SCH3U- UNIT 1 REVIEW: MATTER AND CHEMICAL BONDING

THE PERIODIC TABLE

When Dmitri Mendeleev arranged the first Periodic Table (~1905), its elements were arranged according to periodicity: patterns that repeat at definite intervals

- in chemistry periodicity is the occurrence of similar physical and chemical properties of elements at regular intervals

• elements arranged in order of increasing atomic number show a periodic repetition of properties

• each column of the periodic table (called a group) has similar properties (periodic trends)

• periodic law: many of the physical and chemical properties of elements tend to recur in a systematic manner, when arranged by increasing atomic number

- on the Periodic Table, each element has a

mass number: protons + neutrons

atomic number: protonsatomic symbol

Furthermore, the elements on the Periodic Table are divided into three categories:1. Metals

• Alkali metals (group 1)• Alkaline earth metals

(group 2)• Transition metals (rest)• Inner transition metals

(periods 6 and 7: Lanthanoids and Actinoids)

– left side and centre of periodic table– Solid at room T (except mercury, Hg)– Silver (except Cu and Au)– Shiny, conduct electricity and heat,

malleable and ductile

2. Metalloids• Along zig-zag line dividing metals and non-metals

• Metalloids are: Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium

• Have some metallic and non-metallic properties

• Ex. Si - silicon: shiny, nonmalleable, a semi-conductor

3. Non-metals• Halogens (Group 7)• Noble gases (Group 8)

– Right side of periodic table1

x atomic symbol

Amass number (atomic mass)

zatomic number

– Found in all three states– Variety of colours– Poor conductors of heat/electricity– Usually brittle (having little elasticity: easily cracked or fractured or snapped )

TRENDS IS THE PERIODIC TABLE

1. Atomic Radius- distance measured from the centre of a nucleus to its outermost electrons- increases down a group (due to increase in the number of energy levels), decreases across a period, left to right (increase in the pull of the protons on the electrons)

2. Electronegativity- the measure of an atom’s ability to attract electrons in a chemical bond , relates to its tendency to gain electrons: the stronger the pull of nucleus on surrounding the greater its tendency to gain e- increases up a group, increases across a period from left to right (not including Noble Gases) why?

3. Ionization Energy- energy reguired to remove the outermost electron from a gaseous atom- increases up a group, increases across a period, left to right why?: look at Atomic Radius

CHEMICAL COMPOUNDS AND BONDING

Based on their physical properties, compounds are classified into two groups: ionic compounds or covalent compounds (also called molecules).

PROPERTY IONIC COMPOUND COVALENTstate at room temperature crystalline solid Liquid, gas, or solidmelting point High Lowelectrical conductivity in water

Yes No

solubility in water most have high solubility most have low solubility

1. Ionic Compounds- form as a result of transfer of electrons from a metal to a non-metal- the metal forms a positively charged cation, the non-metal forms a negatively charge anion- held together by electrostatic forces of attraction

2. Covalent Compounds (Molecules)- form as a result of sharing of valence electons between two non-metals

Imagine that you are a chemist. A colleague has just carried out a series of tests on the following compounds: ethanol, carbon tetrachloride, glucose, table salt, water, and potassium permanganate.

You take the results home to organize and analyze them only to discover that your colleague has labelled each test by sample number

only and has forgotten to write down which compound corresponds to each sample number. You realize however, that you can use the physical properties of the compounds to help identify them.

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Thought Lab

IncreasingElectronegativity

IncreasingIonization Energy

Procedure: Complete the following table.

Sample

Compound Name

Dissolves in water

Conductivity

Melting

point

Appearance Covalent or Ionic

1 Yes High 8010C Clear, white crystalline solid

2 Yes Low 0.00C Clear, colourless liquid

3 Yes High 2400C Purple, crystalline solid

4 Yes Low 1460C White powder

5 No Low -230C Clear, colourless liquid

6 Yes Low -1140C Clear, colourless liquid

1. Based on what you know about the properties of ionic and covalent compounds, decide which compound corresponds to each set of properties in your table.

2. Examine the properties of each compound and decide whether it is an ionic compound or a covalent compound.

Analysis:1. Write down the reasoning you used to identify each compound, based on the properties

given. 2. Write down the reasoning you used to decide whether each compound was ionic or

covalent.3. Were you unsure how to classify any of the compounds? Which ones, and why?4. Which property is most useful for deciding whether a compound is ionic or covalent?

