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SCH4CGrade 12

College Chemistry

Version A

SCH4C – Chemistry Introduction

Introduction

Welcome to the Grade 12 College Chemistry Course, SCH4C. This full-credit course is part of the Ontario Secondary School curriculum.

This course enables students to develop an understanding of chemistry through the study of matter and qualitative analysis, organic chemistry, electrochemistry, chemical calculations, and chemistry as it relates to the quality of the environment. Students will use a variety of laboratory techniques, develop skills in data collection and scientific analysis, and communicate scientific information using appropriate terminology. Emphasis will be placed on the role of chemistry in daily life and the effects of technological applications and processes on society and 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 lesson. 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.

Copyright © 2008, Durham Continuing Education Page 2 of 61


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 5 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 Matter and Qualitative Analysis

Lesson 1 Making ObservationsLesson 2 Atoms and ElementsLesson 3 Chemical Bonding and NomenclatureLesson 4 Types of Chemical Reactions

Unit 2 Quantities in Chemistry

Lesson 5 The Mole Lesson 6 Percent CompositionLesson 7 Stoichiometric RelationshipsLesson 8 Preparing Solutions

Unit 3 Organic Chemistry

Lesson 9 HydrocarbonsLesson 10 Functional GroupsLesson 11 Types of Organic ReactionsLesson 12 Polymers

Unit 4 Chemistry in the Environment

Lesson 13 Ions and Water TestingLesson 14 Introduction to Acids and BasesLesson 15 Acid-Base ReactionsLesson 16 Clean Air for Tomorrow

Unit 5 Electrochemistry

Lesson 17 Oxidation and Reduction ReactionsLesson 18 The Activity Series of MetalsLesson 19 Galvanic CellsLesson 20 Electrolytic Cells


SCH4C – Chemistry Unit 1 - Introduction

Unit 1: Matter and Qualitative Analysis

Introduction

Welcome to Grade 12 College Chemistry, SCH4C. Chemistry is the study of matter and energy and the interactions between them. Chemistry focuses on the properties of substances and the interactions between different types of matter.

Matter is anything that has mass and occupies space. Matter will be the central focus of this unit. In order to better understand matter and interactions between various types of matter, chemists must be able to properly describe it by qualititative means such as sight, smell and sound, and quantitative measures such as mass or volume.

Overall Expectations

After completing this unit, you will be able to

evaluate the effects of chemical substances on the environment, and analyse practical applications of qualitative analysis of matter;

investigate matter, using various methods of qualitative analysis; demonstrate an understanding of the basic principles of qualitative analysis of

matter.

There are four lessons in this unit:

Lesson 1 Making ObservationsLesson 2 Atoms and ElementsLesson 3 Chemical Bonding and NomenclatureLesson 4 Types of Chemical Reactions


SCH4CGrade 12

College Chemistry

Lesson 1Making Observations

Lesson 1: Making Observations

This lesson will be an introduction to the course and will provide a review of some chemistry fundamentals such as how to make proper observations, working with numbers and interpretation of experiment results.

After completing this lesson, you will be able to

demonstrate an understanding of the basic principles of qualitative analysis and underlying theories;

carry out qualitative analyses, using flow charts and appropriate laboratory equipment and instruments;

explain the distinction between observation and inference;

Making Observations

When you perform a science experiment, you are often trying to answer a question you have while observing the world around you. You research your problem, then make a prediction or hypothesis and finally perform your experiment. An important step to the scientific method comes while performing your experiment: making observations. A good scientist is observant and notices things in the world around him/her. (S)he sees, hears, or in some other way notices what’s going on in the world and becomes curious about what’s happening. This includes reading and studying what others have done in the past because scientific knowledge builds on the knowledge of what is already known. For example, in order for Neils Bohr and Ernest Rutherford to come up with their now famous model of the atom, the Bohr-Rutherford model, they used information gathered and analyzed from other chemists such as John Dalton. Thus, keeping detailed observations is an important step in scientific research.

Qualitative and Quantitative Observations

There are two types of observations: qualitative and quantitative. Qualitative observations describe matter using words. For example, examine a metal fork at your house. What words could you use to describe the fork? Here are some examples; silver, solid, hard, and shiny. You may have also made some qualitative observations about the form or shape of the fork. Qualitative observations your senses to physically describe matter (sight, smell, sound, taste, and touch).

Quantitative observations use numbers rather than words. Using the fork example, we could describe the fork as having a mass of 10 grams, or that it has a density of 22g/cm3. Table 1.1 Summarizes common quantitative observations that are referred to as SI units. SI is the abbreviation of Systeme International which is the standardized units developed by scientists in France and used today globally.

Measurement SI unitMass grams (g), kilograms (kg)Volume litres (L), millilitres (mL)Length metres (m), centimetres (cm)Density grams per millilitre (g/mL)

grams per litre (g/L) Table 1.1: Common SI units

In previous science courses you have learned the steps in the scientific method. Figure 1.1 below depicts how observations fit into the scientific method:

Figure 1.1: The Scientific Method

Using the chart on the previous page, your observations will allow you draw conclusions about your predictions. These conclusions may or may not support your original hypothesis. When you draw conclusions from your observations you are really making inferences about the data you observed. For example, you may notice that it is cloudy outside, so you infer that it is going to rain. A scientist performing an experiment involving plants and sunlight may infer that plants need sunlight when observing through experimentation that plants grow better in sunlight than when left in darkness.

