Science10ChemistryWorkbook2019 (1).doc

Chemistry

: Energy and Matter in Chemical Change

The five major concepts developed in this unit are:

 Matter is classified on the basis of its properties

 Matter has a well-defined underlying structure

 Elements combine to form a vast array of compounds

 Energy is involved in each change that matter undergoes

 Matter is conserved in chemical reactions

Topic one: Understanding Matter

1. Properties of Matter

2. WHMIS symbols

3. Laboratory Safety 4. Structure of atoms Topic Two: Composition of chemical compounds

1. Ionic Compounds

2. Hydrated Ionic Compounds 3. Solubility of Ionic Compounds 4. Acid forming Compounds

5. Molecular Compounds and Elements

Topic Three: Chemical Change

1. Balancing and Interpreting Equations 2. Formation and Decomposition Reactions 3. Single Replacement Reactions

4. Double Replacement Reactions 5. Combustion Reactions

Understanding Matter -

Properties of Matter

Definition: Matter is anything that has mass and occupies space

Definition: All matter has Physical and Chemical properties.

Physical properties: colour, melting point, boiling point, density, lustre, hardness....

Chemical properties: how 'reactive' it is. For example paper will burn when heated to about 3000C. Iron will not, but iron will rust. Chromium will not do either.

Matter is classified according to their physical and chemical properties.

All Matter

Pure Substances Mixtures

Elements Compounds Solutions Mechanical

Definitions:

Pure Substance: One of the two major classifications of matter. They have fixed properties that do not vary if measured under the same conditions.

Mixtures: One of two major classifications of matter. Their properties vary depending upon the ratio and identity of their components.

Mechanical Mixture: A mixture in which the individual parts or components remain intact and visible. (for example salt and pepper mixed together)

Solution: A homogenous (the same throughout) mixture of solute and solvent . Solutions are usually transparent or clear.

Compound: A combination of two or more elements that can not be separated by ordinary means. (e.g. sodium chloride)

Practice:

1. Sort the following into mechanical mixtures, solutions, compounds or elements.

water, pop, cereal with milk, air, oxygen, spaghetti sauce, sugar, syrup, lead, salt, vinegar, salad dressing, gasoline, wood, glass, copper, soil, baking soda also called sodium bicarbonate, sulfuric acid use to unclog drains, nitrogen, aspirin, clear beef broth, jello, mixed nuts, chocolate bar, toothpaste, glue, tears, saliva, carbon dioxide.

Mixtures Solutions Compounds Elements

Understanding Matter :

Laboratory Safety

• Post laboratory rules, including these safety guidelines, in a conspicuous place in the laboratory.

• Before beginning an experiment, review specific safety rules with the class and demonstrate proper procedures.

• Students should be supervised at all times when working in the lab. No unauthorized investigations should ever be conducted in, and no unauthorized materials should ever be brought into, the laboratory.

• Never leave the lab unattended. Lock the laboratory (and storeroom) when you are not present.

• Mark locations of eyewash stations, safety shower, fire extinguishers (ABC tri-class), chemical spill kit, first aid kit, and fire blanket in the laboratory and storeroom. Make sure all safety equipment is present and in good working order before beginning a lab.

• Post an evacuation diagram and established evacuation procedure by every entrance to the laboratory.

• Outfit your lab with labeled disposal containers for glass, sharp objects, and waste chemical reagents.

• No food or beverages should be allowed in the laboratory. Instruct students to keep

their hands away from their faces and to wash their hands with soap and water before leaving the laboratory.

• Know the location for the master shut-off for laboratory circuits.

• Follow prescribed procedures for any safety incident—including fully documenting

every incident. Remind students that any safety incident, no matter how minor, must be reported.

Personal Protective Equipment

• Chemical goggles [meeting ANSI Standard Z87.1] should be worn when working

with any chemical or chemical solution other than water, open flame, or mechanical device or physical process that could eject an object. If a student must wear contact lenses to correct his or her vision, the student should wear eye-cup safety goggles meeting the same ANSI safety standard.

• Face shield [meeting ANSI Standard Z87.1] should be worn in addition to goggles when working with corrosives.

• Eyewash station [meeting ANSI Standard Z358.1] must be capable of delivering

a steady, gentle flow of water to both eyes for at least 15 minutes and must be within a 30-second walking distance from any spot in the room. A plumbed-in fixture or a perforated spray head on the end of a hose, attached to a plumbed-in outlet, designed for use as an eyewash fountain, is suitable if it meets the ANSI safety standard. Portable liquid supply devices are not satisfactory and should not be used.

• Safety shower [meeting ANSI Standard Z358.1] should be within a 30-second

walking distance from any spot in the room. Students should be instructed in the use of the safety shower in the event of a fire or chemical splash.

Understanding Matter:

Structure of the Atom.

Background Information: Over time, as our understanding of the atom evolved, so did the model that we use to describe it. Early Greeks believed that all matter was made of only four things, Earth, fire, water and air. There were a few Greeks that lived about 500 BC who thought otherwise. They were called the "Atomists".

The Atomists and John Dalton Democritus believed that:

1. all matter was made of very tiny but indivisible particles that could not be further divided. The word atom comes from the Greek word Atomos meaning indivisible.

2. matter could not be created or destroyed only changed from one form to another.

3. all atoms were of the same kind, that is the same substance, but differed in size, shape and position.

4. atoms existed in empty space and were therefore able to move around.

5. atoms were in ceaseless motion although they didn't understand why this would be.

All of these assumptions about atoms were made without the aid of any scientific instruments at all. Microscopes had not been invented; even the clock had not been invented yet.

In 1808, John Dalton added to the

theories of the Atomists by describing the atom as being like a billiard ball and so his model is called the billiard ball model.

Thomson Model of the Atom

Just before 1900, scientists began to believe that matter had an electrical charge. Atoms seemed to have either a positive or a negative charge.

Rutherford Model of the Atom

In 1911, a physicist studying at McGill University (in Canada!) suggested the nuclear model for the atom. His research indicated that most of the mass of the

atom was concentrated in a very tiny

nucleus and the area around the nucleus contained mostly empty space.

He concluded that the nucleus of the atom was where all the positive charges were located and all the negative charges were located at great distances from each other in the empty space surrounding the nucleus.

We now know that this is true and if you were to enlarge the size of the nucleus of uranium to that of a grapefruit and place it over the centre of Canada the electrons in the highest energy level would be orbiting between Vancouver and Ottawa and would be about the size of grains of wheat. In a sphere that size there would only be one grapefruit nucleus and 92 wheat grain sized electrons.

The Bohr model of the Atom

The Electron Cloud model of the Atom

By 1930, several physicists had proposed the electron cloud model or the quantum mechanical model of the atom. They believe that the electrons have distinct energy levels but there is no way to

Understanding Matter -

Periodicity of Elements

Elements are arranged in periods and groups. Each group (also called a family) has similar physical and chemical properties. In the Periodic Table of Elements groups are arranged in columns while periods are rows.

The Periodic Table of Elements contains a great deal of information about each element.

