non-metal(s) I2(s), H2O(s), CO2(s)

Covalent network Metalloids/carbon C(s), SiC(s), SiO2(s)


56/64



Ionic Crystals

Earlier in this unit, you learned that an ionic bond is formed when there is a transfer ofelectrons from a metal atom to a non-metal atom. However, the transfer of electrons isnot what forms the bond. The ions that are formed as a result of electron transfer allow

the resulting ions to form arrangements of in a definite crystal pattern called a crystallattice.

The crystal lattice for sodium chloride (NaCl) is depicted below in figure 4.1

Figure 4.1 Sodium Chloride Crystal Lattice

There is a huge variety of crystal shapes that we willnot discuss here. Ionic compounds are relatively hardbut brittle substances. They conduct electricity inliquid but not in solid state. They also have high

melting points.

Metallic Crystals

You are likely aware of the properties of metals. They are shiny, silvery solids that aregood conductors of heat and electricity, their hardness ranges from hard to soft. The

melting points also range from high to low. Metals have acontinuous compact crystal structure.

The properties of metal are the result of bonding betweenfixed positive nuclei and mobile loose valence electrons. Inmore simplistic terms, the valence electrons float around thenetwork between positive centers. The motility of theelectrons explains why metals are such good conductors ofheat and electricity.

Figure 4.2 Metallic Bonding A sea of negative electrons floats freely amongstpositive nuclei.

Molecular Crystals

Molecular solids may be elements such as iodine and sulphur or compounds such asice or carbon dioxide. They are crystals that have relatively low melting points, they arenot hard, and are not conductors of electricity.

They are arranged by neutral molecules with weak intermolecular forces.


57/64



Covalent Crystals

Massive aggregates of atoms, as distinct from ions, can exist where neighbouringatoms share a pair of electrons to form multiple, single, covalent bonds. This type ofcrystal is usually found in the elements carbon and silicon.

Carbon and Silicon both have four valence electrons. Diamond and Graphite are bothmade up of carbon. Diamond is a huge network of C-C linkages where all four of thecarbon bonds are equal in strength and the whole structure is enormously strong.Graphite, on the other hand, is a carbon crystal inwhich each carbon atom uses only three of its fouravailable valence electrons to bond with adjacentcarbons atoms. This results in layers of carbonthat are not bonded. The unused bondedelectrons can move through layers, explaining whygraphite can conduct electricity, whereas diamond,

does not. The layers of carbon in graphite are heldtogether by Vander Waals forces. The layers ingraphite can slip over one another easily, andalthough graphite is strong in two dimensions, it isweak in third. This makes graphite useful forproducts such as pencils, golf clubs, and as alubricant. Diamond on the other hand is used to

cut glass, and is considered precious due itshardness. Figure 4.3 - Diamond

Covalent crystals tend to exceptionally hard,with very high melting points. They areinvariably insoluble in water, since there areno ions to attract the polar water molecule.

Figure 4.4 Graphite


58/64



Key Question #4

1. Use the Internet to write a short report on X-ray crystallography, describing how X-

rays can be used to give information on structures of solids at the atomic level.(10 marks)

2. How does the melting point relate to the type of particle and forces present?(2 marks)

3. Identify the main type of bonding and the type of solid for each of the following;(6 marks)

a) SiO2 b) CH4 c) Cr d) Na2Se) C f) CaO

4. Water beads on the surface of a freshly waxed car hood. Use your knowledge ofintermolecular forces to explain this observation. (5 marks)

5. All molecular compounds may have London, dipole-dipole, and hydrogen-bondingintermolecular forces, affecting their physical and chemical properties. Indicatewhich intermolecular forces contribute to the attraction between molecules in each ofthe following classes of organic compounds: (8 marks)

a) pentane, C2H5 b) 2-propanol, CH3CHOHc) acetic acid, CH3CHOHCH d) ethybenzoate, C6H5COOCH2CH5

e) dimethylether, CH3OCH3 f) ethylamide, CH3CONH2g) diamond, C h) calcium carbonate, CaCO3

6. Use the theory of intermolecular bonding to explain the sequence of boiling points inthe following alkyl bromides: CH3Br(g)(4

oC), C2H5Br(l)(38oC), and C3H7Br(l)(71

oC).(6 marks 3 marks each)

7. Compare the particles and forces in the following pairs of solids; (4 marks)

a) metallic and covalent b) molecular and ionic


59/64

SCH4UGrade 12


Support Question Answers


60/64

SCH4U Chemistry Unit 1 Support Question Answers


Answer to Support Questions

1.

ElementAtomicNumber

MassNumber

Number ofProtons

Number ofElectrons

Number ofNeutrons

Hydrogen 1 1 1 1 0Helium 2 4 2 2 2

Lithium 3 7 3 3 4

Beryllium 4 9 4 4 5

Boron 5 11 5 5 6

Carbon 6 12 6 6 6

Nitrogen 7 14 7 7 7

Oxygen 8 16 8 8 8

Fluorine 9 19 9 9 10

Neon 10 20 10 10 10

Sodium 11 23 11 11 12

Magnesium 12 24 12 12 12Aluminum 13 27 13 13 14

Silcon 14 28 14 14 14

Phosphorus 15 31 15 15 16

Sulfur 16 32 16 16 16

Chlorine 17 35 17 17 18

Argon 18 40 18 18 22

Potassium 19 39 19 19 21

Calcium 20 40 20 20 20

2.H)1e- IIA IIIA IVA VA VIA VIIA He)2 e-

Li)2 e-)1 e

- Be)2 e

-)2 e

- B)2 e

-)3 e

- C)2 e

-)4 e

- N)2 e

-)5 e

- O)2 e

-)6 e

- F)2 e

-)7 e

- Ne)2 e

-)8 e

-

Na)2 e-)8 e

-

)1 e-

Mg)2 e-)8 e

-

)2 e-

Al)2 e-)8 e

-)3

e-

Si)2 e-)8 e

-

)4 e-

P)2 e-)8 e

-

)5 e-

S)2 e-)8 e

-

)6 e-

Cl)2 e-)8 e

-

)7 e-

Ar)2 e-)8 e

-

)8 e-

K)2 e-)8 e

-

)8 e-) 1 e

-

Ca)2 e-)8 e

-

)8 e-)2 e

-

Trend: The group number indicates the number of valence electrons, the periodindicated the number of shells the atom has.

