CHEMISTRY Science 10 – Unit 1

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SCIENCE 10 CHEMISTRY

Objectives: 1. Atoms are neutral. In ions, the number of electrons and protons differ, giving the ion an electrical charge. 2. Compounds containing a metal and a non-metal usually form ionic compounds in which positive and negative ions are connected by ionic bonds. Compounds containing only non-metals form molecules in which the atoms are connected by covalent bonds. 3. Chemical equations are words or symbols that identify the reactants and products in a chemical reaction. 4. The law of conservation of mass states that the total mass of all the reactants in a chemical reaction is equal to the total mass of all the products. 5. The formula of an acid has an H on the left side. The formula of a base has an OH on the right of a metal. A salt is an ionic compound formed from an acid-base neutralization. 6. The pH scale is a way of measuring the concentration of the H+ ion. A neutral solution has a pH = 7, an acidic solution has a pH < 7, and a basic solution has a pH > 7. 7. Chemical reactions can be classified as synthesis, decomposition, single replacement, double replacement, neutralization (acid-base), or combustion. 8. It is possible to predict the identity of the products of a reaction based on its classification and knowledge of the reactants. 9. Factors that affect the rate of a reaction include temperature, concentration, surface area, and the presence of a catalyst. 10. A catalyst is a substance that speeds up the rate of a chemical reaction but is still present in its original amounts at the end of the reaction.

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Vocabulary: atomic number acids catalyst

atoms alcohol catalytic converter

balanced chemical equation bases combustion

Bohr diagram bromothymol blue decomposition

chemical equation concentration double replacement

chemical reaction indigo carmine neutralization (acid-base)

compound litmus paper precipitate

conservation of mass methyl orange rate of reaction

covalent bonding pH indicators single replacement

covalent compound phenolphthalein surface area

electrons salts synthesis

ionic bonding solvent

ionic compounds

ions

Lewis diagram

molecule

neutron

polyatomic

products

proton

reactants

skeleton equation

subscript

symbolic equation

valence electrons

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Section 1.1 Atomic Theory and Bonding An ____________is the ___________________________of an element

that still has the properties of that element. Atoms join together to form

________________________.

An ________________ is a pure substance made of 1 kind of atom.

_____________ are also pure substances, but are made up of

______________________ elements joined in a molecule. An example of a

compound is ___________________.

A _____________________________occurs when the arrangement of

atoms in compounds change to form new __________________________.

Atoms are made up of smaller particles known as

____________________________.

Table. 1.1 Subatomic particles

Name Symbol Electric Charge

Location in the Atom

Relative Mass (amu)

Proton Neutron Electron

The____________________ is at the _____________ of an atom. The

nucleus is composed of _________________ and _______________.

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Electrons exist in the ________________________ the nucleus. The

number of protons ________________the number of electrons in every atom.

The nuclear charge is the charge on the nucleus, which equals the number of

protons. The _______________________ is equal to the number of

protons, which _________________ to the number of electrons.

Drawing Atoms:

_____________________

show how many electrons appear in each Bohr Diagram

electron shell around an atom. Each shell holds a _________

number of electrons. Electrons in the outermost shell are

called ______________________. There is a maximum of

_____ electrons in the first shell, ______ in the second shell,

and ____ in the third shell.

The _____________________ = # of shells in an ___________.

_____________________only show the electrons Lewis Diagram

that appear in the valence shell.

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Bohr Diagram Example Lewis Diagram Example

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Assignment 1.1 A Atomic Structure (Bohr and Lewis Diagrams) Construct a model of the atom for each of the first 20 elements of the Periodic Table using the Bohr diagram and then the Lewis diagram. Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons - Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons - Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons -

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Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons - Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons - Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons -

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Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons - Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons - Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons -

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Name - Name - Symbol - Symbol - Atomic # - Atomic # - Atomic Mass - Atomic Mass - Protons - Protons - Electrons - Electrons - Neutrons - Neutrons -

Organization of the Periodic Table The ____________________________ organizes all know elements. Elements are

listed in order by their atomic number. ____________________ are on the

_________ (the transition metals range from group 3 to group 12),

________________________ are on the ___________ and the

_______________________ form a “______________________” in the

middle.

Rows of elements (across) are called __________________. The period number is equal

to the number of _____________ in the atom. Columns of elements are called

____________ or _______________. All elements in a family have similar

properties and bond with other elements in similar ways.

Group 1 – ________________ Group 2 – ________________________

Group 17 – ________________ Group 18 – ________________________

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Practice Problems: Use the periodic table in your data booklet to answer the following questions. 1. Based on the patterns of the periodic table, identify the number of occupied shells for each of the following elements. a) calcium, Ca - c) sulphur, S -

b) krypton, Kr - d) iodine, I -

2. Based on the patterns in the periodic table, identify the number of valence electrons for each of the following elements. a) chlorine, Cl - c) strontium, Sr -

b) magnesium, Mg - d) bromine, Br -

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How Atoms Form Compounds

Atoms are electrically neutral. Atoms can gain or lose electrons or share electrons to form

bonds. When atoms gain or lose electrons they become electrically charged particles called

___________. Metals lose electrons and become positive ions (___________), Na+.

Non-metals gain electrons and become negative ions (_____________), Cl-.

