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SCH4U

Grade 12 University Chemistry

Version C

SCH4U – Chemistry Introduction

Copyright © 2008, Durham Continuing Education of 64

Introduction Welcome to the Grade 12 University Chemistry Course, SCH4U. This full-credit course is part of the Ontario Secondary School curriculum. Prerequisite – Grade 11 University Chemistry. This course enables students to deepen their understanding of chemistry through the study of organic chemistry, energy changes and rates of reaction, chemical systems and equilibrium, electrochemistry, and atomic and molecular structure. Students will further develop problem solving and laboratory skills as they investigate chemical processes, 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 evaluating the 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 assignments to be completed. In this course you must complete ALL assignments. Be sure to read through all the material presented in each lesson before trying to complete the assignments. Important Symbols

Questions with this symbol are Key Questions. They give you an opportunity 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 to be submitted to the marker, but they will help you understand the course material more fully. Answers for support questions are included at the end of each unit. Refer to these for suggestions of how to properly structure the answers to questions.

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



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 exam after Unit 3.

Complete and pass a final examination.

After you submit lessons for evaluation, begin work on your next lesson(s) right away! 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 lined paper, 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 each unit.

Term



Table of Contents

Unit 1 – Structure and Properties Lesson 1 Atomic Theories Lesson 2 Quantum Mechanics Lesson 3 Chemical Bonding Lesson 4 Intermolecular Forces Unit 2 – Organic Chemistry Lesson 5 Hydrocarbons Lesson 6 Functional Groups Lesson 7 Types of Organic Reactions Lesson 8 Polymers Unit 3 – Rates of Reactions Lesson 9 Thermochemistry Lesson 10 Enthalpies of Reactions Lesson 11 Energy Options Lesson 12 Chemical Kenetics Unit 4 – Electrochemistry Lesson 13 Oxidation and Reduction Reactions Lesson 14 The Activity Series of Metals Lesson 15 Galvanic Cells Lesson 16 Electrolytic Cells Unit 5 – Chemical Systems and Equilibrium Lesson 17 Introducing Equilibrium Lesson 18 The Equilibrium Constant Lesson 19 Acid and Bases Equilibrium Lesson 20 Solubility Equilibriums

SCH4U


Lesson 1 – Atomic Properties

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 the first computer looked like. It may seem odd and funny to look back at such pictures, but most 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 computer technologies. There have been many modifications along the way to current modern atomic theory. In this unit you will learn about these models and modification. You will also 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 how

types of chemical bonding account for the properties of ionic, molecular, covalent network, and metallic substances;

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

• describe products and technologies whose development has depended on understanding molecular structure, and technologies that have advanced the knowledge of atomic and molecular theory.



Lesson 1: Atomic Theories The atom is the smallest unit of an element that retains the chemical properties of that element. In this lesson, you will learn about the various theories that led up to the current model of the atom. What You Will Learn After completing this lesson, you will; • explain the experimental observations and inferences made by Rutherford and Bohr

in developing the planetary model of the hydrogen atom; • describe some applications of principles relating to atomic and molecular structure in

analytical chemistry and medical diagnosis (e.g., infrared spectroscopy, X-ray crystallography, nuclear medicine, medical applications of spectroscopy);

• describe advances in Canadian research on atomic and molecular theory The Development of the Atomic Model Matter is anything that has mass and takes up space. This means everything around you (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 proposed that the atom was the smallest particle that could not be subdivided. As experimentation and the scientific method gained importance, the model of the atom began to evolve. Let’s take a look at the scientists involved in the development of the atomic model. Making a Model: Key Scientists JOHN DALTON (1809) Dalton was an English schoolteacher came up with his atomic theory based on many years of experimentation by many scientists. Dalton’s Atomic Theory 1. All matter is composed of tiny particles called atoms 2. Atoms can be neither subdivided nor changed into one another 3. Atoms cannot be created or destroyed 4. 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 elements 6. Chemical change is the union or separation of atoms 7. Atoms combine in small whole-number ratios such as 1:1. 1:2, 2:3, etc.