WHAT IS BONDING?

Chemical bonding: involves the interaction between the valence electrons of atoms.

Different atoms are interacting or reacting in order to achieve greater stability. In chemistry, stability is affected by the number of valence electrons. Noble gases are the most stable atoms, because they have a full valence shell: they have eight electrons on their outer most energy shell. This explains why they are so non-reactive. This desire to achieve a full outer valence shell is referred to as the OCTET RULE.

OCTET RULE: atoms bond in order to achieve an electron configuration that is the same as the electron configuration of a noble gas (a full valence shell).

ISOELECTRONIC atoms are two atoms or ions that have the same electron configurations. For example, Cl- is isoelectronic with Ar because they both have 18 electrons and a filled outer energy shell.

TYPES OF CHEMICAL BONDING

There are TWO TYPES OF CHEMICAL BONDING:

IONIC BONDING: involves the transfer of electrons from a metal to a non-metal.- forms ionic compounds

COVALENT BONDING: involves the sharing of electrons between two non-metals.- forms covalent compounds (or molecules)

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PREDICTING BOND TYPE

Recall, electronegativity is the measure of an atom’s ability to attract electrons in a chemical bond.

Electronegativity is a periodic property of atoms involved in chemical bonds. In general, electronegativity increases from left to right across the period and decreases down the groups.

You can use the differences in electronegativity to decide whether the bond between two atoms is ionic or covalent.

Difference in EN

3.3 1.7 0.5 0

Mostly Ionic Polar Covalent Mostly

Covalent

Thus we can offer a new definition for types of chemical bonds: Ionic Bonds: have an EN difference greater than 1.7

Covalent Bonds: have an EN difference less than 0.5

Polar Covalent Bonds: have an EN difference between 1.7 and 0.5

POLAR COVALENT BONDS: are covalent bonds that have polar properties due to the unequal sharing of valence electrons. In polar covalent bonds, electrons spend a greater amount of time around the atoms with the higher electronegativity, resulting in partial charges,

IONIC BONDING

Metals have a low electronegativity while non-metals have a high electronegativity. Thus, when metals and non-metals react, electrons are transferred from metals to non-metals.

When atoms loose electrons they form positively charged ions, called cations. When atoms gain electrons they form negative ions, called anions.

Using Lewis-Dot diagrams to represent IONIC BONDING

STEP 1: Draw the Lewis-dot diagram for each atom.STEP 2: Determine the number of each atom required to achieve a full octet for the most electronegative atom.STEP 3: Draw the resulting ions in square brackets. Include the charge of each ion. ‥ ‥

Na• + •Cl: [Na]+ [:Cl:]-

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¨ ¨ PRACTICE PROBLEMS Determine the electronegativity difference for each bond shown.

Indicate whether the bond is ionic or covalent. Draw the Lewis Structure for each compound.

O – HC – HMg – Cl

B – FCr – OC – N

Na – INa – BrBa – O

Li – BrCa – OK – C

COVALENT BONDING

When electronegativity differences are small (less than 1.7) instead of transferring electrons, electrons are shared between atoms.

A COVALENT BOND consists of a pair of shared electrons, where each atom contributes one electron to the bond.

One pair of electrons= single bondTwo pairs of electrons = double bondThree pairs of electrons = triple bond

PURE COVALENT BONDS result from the equal sharing of electrons between atoms. Atoms that bond in this way are called DIATOMIC ELEMENTS, (Cl2, H2, F2, O2, N2, Br2 , I2,).

A lone pair of electrons consists of two electrons that do no participate in bonding.

PRACTICE PROBLEMS

1. Using Lewis dot diagrams, for covalent compounds, show the formation of a covalent bond between two atoms of each of the following.

a. Iodine and Iodineb. Bromine and

Brominec. Hydrogen &

Hydrogend. Oxygen and Oxygen

e. Fluorine and Fluorinef. Hydrogen and

oxygeng. Chlorine and oxygenh. Carbon and

hydrogen

i. Iodine and hydrogenj. Nitrogen and

rubidiumk. Nitrogen & Hydrogenl. Oxygen and Nitrogen

2. Use Lewis structures to show how the following elements form covalent bonds.a. One silicon and two oxygen atomsb. One carbon, one hydrogen, and three chlorine atomsc. Two nitrogen atoms

3. In general, “the further away two atoms are from each other on the periodic table, the more likely they are to participate in ionic bonding”. Do you agree with this statement? Explain why or why not.