See the key on your periodic table for all that it has to offer.

NOTE: Symbols for elements are:

- one capital letter or

- one capital letter followed by one lower case letter

Na

or

Al

or

Ag

NOT

n

a

or

a

l

or

AG

Elements should be written in lower case unless at the beginning of a sentence

Understanding Matter -

An Elemental Tale

'The Kid' mounted his trusty steed, old (B) ____________. His shooting

(Fe)____________ strapped to his side, he headed out for the bright (NE) _________

lights of Cameras, aiming to rob the Cactus Stage. There was sure to be a load of

precious (U)__________ on board, and probably (K) _________ too. Inhaling a deep

breath of (O)__________, the coughed on the (S)__________ fumes from the nearby gas

plants. Since the (Hg) ___________ was climbing, he quenched his thirst with some

H

O, tasting the (CL)__________ all big cities add to their 'aqua purr'. As he headed

North, his bones ached from (Ca)__________ built up over years of riding the

(Zn)__________ trail. Overhead a (He)_________ filled balloon floated in the breeze

and the sun beat down like burning (P)___________. Soon he spotted the stage, guarded

only by a sheriff with a (Sn)_________ badge. He (Kr)____________, slowly

approaching the stage. "Halt," he yelled, "or I'll fill you full of (Pb)__________!" The

Sheriff drew his gun, but alas, he was too slow. The Kid's gun, blazing like flaming

(Mg)___________ did the (Cu)____________ in. All anyone could do was

(Ba)___________. Anyone who drew on the Kid would know that his life wasn't worth a

plugged (Ni)__________! A (Pt)____________ blonde riding beside the

(Al)___________ framed coach rode for her life when the Kid pulled out some

(N)___________ compounds, preparing to blow the strongbox to atoms. Suddenly a

shout rang out. "Hi Ho (Ag)___________." A masked man on a white horse with an

(In)____________ friend raced across the (Si)___________ sands like

Understanding Matter -

Atomic Structure

Atoms are made of protons, neutrons and electrons

Particle Location Relative Mass Relative Charge

Electron Outside Nucleus 1/1840 -1

Proton Inside Nucleus 1 +1

Neutron Inside Nucleus 1 0

Use the Periodic Table of Elements to find the number of p+_{, n}0 _{and e}

-For an atom:

Number of Protons = Atomic number

Number of Neutrons = (Round the Atomic Mass) - Atomic Number Number of Electrons = Number of protons

Practice:

1. Use your periodic table to complete the table below.

Element symbol Protons Neutrons Electrons nitrogen

cadmium

chlorine

strontium

nickel

antimony

silicon

13

47

Understanding Matter -

Periodic Table Assignment

1. Fill in the group (family) members on the following outline of the periodic table. Also write the common names for the families or series of elements in the appropriate areas on the table.

Complete the following table.

Use/Source English Name AtomicNo. Element Symbol Group No. Number Period SATP State 2. Rich ores at Great Bear Lake,

NWT

Radium

3. Rich deposits at Bernie Lake,

Manitoba 1 (IA) 6

4. Potash deposits in Saskatchewan

19

5. Extracted from Alberta sour Natural gas.

S

6. Radiation source for cancer treatment

9 (VIIIB) 4

7. World scale production in Sudbury.

28

9. Provide an empirical definition of metals.

10. Provide an empirical definition of non-metals.

11. Why do chemists classify elements?

Element Name Element Symbol

Atomic Number

Group Number

Period Number

Metal(m) Non (nm)

State Family Name 1. chlorine

2. magnesium

3. 30

4. N

5. 17 5

6. 79

7. 3 alkali

metals 8. thorium

9. 12 (l)

10 Br

11 argon

12 11 5

13 19

14 calcium

Understanding Matter

: Energy Level Diagrams of Atoms

All the electrons a given element has are arranged in shells around the nucleus. Each shell can only hold a certain number of electrons before it is 'full'.

SHELL FULL NUMBER of

e-1st shell 2 electrons total 2

2nd shell 8 electrons total 10

3rd shell 8 electrons total 18

4th shell 18 electrons total 36

5th shell 18 electrons total 54

6th shell 32 electrons total 86

7th shell 32 electrons (total 118)

 memorize the bolded sections

A diagram of an atom with all its electrons can look like these.

This diagram shows boron

a - electrons

b - 1st electron shell

c - nucleus

d - second electron shell

---2e- _{--- # of electrons in first shell}

5p+ _{number of protons}

6n0 _{number of neutrons}

Understanding Matter -

Ions (Charged Atoms)

Ions are atoms that achieved a filled outer energy level. Because the electron has very little mass (almost none), the mass does not change. However, the charge of the atom does change.

The reason atoms form ions is that certain electron configurations have a lower energy (are more desirable). These configurations are also called noble gas or inert gas configurations.

They are like:

helium 2 electrons neon 10 electrons argon 18 electrons krypton 36 electrons xenon 54 electrons radon 86 electrons

If an element can give up or gain electrons it will try to form an ion with 2, 10, 18, 36, 54 or 86 electrons (the atom will be happy).

Note: The transition elements obey slightly different rules.

For example:

Sodium has 11 electrons and 11 protons it could attain the stability of neon if it lost 1 electron. The only problem is that it would have 1 more protons than electrons. This gives it a charge of +1.

sodium atom (Na) sodium ion (Na+) 8e

2e

-11 protons = 11+ 11 protons = 11+

11 electrons = 11- 10 electrons = 10 - 11p+

_________ _________ 12n0

0 1+

Na+

Chlorine has 17 electrons and 17 protons it could attain the stability of argon if it gained 1 electron. The only problem is that it would have 1 more electron than protons.

This gives it a charge of -1.

8e

-chlorine atom (Cl) chloride ion (Cl-) 8e

-2e

-17 protons = 17+ 17 protons = 17+

17 electrons = 17- 18 electrons = 18 - 17p+

_________ _________ 18n0

-Energy Level Diagrams for

ATOMS

1

2

13

14

15

16

17

18

Energy Level Diagrams for IONS

1

2

13

14

15

16

17

18

Understanding Matter -

Review of Atoms and Ions

1. Write the English names for each of the following elements:

H : P : Na : Cu :

I : Cl : Hg : Ni :

2. Give an empirical (based on observation) definition of a metal.

3. Give a theoretical (based on theory) definition of a metal.

4. Give an empirical definition of a nonmetal.

5. Give a theoretical definition of a nonmetal.

6. What does SATP mean?

7. From the periodic table, list the elements that are 1) gasses and 2) liquids at SATP, including the state. (eg. A(g) would be element A - a gas).