Figure 1.1: Dalton’s “Billiard ball” model

Dalton’s 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 were

subatomic particles within the atom. • J.J. Thompson proposed that the atom is a sphere of uniform positive electricity in

which negative electrons were embedded like raisins in plum pudding (or chocolate chips in a cookie)

Figure 1.2 – Thompson’s raisin bun model for the atom RUTHERFORD (1909) • Radium gives off three different types of radiation (alpha and beta particles, and

gamma rays) • Alpha particles are Helium nuclei (2 protons, 2 neutrons) • Using these principles, Rutherford performed his famous gold foil experiment The Gold Foil Experiment Rutherford hypothesized that If Thompson’s model is true, then the high speed positively charged alpha particles should pass through the gold foil without being deflected. Although most alpha particles passed through the gold foil, some were deflected, and some even reflected back towards the source. Since opposite charges attract, this means that there must be a positive charge present in the centre or nucleus of the atom. From these observations, Rutherford formulated his own nuclear model of the atom.



Rutherford’s 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 Rutherford’s 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 when electricity is applied to elemental gas. Light is then passed through the sample and a prism. The pattern revealed is called a line spectrum. Each element has its own unique spectrum)



Bohr’s 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 n value, 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 Bohr’s terminology can be used to develop 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, is called the principle quantum number.

2. Each energy level can hold up to 2n2 electrons.

The 1st energy level can hold 2 The 2nd energy level can hold 8 The 3rd can 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 below summarizes 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



The number of protons in the nucleus is called the atomic number. This number determines the identity of an atom. Atoms are electrically neutral; therefore the number of 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 electrons orbiting 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. Let’s look at the element oxygen for an example. Oxygen has a mass number of 16. The number of neutrons = mass number – atomic number or 16-8 =8. Thus oxygen has 8 protons, 8 electrons and 8 neutrons. Elements are organized on a chart called the periodic table. You should have received a periodic table at the start of this course. The elements are organized into vertical columns called groups, and horizontal rows called periods.

Figure 1.5: The Periodic Table of Elements Let’s examine how you can retrieve information about subatomic particles from the periodic table. Example 1: Determine the number of protons, neutrons, and electrons a potassium element has.



Solution 1: The atomic number of potassium is 19, and since the atomic number is equal to the number 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 taught and are strongly recommended – DO NOT write in this book). 1. Reproduce this chart in your notes and your periodic table to complete the missing

information:

Element Atomic Number

Mass Number

Number of Protons

Number of Electrons

Number of Neutrons

Hydrogen 2 7 4 5 Carbon 7 16 9 10 Sodium 12 27 14 15 Sulphur 17 40 19 20

19 K 39.10

Atomic number

Mass number



Structure of the Atom: Bohr-Rutherford Diagrams Electrons move around the nucleus in circular paths called shells like planets around the Sun.

Electrons are spinning so fast in their orbits that they seem to form a solid shell around the nucleus. Electrons cannot exist between these orbits, but can move up or down from 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:

1st orbit (K shell) – holds 2 electrons 2nd orbit (L shell) – holds 8 electrons 3rd orbit (M shell) – holds 8electrons 4th orbit (N shell) – 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 electrons

for 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!



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 L

shell holds the remaining 4 electrons. The shells total 6 electrons.

6p+

6n0



Example 3: Nitrogen Draw a Bohr Diagram for an atom of the element nitrogen Solution 3:

14 N 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 on the locations of the elements on the periodic table.



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 one of the following Canadians that made advances in

atomic and molecular theory. (15 marks)

a. Harriet Brooks b. R.J. Leroy c. 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.