Significant Digits

When stating your quantitative observations, it is important to state the numbers to the correct number of significant figures. A significant figure is the number of digits you are certain of. For example, you may be certain that your weight is 165 pounds, but you are uncertain whether your weight is 165.2 or 165.8. When recording data, scientists only state the digits they are certain of. Here are some general rules to follow for significant figures:

Non-zero digits are always significant. Thus, 19 has two significant digits, and 42.3 has three significant digits.

With zeroes, the situation is more complicated: Zeroes placed before other digits are not significant; 0.056 has two significant

digits. Zeroes placed between other digits (“sandwiched zeros” are always significant;

7004 kg has four significant digits. Zeroes placed after other digits but behind a decimal point are significant; 2.90

has three significant digits.

Multiplying and Dividing Significant Digits – In a calculation involving multiplication or division, the number of significant digits in an answer should equal the least number of significant digits in any one of the numbers being multiplied or divided.

Adding and Subtracting Significant Digits – When quantities are being added or subtracted, the number of decimal places (not significant digits) in the answer should be the same as the least number of decimal places in any of the numbers being added or subtracted.

Example:

5.67 J (two decimal places) 1.1 J (one decimal place) 0.9378 J (four decimal places) 7.7 J (one decimal place)

Support Questions (Reminder: these questions are not to be submitted but reinforce the

material taught and are strongly recommended).

1. Classify the following as a qualitative or quantitative observation:

a. Michelle observed five eggs in a basket.b. Rahim determined the objects in the basket have a circumference of ten

centimetres.c. Ian measured an object to have a mass of 21.5 grams.d. The object has a black shiny surface.e. The thermometer indicates that the liquid has a temperature of 320C.f. Jacob observed 3 green geckos sitting on a tree branch.g. Nancy recorded in her journal that the blue block floated lower in the water than

the red block.h. The earth material has a density of 3.2 g/cm3.i. The mineral is greenish-blue in colour.j. This afternoon, the temperature was 8oC.

2. For his science experiment, Mohammed decided to burn a candle, and wrote the following statements down. Classify the statements as an observation or an inference.

a. The candle is redb. The candle is using oxygenc. The flame is yellowd. The candle is cylindricale. The flame is giving off carbon dioxidef. Wax is dripping down the candleg. The candle is 10 centimetres high

3. State the number of significant digits each of following has:

a. 1234 b. 0.023 c. 890 d. 91010 e. 9010.0 f. 1090.0010

4. The mass of empty graduated cylinder was found to be 23.1g. The mass of the graduated cylinder filled with water has 32 g. Calculate the mass of the water. State your answer to the correct number of significant figures.

Key Question #1 (27 marks)

Activity: Mystery Powder Analysis Complete the Student Exploration on the analysis of a mystery powder and answer the questions at the bottom of the next page. This is a GIZMO that can be found on the website www.explorelearning.com. Ask me for the worksheets and your password and login for that website. [15 marks] JUST ANALYZE THE FIRST 10 SUBSTANCES!!!!

You will not prepare a lab report for this activity, but, in general, lab reports should be typed and include the following:

Title Page: Title of Lab, Course Name/Code, Date(s) Performed, Your Name

Purpose Statement: One or two lines which explain the reason for doing the experiment, or what is hoped to be learned by doing the experiment. “The purpose of this experiment is to…”

Hypothesis: State your hypothesis in an “If” (independent variable)… “then” (dependent variable)…”because” (educated guess)… statement.

Materials: Include a list of all equipment and materials (chemicals) that are needed for the experiment.

Apparatus: A list of all the glassware and tools needed to perform the experiment. These could be documented in the form of a diagram.

Procedure/Method: These should be short numbered steps written in the past-impersonal tense. In many cases you may refer to the text or a handout, noting any modifications.

Observations: These should be organized in short numbered statements or preferably in a table. All data tables must be numbered and have a relevant name. Units must be included. The aim is to present the observations as clearly as possible and to aid in forming conclusions. Remember that observations are things that you have seen, heard, smelled, or measured, not things that you have inferred based on your observations.

Analysis: Answer analysis questions/calculation in this section in numbered form. Not all labs will have calculations. Always include the units in your calculations. Indicate what you are trying to calculate at the top of each calculation. If appropriate, the results of the calculations could be presented in a table.

http://www.explorelearning.com/

Discussion: Answer the following in paragraph form:

$ What did you do?$ How did you do it?$ What were your results?$ What do your results mean?$ State any sources of possible experimental error. (Never write there was no error).

Conclusions: These are short numbered statements describing what was learned from your observations

QUESTIONS

1. What is a hypothesis? Write a hypothesis for this lab. [4]

2. Write a list of materials for this lab. [2]

3. List the apparatus needed for this lab. [4]

4. What would be 2 sources of error that would likely affect the results of this lab? [2]

Grade 12College Chemistry

Lesson 2Atoms and Elements

SCH4C – Chemistry Lesson 2

Lesson 2: Atoms and Elements An atom is the smallest particle of an element that still retains the identity and properties of the element. All elements contain one type of atom, and cannot be broken down into smaller particles. In this lesson, you will learn about various scientists’ contributions to the current atomic model as well as how to draw the first 20 elements on the periodic table. You will also learn some practical applications of the current Bohr-Rutherford model of the atom.