Atoms do this in an attempt to have the same number of valence electrons as the nearest noble

gas. The goal of all atoms is to get a complete outer shell. This is why chemical reactions occur.

Atoms can complete their outer shell by ________________electrons,

_____________electrons or ___________________ electrons.

Ionic Bonds Ionic bonds are formed between metals and non-metals. Metals lose electrons and non-metals gain electrons. Bohr diagram Lewis diagram BeF2 MgO

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Covalent Bonds Covalent bonds are formed between non-metals ONLY. Electrons are always shared between non-metals. No ions are formed. Bohr diagram Lewis diagram H2O NH3

Diatomic Elements Gases, when by themselves, form diatomic molecules. H2, N2, O2, F2, Cl2, Br2, I2

Oxidation and Reduction: (Ions are Charged)

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Assignment 1.1 B Modelling Compounds Use these charts to draw Bohr diagrams and Lewis diagrams for modelling compounds. 1) Bohr Diagrams

Hydrogen Lithium Magnesium Oxygen Chlorine Fluorine

2) Li2O Ionic Compounds: 3 a) LiCl b) MgO c) MgCl2 Covalent Compounds: 4 a) HF b) CH4 c) OF2

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5) Lewis Diagrams

Hydrogen Lithium Magnesium Oxygen Chlorine Fluorine

Ionic Compounds: 6 a) Li2O b) LiCl c) MgO d) MgCl2

Covalent Compounds: 7 a) HF b) H2O c) OF2 What did you find out? 1. Describe the information contained in a Bohr diagram compared with the information contained in a Lewis diagram.

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2. a) Which diagram do you find easier to use, Bohr diagram or a Lewis diagram? Why?

Chemical Reactions Exothermic Reaction - Zn + HCl H2 + ZnCl2 H2 + O2 H2O Endothermic Reaction - (NH4)(SCN) + Ba (OH)2 (NH4)(OH) + Ba (SCN)2

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Section 1.1 Review

The atom and the subatomic particles 1. Examine the periodic table for the element below and complete the blanks.

a) atomic number - d) number of protons -

b) average atomic mass - e) name of the element -

c) ion charge - f) number of neutrons -

2. Complete the following table for the different atoms and ions. The first two rows have been completed to help you.

Element

Name

Atomic

Number

Ion Charge Number of

Protons

Number of

Electrons

Number of

Neutrons

potassium 19 1+ 19 18 20

phosphorus 15 0 15 15 16

3 0

2+ 20

nitrogen 3-

5 0

argon 18

13 10

chlorine 0

11 10

Bohr Diagrams 3. Define the following terms: a) Bohr diagram -

b) stable octet -

c) valence shell -

d) valence electrons -

35 -

Br Bromine 79.9

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4. Complete the following table.

Aton/Ion Atomic Number

Number of Protons

Number of Electrons

Number of Neutrons

Number of Electron shells

neon atom

fluorine atom

fluorine ion

sodium atom

sodium ion

5. Use the table above to draw the Bohr model diagram for each of the following atoms and ions.

neon atom fluorine atom fluorine ion sodium atom sodium ion

6. Draw the Bohr model diagram for each of the following compounds.

Ammonia (NH3) Calcium chloride (CaCl2)

Lewis Diagrams 7. Define the following terms: a) Lewis diagram -

b) lone pair -

c) bonding pair -

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8. Draw Lewis diagrams for each of the following elements.

a) boron b) nitrogen c) aluminum d) chlorine

9. Draw Lewis diagrams for each of the following ionic compounds.

sodium oxide potassium chloride

10. Draw Lewis diagrams for each of the following covalent compounds.

Phosphorus trifluoride, PF3 Silicon tetrachloride, SiCl4

11. Draw Lewis diagrams for each of the following diatomic molecules.

Chlorine, Cl2 Nitrogen, N2 Hydrogen, H2

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Section 1.2 Names and Formulas for Ionic and Covalent Compounds

Names and Formulas for Ionic Compounds Naming:

1. Write the name of the positive ion first. 2. Write the name of the negative ion second. 3. Change the ending of the negative ion to ‘ide’, NaCl will be sodium chloride.

BaF2 -

K3P -

Formula: 1. First identify each ion and its charge. 2. Write the symbol for the positive ion and then the symbol of the negative ion. 3. Next, determine the total charges needed to balance the +ve with –ve. You must use

subscripts to show how many of each ion is required to balance the ionic charges. Sodium - Na+ and chlorine – Cl - will be NaCl. Magnesium phosphide - Aluminum bromide -

Assignment 1.2 A Writing Chemical Compounds Write the formula and name of each of the following. 1. magnesium and chlorine

2. hydrogen and iodine

3. boron and fluorine

4. carbon and chlorine

5. lithium and nitrogen

6. barium and sulfur

7. potassium and oxygen

8. magnesium and bromine

9. zinc and chlorine

10. sodium and oxygen

11. zinc and nitrogen

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12. magnesium and phosphorus

13. aluminum and carbon

14. cesium and oxygen

15. lithium and fluorine

Write the name of each of these compounds.