SCH4U


Lesson 2 – Quantum Mechanics



Lesson 2: Quantum Mechanics The development of the Bohr model led to further study of the atomic model based of wave properties of electrons, a branch of chemistry known as Quantum Mechanics. It can be a very complex area of chemistry, so this lesson will cover the fundamental concepts. What You Will Learn After completing this lesson, you will; • describe the quantum mechanical model of the atom (e.g., orbitals, electron

probability density) and the contributions of individuals to this model (e.g., those of Planck, de Broglie, Einstein, Heisenberg, and Schrödinger);

• list characteristics of the s, p, d, and f blocks of elements, and explain the relationship 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 Hund’s rule;

Schrodinger (1924) postulated that sometimes electrons behave as particles, and sometimes like waves. Because of this we cannot measure both the position and velocity of an electron at the same time. This exclusion is referred to as the Pauli Exclusion Principle. What this really means is that we cannot determine the momentum (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 certain place 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 on the mathematical in this course, Regions where electrons are most likely to be found are called orbitals. For every value of n, there are n types of orbitals and n2 actual orbitals

1st energy level has 1 type of orbital and (12) orbital. 2nd energy level has 2 types of orbital and (22) 4 orbitals 3rd energy level has 3 types of orbital and (32) 9 orbitals

Each electron has a set of four numbers, called quantum numbers that specify it completely; no two electrons in the same atom can have the same four. As mentioned above, 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 number represents.



The Principle Quantum Number, n The letter n was used by Bohr to identify the orbits and energies. In other words, n tells you which of the "main" energy levels the electrons are in.

• identifies the energy of 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 difference between the two is released as a photon of light.

Figure 2.1 Energy Levels (Source: Nelson Chemistry 12) The Secondary Quantum Number, l The secondary quantum number, l was introduced by Arnold Somerfield in 1915 to explain the line splitting of the hydrogen spectrum. The secondary quantum number l identifies the electron subshells or sublevels that part of a main energy level.

• identifies the shape of the orbital • values: 0 to (n-1) i.e: 0, 1, 2, 3, …n-1

s p d f



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

Figure 2.2 – Energy Sublevels (Source: Nelson Chemistry 12)



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 energy level

Principle quantum number

Possible secondary quantum numbers

Number of sublevels per primary level

1 2 3 4 5

5. Write a general rule that can be used to predict all possible values from lowest to

highest, of the secondary quantum number for any value of the principal quantum number.

The Magnetic Quantum Number, ml Recall from lesson one that a line spectrum is created by the passing of electricity through a gaseous element sample contained within a gas discharged tube. Light is passed 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 new lines that were not initially present. This occurrence was discovered by Pieter Zeeman is 1897 and is called the Zeeman Effect. For example he observed a single line transformed into three lines when the magnet was applied. This effect was explained using another quantum number, the magnetic quantum number, ml. The magnetic quantum number explains that orbits could exist at various angles.

• identifies direction of orientation (x, y, z) to an external magnetic field • values: -l…0…+ l i.e. l = 2 p orbital split into 3 planes (-1, 0, +1)



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 Number Value of l 0 to n-1 Values of ml -l to +l 0 0 1 -1, 0, +1 2 -2, -1, 0, +1, +2 3 -3,-2,-1,0,+1,+2,+3



The Spin Quantum Number ms Atoms have their own magnetism which is unique to the magnetism of a group of atoms. This magnetism is referred to as paramagnetism. It was suggested that each electron spins on its own axis. It can spin in either a clockwise or counterclockwise direction.

• identifies the spin direction of an electron • only 2 values: +½ (8) and -½ (9)

Table 2.2 Summary of Quantum Numbers Principle Quantum Number, n

Secondary Quantum Number, l

Magnetic Quantum Number, ml

Spin Quantum Number, ms

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

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

3 0 1 2

-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 in some 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, n and l describe electrons that have different energies under normal circumstances in multi-electron atoms. The other two quantum numbers apply 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 with letters 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 of designation

sharp principal diffuse fundamental

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



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. The atomic spectra indicate the energy sublevels, defined by a quantum number. An energy level diagram shows the relative energies of the electrons in various orbitals.