8. Draw the energy level diagrams for the following chemical species:

boron atom Chloride ion aluminum ion neon atom oxide ion

9. Horizontal rows on the periodic table are referred to as _____________________.

11. Predict, where possible, the number of valence electrons for each of the following atoms.

argon - carbon - tin - nitrogen - nickel

-12. Predict, where possible, the most likely charge on the following ions.

alkali metals ____ Alkaline earth metals____ Group 13 ____ Group 15 ____

Group 16 ____ Halogens ____ Group 5 ____

13. Complete the following table: English

Name ChemicalSymbol No. ofProtons ElectronsNo. of # of e

- _Donated

or Accepted ChargeNet

eg. chloride ion Cl- _{17 18 accepted 1}

1-eg. sodium atom Na 11 11 0 0

1. 20 18

2. 1

1-3. Ar

4. Mg 12

5. chlorine atom

6. 9 10

7. 6 0

8. H+

9. sodium ion

10. 1 1

11. 7 0

12. 16

argon atom

13. S

2-14. I

15. uranium atom

16. 10 accepted 2

Composition of Chemical Compounds -

Ionic Compounds

Ionic compounds form because some elements would really like to pick up an extra electron or two or three. They can only do this if they can find another element that is willing to give them up.

Take lithium and fluorine below. Lithium could attain the electron configuration of helium if it could give away its valence electron. Fluorine could become like neon if to could only gain an electron.

Put both elements together and they make a trade.

Lithium fluoride has formed

Li

F

Li

_F

-In this exercise, circle all the metal ion names, and underline all the non-metal names.

zinc sulfide magnesium chloride potassium iodide strontium nitride

What generalization can you make about the names of each of these binary ionic compounds?

Ionic compounds are made of just one metal and one non-metal ion. Ionic compounds are always named with the positive ion (cation) first followed by the anion (a negative ion). In the example that we use so often, NaCl is named sodium chloride.

When the formula for the compound is written, it must be written in such a way that it has a neutral charge. eg. the sodium ion, Na+_{, plus the chloride ion, Cl}- _{combine to form NaCl. The charge is}

neutral because the metal ion has a one positive charge and the non-metal has a one negative charge. Not so with magnesium iodide,

The magnesium ion, Mg2+ _{and the iodine ion, I}-_{do not produce a neutral compound when they are}

put together. - MgI. Instead, we must balance the compound by adding an additional negative ion to produce a compound with a neutral charge.

Thus the compound requires one Mg2+_,I- _{, I}- _{. This forms the compound magnesium iodide, MgI} 2.

Two iodide ions are needed to neutralize the charge of the one magnesium ion.

Chemical Formula

eg. CaCl2 Ca2+ , Cl-, Cl- calcium chloride

1. potassium iodide

2. MgO

3. aluminum chloride

4. NaBr

5. CaO

6. lithium nitride

7. Al2O3

8. barium chloride

9. sodium chloride

10. ZnO

11. silver bromide

12. magnesium hydride

13. magnesium chloride

14. zinc chloride

15. Ag2S

16. potassium chloride

17. CaF2

18. sodium sulfide

19. CaH2

Composition of Chemical Compounds –

Ionic compounds with multivalent metals

Some ions can exist with different ion charges in different situations. For example, iron can exist as Fe3+ _{and Fe}2+_{. To distinguish one from the other, Roman numerals are used in the}

name of compounds that are made with ions that can have different charges.

 Fe3+ _{is called the iron(III) ion and}

 Fe2+_{. is called the iron (II) ion.}

Write the chemical formula for iron (II) oxide. ________, for iron(III)oxide. ________

Complete the following table using the Stock System for naming ionic compounds. Remember that all transition metals except for aluminum, zinc and silver require the Roman numeral to specify ion charge.

Chemical Formula IUPAC Name Summary of Charges

1. Cu3N2 copper(II) nitride Cu2+, Cu2+, Cu2+, N3-, N

3-2. gold (I) chloride

3. Al2S3

4. tin(II) oxide

5. antimony(V) bromide

6. SbF3

7. CuCl

8. Fe2Se3

9. ZnCl2

10. aluminum sulfide

11. mercury(I) nitride

12. AlP

13. Ni2S3

14. VI5

15. cobalt(II) phosphide

16. chromium(II) hydride

Composition of Chemical Compounds -

Ternary Ionic Compounds

Ternary ionic compounds are usually compounds made from polyatomic ions. These polyatomic ions are identified on your periodic table in the box at the top of the page. Generally, if you don't recognize the name of a substance it is a polyatomic ion.

For example: If asked to write the formula for sodium sulfate. You can recognize the sodium part but sulfate may be a word you don't recognize. In that case check the box containing the polyatomic ions. Sulfate shows up as SO42-. This means it is a polyatomic ion with a charge of -2.

To write the formula for the compound sodium sulfate, you need two sodium ions and one sulfate ion. Na+ _{+ Na}+ _{+ SO}

42- makes Na2SO4 and it has a neutral charge so you are done.

Whenever you need multiple polyatomic ions, you must put parentheses around the polyatomic ion and add the subscript outside the parentheses.

E.g. Write the name of Mg(ClO)2 ____________________________________.

The following table that requires you to name some common polyatomic ions. Remember that polyatomic ions are not molecules and cannot exist by themselves as they are on the periodic table or in this exercise.

Ion Name Formula Ion Name Formula

1. hydrogen sulfate 6. sulfite

2. ClO3- 7. NO3

3. NH4+ 8. hydrogen sulfide

4. Dichromate 9. HPO4

5. OH- ₁₀

.

-Naming Compounds with Polyatomic Ions

Use the table of polyatomic ions to complete the following exercise. Only those polyatomic ions listed on the periodic table are used in science 10.

International Formula Summary of Charges IUPAC English Name 1. Na2CO3

2. (NH4)2CO3

3. FeSO4

4. lithium hydroxide

5. aluminum hydroxide

6. NaClO

7. potassium dichromate

8. LiC6H5COO

9. NaNO2

10. ammonium sulfate

11. sodium hydrogen carbonate

12. Na3PO4

13. calcium dihydrogen phosphate

14. PbCrO4

15. sodium hydrogen sulfate

16. KMnO4

17. aluminum silicate

18. Li2CO3

Complete the following table using the stock system and polyatomic ions for naming compounds.

Chemical Formula Summary of Charge Name of Compound eg. Cu2SO3 Cu+, Cu+, SO32- copper(I) sulfite

1. uranium(IV) oxide

2. lead(IV) sulfate

3. Sn(HPO4)2

4. Al2O3

5. manganese(IV) iodate

6. Sb2S3

7. thallium(III) hydroxide

8. HgS

9. MoS3

10. polonium (II) thiosulfate

11. FeSO4

12. lead(IV) chlorate

13. Hg(NO3)2

14. ZnSe

15. V2O5

16. tin (II) borate

17. chromium (III) phosphate

18. TiO2

19. Ag2SO3

20. AuCl3

21. uranium(IV) cyanide

22. NiBr2

Composition of Chemical Compounds -

Hydrated Compounds

Hydrated compounds are compounds that contain water as part of their structure. Some compounds are water seeking and are most stable when they are attached to many water molecules.

To name these substances we need to know the common prefixes so that the compounds can be described correctly.

The common substance, Epsom salts, is named magnesium sulfate heptahydrate and its formula is MgSO47H2O. If you

remove the water by heating or by some other means, the compound can be used to soak up water as it gets back the water you have taken away. Anti-perspirants are made in this way.