In this lesson you will:

use appropriate symbols and language to represent atoms and their structure describe and explain basic processes involved in qualitative analysis, including

flame tests and atomic line spectra

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 who 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 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 properties5. 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.



Figure 2.1: 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$ 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 2.2 – 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 repel, this meant 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 2.3 Rutherford beehive model

Weaknesses of the Rutherford model

Why are negative electrons not pulled into the positive nucleus by the attraction of unlike charges?

A clue to the problem is obtained from the study of the light given off from high energy substances. If a Tungsten wire is in a high energy state due to the heating effect caused by an electric current, it releases the extra energy in the form of white light. White light is composed of all colours or frequencies of light. Each frequency has a characteristic energy.

The relationship between energy and frequency was shown by the German physicist, Max Planck. He assumed that the light is made up of discrete (separate) packages of energy. Each package is called a quantum or a photon. The relationship between energy and frequency is E=hv where E is the amount of energy possessed by a quantum, h is Plank’s constant, and v is the frequency of the quantum. According to the theory, red light consists of a stream of quanta, each of which is a certain amount of energy and has a characteristic frequency. Violet light also consists of a stream of quanta. However, quanta of violet light have more energy and a higher frequency than do quanta of red light.



White light is composed of all frequencies and all energies of visible light. White light can be broken down into its component colours by a prism or a diffraction grating. A spectrum of colours is obtained. It is called a continuous spectrum because it consists of all the colours or frequencies of visible light. In order of increasing frequencies (and increasing energies) the colours are Red, Orange, Yellow, Green, Blue, Indigo, and Violet or ROYGBIV for short.

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

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.

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

Each energy level can hold up to 2n2 electrons.

The 1st energy level can hold 2 The 2nd energy level can hold 8The 3rd energy level can hold 18* (* holds up to 8 in the first 20 elements – however, increases to 18 after the Calcium element)

Energy Level Population = electrons in ground state

Example:

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

Therefore, sodium holds 2 electrons in the first level, 8 electrons in the second level and 1 electron in the third level. Phosphorus is 2, 8, 5.



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).

Subatomic particle Charge Location in atom Relative massProton Positive (+1) nucleus 1

Neutron Neutral (0) nucleus 1Electron Negative (-1) Orbits nucleus 1/1800

Table 2.1: Subatomic particles

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.

Note: Get your own copy of the periodic table from McGill University at http://www.mcgill.ca/ehs/radiation/basics/periodic/ - you can download and print a coloured periodic table from the .pdf file available on this site.


http://www.mcgill.ca/ehs/radiation/basics/periodic/


Figure 2.1: The Periodic Table of Elements

Let’s look at 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 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 39 -19 or 20 neutrons (note: you must round the mass number to the nearest whole number before calculating the neutrons).


19

K

39.10

Atomic number

Mass number


Support Questions(Reminder: these questions are not to be submitted but reinforce the material taught and are strongly recommended).

1. Recreate this table in your notes. Use your periodic table to fill in 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

Sulfur

17

40

19

20

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 8 electrons (remember – holds up to 8 in the first 20 elements then can increase to 18 after the 21st element)

4th orbit (N shell) – holds 2 electrons

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.

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.



Example 3 – Nitrogen

Draw a Bohr Diagram for an atom of the element nitrogen

Solution 3

14N 7

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. You may use the short-hand version if you like.

Applications of the Bohr Model

As you have learned, every atom consists of a nucleus with tiny electrons rotating around it. The further away from the nucleus the electrons are, the more energy they possess. If a metal atom is heated, it jumps higher away from the nucleus. When they fall back closer to the nucleus, they give off energy as light.



Different metals produce different colours of light. If we look at the colour of the light when a solution of metal is heated in a flame, we can tell which metal is present.

When a firework explodes into a beautiful array of colours, we see evidence of the electrons jumping to different energy levels and releasing light when they fall back.

Other practical applications include counterfeit money detection. Did you ever wonder why the cashier puts a money bill under an ultraviolet light? The light causes the electrons to jump and produce a colourful image when they fall back. This lets the cashier know that the money is real.

Common characteristic element colours are summarized below. Chemists perform flame tests, which involves heating the element and observing its characteristic colours. This is one reason that scientist know which elements are present in the Sun.

Flame colour ElementRed Lithium

StrontiumOrange CalciumYellow SodiumBlue CopperPurple Potassium

Table 2.2: Flame Colour of common elements



Key Question #2

1. On a separate piece of paper, create a chart similar to the one following. Fill in the missing information. (12 marks)

Standard Notation

Atomic number

Atomic mass

Number of protons

Number of electrons

Number of neutrons

Bohr Diagram

10

35 Cl17

2. Complete the Student Exploration on Electron Configuration. This is a GIZMO that can be found on the website www.explorelearning.com. Ask me for the worksheets and your password and login for that website. [15 marks]


http://www.explorelearning.com/

SCH4CGrade 12

College Chemistry

Lesson 3Chemical Bonding and Nomenclature


Lesson 3: Chemical Bonding and Nomenclature

In the previous lesson you learned what an atom is and how to draw an atom. This lesson will deepen your understanding of how atoms interact with each other to form compounds. You will learn about the types of chemical bonding, and how to name the resulting compounds.