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

Each electron orbital is represented by a circle and can contain two electrons. The energy of the electrons increases with increasing principle quantum number, n. For a given value of n, the sublevels increase in energy, in order, s<p<d<f. The restrictions on the quantum numbers require that there can only be one s orbital, three p orbitals. Five d orbitals, and seven f orbitals. The energy diagram for an atom is directly related to its chemical properties and position on the periodic table.



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 sublevel

have at least 1 electron (Hund’s Rule)

Figure 2.5: Energy-level diagrams for lithium, carbon and fluorine. (Source:

Nelson Chemistry 12) Aufbau Diagram for Filling Orbitals

Figure 2. Memory Tip for Filling Orbitals (Source Nelson Chemistry 12) To use this diagram to help with drawing energy level diagrams, start at the bottom left side and add electrons in the order shown by the diagonal arrows. Work your way up to the upper right-hand corner.



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 in energy 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)



Example 2: Draw the energy level diagram for nitrogen anion. Solution 2: Nitrogen (N) 1s22s22p3 Nitrogen anion (N-3) 1s22s22p6 Use the same procedure except remember that a nitrogen anion will have 7 + 3 additional electrons. Electron Configurations Electron Configurations provide the same information as energy level diagrams, but in a more concise format. An electron configuration depicts the type of electron in increasing energy level. For example the electron configuration for an oxygen atom is 1s22s22p2. Example 3: Write the electron configuration for the following elements:

a) sodium atom b) fluorine ion

Solution 3: Sodium Atom 1s22s22p63s1 Fluorine ion 1s22s22p6



Shorthand Form of Electron Configuration The shorthand form abbreviates writing the electron configuration by using the nearest Nobel Gas in its written format. For example: Cl: 1s22s22p63s23p5 becomes [Ne]3s23p5 Sn:1s22s22p63s24s23d104p65s24d105p2 becomes [Kr]5s24d105p2

Support Questions

6. Draw an energy level diagram for the following

a) Carbon b) Cl-1

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

a) lithium b) oxygen c) calcium d) titanium e) rubidium f) lead g) erbium

8. Write the abbreviated (shorthand method) ground state electron configurations for

the following: a) helium b) nitrogen c) chlorine d) iron e) zinc f) barium g) polonium



Key Question #2

1. Calculate the maximum number of electrons with principal quantum number; (8 marks) a) 1 b) 2 c) 3 d) 4

2. Show the permissible values of l and m for; (8 marks) a) n=1 b) n=2 c) n=3

3. Draw energy level diagrams for beryllium, magnesium and calcium ions. What is the

similarity 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 electron configuration). (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 an

information pamphlet including;

• uses of laser technology • how it applies quantum mechanics • advantage/disadvantages of laser technology NOTE: Don’t forget to include the references for your information (Wikipedia or search engine references are not acceptable – look for professional sites or journal articles) (25 marks – 20 for content, references and design, 5 for grammar/spelling)

SCH4U


Lesson 3 – Chemical Bonding



Lesson 3: Chemical Bonding Chemical compounds are formed by the joining of two or more atoms. You will review some 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 After completing this lesson, you will; • predict molecular shape for simple molecules and ions, using the VSEPR model; • use appropriate scientific vocabulary to communicate ideas related to structure and

bonding (e.g., orbital, absorption spectrum, quantum, photon, dipole); • predict the polarity of various substances, using molecular shape and the

electronegativity values of the elements of the substances Ionic and Covalent Bonding: The Octet Rule Atoms form bonds to become more chemically stable. The most chemically stable elements on the periodic table are the noble gases. We know this because they are extremely unreactive and tend not to form compounds. According to the octet rule, atoms bond in order to achieve the same electron configuration as a noble gas. This rule is called the octet rule because all the noble gases (except helium) have eight valence electrons. Generally, when an atom tends to gain or lose electrons, an atom will have the same electron configuration (arrangement of electrons) as a noble gas. Atoms that have identical 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 atom loses electrons, it becomes positively charged and is referred to as a cation.