1. Li2O

2. Al2O3

3. H3N

4. RbF

5. AgI

6. AlBr3

7. MgSe

8. Ca3P2

9. Ca2C

10. KH

11. ZnTe

12. Cs2Se

13. Be3As2

14. SrCl2

15. NaF

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Naming Ionic Compounds Where the Metal has More Than One Combining Capacity A metal that has more than one combining capacity is called ____________________.

The rules for writing the formulas for compounds containing multivalent metals are the same

as we have just learned. Except that the metal will need a _______________________

to indicate the charge of the metal ion.

Roman Numerals: I II III IV V VI VII VIII IX X

Formula: 1. First identify each ion and its charge. For the metal use the Roman Numeral. 2. Write the symbol for the positive ion and then the symbol of the negative ion. 3. Next, determine the total charges needed to balance the +ve with –ve. You must use

subscripts to show how many of each ion is required to balance the ionic charges.

For Example:

Chromium III oxide -

Tin IV sulfide -

Writing the name for these ionic compounds begins with_______________________.

Verifying that it has more than one combining capacity. Then we must determine the

combining capacity of the metal by looking at the combining capacity of the negative ion.

Recall that a formula must have an overall charge of ___________. Determine the charge

on the positive ion (the metal) that __________________the charge on the

negative ion. Write the name including a roman numeral to indicate the charge of the

metal ion.

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Naming: 1. Write the name of the positive ion first. (check the metal) 2. Write the name of the negative ion second 3. Change the ending of the negative ion to ‘ide’. 4. Add Roman Numeral if needed.

Cu3N - PtS2 -

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Names and Formulas of Compounds Containing Polyatomic Ions A polyatomic ion is an ion that contains many atoms. It’s a group of covalently bonded atoms that carry an overall charge. (See data table) Ex: phosphate PO43- ammonium NH4+

These polyatomic ions behave as a single unit when combining with other elements to form compounds.

Writing Names of Compounds with Polyatomic Ions Check to see if the first ion in the formula has more than one combining capacity. If so you

must use the ____________________________. Pick out the polyatomic ion (most

are –ve). You will know if there is a polyatomic ion involved if there are more than two capital

letters in the formula. Write the name of the compound using the name of the polyatomic.

_____________________change the ending of the polyatomic ion.

FeSO4 -

Pb(OH)2 -

Writing Formulas of Compounds with Polyatomic Ions When writing the formula for compounds with polyatomic ions you start by identifying each

ion and its charge. Then you determine the total charges needed to balance + ve and – ve.

Next you use subscripts and __________________________to write the formula.

Sodium bicarbonate - Calcium hydroxide - Mercury II bicarbonate -

Assignment 1.2 B Multivalent elements and Polyatomic Compounds Write the formula for each of these names. 1. Hydrogen sulfate 2. Zinc hydroxide 3. Magnesium oxide 4. Gold I chloride 5. Calcium chloride 6. Barium nitrate

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7. Copper II bicarbonate 8. Chromium III sulfate 9. Aluminum oxide 10. Lead IV phosphate 11. Mercury II carbonate 12. Copper II fluoride 13. Nickel II hydroxide 14. Sodium iodide 15. Tin II chloride Write the name of the following compounds. 1. NaCl 2. CuSO4 3. K2O 4. Ca3N2 5. Fe(OH)2 6. PbS2 7. (NH4)2S 8. HgNO3 9. PbO2 10. ZnO 11. Pb3(PO4)4 12. Fe(HCO3)2 13. Na2CO3 14. Al(OH)3 15. Pb(NO3)2

Names and Formulas of Covalent (Molecular) Compounds The formulas of covalent compounds show the _____________________________ of each element in a molecule. Subscripts are not used to show the ratio of atoms the same way as ionic formulas. First you write the name of the first element. Then you name the second element and change the ending to ‘ide’. Next, you add _____________________to indicate the number of atoms of each element. (See data booklet) Mono – one, di – two, tri – three etc. H2O2 -

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NH3 - N2O3 - CO - CH4 -

Writing Formulas for Covalent Compounds When writing the formulas for covalent compounds you write the symbols of the elements and then use the prefixes as the subscripts. DO NOT _______________________the formula. carbon tetrachloride - diiodine hexafluoride - sulphur trioxide - dinitrogen tetrasulphide - oxygen difluoride -

How do you know if a compound is ionic or covalent? 1. Examine the formula to see if a __________________ or the _______________________ is present in the front, then the compound is ionic. If the compound is ionic then we follow the ionic rules. 2. If the compound is covalent use prefixes. Indicate whether these compounds are ionic or covalent. (NH4)2S - FeF2 - NBr3 - SCl2 - NaNO3 -

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Practice Problems: Write the name of each covalent compound. 1. N2O - 2. PI3 - 3. N2O4 - 4. P4S10 - 5. CO2 - 6. OCl2 - Write the formulas for each covalent compound. 1. nitrogen tribromide - 2. xenon hexafluoride - 3. dichlorine monoxide - 4. phosphorus pentabromide - 5. carbon tetraiodide - 6. dinitrogen trioxide - Indicate whether these compounds are ionic or covalent. 1. Na2Cr2O7 - 2. N2O3 - 3. SO2 - 4. Li2SO4 - 5. PbCO3 - 6. SBr2 -

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Section 1.2 Review: Multivalent metals and Polyatomic ions 1. Define the following terms:

a) ionic compound -

b) multivalent metal -

c) polyatomic ion -

2. Write the formulae and names of the compounds with the following combination of ions. The first row is completed to help guide you.