After completing this lesson, you will be able to

represent ionic and molecular compounds by their accepted formulae and names explain covalent bonding in simple molecules, using Lewis structures demonstrate an understanding of the formation of ionic bonds demonstrate an understanding of the formation of polar covalent bonds

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.

For example, consider the formation of the calcium cation. The calcium atom has the following configuration:

Example 1 – Calcium

Explain the formation of a calcium ion

Solution 1

Ca)2e-)8e-)8e-)2e-

If we examine the outermost shell of the calcium atom, we notice that there are two electrons (shown enlarged). The electrons in the outermost shell are referred to as valence electrons, and are the most reactive when atoms combine to form



compounds. This calcium atom will lose two electrons to become isoelectronic (meaning to have the same electron configuration as argon:

[Ca)2e-)8e-)8e-]+2

Since the calcium cation lost two electrons, it has a charge of positive two (+2). Notice how the ion is written with square brackets and that the overall charge is indicated in the upper right hand corner.

Example 2 – Nitrogen

Show the formation of the nitrogen ion

Solution 2N)2e-)5e-

In this case, nitrogen has five valence electrons, and must gain three to obtain a full octet (8 valence electrons). Thus nitrogen will form a negatively charged ion or anion. Since the ion gained three electrons, it now has a charge of negative three (-3)

[N)2e-)8e-]-3

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 examples following show 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 3 – Chlorine

Draw the Lewis structure for a chlorine atom

Solution 3

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 4 – Lewis Chlorine Ion

Draw the Lewis structure for a chlorine ion

Solution 4

Support Questions

1. Reproduce and complete the following table;

Element Atomic Symbol Number of valenceelectrons

Valence

OxygenChlorineSodiumPhosphorus

2. Reproduce and complete the following table. You may use the shorthand method for drawing Bohr Diagrams.

Element Bohr Atom Bohr Ion Lewis Atom Lewis IonLithium

Magnesium

Oxygen



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 5 – Potassium and Fluoride

Using Lewis structures draw the formation of a bond between potassium and fluorine.

Solution 5

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

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

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.



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 6 – Magnesium and Nitrogen

Using Lewis structures draw the formation of a bond between magnesium and nitrogen.

Solution 6

Again, first draw the Lewis structures

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 following.



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

3. Reproduce and complete the following table. Use Lewis structures to depict atoms.

Bond formation

Name of compound

Chemical formula

Anion Cation

lithium and fluorine

calcium and phosphorus

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.

Examples 7 and 8

The nitrate ion The sulphate ion

When a compound containing a polyatomic ion is dissolved in water, the metal ion separates from the polyatomic ion, but the atoms of the polyatomic ion stay together as a unit.



Name of polyatomic ion Ion formula Ionic chargenitrate NO3

– 1–hydroxide OH– 1–

bicarbonate HCO3– 1–

chlorate ClO3– 1–

carbonate CO32– 2–

sulphate SO42– 2–

phosphate PO43– 3–

Table 3.1 – Common Polyatomic Ions and Their Ionic Charges

Polyatomic ions have many practical everyday uses and sources. These are summarized in table 3.2 below:

Compound Formula Use or sourcecalcium carbonate CaCO3 chalk and building materials

magnesium hydroxide Mg(OH)2 stomach antacidssulphuric acid H2SO4 car battery acid

copper (II) sulphate CuSO4 fungicidesodium carbonate Na2CO3 laundry detergentsammonium nitrate NH4NO3 fertilizer

Table 3.2 – Common Polyatomic Compounds

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 9

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

Solution 9

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 10

Draw the formation of chlorine gas using Lewis structures

Solution 10



Other examples of covalently bonded molecules include:

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

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 EN stands for the difference in electronegativity between two values

3.3 1.7 0.5 0

MOSTLY IONIC POLAR COVALENT COVALENT



Element ElectronegativityH 2.1

Metals Li 1.0Be 1.5Na 0.9Mg 1.2K 0.8Ca 1.0

Non-metals C 2.5N 3.0O 3.5F 4.0P 2.1S 2.5Cl 3.0

Table 3.3 – Electronegativities of Various Elements

Example 11

Determine if the elements below would form ionic, covalent or polar covalent bonds:

Solution 11

Substance EN Element 1 EN Element 2 EN Ionic or Covalent?

KF 0.8 4.0 3.2 ionic

O2 3.5 3.5 0 covalent

HCl 2.1 3.0 1.1 Polar covalent

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 12Draw the polar bond formed in a molecule of carbon tetrachloride (CCl4) between the elements carbon and chlorine



Solution 12 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+).

Support Questions

4. Reproduce and complete the table below. If the molecule is covalent, indicate if it is polar covalent or not.


NaCl

Cl2

HF



5. Compare and contrast covalent and polar covalent compounds.

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.