Lewis Structures Another way to draw atoms and ions is to use Lewis structures. Lewis structures depict only the valence electrons an element has. The figure below shows the Lewis structures 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 is

considered 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:

Clx 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 electrons an atom has. Thus, lithium is in group one and has one valence electron, while fluorine is in group seven and has seven valence electrons.



Ions can also be represented by Lewis symbols. The Lewis symbol is enclosed by a bracket as with the Bohr diagram: Example 2: Lewis Chlorine Ion - Draw the Lewis structure for a chlorine ion Solution 2:

Clx 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 and the six electrons in the p orbital’s.

Figure 3.2 - Correlation between Atomic and Bonding Theories

(Source: Nelson Chemistry 12)



Support Questions

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

Element Atomic Symbol

Electron Configuration

Number of valence

electrons

Valence

Oxygen Chlorine Sodium Phosphorus

10. Identify the elements and write the Lewis structure for the following electron

configurations:

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

Electronegativity Electronegativity is a measure of an atom’s ability to attract electrons in a chemical bond. It is a periodic property.

Predicting Bond Type Using Electronegativity You 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. The symbol ΔEN stands for the difference in electronegativity between two values 3.3 1.7 0.5 0 MOSTLY IONIC POLAR COVALENT COVALENT



Table 3.1– Electronegativities of Various Elements Element Electronegativity H 2.1 Metals Li 1.0

Be 1.5 Na 0.9 Mg 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 polar covalent bonds: Solution 3:

Substance EN Element 1 EN Element 2

ΔEN Ionic or

Covalent?

KF

0.8 4.0 3.2 ionic

O2

3.5 3.5 0 covalent

HCl

2.1 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 or gains an electron in conjunction with another atom forming a chemical bond. There are three main types of intramolecular bonds we will explore in this lesson, ionic, covalent and polar covalent. Ionic Bonding

• electrons are transferred from one atom to another • usually occurs between metals and non-metals.



Example 4: Potassium and Fluoride - Using Lewis structures draw the formation of a bond between potassium and fluorine. Solution 4: Step 1: Write out the correct number of valence electrons for each atom.

K + F

x x

x xxxx

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

K F

x x

x xxxx

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

K F

x x

x xxxx

+1 -1

Notice how the ions are written with square brackets and the overall charge is indicated in the upper right hand corner. When naming ionic compound, we name the cation first, then the anion. The ending of the anion is replaced with “ide”. Thus, the name of this compound is potassium fluoride. The positive potassium cation (+) is attracted to the negative fluorine anion (-). This is what forms the ionic bond.



Example 5: Magnesium and Nitrogen - Using Lewis structures draw the formation of a bond 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, since this does not balance we need three magnesium atoms and two nitrogen atoms to make this balance. The transfer of the electrons from magnesium to nitrogen is shown below.

Mg N xx

x

x

x

Mg N xx

x

xx

Mg

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



This compound is named magnesium nitride.

Support Questions

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

formation Name of compound

Chemical formula

Anion Cation

a) lithium and fluorine

b) calcium and phosphorus

Covalent Bonding Covalent bonds form when two or more non-metals share one or more pairs of electrons. As a result of forming covalent bonds through sharing electrons, the atoms end up with a stable electron arrangement in their outer orbit similar to that of a noble gas.



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 to

become stable. → Two chlorine atoms share a pair of electrons to form a covalent bond. Each chlorine

atom 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)



Drawing Lewis Structures for Polyatomic Ions When an ion is composed of more that one atom, it is termed a polyatomic ion. Polyatomic ions are groups of atoms that tend to stay together and carry an overall ionic 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– 2– sulfate SO4

2– 2– phosphate PO4

3– 3– Drawing Lewis Structures for polyatomic Ions can be tricky. There are some general rules for 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, and singular although not usually hydrogen or oxygen). In this case, nitrogen is our central atom.

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



Step 3: Place a bonding pair of electrons between the central atom and each of The 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 surrounding atoms to form double or triple bonds.