Positive ion

Negative ion

Formula Compound name

a) Pb2+ O2- PbO Lead II oxide

b) Sb4+ S2-

c) Mo2S3

d) Rh4+ Br-

e) Copper I telluride

f) NbI5

3. Write the chemical formula for each of the following compounds. a) manganese II chloride -

b) chromium III sulphide -

c) vanadium V oxide -

d) platinum IV nitride -

e) nickel II cyanide -

4. Write the formulae for the compounds formed from the following ions. Then name the compounds.

Ions Formula Compound Name

a) K+ NO3- KNO3 Potassium nitrate

b) Ca2+ CO3-

c) Li+ HSO4-

d) Sr2+ CH3COO-

e) NH4+ Cr2O72-

f) Ba2+ CrO42-

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5. Write the chemical formula for each of the following compounds. a) barium bisulphate -

b) potassium chromate -

c) potassium hydroxide -

d) aluminum sulphate -

e) silver nitrite -

f) calcium phosphate -

g) ammonium hydrogen carbonate -

h) calcium cyanide -

Chemical names and formulas of Ionic compounds. 1. Write the name for each of the following compounds.

a) Hg3N2 -

b) CoBr2 -

c) Cr2(SO4)3 -

d) K2Cr2O7 -

e) Pb(CO3)2 -

f) Mg(CN)2 -

g) Bi3(PO4)5 -

h) RbClO2 -

2. Write the chemical formula for each of the following compounds.

a) platinum II sulphide -

b) titanium IV nitrite -

c) ammonium sulphate -

d) sodium acetate -

e) tin II permanganate -

f) chromium II chloride -

g) iron III bisulphate -

h) gold III sulphate -

i) lead II hydrogen sulphate -

j) zinc phosphate -

k) aluminum bromide -

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Chemical names and formulas of Covalent compounds.

1. What is a covalent compound?

2. What type of bond is formed in a covalent compound?

3. What is used in naming covalent compounds?

4. Write the chemical formula for each of the following compounds.

a) silicon dioxide -

b) arsenic trichloride -

c) chlorine heptaoxide -

d) dinitrogen monoxide -

e) arsenic pentachloride -

f) disulphur pentaoxide -

g) diphosphorus octaoxide -

h) dinitrogen trioxide -

i) diiodine hexachloride -

Ionic or Covalent?

a) FeF2 -

b) SCl2 -

c) NaNO3 -

d) NBr3 -

e) Cr2O3 -

f) N2O3 -

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Section 1.3: Writing and Balancing Chemical Reactions

Law of Conservation of Mass Matter is not gained or lost in a chemical reaction. Atoms rearrange themselves to form

products, they do not disappear or appear. This means the mass of the reactants in a

chemical change _______________________________the mass of the

products.

Demonstration:

Writing and Balancing Reactions Like formulas, chemical reactions need to be “balanced”. The number of atoms and molecules on _______________________________ of the equation must be in correct proportions. Chemical equation Word equation: sodium + chlorine gas sodium chloride

reactants products

Symbolic equation: Skeleton equation Na (s) + Cl2 (g) NaCl (s)

Balanced equation 2 Na (s) + Cl2 (g) 2 NaCl (s)

Coefficients

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General Steps for Balancing Equations

1) Write the correct formulas for reactants and products.

2) Count the number of atoms of each element on each side of the arrow.

3) Make one change to the coefficients.

4) Recount after each change. Keep repeating steps 3 and 4 until the same

number of atoms

of each element appear on each side of the arrow.

HINTS:

* NEVER change subscripts in formulas.

* Balance polyatomic ions as a unit only if the same ion appears on both sides.

* Often oxygen and hydrogen appear in more than one compound. If so save

them until last.

When balancing Hs and Os, do the hydrogen’s first.

H2SO4 H2 + SO4 KCl + O2 KClO3 Diatomic molecules - gases that when by themselves come in pairs. H2, N2, O2, F2, Cl2, Br2, I2 Example: When zinc is placed in hydrochloric acid, zinc chloride and hydrogen gas are formed. Practice Problems: 1) Fe + Br2 FeBr3 2) Al2O3 Al + O2 3) Ca(OH)2 + HCl CaCl2 + H2O 4) NH4Cl + CaO CaCl2 + NH3 + H2O

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Assignment 1.3 A Balancing Reactions 1. Na + Cl2 NaCl

2. Ca + O2 CaO

3. O2 + K K2O

4. Fe + Br2 FeBr3

5. P4 + Mg Mg3P2

6. BeO O2 + Be

7. NiF3 F2 + Ni

8. Al2O3 Al + O2

9. MnBr7 Br2 + Mn

10. SnN Sn + N2

11. NaCl + K KCl + Na

12. CaO + Cl2 CaCl2 + Br2

13. LiBr + O2 Li2O + Br2

14. I2 + CaS CaI2 + S8

15. AlF3 + N2 AlN + F2

16. LiF + MgO MgF2 + Li2O

17. CaCl2 + MgO CaO + MgCl2

18. Al2O3 + KBr AlBr3 + K2O

19. Fe2O3 + CuI2 FeI3 + CuO

20. Cu(OH)2 + Li2S LiOH + CuS

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Assignment 1.3 B: Balancing Chemical Equations