Generally the chemical formula ends with __H2O

- where the blank is the correct value for the number of water molecules that are bonded to the ionic compound.

The IUPAC name can be written in two different ways. For the above example MgSO47H2O would

be written either;

- using the prefix system as magnesium sulfate heptahydrate or - using the number system as magnesium sulfate-7-water.

***You must be able to read and write the names of hydrated compounds using either system.

Eg. – Cu(NO3)2  4 H2O

  tetrahydrate

copper (II) nitrate

Full Name :

copper (II) nitrate tetrahydrate

1 - mono 2 - di 3 - tri 4 - tetra 5 - penta

Naming Hydrated Compounds. Complete the following table.

Name of Hydrate Common Name, Use or Discrition Formula e.g. copper(II) sulfate

pentahydrate

blue vitriol, bluestone, copper plating,

blue solid CuSO45H2O(s)

1. Epsom salts, white solid explosives, matches MgSO47H2O(s)

2. sodium carbonate decahydrate

washing soda, soda ash, water softener, white solid

3. white solid, fireproofing wood,

disinfectants, parchment paper MgCl26H2O(s) 4. barium chloride

dihydrate

white solid, pigments, dyeing fabrics, tanning leather

5. white solid, photographic emulsions Cd(NO3)24H2O(s)

6. white solid, embalming material,

fireproofing lumber, vulcanizing ZnCl25H2O(s) 7. zinc sulfate

heptahydrate

white solid, clarifying glue, preserving wood and skins

8. lithium chloride tetrahydrate

white solid, soldering aluminium in fireworks

9. photographic hypo, antichlor, white solid Na2S2O35H2O(s)

10 cobalt(II) chloride hexahydrate

pink solid, humidity and water indicator, foam stabilizer in beer

11 white solid, antiperspirant AlCl36H2O(s)

12 de-icer used on icy highways, added to

cement mixtures to prevent freezing CaCl22H2O(s) 13 barium hydroxide

octahydrate

white solid, manufacture of glass, water softener

14 nickel(II) chloride hexahydrate

green solid, manufacture of glass, water softener

15 Glauber's salt (a medicine), white solid,

drying agent Na2SO410H2O(s)

Ionic compounds are determined to have either high solubility or low solubility when placed into water. Since it is not possible for you to know which is which, a solubility chart is provided for you on the periodic table. The top row of ions on the Solubility Chart contains the high solubility (H) compounds, the bottom row contains the low solubility (L) compounds.

To determine the solubility of an ionic compound you first look for the anion on the chart (the non-metal ion). Once you find the anion, there are two choices for the cation in the compound. If it appears in the top box below the anion name, it has high solubility and is therefore said to be aqueous

(aq). If it appears in the bottom box it has low solubility and the compound is a solid (s).

Ionic compounds are deemed to be high solubility (H) if large amounts of the solid will dissolve in water. Low solubility (L) substances dissolve only very slightly in water.

Predict whether the following Ionic Compounds are high solubility or low solubility.

IUPAC Name

State of the Pure substance (s, l, or g)

Solubility H/L

State in a Water Enviroment (s or aq)

1. AgNO3 ( ) ( )

2. NH4OH ( ) ( )

3. PbS ( ) ( )

4. Ag2SO4 ( ) ( )

5. CaCO3 ( ) ( )

6. Mg(CH3COO)2 ( ) ( )

7. Al2(SO4)3 ( ) ( )

8. Na2S ( ) ( )

The following ionic species (ions) were placed in an aqueous environment (in water). Describe what will happen when each pair of ions reacts to form an ionic compound. Write the chemical formula for the ionic compound, including the state in the aqueous environment, then indicate whether they are high or low solubility and whether a precipitate forms.

If a low solubility substance forms when the ions combine, the evidence will be the formation of a precipitate. You can tell when a precipitate forms because it becomes a solid (usually white and cloudy). In this case the chemical formula will be written with a (s).

High solubility substances that form will remain in solution (no cloudiness), are written as (aq).

Complete the table below by writing the name of the chemical formula including its state in a water environment, its solubility and whether a precipitate will form.

IONS Solubility? Precipitate? Chemical Formula & State 1. Ba2+ _{and SO}

4

2-2. Mg2+ and S 2-3. Fe3+ _{and OH}

-4. K+ _{and CO} 3

2-5. Sr2+ _{and OH}

-6. Na+ _{and OH}

-7. NH4+ and PO4

-Composition of Chemical Compounds -

Naming of Acids

Acids are a special kind of ionic compound. Acids are made up of positive and negative ions but the positive ion in this case is only the hydrogen ion (H+_{). Acids are unique also because they do not}

behave as acids until they are dissolved in water.

Acids have the following properties:

 are solids, liquids or gases at SATP. (when they are not in a water environment)

 are highly soluble in water.

 form conducting solutions.

 turn indicators like blue litmus paper red and bromothymol blue solution yellow.

 react with metals such as zinc to produce hydrogen gas (H2(g)).

Acids naming appears to be complex at first but there are really only three different naming conventions for acids. As ionic compounds (not in water ) they have different names than as acids when they are dissolved in water. This table illustrates the way the different acids are named.

Ionic Name Acid Name

1. hydrogen ________ide

eg. hydrogen nitride

becomes

hydro_________ic acid

hydronitric acid H3N(aq)

2. hydrogen ________ate

eg. hydrogen nitrate becomes

_____________ic acid

nitric acid HNO3(aq)

3. hydrogen ________ite

eg. hydrogen nitrite

becomes

_____________ous acid

nitrous acid HNO2(aq)

Chemical Formula Name as an Ionic Name as an Acid

1. H2 SO4 (aq)

2. H3 PO4 (aq)

3. boric acid

More Acid Naming

Chemical Formula Name as an Ionic Name as an Acid

1. H2 SO4 (aq)

2. H3 PO4 (aq)

3. boric acid

4. carbonic acid

5. hydrogen fluoride

6. hydrosulfuric acid

7. chloric acid

8. H2 S (aq)

9. hydrogen nitrate

10. sulfurous acid

11. HNO3 (aq)

12. HCl (aq)

13. hydrogen nitrite

14. hydrogen benzoate

15. HCNq)

16. oxalic acid

17. HBr (aq)

18. HI (aq)

Distinguishing between ionic compounds and acids

 In the following table, indicate the state of the compound in a water environment as (s), (l) or (aq).

Acid/Ionic Chemical Formula Name of Compound

1. Al(OH)3 ( )

2. aluminum sulfate

3. H3 BO3 ( )

4. sulfuric acid

5. NH4 NO3 ( )

6. potassium carbonate

7. H2 SO4 ( )

8. phosphoric acid

9. CuSO4 ( )

10. hydroiodic acid

11. lead (II) acetate

12. H2 CO3 ( )

13. hydrosulfuric acid

14. sodium chlorate

15. KI ( )

16. KMnO4 ( )

17. ammonium nitrate

18. HNO2 ( )

19. barium hydroxide

20. hydrochloric acid

21. Au(NO3 )3 ( )

22. nitric acid

23. HOOCCOOH ( )

24. aluminum phosphate

Composition of Chemical Compounds -

Molecular Elements

S

H

P

O

N

Molecular Compounds

Some molecular compounds must be memorized because there are no simple rules for naming them.