POLAR COVALENT BONDS

Recall, covalent compounds display a wide variety of physical properties; they can be solid, liquid or gas at room temperature. How can we account for such a wide variety of differences?

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PRACTICE PROBLEMS

1. Predict whether each bond will be covalent, polar covalent, or ionic.

a. C—F e. Na—Fb. Cl—Cl f. Fe—Oc. O—N g. Mn—Od. Cu—O h. Si—H

Many of these properties can be explained in part by the electronegativity difference between the atoms. In other words, by looking at what type of covalent bond actually exists.

POLAR COVALENT BONDS result when two bonding atoms have an electronegativity difference greater than 0.5, but less than 1.7. The atom with the higher electronegativity will hold onto the bonding pair of electrons for more of the time, resulting in an unequal sharing.

This unequal sharing, results in slight charge on both atoms. The atom with the higher electronegativity (the one that has the pair of electrons for a greater amount of the time) will develop a slight negative charge, δ- and the less electronegative atom will develop a slight positive charge, δ+

δ+ δ-

O H H Cl

δ- δ+

These partial charges have great consequence on the physical and chemical properties of the compound. The most important consequence of these polar bonds is the creation of strong intermolecular attractions called dipole-dipole interactions.

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An Aside: DIFFERENT TYPES OF MOLECULAR MODELS

Different models can communicate different aspects of chemical bonding. Below are four different types of molecular models that you should become familiar with.

Lewis Structure Structural Diagram

Ball-and-stick Space-filling model

Shows you exactly how many electrons are involved in each bond in a compound.

Shows single bonds as single lines, and multiple bonds as multiple lines.

Shows atoms as spheres and bonds as sticks. It accurately shows how the bonds within a molecule are oriented in three-dimensional space. This type of model usually also shows the shape of the molecule.

Shows atoms as spheres. It is the most accurate representation of the shape of a real molecule.

POLAR BONDS AND MOLECULAR SHAPE* don’t forget to review intermolecular bonding from unit 3!!!*The water molecule serves as a great model molecule for polar covalent compounds. In water, each oxygen-hydrogen bond is polar, since each oxygen develops δ – charge and every hydrogen develops a δ+ charge.

The overall three-dimensional shape of the molecule also plays a significant role in determining whether or not a molecule is polar. The slight bent shape of the water molecule causes these charges to be asymmetrical or on opposite ends of the molecule; it is a polar molecule.

On the other hand, carbon dioxide is a straight symmetrical molecule. So although it contains polar bonds, the direction of the polarity cancels each other out; it is a non-polar molecule.

δ_ δ+ δ-

WRITING CHEMICAL FORMULA FOR IONIC COMPOUNDS

Since water is made up of polar molecules with positive and negative ends that attract one another, water tends to “stick” to itself. This means that it has a high melting point and boiling point, relative to other covalent compounds.

Lewis dot diagrams are useful for: keeping track of electron transfers in bonding and for making sure that the octet rule is

obeyed. determining the ratio of the atoms in the compound.

Another way of communicating the ratio of atoms in a chemical compound is to use a chemical formula. A chemical formula provides two important pieces of information:

1. the elements that make up the compound2. the number of atoms of each element that are present in a compound.

GENERAL RULES FOR WRITING CHEMICAL FORMULAS1. The symbol for the least electronegative (or metal) atom or ion is listed first.2. subscripts are used to indicate the number of each atom present in the compound.

Do not write the subscript “1” (it is

implied)

Symbol of positive ion CaBr2

Subscript indicating that there are TWO bromine ions.

Symbol of negative ion

Is there an easier way of determining the chemical structure of a compound without having to draw Lewis Dot diagrams?

Yes. Every element has a certain capacity to combine with other atoms. We already said that metals tend to produce stable positive (cations) ions. Why?

When we are considering ionic bonding, we need to consider the stable ions that are produced when bonding takes place.