A. Balance the following equations:

1) K + H2O KOH + H2

2) H2CO3 + Mg(OH)2 MgCO3 + H2O

3) H3PO4 + KOH K3PO4 + H2O

4) Na + H2O NaOH + H2

5) NH4OH + H3PO4 (NH4)3PO4 + H2O

6) Al(OH)3 + H2CO3 Al2(CO3)3 + H2O

B. Write the formula for the equation, then balance it:

1) Calcium oxide + Water Calcium hydroxide

2) Phosphorus + Chlorine Phosphorus trichloride

3) Mercury + Oxygen Mercury I oxide

4) Mercury + Iodine Mercury I iodide

5) Silver nitrate + Sodium chloride Silver chloride + Sodium nitrate

6) Potassium bromide + Chlorine Potassium chloride + Bromine

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7) Potassium + Water Potassium hydroxide + Hydrogen

8) Sodium hydroxide + Hydrochloric acid Sodium chloride + Water

9) Sodium chloride + Hydrogen sulphate Sodium bisulphate + Hydrochloric acid

10) Hydrogen sulphide + Oxygen Water + Sulfur dioxide

11) Iron II sulphide + Hydrogen sulphate Iron II sulphate + Hydrogen sulfide

12) Copper II oxide + Hydrogen Copper + Water

13) Silicon dioxide + Hydrogen fluoride Silicon tetrafluoride + Water

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Assignment 1.3 C: Writing and Balancing Reactions

Write and balance the following chemical reactions.

1) Heating copper and sulfur forms copper I sulfide

2) Heating potassium chlorate forms potassium chloride and oxygen

3) When carbon monoxide comes in contact with iron III oxide, iron II oxide and

carbon dioxide are formed.

4) Zinc sulfide and oxygen forms zinc oxide and sulfur dioxide.

5) When barium oxide comes in contact with carbon and nitrogen, barium cyanide and

carbon monoxide are formed.

6) Heating ammonium nitrate forms nitrogen and water and oxygen.

7) Calcium hydroxide neutralizes hydrochloric acid (HCl) and then forms calcium chloride

and water.

8) Adding sodium nitrite to ammonium chloride forms nitrogen gas, sodium chloride and

water.

9) By adding ammonium chloride to calcium oxide, the result is calcium chloride,

ammonia (NH3) and water.

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Section 1.3 Review:

Balancing Equations – Starting with the skeleton equations, balance the following equations by adding coefficients where appropriate. 1. H2 + F2 HF

2. Sn + O2 SnO

3. MgCl2 Mg + Cl2

4. KNO3 KNO2 + O2

5. BN + F2 BF3 + N2

6. CuI2 + Fe FeI2 + Cu

7. Li + H2O LiOH + H2

8. NH3 + O2 N2 + H2O

9. V2O5 + Ca CaO + V

10. C9H6O4 + O2 CO2 + H2O

11. H2S + PbCl2 PbS + HCl

12. C3H7OH + O2 CO2 + H2O

13. Zn + CuSO4 Cu + ZnSO4

14. C6H12O6 + O2 CO2 + H2O

15. C2H5OH + O2 CO2 + H2O

16. Al + H2SO4 H2 + Al2(SO4)3

17. FeCl3 + Ca(OH)2 Fe(OH)3 + CaCl2

18. Pb(NO3)2 + K2CrO4 PbCrO4 + KNO3

19. Cd(NO3)2 + (NH4)2S CdS + NH4NO3

20. Ca(OH)2 + NH4Cl NH3 + CaCl2 + H2O

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Word Equations Write the skeleton equation for each of the following reactions, and then balance each. 1. hydrogen + oxygen water 2. iron III oxide + hydrogen water + iron 3. sodium + water sodium hydroxide + hydrogen 4. calcium carbide + oxygen calcium + carbon dioxide 5. potassium iodide + chlorine potassium chloride + iodine 6. chromium + tin II chloride chromium III chloride + tin 7. magnesium + copper II suphate magnesium sulphate + copper 8. zinc sulphate + strontium chloride zinc chloride + strontium sulphate 9. ammonium chloride + lead III nitrate ammonium nitrate + lead III chloride 10. iron III nitrate + magnesium sulphide iron III sulphide + magnesium nitrate 11. aluminum chloride + sodium carbonate aluminum carbonate + sodium chloride 12. sodium phosphate + calcium hydroxide sodium hydroxide + calcium phosphate