Many molecular compounds can be named using the prefix system as with hydrates.

In molecular compounds, the prefixes refer to the subscript number. This is different than hydrates where the prefix referred to the coefficient (the number in front of the H2O). Note: The mono

prefix is only used for the second element where necessary.

E.g. - CO is named carbon monoxide not monocarbon monoxide. - CO2 is named carbon dioxide.

The molecular compounds that must be remembered are: 1 - mono 2 - di 3 - tri 4 - tetra 5 - penta

Molecular Naming

Molecular Formula (inc. SATP state) IUPAC English Name

1. dinitrogen monoxide

2. nitrogen dioxide

3. dinitrogen trioxide

4. NO(g)

5. N2O4(g)

6. N2O5(g)

7. P2O5(s)

8. sulfur hexafluoride

9. sulfur trioxide

10. phosphorus trichloride

11. PCl5(s)

12. CH3OH(l)

13. O2F2(g)

14. CO(g)

15. NH3(g)

16. sucrose

167 .

Composition of Chemical Compounds –

More Molecular Compounds

1. List the molecular prefixes from one to ten.

2. For which type of molecular substance are these prefixes used ?

3. Why is memorisation required for the nomenclature of many molecular substances in this unit?

Molecular Formula (inc. SATP state) IUPAC English Name

4. oxygen

5. P2O5(s)

6. hydrogen monochloride

7. NH3(g)

8. dinitrogen tetrahydride (liquid)

9. PCl5(g)

10. methane

11. NI3(l)

12. CH3OH( )

13. sucrose

14. S4N2(s)

15. ethanol

16. CO( )

17. H2O2( )

18. SO3(g)

19. sulfur

Science 10 – Naming Summary

_{PURE SUBSTANCES}

IONIC COMPOUNDS Empirical evidence

- Form white solids at room temperature. - Can be soluble in water to become (aq) ions.

- If not soluble, form solids (s).

- Form conducting solutions if soluble in water Theoretical evidence

- They transfer electrons to form + and - ions. - There is an attraction between + and - ions. - Cations - positively charged ions.

- Anions - negatively charged ions. - Made of: a metal and a non-metal or

a metal and a polyatomic ion or two polyatomic ions.

Naming

- Positive ion is named first then negative ion. - The negative ion ends in -ide.

- Formula must be written so that the positive and negative charges are balanced. Stock system

- use Roman numerals after the name of any transition metal to specify the ion charge. 3 exceptions are Al(s), Zn(s) &

Ag(s).

Hydrates

- are compounds with water molecules as part

ACIDS Empirical evidence

- Can be solid, liquid or gas in pure form. - Form conducting aqueous solutions.

- Form solutions which turn blue litmus pink. - React with active metals and release H2(g).

- They taste sour (like vinegar).

Theoretical evidence

- the positive ion in acid formulas is the H+ _ion.

- they dissociate in water to form hydrogen ions (H+

(aq)).

Naming Acids

1) hydrogen ____ide  hydro_____ic acid

2) hydrogen _____ate  ______ic acid

3) hydrogen _____ite  _____ous acid

- all acids are written as aqueous (aq).

6 Strong Acids

HClO4(aq) perchloric acid

HNO3(aq) nitric acid

HCl(aq) hydrochloric acid

HBr(aq) hydrobromic acid

HI(aq) hydroiodic acid

H2SO4(aq) sulfuric acid

MOLECULAR COMPOUNDS Empirical evidence

- Can be solids, liquids and gases at SATP. - Form non-conducting solutions.

- Most are not soluble in water. Theoretical evidence

- elements in molecular compounds share electrons.

- molecular compounds form covalent bonds. - Made of two or more non-metals.

Naming

The following molecular compounds including state and name must be memorized.

H2O(l) water H2O2(1)hydrogen peroxide

NH3(g) ammonia C12H22O11(s) sucrose

CH4(g) methane C3H8(g) propane

CH3OH(l) methanol C2H5OH(l) ethanol

C6H12O6(s) glucose O3(g) ozone

- Other molecular compounds are named using the prefix system (mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca.)

Polyatomic Molecular Elements

Composition of Chemical Compounds

- Naming Review 1

 Classify each substance as ionic, molecular or acid. Predict the international chemical formula (including state of matter at SATP) or the IUPAC English name. Communicate the solubility of the substance in a water environment using (aq) for high solubility and (s, l or g) for low solubility substances.

i,m,a Chemical Formula Solubility IUPAC English Name

1. PbI2(s)

2. (l) ethanol

3. NaHS(s)

4. sulfurous acid

5. H2O2(l) (l)

6. titanium(IV) oxide

7. Co(NO3)26H2O(s)

8. H2S(g)

9. gallium sulfide

10. sulfuric acid

11. CH4(g) (g)

12. ammonium chromate

13. SO3(g) (g)

14. H2CO3(aq)

15. (g) dinitrogen tetraoxide gas

16. Al2(SO4)3(s)

17. Na2SO3(s)

18. (g) ammonia

19. sodium thiosulfate pentahydrate

Composition of Chemical Compounds –

Review of Naming 2

 Complete the following table for the binary ionic compounds and those containing polyatomic ions.

Chemical Formula IUPAC Name Summary of Charges

1. MnCl2(s)

2. AlBr3(s)

3. zinc oxide

4. iron(II) fluoride

5. Ni2O3(s)

6. Cu2S(s)

7. cobalt(II) chloride

8. Al(NO3)3(s)

9. tin(II) chlorite

10. chromium(III) hydroxide

11. Fe2(SO4)3(s)

12. Cu(NO2)2(s)

13. iron(II) dichromate

14. Ag2SO3(s)

15. Ca(HCO3)2(s)

16. aluminum chlorate

Composition of Chemical Compounds –

Review of Naming 3

Predicted Formula IUPAC Name Summary of Charges

1. K2SO4(s)

2. Ca(HS)2(s)

3. magnesium hypochlorite

4. tin(II) acetate

5. chromium(II) sulfite

6. iron(III) acetate

7. Co(OH)2(s)

8. Cu2SO4(s)

9. Cr3N2(s)

10. Ni2(S2O3)3(s)

11. vanadium(V) silicate

12. aluminium sulfate

13. nickel(III) chlorate

14. Cr(CN)2(s)

15. CuSO45H2O(s)

16. calcium chloride hexahydrate

17. Na2CO310H2O(s)

Composition of Chemical Compounds –

Review of Naming 4

 Complete the following table of ionic compounds.