If we look at the valence electrons of the alkali metals (Group 1) we will notice that all atoms in group 1 have one valence electron. In order to achieve a stable octet, each alkali metal would either have to lose one electron or gain seven electrons. Since it is much easier to lose one electron than it would be to gain seven, the alkali metals form positive ions with a charge of +1.

If we were to look at the valence electrons of Group 17 (the halogens) we would notice that all the atoms in Group 17 have 7 valence electrons. Since it is much easier to gain one electron than it would be to lose seven, the halogens form negative ions with a charge of -1.

When we are considering covalent bonding, the valence of the non-metals tells you how many electrons the atom must find to share. This also corresponds to the number of bonds it must form in order to share the required number of electrons.

FORMING BINARY IONIC COMPOUNDSBinary Compounds are inorganic compounds that contain only two elements.When writing the chemical formula for ionic compounds we can use the valence (the charge of the stable ion) to form the compound. The zero sum rule states that “for a neutral chemical formula containing ions, the sum of the positive charges plus the negative charges of the atoms in a compound must equal zero”.1. Write the symbols for the atoms;

remember that the metal is written first, Mg Cl

the non-metal is written second.2. place the valence charges of each atom

on top of the appropriate symbol 2+ 1-

Mg Cl

3. Using arrows bring the numbers (without the signs) down and write as a subscript by criss-crossing over.

2+ 1-

Mg Cl

4. Check the subscripts. Any “ones” can be removed. Reduce subscripts to their lowest terms. For example, Mg2O2 becomes MgO

MgCl2

WRITING FORMULAS WITH MULTIVALENT IONSMany of the transitional metals and non-metals can form more than one stable ion. For example, Copper can form Cu1+ or Cu2+ ions. When forming compounds it is important to determine which ion is forming the compound and to respect the charges of that particular ion. Therefore, when naming copper-containing compounds you need to state whether it’s copper (I) or copper (II).

WRITING FORMULAS WITH POLYATOMIC IONSPolyatomic Ions are ions that are made up of more than one atom. In fact, polyatomic ions are charged molecules. The charge associated with a polyatomic ion belongs to the whole molecule and thus behave as a single unit and should be treated as a single ion.

1. Write the symbols for the atom and polyatomic ion; remember that the metal is written first, the polyatomic ion is written second. Place brackets around any polyatomic ion present.

K (PO4)

2. place the valence charges of each atom on top of the appropriate symbol

1+ 3- K (PO4)

3. Using arrows bring the numbers (without the signs) down and write as a subscript by criss-crossing over.

1+ 3-

K (PO4)

4. Check the subscripts. Any “ones” can be removed. Reduce subscripts to their lowest terms. Remove any unneeded brackets

K3PO4

NAMING IONIC (metal + non-metal) COMPOUNDS: BINARY, MULTIVALENT, & POLYATOMIC

Chemical nomenclature is the system that is used in chemistry for naming chemical compounds. In 1919, the International Union of Pure and Applied Chemistry (IUPAC) was formed by a group of chemists. The main aim of the group was to establish an international standard for masses, measurements, chemical symbols, and chemical names.

NAMING BINARY IONIC COMPOUNDS1. The metal in a binary compound is always named first by using the name of the element.2. The name of the non-metal is written by changing the ending of the name to “ide”.

For example, NaCl would be named sodium chloride.

NAMING MULTIVALENT IONIC COMPOUNDS1. The multivalent element’s name is followed by the valence written in Roman numerals in

brackets.For example: Sn4+ is tin(IV) and Fe3+ is iron(III)

2. The name of the non-metal is written by changing the ending of the name to “ide”.For example, the compound of Sn4+ and oxygen (Sn2O4 SnO2) would be named tin(IV) oxide

NAMING COMPOUNDS WITH POLYATOMIC IONSIonic compounds formed from polyatomic ions are called tertiary compounds because they contain three elements.

1. The metal in a binary compound is always named first by using the name of the element.2. The name of the polyatomic ion is written.

For example, NaCl would be named sodium chloride.When naming polyatomic ions, prefixes and suffixes are sometimes used:

hypo ite 2 less oxygen atoms

ite 1 less oxygen atom

ate the base polyatomic ion

per ate 1 extra oxygen atom

NAMING COVALENT COMPOUNDS (two non-metals)

In order to indicate that a binary compound is made up of two non-metals (that it is a molecular compound), a prefix is added to both non-metals in the compound. This prefix indicates the number of atoms of each element in one molecule or formula unit of the compound.