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Investigation Activity Lab 1.3: Balancing Chemical Reactions Purpose: Carry out 3 chemical reactions. Write the products and balance the chemical reactions. Apparatus and Materials: - sodium sulfide solution - cobalt II chloride solution - petri dish - lead II nitrate solution - potassium iodide solution - droppers - potassium sulfate solution - barium chloride solution -safety glasses Caution Some of these chemicals are poisonous. Wear safety glasses. If any of these chemicals get on the skin or in the eyes, rinse the areas immediately with water and inform the teacher! Procedure: Part 1: sodium sulfide and cobalt II chloride 1. Place a few drops of sodium sulfide into a petri dish. 2. Add a few drops of cobalt II chloride. Record observations. (2) When a substance cannot dissolve in water, it is called insoluble. An insoluble substance that is formed when solutions are mixed is called a precipitate. Part 2: potassium iodide and lead II nitrate 1. Place a few drops of potassium iodide into a petri dish. 2. Add a few drops of lead II nitrate. Record observations. (2) Part 3: potassium sulfate and barium chloride 1. Place a few drops of potassium sulfate into a petri dish. 2. Add a few drops of barium chloride. Record observations (2) Questions 1. What evidence is there of a chemical reaction in each case? (2) 2. Write a balanced chemical equation for each reaction. (6) Conclusion What is a precipitate? (2) What evidence is there that a chemical reaction has occurred? (2) What is a reactant in a chemical equation? (2) What is a product in a chemical equation? (2)

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Section 2.1 Acids and Bases Acids and bases are classified on ___________________________________and should

never be identified by taste or touch.

Acids and bases are dissolved in water so are often written with the _____________ (aq).

The strength of an acid or base is indicated on the____________________. The pH is

measured by other chemicals called ___________________________, or by a pH

meter. pH indicators ____________________________ to indicate whether a

solution is acidic or basic.

Commonly used indicators include:

Litmus paper – blue litmus turns red in acid and red litmus turns blue in a base

Phenolphthalein – turns pink in a base

Purple cabbage juice – turns red in an acid, blue in a neutral, and green in a base

Acids generally begin with an ______and release hydrogen ions, H+ (aq).

Bases generally end with an _______ and release hydroxide ions, OH- (aq).

Acids and bases react to form salt and water, they cancel out each other or

____________________________each other to form water,

H + OH HOH or H2O.

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Demonstration

pH Blue litmus Red litmus

HCl

Water

NaOH

Purple Cabbage Juice

Control Acid Neutral Base

Other Chemicals Tested

Chemical Acid Neutral Base

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Acids 1. Turn blue litmus paper red, purple cabbage juice turns red, pH scale: below 7

2. Acids taste sour: lemons

3. Acids corrode active metals and produce hydrogen gas

Example: Zn + 2 HCl ZnCl2 + H2 (g)

4. Acid and water solutions conduct electricity

Common Acids Name Formula Use Hydrochloric HCl (aq) Produced in the stomach

Acetic HC2H3O2 (aq) Present in vinegar

Nitric HNO3 (aq) Used to make fertilizer

Carbonic H2CO3 (aq) Used in sodas and soft drinks

Sulfuric H2SO4 (aq) Used in automobile batteries

Phosphoric H3PO4 (aq) Used in dental products and cosmetics

Naming Acids 1. Naming acids containing halogens. (N, O, F, Cl, Br, I) Ex: HCl hydrogen chloride hydrochloric acid

HI hydrogen iodide hydroiodic acid

HBr 2. Naming acids containing polyatomics with an ‘ate’ ending Ex: H2SO4 hydrogen sulphate sulfuric acid

H2C2O4 hydrogen oxalate oxalic acid

H2CrO4 3. Naming acids containing polyatomics with an ‘ite’ ending Ex: H2SO3 hydrogen sulphite sulphurous acid

HNO3 4. Naming other acids - just name the compounds as normal

Ex: HTe hydrogen telluride

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Bases:

1. Turn red litmus paper blue, phenolphthalein turns pink, purple cabbage just turns green, pH scale: more than 7

2. Feel slippery

3. Bitter taste: baking soda

4. Basic solutions conduct electricity

5. Neutralize acids

Acid + Base Salt + Water Common Bases Name Formula Use

Sodium hydroxide NaOH (aq) Oven cleaner

Potassium hydroxide KOH (aq) Drain cleaner

Ammonium hydroxide NH4OH (aq) Household cleaner

Calcium hydroxide Ca(OH)2 (aq) Lime

Magnesium hydroxide Mg(OH)2 (aq) Laxative

Aluminum hydroxide Al(OH)3 (aq) Antacid

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Investigation Activity Lab 2.1: Acid Versus Base Purpose: To neutralize an acid with a base. To determine the products of an acid/base reaction. Apparatus and Materials: - evaporating dish - acid solution - base solution - stirring rod - hot plate - phenolphthalein - tons -safety glasses Caution Wear safety glasses. If any of these chemicals get on the skin or in the eyes, rinse the areas immediately with water and inform the teacher! Procedure: Part 1: sodium sulfide and cobalt II chloride 1. Measure 20 drops of acid into an evaporating dish. Add a drop of phenolphthalein. Note and record the color change. (2) 2. Add the base solution, one drop at a time, stirring after each drop. Continue to add base, one drop at a time until a faint pink color appears. Record the number of drops of base needed. (2) 3. Add the smallest amount of acid possible with the dropper to just remove the pink color. 4. Place the evaporating dish on a hot plate and evaporate the solution. Examine the solid residue. Describe and sketch the residue. (4) Questions 1. Consider phenolphthalein’s reaction with acids and bases, what kind of substance is it? (2) 2. The reaction in this activity yields hydrogen ions from the acid (HCl) and hydroxide ions from the base (NaOH)? What is the probable formula for this compound when these ions combine? (2) What is the name of this substance? (2) What is the state of this substance at room temperature? (2) 3. What is the probable formula for the other compound formed? (Sodium ions from the base and the chloride ions from the acid?) (2) What is the name of this substance? (2) What is the state of this substance at room temperature? (2) Conclusion Why did we use phenolphthalein in this experiment? (2) Give a possible equation to describe the neutralization of an acid with a base. (4)