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 3.5 below.

Force Description ExampleLondon 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

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

Table 3.4 Vander Waals forces

Support Questions

6. Determine the type of Vanderwaals forces that would occur between the following moleculesa. water, H2Ob. butane C4H10



c. hydrogen chloride, HCl



Key Question #3

1. Recreate and complete this table on a separate piece of paper. Use Lewis structures to depict atoms. (10 marks)

Bond formation

Name of compound

Chemical formula

Anion Cation

sodium and oxygen

lithium and nitrogen

2. Draw Lewis structures for the each of the following covalent bonds, then state if they are polar covalent or non-polar covalent molecules.

a. water, H2O

b. Iodine gas, I2

c. methanol, CH3OH


SCH4CGrade 12

College Chemistry

Lesson 4 Types of Chemical Reactions


Lesson 4: Types of Chemical Reactions

Chemical reactions occur everyday in our world around us. For example, a bike chain rusts when left outside for long periods of time, milk sours if it is left too long in the fridge, or even when we burn the food we each, thousands of chemical reactions are occurring in the cells in our bodies. In this lesson, you will learn about the types of chemical reactions and practical applications of this topic.

After completing this lesson, you will be able to;

Use appropriate scientific vocabulary to communicate ideas related to qualitative analysis, such as double displacement reactions

Write and identify balanced chemical equations Predict the formula of the precipitate formed in a chemical reaction by writing

double-displacement equations Predict the precipitate formed in a chemical reaction by writing net ionic equations

and using a table of solubility rules Carry out and explain applications of qualitative analysis in the field of water

treatment

Chemical Reaction Types

A chemical reaction has occurred when substances change into new substances with new physical and chemical properties. There are four main types of chemical reactions:

Synthesis Decomposition Single Displacement Double Displacement

Synthesis

Synthesis reactions involve direct combinations of 2 substances to produce 1 new substance:

A + B AB

You can often predict the formula of the product (right hand side of the equation) of a synthesis reaction involving elements if you know their valences.

Example 1 – A metal and a non-metal form an ionic bond

Solution 12Na(s) + Cl2(g) 2NaCl(s)


Subscripts indicate the form:(s) = solid(l) = liquid(g) = gas(aq) = dissolved in H2O


Example 2 – A metal or a non-metal combines with oxygen to form an oxide

Solution 22Mg(s) + O2(g) 2MgO(s)

C(s) + O2(g) CO2(g)

Decomposition

Decomposition reactions involve the breaking up of a compound into simpler substances:

AB A + B

A compound will separate into either elements or compounds. These reactions are much harder to predict than synthesis reactions.

Example 3 – The decomposition of water

Solution 3

H2O(l) O2(g) + H2(g)

Example 4 – Decomposition of hydrogen peroxide

Solution 4

2H2O2(aq) O2(g) + 2H2O(l)

Example 5 – Decomposition of calcium carbonate

Solution 5

CaCO3(s) CaO(s) + CO2(g)

Single Displacement Reactions

These reactions are also referred to as substitution reactions because one element replaces another in a compound.

X + AB A + XB

OR

Y + AB B + AY



Example 6 – A metal replaces a metal

Solution 6

Mg(s) + 2AgNO3(aq) 2Ag(s) + Mg(NO3)2(aq)

Example 7 – A non-metal replaces a non-metal

Solution 7

Br2(l) + CaI2(aq) I2(s) + CaBr2(aq)

Double Displacement Reactions

These reactions are also referred to as substitution reactions because one element replaces another in a compound:

AB + CD AD + CB

Example 8 – Lead (II) nitrate + potassium iodide lead (II) iodide + potassium nitrate

Solution 8 Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq)

NOTE: In both single and double displacement reactions, metals exchange places with metals,

and non-metals exchange places with non-metals.

Balancing Equations

Chemical equations do not come already balanced. This must be done before the equation can be used in a chemically meaningful way. A balanced equation has equal numbers of each type of atom on each side of the equation.

The Law of Conservation of Mass is the rationale for balancing a chemical equation. The law was discovered by Antoine Laurent Lavoisier (1743-94). The Law of Conservation of Mass states that "Matter is neither created nor destroyed."


Many double displacement reactions are also called Precipation reactions. A solid precipitate will often form when two aqueous solutions are mixed. In order to predict this, a solubility table must be used. This will be discussed in the section that follows


Therefore, we must finish our chemical reaction with as many atoms of each element as when we started.

Here is an example unbalanced equation for this lesson:

H2 + O2 H2O

It is an unbalanced equation (sometimes also called a skeleton equation). This means that there are UNEQUAL numbers at least one atom on each side of the arrow.

In the example equation, there are two atoms of hydrogen on each side, BUT there are two atoms of oxygen on the left side and only one on the right side.

Remember this: A balanced equation MUST have EQUAL numbers of EACH type of atom on BOTH sides of the arrow.

An equation is balanced by changing coefficients in a somewhat trial-and-error fashion. It is important to note that only the coefficients can be changed, NEVER a subscript.

The coefficient times the subscript gives the total number of atoms.