For example, P2O5 is named diphosphorous pentoxide

NOTICE: the ending of the name of the second non-metal is changed to “ide” and the prefix mono is left out when there is only one atom of the first element in the name.

For example. CO2 is NOT monocarbon dioxide, but rather carbon dioxide.

Numerical Prefixes for molecular binary compounds:1 2 3 4 5 6 7 8 9 10

mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca-

PRACTICE PROBLEMS The following pages are worksheets for you to complete. Answers to these worksheets are available upon request.

IONIC COMPOUNDS: Names and Formulas

1. Write the formulas for the following compounds.

a. magnesium oxide ___________________ k. tin (II) fluoride _________________

b. aluminum nitride ___________________ l. lead (IV) nitride __________________

c. potassium sulfide __________________ m. iron (III) chloride _________________

d. calcium bromide ___________________ n. copper (I) oxide _________________

e. aluminum sulfide ___________________ o. antimony (III) sulfide _______________

f. beryllium oxide __________________ p. mercury (II) oxide _________________

g. strontium phosphide ________________ q. tin (IV) iodide ___________________

h. sodium fluoride __________________ r. arsenic (III) phosphide ______________

i. lithium selenide _________________ s. cobalt (II) sulphide _________________

j. barium oxide __________________ t. tin (IV) sulphide ___________________

2. Write the names for the following compounds.

a. Li2O _____________________________ k. PbS _____________________________

b. AlCl3 ____________________________ l. SnO2 _____________________________

c. MgS _____________________________ m. NiO _____________________________

d. CaF2 ____________________________ n. CuI2 _____________________________e. Al2O3 _____________________________ o. PbCl4 _____________________________

f. BeF2 _____________________________ p. FeP ______________________________

g. K3P ______________________________ q. AuBr3 _____________________________

h. Mg3P2 ____________________________ r. Hg2S _____________________________

i. CaO ______________________________ s. SbF3 _____________________________

j. Ag2S _____________________________ t. MnO2 _____________________________POLYATOMIC COMPOUNDS:

Names and Formulas

3. Write the formulas for the following compounds.

a. magnesium carbonate ________________ k. tin (II) chlorate _________________

b. aluminum nitrate __________________ l. lead (IV) nitrate __________________

c. potassium sulfate __________________ m. iron (III) carbonate _________________

d. calcium chlorate __________________ n. copper (II) hydroxide ________________

e. aluminum sulfate____________________ o. lead (II) nitrate ___________________

f. sodium carbonate __________________ p. mercury (II) chlorate ________________

g. strontium phosphate ________________ q. tin (IV) phosphate _________________

h. sodium chlorate __________________ r. lead (IV) hydroxide _________________

i. lithium nitrate __________________ s. potassium nitrate _________________

j. aluminum hydroxide __________________ t. tin (IV) sulphate ___________________

4. Write the names for the following compounds.

a. Li2SO4 ___________________________ k. PbSO4 ___________________________

b. Al(ClO3)3 __________________________ l. AuOH ___________________________

c. MgSO4 ___________________________ m. GaPO4 ___________________________

d. K2CO3 ____________________________ n. Cu(NO3) __________________________e. Na2SO4 ___________________________ o. Pb(ClO3)4 _________________________

f. AgNO3 __________________________ p. Fe(ClO)3 __________________________

g. K3PO4 ____________________________ q. Au2CO3 ___________________________

h. Sr(ClO3)2 __________________________ r. HgOH ____________________________

i. RbOH _____________________________ s. Sb2(SO4)3 __________________________

j. HClO3 ____________________________ t. MnSO4 __________________________

MOLECULAR COMPOUNDS:Names and Formulas

5. Write the formulas for the following compounds.

a. carbon dioxide __________________ k. diphosphorus trisulphide ____________

b. silicon dioxide __________________ l. dinitrogen monoxide _______________

c. water __________________ m. dichlorine monoxide _______________

d. carbon disulphide __________________ n. bromine gas ________________

e. sulphur trioxide __________________ o. carbon monoxide ________________

f. carbon tetrachloride ________________ p. xenon tetrafluoride ________________

g. sulphur dioxide __________________ q. neon gas ________________

h. dinitrogen tetroxide _________________ r. silicon tetrahydride ________________