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Section 2.1 Review: Acids and Bases

pH scale and pH indicators

1. Define the following terms: a) pH indicator -

b) pH scale -

2. Complete the following table by using page 2 of your data book. Identify whether the substance is an acid or base and indicate what color the pH indicator will turn. a)

Substance pH value Acid or base Methyl orange Bromothymol blue Litmus

lemon

ammonia

milk

b)

Substance pH value Acid or base Methyl red Phenolphthalein Indigo carmine

tomato

oven cleaner

egg

3. Complete the following table. Identify whether the substance is an acid or a base and indicate what color the pH indicator will turn.

Substance pH Value Acid or Base pH indicator Color of pH indicator

black coffee 5 litmus

milk of magnesia

10 phenolphthalein

battery acid 0 bromothymol blue

sea water 8 indigo carmine

orange juice 3 Methyl orange

liquid drain cleaner

14 Methyl red

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Names of Acids

1. An acid will have the suffix “-ic acid” at the end of its name when the negative ion has a suffix . For example, “hydrogen carbonate (H2CO3)” is called “carbonic acid”. 2. An acid will have the suffix “-ous” at the end of its name when the negative ion has a suffix . For example, “hydrogen sulphite (H2SO3)” is called “sulphurous acid”. 3. What is the name of each of the following acids? a) H2CO3 -

b) CH3COOH -

c) H3PO4 -

d) HClO2 -

e) H2SO3 -

f) HNO3 -

g) HF -

h) HCl -

4. What is the chemical formula for each of the following acids? a) hydroiodic acid -

b) sulphuric acid -

c) perchloric acid -

d) nitrous acid -

e) chloric acid -

f) hydrobromic acid -

g) phosphorous acid -

h) hypochlorous acid -

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Acid versus bases 1. Compare and contrast acids and bases by completing the following table.

Acids Bases

definition

pH

what to look for in chemical formula

production of ions

electrical conductivity

taste

touch

examples

2. Classify each of the following as an acid or a base. a) H3PO4 -

b) NH4OH -

c) Mg(OH)2 -

d) has a pH of 4 -

e) has a pH of 9 -

f) sulphurous acid -

g) hydrogen bromide -

h) potassium hydroxide -

i) causes methyl orange to turn red -

j) causes phenolphthalein to turn pink -

k) causes indigo carmine to turn yellow -

l) causes bromothymol blue to turn yellow -

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Section 2.2: Salts Salts are _____________________________ formed when acids and bases react.

Table salt, _________, is found in seawater, salt lakes or rock deposits. NaCl is only

_______________ of salt.

A salt is made up of a ______________________from a base and

________________________ from an acid.

Neutralization reactions occur when an acid and a base

__________________________________a salt and water.

HCl (aq) + NaOH (l) NaCl (s) + H2O (l)

Demonstration:

NaHCO3 + CH3COOH H2O + NaCH3COO + CO2

Section 2.2 Review:

Recognizing acids, bases, and salts 1. State whether each of the following is an acid, a base, or a salt. a) HI - l) Al2(SO4)3 -

b) HBr - m) CH3COOH -

c) KOH - n) Mg(CH3COO)2 -

d) HNO3 - o) calcium nitrate -

e) NaOH - p) sodium chloride -

f) H2SO4 - q) hydrocyanic acid -

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g) H2CO3 - r) hydrogen fluoride -

h) H3PO4 - s) barium hydroxide -

i) Na3PO4 - t) hypochlorous acid -

j) Sr(OH)2 - u) aluminum hydroxide -

k) Ca(OH)2 - v) magnesium carbonate -

2. What acid is present in vinegar?

3. What is the chemical name for table salt?

4. What acid is used in automobile batteries?

5. What base is found in drain and oven cleaners?

6. What base is the active ingredient in some antacids?

7. What acid is produced in the stomach to help digest food?

Acid-base neutralization reactions

Acid + Base Water + Salt

1. Complete and balance the following neutralization reactions. a) H2SO4 + NaOH

b) HNO3 + KOH

c) HCl + Ca(OH)2

d) H3PO4 + Ba(OH)2

e) CH3COOH + NaOH

f) HNO3 + Sr(OH)2

g) HF + Fe(OH)3

h) HBr + Sn(OH)4

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Section 3.1: Types of Chemical Reactions Chemical reactions can be classified as one of_____________________________:

synthesis, decomposition, single replacement, double replacement,

neutralization (acid-base), and combustion. You can identify each type of reaction by

examining the reactants. This makes it possible to classify a reaction and then predict the

identity of products.

1) Synthesis is a reaction where two or more reactants combine to produce a single product.