Predicted Chemical Formula IUPAC English Name

1. SrCl2(s)

2. RbBr(s)

3. Na2O(s)

4. aluminum sulfide

5. magnesium iodide

6. TiO2(s)

7. Cu2O(s)

8. tin(II) sulfide

9. chromium(III) oxide

10. iron(II) sulfide

11. KC6H5COO(s)

12. Na2S2O3(s)

13. NH4HCO3(s)

14. ammonium sulfide

15. barium sulfite

16. magnesium hydroxide

17. FeSO47H2O(s)

18. LiCl4H2O(s)

19. sodium sulfate decahydrate

20. Au(NO3)3(s)

21. bismuth(III) sulfate

22. lead(II) acetate trihydrate

 Give the state of each compound as if they were in a water environment.

Predicted Chemical Formula IUPAC English Name

1. CrCl2

2. V2O5

3. Au2S

4. aluminum chloride

5. gallium sulfide

6. barium bromide

7. CaCl2

8. KI

9. Ag2O

10. calcium hydroxide

11. zinc carbonate

12. ammonium phosphate

13. NaCH3COO

14. K2SiO3

15. NH4HSO4

16. lead(II) oxide

17. nickel(II) sulfate

18. manganese(II) cloride

19. MgSO4H2O

20. BaCl24H2O

21. sodium thiosulfate pentahydrate

22. nickel(II) chloride hexahydrate

23. Sb2(SO4)3

24. calcium hydrogen carbonate

Composition of Chemical Compounds

– Naming Review 6

Complete the following table. Classify the substance as ionic, molecular or acid (i, m, or a) in the first column. Use the subscript to indicate the state of each substance (s, l, g or aq at room temp.) RULES: 1. Pure ionic compounds -- (s) only

2. Pure molecular compounds -- (s), (l) or (g)

3. Acids -- assume acids are in a water environment so label as (aq) only.

Chemical Formula Name of Compound i, m or a 1. Al(OH)3( )

2. sodium sulfate decahydrate

3. sodium nitrate hexahydrate

4. Al2(SO4)3( )

5. calcium chloride hexahydrate

6. NH4NO3( )

7. (g) phosphorous trihydride

8. N2O3(g)

9. (g) methane

10. H2SO4( )

11. H3PO4( )

12. boric acid

13. (NH4)2SO4( )

14. SnF2( )

15. carbonic acid

16. PbO2(s)

17. (s) silicon dioxide

18. NaClO( )

19. potassium permanganate

Chemical Formula Name of Compound i, m,or a

21. K2CO32H2O( )

22. hydrofluoric acid

23. H2S( g )

24. sodium hydroxide

25. NaHSO4( )

26. magnesium sulfate heptahydrate

27. Ca(OH)2( )

28. sodium thiosulfate

29. CaO( )

30. copper(II) sulfate pentahydrate

31. sulfur

32. B2H8(g)

33. KI( )

34. phosphorus

35. SO3(g)

36. sodium chlorate

37. Na2SiO3( )

38. methanol

39. chloric acid

40. lead(II) sulfate

41. Ca(HCO3)2( )

42. (g) nitrogen trichloride

43. sodium hydrogen sulfite

44. CS2(s)

45. H2S(aq)

46. water

Chemical Change

- Balancing and Interpreting Equations

When writing and balancing chemical equations you must do all the following steps.

1. Write all the reactants on the left side of arrow and all the products on the right side of arrow using the correct formulas for the compounds.

2. Matter can not be created nor destroyed, therefore we must balance both sides of the reaction arrow with the correct number of atoms.

3. Generally, you start with atom that is most numerous.

4. When you are finished, you can always check your answer. If you have the correct formula for each compound and the same number of each atom on both sides of the equation, you have it balanced.

Example: The correct equation for the combustion of methane gas in a furnace.

CH4(g) + O2(g) CO2(g) + H2O(g)

General Rule for Combustion Reactions : Balance the oxygen atoms last!!!

The atom that is most numerous is the hydrogen atom. There are hydrogen atoms on the left of the reaction and 2 on the right. We balance this by placing a coefficient (2) in front of the water molecule.

CH4(g) + O2(g)  CO2(g) +

2

Now that the hydrogen atoms are balanced we move on the carbon. (Remember oxygen last.) There is one carbon atom on the left and one on the right – the carbons are balanced.

Finally, we look at the oxygen atoms. There are two on the right of the arrow and four on the right. By placing the coefficient 2 in front of the O2(g) the equation will be balanced.

CH4(g) +

2

This activity has all the correct reactant and product formulas written for you. Balance the following equations.

1. ___K(s) + ___ Cl2(g)  ___ KCl(s)

4. ___NaCl(s)  ___Na(s) + ___Cl2(g)

5. ___AsCl3(aq) + ___H2S(aq)  ___As2S3(s) + ___HCl(aq)

6. ___CuSO4 5H 2O(s)  ___CuSO4(s) + ___H2O(aq)

7. ___Na(s) + ___O2(g)  ___Na2O(s)

8. ___H2S(aq) + ___KOH(aq)  ___HOH(l) + ___K2S(aq)

9. ___Fe(s) + ___H2O(g)  ___H2(g) + ___Fe3O4(s)

10. ___Al(s) + ___H2SO4(aq)  ___H2(g) + ___Al2(SO4)3(aq)

11. ___H3PO4(aq) + ___NH4OH(aq)  ___HOH(l) + ___(NH4)3PO4(aq)

12. ___C3H8(g) + ___O2(g)  ___CO2(g) + ___H2O(g)

13. ___CH4(g) + ___O2(g)  ___CO2(g) + ___H2O(g)

14. ___AlCl3(aq) + ___NaOH(aq)  ___Al(OH)3(s) + ___NaCl(aq)

15. ___Na2CO3(aq) + ___HCl(aq)  ___NaCl(aq) + ___H2O(l) + ___CO2(g)

16. ___Fe(s) + ___CuSO4(aq)  ___Cu(s) + ___Fe2(SO4)3(aq)

17. ___H2SO4(aq) + ___KOH(aq)  ___HOH(l) + ___K2SO4(aq)

Chemical changes -

Formation and Decomposition Reactions

Many of the chemical reactions you encounter in high school chemistry can be classified as one of the following five types.

1. formation (f)

2. simple decomposition (sd) 3. single replacement (sr) 4. double replacement (dr) 5. combustion or oxidation(c)

Not every chemical reaction will fit into this empirical classification scheme. For the time being, any reactions, which do not fit these categories, are classified as other (o). 1. Formation reactions (forming a compound)

element + element  compound

Eg. 2 Mg(s) + O2(g)  2 MgO(s)

Eg. S8(s) + 8 O2(g)  8 SO2(g)

1. A frequent technological problem associated with the operation of swimming pools it that chlorine corrodes copper pipes.

___Cu(s) + ___Cl2(g)  ___CuCl (s)

2. When aluminum reacts with air a tough, protective coating forms. ___Al(s) + ___O2(g) 

3. When zinc is exposed to oxygen a protective coating of zinc oxide forms on the surface of the metal. This reaction makes zinc coating (galvanizing) metal technologies desirable.