i. nitrogen monoxide _________________ s. iodine heptachloride _______________

j. arsenic tribromide __________________ t. krypton difluoride _________________

6. Write the names for the following compounds.

a. CF4 _____________________________ k. NF3 _____________________________

b. NH3 ____________________________ l. P2S5 _____________________________

c. PBr3 _____________________________ m. PF5 _____________________________

d. F2 gas _____________________________ n. ICl _____________________________e. CS2 _____________________________ o. SeCl2 ____________________________

f. CO ______________________________ p. Cl2O _____________________________

g. SiC ______________________________ q. AsBr3 ____________________________

h. N2O4 ____________________________ r. H2S ______________________________

i. P2O5 _____________________________ s. B2H8 ____________________________

j. SF4 ______________________________ t. TeCl2 ____________________________

PUTTING IT ALL TOGETHER:Names and Formulas

7. Write the formulas for the following compounds.

a. calcium fluoride _______________ k. potassium sulphate _______________

b. carbon disulfide _______________ l. barium nitride _______________

c. nitrogen triodide _______________ m. aluminum hydroxide ______________

d. sodium phosphide _______________ n. fluorine gas _______________

e. dichlorine monoxide _______________ o. silicon dioxide _______________

f. iron (III) carbonate _______________ p. calcium hydroxide _______________

g. manganese (IV) sulphate ______________ q. xenon gas _______________

h. diphosphorus pentasulphide __________ r. gold (I) nitrate _______________

i. tin (IV) chloride _______________ s. sulphur trioxide _______________

j. magnesium chlorate _______________ t. iron (II) phosphate _______________

8. Write the names for the following compounds.

a. CCl4 _____________________________ k. NaNO3 ___________________________

b. Mg(ClO3)2 ________________________ l. PCl5 _____________________________

c. PBr3 _____________________________ m. BiF5 ____________________________

d. H2 gas ____________________________ n. Al(ClO3)3 _________________________

e. PbS2 _____________________________ o. FeCl2 ____________________________

f. Al2(CO3)3 _________________________ p. N2O ____________________________

g. Na2SO4 __________________________ q. CuClO3 ___________________________

h. Na2O ____________________________ r. Li3PO4 __________________________

i. Al2(SO4)3 __________________________ s. SnO ____________________________

j. Li2SO4 _____________________________ t. SeCl2 ____________________________

CHEMICAL REACTIONS

A chemical equation describes what happens when elements and compounds interact with one another to form new substances. These interactions that involve the production of a new substance are called chemical reactions.

A substance that undergoes a chemical reaction is called a reactant. A substance that is formed in a chemical reaction is called a product.

The general structure of a chemical equation is to write all reactants on the left side of the equation and all products on the right side; reactants and products are separated by an arrow, which stands for “yield” or “produces”.

Reactants Products

Chemical equations come in different forms; each condenses a large amount of information into a short, concise statement.

I. WORD EQUATIONS - identifies the reactants and products of a chemical reaction by name.

sodium + chlorine sodium chloride

In this equation, “+” means “reacts with” and “” means “to form/yield”.

II. SKELETAL EQUATIONS - identifies the reactants and products of a chemical reaction by replacing chemical names with chemical formulas. A skeletal equation also shows the state of each reactant and product in the chemical equation.

Na(s) + Cl2(g) NaCl(s)Skeletal equations, although simple and useful, do not fully describe chemical reactions because they are unbalanced. A complete chemical equation accounts for the law of conservation of mass, which states that in any chemical reaction, the mass of the products is always equal to the mass of the reactants. III. BALANCED CHEMICAL EQUATIONS - shows that there is the same number of each kind of atom on both sides of the equation. In other words, a balanced chemical equation reflects the law of conservation of mass.

Consider the following reaction: Na(s) + Cl2(g) NaCl(s)

What has happened to the other chlorine atom ?

In order to balance chemical equations you must add numbers (coefficients) in front of the appropriate formulas so that the number of atoms on the reactant side equals the number of atoms on the product side of the reaction.YOU CANNOT balance a chemical equation by changing ANY of the chemical formulas. The only way to balance a chemical equation is through the use of coefficients.