Three quick examples before balancing the equation.

a. 2 H2 - there are 2 x 2 atoms of hydrogen (a total of 4).

b. 2 H2O - there are 2 x 2 atoms of hydrogen (a total of 4) and 2 x 1 atoms of oxygen (a total of 2).

c. 2 (NH4)2S - there are 2 x 1 x 2 atoms of nitrogen (a total of 4), there are 2 x 4 x 2 atoms of hydrogen (a total of 16), and 2 x 1 atoms of sulphur (a total of 2).

So, now to balancing the example equation:

H2 + O2 H2O

The hydrogens are balanced, but the oxygens are not. We have to get both balanced. We put a two in front of the water and this balances the oxygen.

H2 + O2 2 H2O

However, this causes the hydrogen to become unbalanced. To fix this, we place a two in front of the hydrogen on the left side.

2 H2 + O2 2 H2O

This balances the equation.



Two things you CANNOT do when balancing an equation.

1. You cannot change a subscript.

You cannot change the oxygen's subscript in water from one to two, as in:

H2 + O2 H2O2

True, this balances the equation, but you have changed the substances in it. H2O2 is a completely different substance from H2O.

2. You cannot place a coefficient in the middle of a formula.

The coefficient goes at the beginning of a formula, not in the middle, as in:

H2 + O2 H22O

Water only comes as H2O and you can only use whole formula units of it.

There is another thing you should avoid. Make sure that your final set of coefficients are all whole numbers with no common factors other than one. For example, this equation is balanced:

4 H2 + 2 O2 4 H2O

However, all the coefficients have the common factor of two. Divide through to eliminate common factors like this.

Example 9 – Balance this equation

H2 + Cl2 HCl

Solution 9

Remember that the rule is: A balanced equation MUST have EQUAL numbers of EACH type of atom on BOTH sides of the arrow.

The correctly balanced equation is:

H2 + Cl2 2 HCl

Placement of a two in front of the HCl balances the hydrogen and chlorine at the same time.



Support Questions

1. Recreate these equations in your notebook – DO NOT WRITE IN THIS BOOKLET. Balance the following reactions, and then classify the type of reaction that has occurred:

a. ____ NaBr + ____ H3PO4 ____ Na3PO4 + ____ HBr

Type of reaction: ____________________

b. ____ Ca(OH)2 + ____ Al2(SO4)3 ____ CaSO4 + ____ Al(OH)3

Type of reaction: ____________________

c. ____ Mg + ____ Fe2O3 ____ Fe + ____ MgO

Type of reaction: ____________________

d. ____ C2H4 + ____ O2 ____ CO2 + ____ H2O

Type of reaction: ____________________

e. ____ PbSO4 ____ PbSO3 + ____ O2

Type of reaction: ____________________

f. ____ NH3 + ____ I2 ____ N2I6 + ____ H2

Type of reaction: ____________________

g. ____ H2O + ____ SO3 ____ H2SO4

Type of reaction: ____________________

h. ____ H2SO4 + ____ NH4OH ____ H2O + ____ (NH4)2SO4

Type of reaction: ____________________The Solubility Rules



Solubility is a measure of the amount of a substance that dissolves in water at a given temperature and pressure. A substance is considered soluble if it has solubility greater than 0.1 mol/L, this is often indicated in the reaction as aqueous (aq). If a solute has solubility less than 0.1 mol/L then it is considered insoluble. Often when two substances are reacted and an insoluble product is formed, a precipitate will form, which is the sudden appearance of a solid in a liquid. A precipitate is indicated in the reaction as a solid (s).

Solubility Rules for Ionic Compounds

Anions Cations Solubility of Compoundsmost alkali ions (Li+, Na+, K+, Rb+,

Cs+, Fr+)soluble

most hydrogen ion, H+ solublemost ammonium ion, NH4

+ solublenitrate (NO3

-) most solubleacetate (C2H3O2

-) Ag+ low solubility

most others solublechloride, Cl-bromide, Br-

iodide, I-

Ag+, Pb2+, Hg22+, Cu+, TI+ low solubility

all others soluble

sulphate (SO42-) Ca2+, Sr2+, Ba2+, Pb2+, Ra2+ low solubility

all others soluble

sulphide, S2- alkali ions, H+(aq), NH4+, Be2+, Mg2+, Ca2+, Sr2+, Ba2+, Ra2+

soluble

all others low solubility

hydroxide (OH-) alkali ions, H+(aq), NH4

+, Sr2+, Ba2+, Ra2+, TI+

soluble

all others low solubilityphosphate (PO4

3-)carbonate (CO3

2-)sulphite (SO3

2-)

alkali ions, H+ (aq), NH4

+ soluble

all others low solubility

Formation of Precipitates

A precipitate can result from the formation of a new insoluble compound. A precipitate will form when positive ions and negative ions that make up a compound of low solubility are added together from two different sources. For example, if acetate (negative ion) is mixed with silver (positive ion), a precipitate will form.



Ionic Equations

Ionic equations show the ions that are present in the reactions that produce precipitates. They are used to describe the dissociation of a solute into ions in water.

These equations consist of ;

$ Spectator Ions (ions that do not form the precipitate)$ Participatory Ions (ions that do form the precipitate)

Net Ionic Equations are written to show the precise reaction that forms the precipitate.