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



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 theory explains why and how electrons are shared between atoms. The VB theory postulates how individual atoms, each with its own orbitals and electrons come 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 overlap and 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)



According to Valence Bond Theory: • A half-filled orbital in one atom can overlap with another half-filled orbital of a

second 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 they change shape. They form new orbitals called hybrid orbitals. Hybridization is the process of combining two or more atomic orbitals to create new orbitals, called hybrids that will fulfill the geometric demands of the system. When an atom hybridizes, it will restructure its original set of s and p atomic orbitals into a new set of hybrid orbitals. The process of hybridization is driven by the needs of the atoms to produce specific geometric patterns. Example 9: What are the bonding orbitals of the BF3 molecule? 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 this case the 2s has two electrons, and p orbital has one. B: 1s22s22p1

2s2 2px

1 2py 2pz In this case, one of the 2s2 electrons can be promoted to an empty p orbital 2s1 2px

1 2py 2pz



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

More forms of hybridization are summarized in the table following. Table 3.3: Forms of hybridization



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 (σ)

b) Side by side overlap forms pi (π) 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 one “normal” 2pz orbital.



This is shown in figure 3.4 following:

Figure 3.4: Partial hybridization (Source: Nelson Chemistry 12) Partial hybridization can also be used to explain how triple bonds form Consider ethyne gas C2H2

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



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 following atoms? a) H b) F c) S d) Br

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

following atoms and indicate the type of hybridization involved when each atom forms a compound; a) carbon in CH4 b) boron in BH3 c) beryllium in BeH2

15. When are π bonds formed? 16. Provide an explanation for bonding in each of the following;

a) C2Cl4 b) C2F2



The Valence Shell Electron Pair Repulsion (VSEPR) model:

One of the more important properties of any molecule is its shape. It is very important to know the shape of a molecule if one is to understand its reactions. It is also desirable to have a simple method to predict the 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 called bonding pairs. More than one set of bonding pairs of electrons may bind any two atoms together (multiple bonding).

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

• The bonding pairs and lone pairs around any particular atom in a molecule adopt positions in which their mutual interactions are minimized. The logic here is simple. Electron pairs are negatively charged and will get as far apart from 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 the general formula AX2 and is considered linear.



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 electrons totalling 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 following molecules: a) CO2 b) HCN c) BF3 d) SiCl4 e) CH4 f) OCl2



Polar Covalent Bonds A polar covalent intramolecular bond is formed when there is unequal sharing between valence 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, ΔEn is 0.5, and the intramolecular bond is considered polar covalent. Since the chlorine has a higher electronegativity, the electrons are pulled closer to the chlorine’s atomic nucleus. This makes the chlorine ends or dipoles slightly negative (δ-). Conversely, the electrons are farther away from carbons nucleus, making it slightly positive (δ+).

Molecular Polarity Covalent bonds can be polar or nonpolar. However the polarity of a molecule as a whole is dependent on bond polarity and molecular 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 your reasoning. Solution 12: First, draw the Lewis structure



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

polar covalent or not.

Substance EN Element 1 EN Element 2

ΔEN Ionic or

Covalent?

NaCl

Cl2

HF

19. Predict the polarity of following molecules.

a) BF3 b) OF2 c) CI2 d) PCl3



Key Question #3

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

Molecule or Compound ion

Lewis structure

Name of Shape

Bond Polarity

Molecular Polarity

HCl CH4 CH3Cl CO2 H2O NH3

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) H2O b) C2H2 c) C2H4 d) C2H6

SCH4U


Lesson 4 – Intermolecular Forces



Lesson 4: Intermolecular Forces In this lesson, we will examine the other types of bonding that have not yet been covered 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 After completing this lesson, you will; • 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 of forces between them (e.g., covalent bonds, Van der Waals forces, dipole forces, and metallic bonds)

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

• conduct experiments to observe and analyse the physical properties of different substances, and to determine the type of bonding present

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

Bonds between Molecules: Intermolecular Bonds Intermolecular bonds are the chemical bonds between molecules. These bonds determine the physical state of molecular substances. These bonds are broken as a substance undergoes a change of state. Intermolecular forces are much weaker than covalent bonds. There are generally three types of intermolecular bonds: London forces, dipole-dipole forces, and hydrogen bonds. These intermolecular forces are collectively called Van der 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 weak forces of attraction. Momentary dipoles are created by the electrons contained within the compound, which are constantly in motion

All molecules Methane gas (CH4)

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

Hydrogen Chloride(HCl)

Hydrogen bonding

Formed between the electropositive Hydrogen dipole and an electronegative dipole of Oxygen, Chlorine, or Fluorine.