So the above equations when balanced becomes: Na(s) + Cl2(g) 2 NaCl(s) STEPS FOR BALANCING CHEMICAL EQUATIONSStep 1 Write the word equation.

Step 2 Write out the skeletal equation. Ensure that you have copied all the chemical formulas correctly and respected all the charges and combing capacities of each atom.Make sure to include the states of each compound.

Step 3 Balance the atoms that occur in the highest number. Leave hydrogen, oxygen, and other elements until later.

Step 4 Balance any polyatomic ions that occur on both sides of the equation.

Step 5 Balance any hydrogen and oxygen atoms that occur in a combined and uncombined state

Step 6 Balance any other element that occurs in the uncombined state.

Step 7 Check to make sure that the total number of each atom is equal on both sides of the equation and that the coefficients of the equation are written in lowest terms

Copper(II) nitrate reacts with potassium hydroxide to form potassium nitrate and solid copper(II) hydroxide. Write the balanced chemical equation for this reaction.

STEP 1: Write the word equation:

Copper(II) nitrate + potassium hydroxide potassium nitrate + copper(II) hydroxide

STEP 2 : Write out the skeletal equation. If required, divide this step into two:

In the first part, write all symbols for each element along with their valence charge or combining capacity.

Cu2+ NO3- + K+ OH- K+ NO3- + Cu2+ OH-

In the second part, form the compounds by criss-crossing valence charges.

Cu(NO3)2 (aq) + KOH(aq) KNO3(aq) + Cu(OH)2(s)

STEP 3: Balance the atoms that occur in the highest number first. Leave hydrogen, oxygen, and other elements until later. In the case of this reaction, these are the polyatomic ions.

STEP 4: Balance any polyatomic ions that occur on both sides of the equation.

There are two NO3- ions on the left side, so we will also need two on the right side.Cu(NO3)2 (aq) + KOH(aq) 2 KNO3(aq) + Cu(OH)2(s)

There are two OH- ions on the rights side, so we will also need two on the left side.Cu(NO3)2 (aq) + 2 KOH(aq) 2 KNO3(aq) + Cu(OH)2(s)

STEP 5: Balance any hydrogen, and oxygen atoms that occur in the combined and uncombined state. There are none to consider for this reaction.

STEP 6: Balance any other elements that occurs in uncombined state. There are none to consider.

STEP 7: Check to make sure that the total number of each atom is equal on both sides of the equation and that the coefficients of the equation are written in lowest terms.

Cu(NO3)2 (aq) + 2 KOH(aq) 2 KNO3(aq) + Cu(OH)2(s)

Left Side Right SideCu 1 1NO3 2 2

K 2 2OH 2 2

SAMPLE PROBLEM

WRITING CHEMICAL EQUATIONS

The following elements exist only in molecular form and their formulas need to be memorized:

P4 (s) PhosphorousS8(s) SulfurH2(g) Hydrogen gasO2(g) Oxygen gasN2(g) Nitrogen gasF2(g) Fluorine gasCl2(g) Chlorine gasBr2(l) BromineI2(s) Iodine

Write a chemical equation for each of the following and then balance.

(a) lithium oxide + water lithium hydroxide(b) iron(III) hydroxide iron(III) oxide + water(c) lead(II) nitrate lead oxide + nitrate + oxygen (d) barium oxide + water barium hydroxide(e) calcium + aluminum chloride calcium chloride + aluminum(f) ammonia + nitrogen monoxide nitrogen + water(g) hydrogen chloride + potassium carbonate potassium chloride + bicarbonate(h) gold(III) oxide gold + oxygen(i) iodine + nitric acid iodic acid + nitrogen(j) barium chloride + aluminum sulphate barium sulphate + aluminum chloride(k) ammonium chloride + calcium hydroxide calcium chloride + ammonia + water(l) potassium oxide + carbon dioxide potassium carbonate + oxygen

TYPES OF CHEMICAL REACTIONS

Different types of compounds are formed from different types of reactions. Chemists have classified reactions into different classes based on key characteristics and they use this information to predict the products of chemical reactions.The four major classes of reactions are:

1. SYNTHESIS REACTIONS- two or more elements or compounds combine to form a single new substance.

A + B C

Predicting the product of a synthesis reaction is often difficult and most of the time requires that the reaction be carried out in the lab and the product be isolated and analyzed. Thus, you cannot predict the products of a synthesis reaction with certainty; rather you can only give options.