___Zn(s) + ___O2(g) 

2. Simple decomposition reactions (decomposing a compound into its elements)

compound  element + element

E.g. 2 H2O(l)  2 H2(g) + O2(g)

Eg. 2 NaCl(s)  2 Na(s) + Cl2(g)

1. Since the Bronze Age (about 3000 BC), copper has been produced by heating the mineral, copper(II) oxide.

___CuO(s)  ___Cu(s) + ___O2(g)

2. Another ancient metallurgical process involves heating the mineral, lead(II) oxide, to produce lead.

___PbO(s)  ___Pb(s) + ___O2(g)

3. A major technological breakthrough occurred in 1807 when Humphry Davy isolated potassium by passing electricity through molten potassium oxide.

___K2O(s) 

4. Another similar major development happened in 1886. Charles Martin Hall discovered, after hundreds of trials, a technological process of electrolyzing alumina.

___Al2O3(s) 

5. A current technology for producing sodium involves passing electricity through molten lye to decompose it into its elements.

___NaOH(s) 

6. Magnesium metal is produced by our industrialized society by passing electricity through molten magnesium chloride.

___Mg(ClO3)2(s) 

Chemical Change -

Single Replacement Reactions

 If the element is a metal, it will exchange places with the positive ion (metal) in the compound by attaching itself to the negative ion.

 If the element is a non-metal, it will exchange places with the negative ion (non-metal) by attaching itself to the positive ion.

Element + compound  new compound + new element

A + BC  AC + B

E.g. Cu(s) + 2 AgNO3(aq)  2 Ag(s) + Cu(NO3)2(aq) - Metal replaces with the metal ion

Eg. Cl2(aq) + 2 NaI(aq) I2(s) + 2 NaCl(aq)

- Non-metal replaces with the non-metal ion

Notes: - Polyatomic ions generally remain intact in a chemical reaction.

- Water is generally written as HOH(l) (to represent a H+ ion and a OH- ion).

- If the reaction occurs in an aqueous environment use your Solubility Chart to predict the solubility of the ionic products of single replacement reactions in water.

Single replacement reactions are used by industry to produce elemental metals from their ores or oxides.

1. Silver metal can be produced by placing aluminum foil in an aqueous solution of silver nitrate.

Al(s) + 3 AgNO3(aq) Al(NO3)3(aq) + 3 Ag(s) sr

2. Chlorine is found to replace bromine from an aqueous solution of sodium bromide.

___ Cl2(g) + ___ NaBr(aq) 

3. Chlorine will replace fluorine to from an aqueous solution of potassium chloride. ___ Cl2(g) + ___ KF(aq) 

4. Hydrogen is produced in the laboratory by reacting zinc metal with sulfuric acid. Hydrogen is also obtained using new technologies from water or methane gas.

___ Zn(s) + ___ H2SO4(aq) 

5. A reaction occurs when a zinc strip is placed in an aqueous solution of lead(II) nitrate. ___ Zn + ___ Pb(NO )

6. A single replacement reaction occurs when hydrogen gas is passed over hot copper(II) oxide.

___ H2(g) + ___ CuO(s) 

7. A chemical reaction occurs when aqueous bromine is added to added to aqueous sodium iodide.

___ Br2(aq) + ___ NaI(aq) 

8. Sodium metal reacts vigorously with water to form a basic solution and a gas that burns vigorously.

___ Na(s) + ___ HOH(l) 

9. Molten iron is produced by the reaction between powdered aluminium and iron(III) oxide. In this process a huge amount of heat is produced. Railways use this process to weld steel rails together.

___ Al(s) + ___ Fe2O3(s) 

For the following write the chemical equation from the word equation.

10. Sodium solid reacts with a aqueous silver nitrate solution.

11. Chlorine gas is bubbled through a solution of ammonium sulfide.

12. Strontium metal reacts with water.

13. Cesium metal reacts with water,

Chemical Change -

Double Replacement Reactions

Double replacement reactions involve the exchange of ions between compounds. Usually, the reactants are two compounds and the products are two compounds.

 In the reaction the negative ions become attached to the other positive ions to form new compounds.

 Double replacement reactions are further subdivided into neutralization reactions or precipitation reactions.

 Neutralization reactions usually produce water as one of the products and the other is often a salt of some kind.

 Precipitation reactions produce one low solubility compound and one high solubility compound from two reactant compounds that are usually high solubility.

Compound + compound new compound + new compound AB + CD  AD + CB

E.g. 1 AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)

Example 1 is a precipitation reaction (AgCl(s) is the solid formed – precipitate)

E.g. 2 H2SO4(aq) + 2NaOH(l)  2HOH(l) + Na2SO4(aq)

Example 2 is a neutralization reaction (note: HOH(l))

1. The presence of the chloride ion in a sample is detected by the formation of a white precipitate when silver nitrate is added to the sodium chloride solution.

___AgNO3(aq) + ___NaCl(aq) 

2. The formation of a white gelatinous precipitate when sodium hydroxide is added to a solution indicates the possible presence of the aluminum ion.

___NaOH(aq) + ___Al2(SO4)3(aq) 

3. An analytical chemist uses sodium oxalate to precipitate calcium ions in a water sample from an acidic lake.

___Na2OOCCOO(aq) + ___CaCl2(aq) 

4. A potassium hydroxide solution can be used to determine the concentration of an aqueous hydrogen fluoride solution used to etch (frost) light bulbs in industry.

___KOH(aq) + ___HF(aq) 

5. A white precipitate of lead(II)chloride forms when hydrochloric acid is added to a toxic aqueous lead(II) nitrate solution. The toxic precipitate is placed in a waste container for disposal.

___HCl(aq) + ___Pb(NO3)2(aq) 

dr

6. The presence of carbonate ions in a washing detergent is indicated by the formation of a white precipitate when aqueous barium chloride is added.

___Na2CO3(aq) + ___BaCl2(aq) 

7. Aqueous sodium hydroxide solution is used to determine the concentration of acetic acid in a commercial vinegar sample.

___NaOH (aq) + ___CH3COOH(aq) 

8. When sodium hydroxide is added to well-water sample, the formation of a rusty brown precipitate indicates the presence of an iron(III) compound in the sample.

___NaOH(aq) + ___FeCl3(aq) 

9. The concentration of sulfurous acid in an acid rain sample can be determined by using aqueous potassium hydroxide. (Environmental chemists do this kind of work.)

___H2SO3(aq) + ___KOH(aq) 

Write the chemical equation from the following word equations.

10. An aqueous ammonium chloride solution reacts with an aqueous silver nitrate solution.

11. An aqueous solution of mercury (I) chlorate is added to an aqueous solution of barium hydroxide.

12. An aqueous calcium acetate solution reacts with an aqueous thallium (I) hydroxide solution.

13. Sulfurous acid reacts with an aqueous sodium hydroxide solution.

Chemical Change -

Combustion Reactions

In this reaction a compound of element react with oxygen to form the most common oxides. The most common oxides are:

- If a metal is combusted, use the ionic formula for that metal oxide - If carbon is combusted, carbon dioxide CO2(g) is produced

- If hydrogen is combusted, water H2O(g) is produced

- If nitrogen is combusted, nitrogen dioxide NO2(g) is produced

- If sulfur is combusted, sulfur dioxide SO2(g) is produced

NOTE: ALWAYS BALANCE THE OXYGEN LAST!!

compound + oxygen  most common oxides

(Common oxides like CO2(g), H2O(g), NO2(g), SO2(g))

E.g. 1 4 C2H7N(l) + 19 O2(g)  8 CO2(g) + 14 H2O(g) + 4 NO2(g)

E.g. 2 4 Fe(s) + 3 O2(g)  2 Fe2O3(s)

1. In Western Canada natural gas, assume CH4(g), is the most common fuel for home heating.

___CH4(g) + ___O2(g) 

2. Propane is used as an alternate fuel for automobiles, for home barbecues, for heat in recreational vehicles and for home heating.