Pure distilled water



Using Vander Waals forces to predict Boiling Points As the number of electrons in the molecule increases, the boiling point increases as well. 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 following molecules;

a) water, H2O b) butane C4H10 c) hydrogen chloride, HCl

21. Which of the following pure substances has a stronger dipole-dipole force than the

other? Discuss your reasons for your conclusion.

a) hydrogen chloride or hydrogen fluoride b) chloromethane or iodomethane c) 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 greatly in such physical properties such as hardness, melting point, mechanical characteristics and 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 Solids Class 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)



Ionic Crystals Earlier in this unit, you learned that an ionic bond is formed when there is a transfer of electrons from a metal atom to a non-metal atom. However, the transfer of electrons is not 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 crystal lattice. 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 will not discuss here. Ionic compounds are relatively hard but brittle substances. They conduct electricity in liquid 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 are good conductors of heat and electricity, their hardness ranges from hard to soft. The

melting points also range from high to low. Metals have a continuous compact crystal structure. The properties of metal are the result of bonding between fixed positive nuclei and mobile “loose” valence electrons. In more simplistic terms, the valence electrons float around the network between positive centers. The motility of the electrons explains why metals are such good conductors of heat and electricity.

Figure 4.2 – Metallic Bonding – A sea of negative electrons floats freely amongst positive nuclei. Molecular Crystals Molecular solids may be elements such as iodine and sulphur or compounds such as ice or carbon dioxide. They are crystals that have relatively low melting points, they are not hard, and are not conductors of electricity. They are arranged by neutral molecules with weak intermolecular forces.



Covalent Crystals Massive aggregates of atoms, as distinct from ions, can exist where neighbouring atoms share a pair of electrons to form multiple, single, covalent bonds. This type of crystal is usually found in the elements carbon and silicon. Carbon and Silicon both have four valence electrons. Diamond and Graphite are both made up of carbon. Diamond is a huge network of C-C linkages where all four of the carbon bonds are equal in strength and the whole structure is enormously strong. Graphite, on the other hand, is a carbon crystal in which each carbon atom uses only three of its four available valence electrons to bond with adjacent carbons atoms. This results in layers of carbon that are not bonded. The unused bonded electrons can move through layers, explaining why graphite can conduct electricity, whereas diamond, does not. The layers of carbon in graphite are held together by Vander Waals forces. The layers in graphite can slip over one another easily, and although graphite is strong in two dimensions, it is weak in third. This makes graphite useful for products such as pencils, golf clubs, and as a lubricant. Diamond on the other hand is used to

cut glass, and is considered precious due its hardness. Figure 4.3 - Diamond Covalent crystals tend to exceptionally hard, with very high melting points. They are invariably insoluble in water, since there are no ions to attract the polar water molecule.

Figure 4.4 – Graphite



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) Na2S e) C f) CaO

4. Water beads on the surface of a freshly waxed car hood. Use your knowledge of

intermolecular forces to explain this observation. (5 marks) 5. All molecular compounds may have London, dipole-dipole, and hydrogen-bonding

intermolecular forces, affecting their physical and chemical properties. Indicate which intermolecular forces contribute to the attraction between molecules in each of the following classes of organic compounds: (8 marks) a) pentane, C2H5 b) 2-propanol, CH3CHOH c) acetic acid, CH3CHOHCH d) ethybenzoate, C6H5COOCH2CH5 e) dimethylether, CH3OCH3 f) ethylamide, CH3CONH2 g) diamond, C h) calcium carbonate, CaCO3

6. Use the theory of intermolecular bonding to explain the sequence of boiling points in

the following alkyl bromides: CH3Br(g) (4oC), C2H5Br(l) (38oC), and C3H7Br(l) (71oC). (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

SCH4U Grade 12

University Chemistry

Support Question Answers

SCH4U – Chemistry Unit 1 – Support Question Answers


Answer to Support Questions 1.