Example 10 – Write the net ionic equation for the following reaction:

Potassium chromate + Silver nitrate Silver chromate + Potassium nitrate

Solution 10

Molecular Equation (don’t forget to balance your equation!!)

K2CrO4(aq) + 2AgNO3(aq) Ag2CrO4(s) + 2KNO3(aq)

Total Ionic Equation

2K+1(aq) + (CrO4)-2

(aq) + 2Ag+1(aq) + 2(NO3)-1

(aq) Ag2CrO4(s) + 2K+1(aq) + 2(NO3)+1

(aq)

The potassium (K+) and nitrate (NO3-) ions cancel each other out since they are soluble

(aqueous) both as reactants and products. The ions remaining should be the ions that form the precipitate. This is the net ionic equation.

Net Ionic Equation

2Ag+1(aq) + (CrO4)-2

(aq) Ag2CrO4(s)

Example 11 – Silver nitrate was reacted with calcium chloride. Determine the net ionic equation for this reaction.


Spectator Ions(do not form precipitate)


Solution 11

First complete the word equation.

Silver Nitrate + Calcium Chloride Silver Chloride + Calcium Nitrate

Then complete the balanced molecular equation. Check your solubility table for precipitates. Using the solubility table, silver chloride will form a precipitate.

2AgNO3(aq) + CaCl2(aq) 2AgCl (s) + Ca(NO3)2(aq)

Now, dissociate your aqueous ions. The solid will not form ions.

2Ag+(aq) + 2NO3

-(aq) + Ca2+

(aq) + 2Cl-(aq) 2AgCl(s) + Ca2+(aq) + 2NO3

-(aq)

Next, cancel your ions that are aqueous on both sides (spectator ions). If correct, you will have equal amounts. Do not cancel participatory ions.

2Ag+(aq) + 2NO3

-(aq) + Ca2+

(aq) + 2Cl-(aq) 2AgCl(s) + Ca2+(aq) + 2NO3

-(aq)

This results in your net ionic equation.

2Ag+(aq) + 2Cl-(aq) 2AgCl(s)

Qualitative Chemical Analysis

If the chemicals in a solution are not known, qualitative chemical analysis can be performed to determine which ions are present in the solution.

Example 12

We suspect that a solution contains either silver nitrate or potassium nitrate. How can we determine which of the two compounds it contains?

Solution 12

We know that silver chloride is insoluble (solubility rules) and potassium chloride is soluble. If we can react the unknown solution with a compound containing chloride we can determine which of the two the unknown sample actually contains.

We add silver nitrate to the sodium chloride.

Sodium Chloride + Silver Nitrate Sodium Nitrate + Silver Chloride

Looking at the products above and referring to the solubility rules, it can be determined that silver chloride would form a precipitate in solution.



Now let’s look the reaction of sodium chloride with potassium nitrate.

Sodium Chloride + Potassium Nitrate sodium nitrate + potassium chloride In this scenario, both products (sodium nitrate and potassium chloride) are aqueous in solution.

Conclusion – if a precipitate forms, then the unknown solution must be silver nitrate. If no precipitate forms, then the solution must be potassium nitrate! Chemists use solubility to test for the presence of specific ions, and also to remove ions from solution. If ions are to be removed from a solution, they are reacted to form a precipitate and then the precipitate is removed using filter paper or a centrifuge (spinning).

Support Questions

2. Show the total ionic and net ionic forms of the following equations. If all species are spectator ions, please indicate that no reaction takes place. Note. You need to make sure the original equation is balanced before proceeding.

a. AgNO3 (aq) + KCl (aq) AgCl (s) + KNO3 (aq)

b. Mg(NO3)2 (aq) + Na2CO3 (aq) MgCO3 (s) + NaNO3 (aq)

Key Question #4

1. Recreate this chart in your notes and balance the following equations. Classify the reaction type as synthesis, decomposition, single displacement, or double displacement. (10 Marks)

Equation Reaction Type Ag + S Ag2S

Cl2 + AlBr3 Br2 + AlCl3

KNO3 KNO2 + O2

NH3 N2 + H2

Na2S + Pb(NO3)2 NaNO3 + PbS



2. When sodium hydroxide and copper (II) sulphate are reacted, a blue precipitate forms. The unbalanced equation for this reaction is: (10 marks)

NaOH + CuSO4 Na2SO4 + Cu(OH)2

Write the;

a. Balanced equation (including status of matter of products and reactants)

b. Net ionic equation


SCH4CGrade 12

College Chemistry

Support Question AnswersFor

Lessons 1 – 4

SCH4C – Chemistry Support Question Answers

Lesson 1

1. Classify the following as a qualitative or quantitative observation.

a. qualitative (stated in words) – but may argue for quantitative due to “five”b. quantitativec. quantitatived. qualitativee. quantitativef. quantitativeg. qualitativeh. quantitativei. qualitativej. quantitative

2. For his science experiment, Mohammed decided to burn a candle, and wrote the following statements down. Classify the statements as an observation or an inference.

a. observationb. inferencec. observationd. observatione. inferencef. observationg. observation

3. State the number of significant digits each of following has:

a. 4 b. 2 c. 2d. 4e. 5f. 8

4. The mass of empty graduated cylinder was found to be 23.1g. The mass of the graduated cylinder filled with water has 32 g. Calculate the mass of the water. State your answer to the correct number of significant figures.