___C3H8(g) + ___O2(g) 

3. The first step in the production of sulfuric acid is to burn sulfur.

___S8(s) + ___O2(g) 

4. Rocket fuel burns to propel a satellite into space.

___H2(g) + ___O2(g) 

5. Barbecue charcoal undergoes incomplete combustion to produce deadly carbon monoxide gas

___C(s) + ___O2(g) 

6. In cellular respiration glucose combines with oxygen.

7. Nitromethane, CH3NO2(l), is a fuel commonly burned in drag-racing vehicles.

8. As a safety precaution, small amounts of sulfur-containing compounds called mercaptans, assume C2H5SH(g), are placed in natural gas to give it an odor that can be detected. These

mercaptans burn in furnaces and hot water heaters in Canadian homes.

9. Sucrose burns.

11. At the high temperatures in the combustion cylinder of an automobile, relativity inert nitrogen gas undergoes combustion. (What other class of reaction is this reaction?)

12. Ammonia is burned in fertilizer plants serving Canadian agriculture. This is an initial step in the technological process to produce nitrate fertilizers used to increase agricultural yields.

3. Gasoline is mixed with air in the carburetor and then exploded by a spark in the cylinder of a car motor.

___C8H18(l) + ___O2(g) 

4. A rock may be tested for limestone content by adding muriatic acid.

___CaCO3(s) + ___HCl(aq )  CaCl2(aq) + ___H2O(l) + ___CO2(g)

5. Kerosene (C14H30) is a mixture of hydrocarbons used as a fuel for stoves and

Chemical Change -

Chemical Equations Practice

1. In 1774 Joesph Priestley discovered oxygen by decomposing the oxide of mercury. ___HgO(s) 

2. Molten table salt is industrially decomposed to produce molten sodium. ___NaCl(s) 

3. Nitrogen from the air reacts with hydrogen to produce ammonia for fertilizers. ___N2(g) + ___H2(g)  ___NH3(g)

4. Copper ore is decomposed to remove the copper metal. ____CuO(s) 

5. Freshly cut lithium reacts with nitrogen from the air.

6. A silver spoon or coin tarnishes when exposed to sulfur.

7. Molten lye (NaOH)(l) is decomposed industrially into its elements.

8. Sodium metal reacts vigorously with water. ___Na(s) + ___HOH(l) 

9. Hydrogen chloride gas is produced in the laboratory from table salt. ___NaCl(s) + ___H2SO4(aq) 

10. Molten iron is produced in the highly exothermic reaction ___Al(s) + Fe2O3(s) 

11. Slaked lime precipitates magnesium ions from hard water. ___Ca(OH)2(aq) + ___Mg(HCO3)2(aq) 

___Cu(s) + ___Ag2SO4(aq) 

14. Phosphoric acid is produced at a fertilizer plant. ___H2SO4(aq) + ___Ca3(PO4)2(s) 

15. Bromine is commercially produced from MgBr2 found in sea water.

___Cl2(g) + ___MgBr2(aq) 

16. Hydrogen sulfide (sour) gas from a wild (untamed or uncontrolled) sour natural gas well reacts with the lead(II) chromate pigment in paint on homes.

___H2S(g) + ___PbCrO4(s) 

17. Hydrogen sulfide gas from a wild sour natural gas well with silver in cutlery and ornaments in homes.

___H2S(g) + ___Ag(s) 

18. A Bunsen burner, gas furnace and gas hot-water tank all burn natural gas. ___CH4(g) + ___O2(g) 

19. This is a fuel used for trailers and where natural gas is not available. ___C3H7NH2(g) + ___O2(g) 

20. Oxygen gas may by produced in the laboratory by heating potassium chlorate. ___KClO3(s)  _KCl(s) + ___O2(g)

21. Limestone mined in Alberta is decomposed by heating to produce lime. ___CaCO3(s)  ___CaO(s) + ___CO2(g)

22. The reaction of magnesium metal and oxygen to form magnesium oxide is used to produce light in disposable flash bulbs.

23. Chlorine gas reacts with an aqueous solution of sodium iodide. The products are aqueous iodine and aqueous sodium chloride.

25. Sulfuric acid, spilled from a battery, reacts with baking soda to produce sodium sulfate, carbon dioxide and water.

26. The roasting of zinc sulfide ore in a smelter involves the heating of the ore in the presence of oxygen to produce zinc oxide and sulfur dioxide gas.

27. Once the protective oxide coating is removed, aluminum metal reacts readily with water to form hydrogen and aluminum hydroxide.

Starting with nitrogen and hydrogen, millions of kilograms of ammonia are produced in Alberta every year for use as a fertilizer. Use this information to answer the next two questions.

28. Write a balanced chemical equation using molecular models of the Dalton Theory.

29. Write the balanced chemical equation using international symbols and states of mater at SATP.

Predict the products and balance the following chemical equations. 30. Ni(s) + ___HCl(aq) 

31. ___Ca(OH)2(s) + ___HCl(aq) 

32. ___I2(aq) + ___NaBr(aq) 

33. ___Cr2O3(s) 

34. ___Fe(s) + ___HCl(aq) 

35. ___C3H6(g) + ___O2(g) 

38. ___KHCO3(s)  ___K2CO3(s) + ___H2O(l) + ___CO2(g)

39. ___H3PO4(aq) + ___NaOH(aq) 

40. ___Ca(NO3)2(aq) + ___Na3PO4(aq) 

Chemical Change -

Evidence of Chemical Change

In this activity you are going to study evidence of chemical change by observing the effects of combining chemicals. All chemical reactions that occur on their own give off or absorb varying, sometimes very small, amounts of energy. A reaction that gives off energy to its surroundings is called exothermic and a reaction that absorbs energy is called endothermic.

Problem: To determine what the evidence of chemical reactions is for each of the following activities.

Safety precautions:

 In this activity you will be working with concentrated NaOH and H2SO4. These chemicals are

fairly strong and can burn your skin and clothes and cause severe eye injuries if they are not handled properly. Don't touch, and if you spill any on your person, wash off immediately with plenty of cool water.

 Wear goggles and protective aprons throughout this activity.

 Do not talk at anything higher than a whisper so you can tell if your teacher is trying to get your attention or someone in the lab needs help.

Design:

List the chemical reaction for each of these activities in the first column. Record all your observations in the second column. Answer questions on a separate sheet of paper.