C(s) + O2(g) CO2(g)2C(s) + O2(g) 2 CO(g)

2. DECOMPOSITION REACTIONS- a large compound breaks down into elements or smaller compounds. Thus, it is the opposite of a synthesis reaction.

C A + BThe simplest type of decomposition reaction involves breaking down a compound into its component elements.

2H2O 2H2 + O2

More complex decomposition reactions occur when compounds break down into other smaller compounds. NH2NO3(s) N2O(g) + 2H2O(q)

3. SINGLE DISPLACEMENT REACTIONS- one element in a compound is displaced (or replaced) by another element. There are two general types of single displacement reactions:

1. Metal replacing a metal:A + BC AC + B

2. non-metal replacing a non-metal:

DE + F DF + E

Most single displacement reactions involve the displacement of a metal by another metal.

Through experimentation, chemists have ranked the relative reactivity of the metals, including hydrogen (in acids and in water), in an ACTIVITY SERIES. The position of the metal on the activity series must be considered in order to determine if a reaction will take place.4. DOUBLE DISPLACEMENT REACTIONS- involves the exchange of cations (most times these are metals) between two ionic compounds, usually in an aqueous solution.Generally, this can be represented by:

AB + CD CB + ADThere are three major types of double displacement reactions. 1. a solid forms (precipitate) - in order to determine if a precipitate will form, you must use the solubility rules to identify any possible insoluble compounds. 2. gas is produced3. water is produced (neutralization)

THE ACTIVITY SERIES OF METALSLi K Ba Ca Na Mg Al Zn Fe Ni Sn Pb H Cu Hg Ag Au

Metal metal ion most reactive least reactiveAny metal to the left will displace a metal ion to the right from a compound.PREDICTING REACTIONS WITH ACIDS (to form H2 gas)

Li K Ba Ca Na Mg Al Zn Fe Ni Sn Pb H Cu Hg Ag Au react with acids do not react with acids

Another important activity series is the HALOGEN ACTIVITY SERIES, which indicates the order of reactivity amongst the halogens from most reactive (fluorine) to least reactive (Iodine).

F Cl Br I

Double displacement reactions can be illustrated using total and net ionic equations:

WRITING TOTAL AND NET IONIC EQUATIONS

Step 1 Identify Type of Reaction and Possible ProductsStep 2 Look Up Solubility of Both ProductsStep 3 Indicate States of Reactants and ProductsStep 4 Write Chemical Equation for ReactionStep 5 Balance EquationStep 6 Write Total Ionic EquationStep 7 Write Net Ionic Equation

ex 1. Write the total ionic equation and net ionic equation for the reaction of barium sulfide with sodium sulfate.Step 1Double displacement reaction:barium sulfide + sodium sulfate barium sulfate + sodium sulfideStep 2barium sulfate: low solubility sodium sulfide: solubleStep 3aqueous barium sulfide + aqueous sodium sulfate solid barium sulfate + aqueous sodium sulfideStep 4BaS(aq) + Na2SO4(aq) BaSO4(s) + Na2S(aq) (unbalanced)Step 5BaS(aq) + Na2SO4(aq) BaSO4(s) + Na2S(aq) (balanced)Step 6Ba2+ (aq) + S2–(aq) + 2 Na+(aq) + SO4

2– (aq) BaSO4(s) + 2 Na+ (aq) + S2– (aq)Step 7Ba2+ (aq) + SO4

2– (aq) BaSO4(s)

PRACTICE PROBLEMSWrite the total and net ionic equations for the following

double displacement reactions:

a) potassium phosphate + magnesium chloride

b) sodium bicarbonate + sulfuric acid

c) ammonium sulfide and lead II nitrate

d) calcium chloride and chromium III nitrate

e) NaCl ( ) + KBr ( ) NaBr ( ) + KCl ( )

f) NaCl ( ) + HOH ( ) NaOH + HCl

g) CaSO4 ( ) + LiF ( ) CaF2 + Li2SO4

h) H2SO4 ( ) + KOH ( )