32 – 23.1 = 9 (1SF)



Lesson 2

1. Use your periodic table to complete this table:

Element Atomic Number

Mass Number

Number of Protons

Number of Electrons

Number of Neutrons

Hydrogen 1 1 1 1 0Helium 2 4 2 2 2Lithium 3 7 3 3 4Beryllium 4 9 4 4 5Boron 5 11 5 5 6Carbon 6 12 6 6 6Nitrogen 7 14 7 7 7Oxygen 8 16 8 8 8Fluorine 9 19 9 9 10Neon 10 20 10 10 10Sodium 11 23 11 11 12Magnesium 12 24 12 12 12Aluminum 13 27 13 13 14Silcon 14 28 14 14 14Phosphorus 15 31 15 15 16Sulfur 16 32 16 16 16Chlorine 17 35 17 17 18Argon 18 40 18 18 22Potassium 19 39 19 19 20Calcium 20 40 20 20 20

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.

IA VIIIAH)1e- IIA IIIA IVA VA VIA VIIA He)2e-

Li)2e-)1e- Be)2e-)2e- B)2e)3e- C)2e-)4e- N)2e-)5e- O)2e-)6e- F)2e-)7e- Ne)2e-)8e-

Na)2e-)8e-)1e-

Mg)2e-)8e-)2e- Al)2e-)8e-)3e- Si)2e)8)4e- P)2e-)8e-)5e- S)2e-)8e-)6e- Cl)2e-)8e-)7e- Ar)2e-)8e-)8e-

K)2e-)8e-)8e-) 1e-

Ca)2e-)8e-)8e-)2e-

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



Lesson 3

1. Complete the following table.

Element Atomic Symbol Number of valenceelectrons

Valence

Oxygen O 6 -2Chlorine Cl 7 -1Sodium Na 1 +1Phosphorus P 5 -3

2. Complete the following table. You may use the shorthand method for drawing Bohr Diagrams.

Element Bohr Atom Bohr Ion Lewis Atom Lewis IonLithium Li)2e-)1e- [Li)2e-]+1

Magnesium Mg)2e-)8e-)2e- [Mg)2e-)8e]+2

Oxygen O)2e-)6e- [O)2e-)8e-]-2



3. Complete the following table. Use Lewis structures to depict atoms.

Bond formation Name of compound

Chemical formula

Anion Cation

lithium and fluorine

Lithium fluoride

LiF fluorine lithium

calcium and phosphorus

Calcium phosphide

Ca3P2 phosphorus calcium

4. Complete the table below. If the molecule is covalent, indicate if it is polar covalent or not.


NaCl Na =0.9 Cl=3.0 2.1 Polar covalent

Cl2 Cl =3.0 Cl =3.0 0 covalent

HF H = 2.1 F = 4.0 1.9 Polar covalent

5. Compare and contrast covalent and polar covalent compounds.

Similar DifferencesBoth involve sharing of electronsBoth contain mostly non-metals

Covalent compounds share electrons equally, whereas polar covalent molecules do not share electrons equally



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

a. Hydrogen bondingb. London forcesc. Dipole-Dipole

Lesson 4

1. Balance the following reactions, and then classify the type of reaction that has occurred.

a. 3 NaBr + 1 H3PO4 1 Na3PO4 + 3 HBr

Type of reaction: double displacement

b. 3 Ca(OH)2 + 1 Al2(SO4)3 3 CaSO4 + 2 Al(OH)3


c. 3 Mg + 1 Fe2O3 2 Fe + 3 MgO

Type of reaction: single displacement

d. 1 C2H4 + 3 O2 2 CO2 + 2 H2O

Type of reaction: combustion

e. 2 PbSO4 2 PbSO3 + 1 O2

Type of reaction: decomposition

f. 2 NH3 + 3 I2 1 N2I6 + 3 H2


g. 1 H2O + 1 SO3 1 H2SO4

Type of reaction: decomposition

h. 1 H2SO4 + 2 NH4OH 2 H2O + 1 (NH4)2SO4




2. Show the total ionic and net ionic forms of the following equations. If all species are spectator ions, please indicate that no reaction takes place. Note! You need to make sure the original equation is balanced before proceeding.

a. AgNO3 (aq) + KCl (aq) AgCl (s) + KNO3 (aq)

Total Ionic: Ag+ (aq) + NO3¯ (aq) + K+ (aq) + Cl¯ (aq)

AgCl (s) + K+ (aq) + NO3¯ (aq)

Net Ionic: Ag+ (aq) + Cl¯ (aq) AgCl (s)

b. Mg(NO3)2 (aq) + Na2CO3 (aq) MgCO3 (s) + 2 NaNO3 (aq)

Total Ionic: Mg2+ (aq) + 2 NO3¯ (aq) + 2 Na+ (aq) + CO32- (aq)

MgCO3 (s) + 2 Na+ (aq) + 2 NO3¯ (aq)

Net Ionic: Mg2+ (aq) + CO32- (aq) MgCO3 (s)


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