3. Bohr- orbits are 2D, fixed distance from nucleus, circular or elliptical path, 2n2electrons per orbit, Sommerfield- 3D region in space, variable distance from nucleus,

no path, 2 electrons per orbital


61/64



4.

Primary energylevel

Principle quantumnumber

Possiblesecondaryquantum numbers

Number ofsublevels perprimary level

1 1 0 1

2 2 0,1 23 3 0,1,2 3

4 4 0,1,2,3 4

5 5 0,1,2,3,4 5

5. l=n-1

6. Draw an energy level diagram for the following;

a) Carbon

1s

2s

2p

b) Cl-1

1s

2s

2p

3s

3p

7. Write the complete ground state electron configurations for the following;

a) 1s22s1b) 1s22s22p4c) 1s22s22p63s23p64s2d) 1s22s22p63s23p64s23d2e) 1s22s22p63s23p64s23d104p65s1f) 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2g) 1s22s22p63s23p64s23d104p65s24d105p66s24f12


62/64



8. Write the abbreviated (shorthand method) ground state electron configurations forthe following;

a) [He] b) [He]2s22p

3

c) [Ne]3s23p5 d) [Ar]4s23d6

e) [Ar]4s

2

3d

10

f) [Xe]6s

2

g) [Xe]6s24f145d106p4

9.

Element AtomicSymbol

ElectronConfiguration

#of valenceelectrons

Valence

Oxygen O 1s 2s 2p 6 -2

Chlorine Cl 1s 2s 2p 3s 3p 7 -1

Sodium Na 1s 2s 2p 3s 1 +1

Phosphorus P 1s 2s 2p 3s 3p 5 -3

10. Identify the elements and write the Lewis structure for the following electronconfigurations;

a) 1s22s22p4 b) 1s22s22p63s23p3 c) [Ar]4s23d104p5 d) [Kr]5s1

a) oxygen b) phosphorus c) bromine d) rubidium

O

x x

x x x x

P

x x

Br

x x

Rb

11.

Bond formation Name ofcompound

Chemicalformula

Anion Cation

a) lithium andfluorine

LiF

x x

x x

xxx

+1 -1 Lithiumfluoride

LiF fluorine lithium

b) calciumand

phosphorus

Calciumphosphide

Ca3P2 phosphorus calcium


63/64



12. Draw Lewis Structures for each of the following molecules or polyatomic ions.

a) ClO4-1 b) CN-1 c)HCO3

-2

Cl O

O

O

O

-1

N C

-1

HO C

O

O -2

13. What atomic orbital or orbitals are available for bonding for each of the followingatoms?

a) H -1s b) F -2s, 2p c) S- 3s,3p d) Br 4p

14. Provide ground state and promoted state electron configurations for each of the

following atoms and indicate the type of hybridization involved when each atomforms a compound;

a) carbon in CH4, 1s22s22p2promoted to1s22s12px

12py12pz

1(sp3)b) boron in BH3 1s

22s22p1promoted to 1s22s12px12py

1(sp2)c) beryllium in BeH21s

22s2promoted to 1s22s12px1(sp)

15. When are bonds formed?

Pi bonds are created when there is overlap sideways of non hybridized orbitals,usually p orbitals

16. a) C2Cl4 - sp2 There will be a double bond between the carbon atoms, one pi bondsand one sigma. The two half-filled p orbitals of the adjacent atom overlap sideways.

b) C2F2 - sp The sigma bonds use sp orbital, the two pairs of half-filled p of theadjacent carbon overlap sideways.

17. a) CO2(g) - AX2- Linearb) HCN - AX2Linearc) BF3 - AX3- Trigonal planard) SiCl4- AX4-tetrahedral

e) CH4 - AX4- tetrahedralf) OCl2 - AX2E2 V-shaped or bent


64/64


18.

Substance EN Element 1 EN Element 2 EN Ionic orCovalent?

NaCl Na = 0.9 Cl = 3.0 2.1 Ionic

Cl2 Cl = 3.0 Cl = 3.0 0 CovalentHF H = 2.1 F = 4.0 1.9 Polar covalent

19. Predict the polarity of following molecules.

a) BF3 -polarb) OF2 non-polarc) CI2 non-polard) PCl3-polar

20. Determine the type of Vanderwaals forces that would occur between the following

molecules;

a) water, H2O hydrogen bondingb) butane C4H10London forcesc) hydrogen chloride, HCl dipole-dipole

21. Which of the following pure substances has a stronger dipole-dipole force than theother? Provide your reasoning.

a) hydrogen chloride or hydrogen fluoride: HF greater electronegativity differencebetween atoms

b) chloromethane or iodomethane: chloromethane - greater electronegativitydifference between atomsc) nitrogen tribromide or ammonia: ammonia - greater electronegativity difference

between atomsd) water or hydrogen sulphide: water - greater electronegativity difference between

atoms

SCH4U - Unit 1 - Version C.pdf

Documents

Transcript of SCH4U - Unit 1 - Version C.pdf