Element Atomic Number

Mass Number

Number of Protons

Number of Electrons

Number of Neutrons

Hydrogen 1 1 1 1 0 Helium 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 12 Aluminum 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 period indicated the number of shells the atom has. 3. Bohr- orbits are 2D, fixed distance from nucleus, circular or elliptical path, 2n2

electrons per orbit, Sommerfield- 3D region in space, variable distance from nucleus, no path, 2 electrons per orbital



4. Primary energy level

Principle quantum number

Possible secondary quantum numbers

Number of sublevels per primary level

1 1 0 1 2 2 0,1 2 3 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) 1s22s1 b) 1s22s22p4 c) 1s22s22p63s23p64s2 d) 1s22s22p63s23p64s23d2 e) 1s22s22p63s23p64s23d104p65s1 f) 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2 g) 1s22s22p63s23p64s23d104p65s24d105p66s24f12



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

a) [He] b) [He]2s22p3 c) [Ne]3s23p5 d) [Ar]4s23d6 e) [Ar]4s23d10 f) [Xe]6s2 g) [Xe]6s24f145d106p4

9.

Element Atomic Symbol

Electron Configuration

#of valence electrons

Valence

Oxygen O 1s22s22p4 6 -2 Chlorine Cl 1s22s22p63s23p5 7 -1 Sodium Na 1s22s22p63s1 1 +1 Phosphorus P 1s22s22p63s23p3 5 -3 10. Identify the elements and write the Lewis structure for the following electron

configurations;

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

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

Ox x

x x x x

Px x

Brx x

Rb 11. Bond formation Name of

compound Chemical formula

Anion Cation

a) lithium and fluorine

Li F

x x

x xxxx

+1 -1 Lithium fluoride

LiF fluorine lithium

b) calcium and phosphorus

Calcium phosphide

Ca3P2 phosphorus calcium



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 following

atoms?

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 atom forms a compound;

a) carbon in CH4, 1s22s22p2 promoted to1s22s12px

12py12pz

1 (sp3) b) boron in BH3 1s22s22p1 promoted to 1s22s12px

12py1 (sp2)

c) beryllium in BeH2 1s22s2 promoted to 1s22s12px1 (sp)

15. When are π 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 bonds

and 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 the adjacent carbon overlap sideways.

17. a) CO2(g) - AX2- Linear

b) HCN - AX2 –Linear c) BF3 - AX3- Trigonal planar d) SiCl4 - AX4 -tetrahedral e) CH4 - AX4 - tetrahedral f) OCl2 - AX2E2 – V-shaped or bent



18. Substance EN Element 1 EN Element 2 ∆ EN Ionic or

Covalent? NaCl Na = 0.9 Cl = 3.0 2.1 Ionic Cl2 Cl = 3.0 Cl = 3.0 0 Covalent HF H = 2.1 F = 4.0 1.9 Polar covalent 19. Predict the polarity of following molecules.

a) BF3 -polar b) OF2 – non-polar c) CI2 – non-polar d) PCl3 -polar

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

molecules;

a) water, H2O – hydrogen bonding b) butane C4H10 –London forces c) hydrogen chloride, HCl –dipole-dipole

21. Which of the following pure substances has a stronger dipole-dipole force than the

other? Provide your reasoning.

a) hydrogen chloride or hydrogen fluoride: HF – greater electronegativity difference between atoms

b) chloromethane or iodomethane: chloromethane - greater electronegativity difference between atoms

c) nitrogen tribromide or ammonia: ammonia - greater electronegativity difference between atoms

d) water or hydrogen sulphide: water - greater electronegativity difference between atoms

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