A + B AB ex: C + O2 CO2 If the reaction gives off energy, it is_________________________. If the reaction requires energy, it is _________________________. 2) Decomposition is the breaking down of a compound into smaller compounds or separate elements. A decomposition reaction is the reverse of a synthesis reaction. AB A + B ex: 2 NaCl 2 Na + Cl2 3) Single Replacement is an element (a metal or a non-metal) and a compound that react to produce another element and another compound. An element replaces an element that is part of a compound. A + BC B + AC

ex: Fe + CuSO4 Cu + FeSO4 ex: F2 + 2 NaI I2 + 2 NaF

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4) Double Replacement reactions involve two ionic solutions that react to produce two other ionic compounds. One of the new compounds forms __________________. (Lab 1.3) AB + CD AD + CB

ex: 2 KI + Pb(NO3)2 2 KNO3 + PbI2 5) Neutralization reactions occur when an acid and a base react to produce a salt and water. HA + BOH BA + H2O

ex: HCl + NaOH NaCl + H2O 6) Combustion is the rapid reaction of a compound or element with oxygen to form an oxide and water and produce heat. CxHy + O2 CO2 + H2O

ex: CH4 + O2 CO2 + H2O

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Practice Problems:

Classifying chemical reactions Classify each of the following reactions as synthesis (S), decomposition (D), single replacement (SR), double replacement (DR), neutralization (N), or combustion (C). 1. 2H2 + O2 2H2O -

2. 2 Al + 3 CuCl2 2AlCl3 + 3Cu -

3. CaI2 + Na2CO3 2 NaCl+ CaCO3 -

4. 2 C6H6 + 15 O2 12 CO2 + 6 H2O -

5. 2 FeBr3 + 2 Fe + 3Br2 -

6. 2 HCl + Pb(NO2)2 2 HNO2 + PbCl2 -

7. 2 Cr + 3 F2 2 CrF3 -

8. 2 NH4Br + Cl2 2 NH4Cl + Br2 -

9. Ca(OH)2 + H2SO4 2 H2O + CaSO4 -

10. 2 N2O 2 N2 + O2 -

11. H2SO4 + Ca(OH)2 CaSO4 + 2 H2O -

12. C2H6O + 3 O2 2 CO2 + 3 H2O -

Predict the product for each reaction and then balance the equation. 1. Mg + N2

2. K3PO4 + MgI2

3. PbCl4 + Al

4. AuCl3

5. H3PO4 + Mg(OH)2

6. C2H4 + O2

7. Cl2 + CsBr

8. Cs + P4

9. C6H12O6 + O2

10. Ca3N2

11. Al(OH)3 + HClO4

12. AgNO3 + Na2CrO4

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Section 3.1Review: Classifying Chemical Reactions

Classifying chemical reactions Classify the following reactions as synthesis (S), decomposition (D), single replacement (SR), double replacement (DR), neutralization (N), or combustion (C). Place the correct letter representing the reaction type in the space provided. Then balance the chemical equation by placing the correct coefficients in the equation. 1) N2 + F2 NF3

2) KClO3 + KCl + O2

3) C12H22O11 + O2 CO2 + H2O

4) CuSO4 + Fe Fe2(SO4)3 + Cu

5) MgF2 + Li2CO3 MgCO3 + LiF

6) H3PO4 + NH4OH H2O + (NH4)3PO4

7) NaF + Br2 NaBr + F2

8) CH3OH + O2 CO2 + H2O

9) HI + H2 + I2

10) H2SO4 + KOH H2O + K2SO4

11) RbNO3 + BeF2 Be(NO3)2 + RbF

12) S8 + H2 H2S

Types of chemical reactions – Word Equations Classify the following reactions as synthesis (S), decomposition (D), single replacement (SR), double replacement (DR), neutralization (N), or combustion (C). Place the correct letter representing the reaction type in the space provided. Then write a balanced equation for each word equation. 1. magnesium + Sulphur magnesium sulphide 2. potassium hydroxide + sulphuric acid water + potassium sulphate 3. aluminum chloride + sodium hydroxide aluminum hydroxide + sodium chloride

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4. chlorine + potassium iodide potassium chloride + iodide 5. lead II oxide lead + oxygen 6. methane (CH4) + oxygen carbon dioxide + water

Predicting the products For each of the following: 1) classify the reaction type 2) predict the products 3) balance the equation 1. H2O 2. H2 + Cl2 3. NaI + F2 4. AgNO3 + Na3PO4 5. Ba(OH)2 + H2SO4 6. CH3OH + O2

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Section 3.2 Rates of Reactions

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Reaction Rates Demonstrations

Temperature, Surface Area, Concentration

Products Type

1. CuCl2 + Al -

a) Cold water / low concentration / low surface area

b) Hot water / high concentration / high surface area

Catalyst

Products Type

2. Al + HCl -

a) Without copper

b) With copper

How does increasing the temperature, concentration and surface are affect the reaction rate?

What was the catalyst in the second reaction and how does the catalyst increase the rate of the reaction?

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Factors Affecting the Rate of Reactions

Rate of reaction is how quickly or slowly reactants turn into products, in a chemical reaction. For products to be formed the particles must collide. Factors:

1) __________________________ – the faster the molecules move the more collisions that will occur between the particles. - the higher the temperature the faster the reaction (Rx) goes.

2) ___________________________- is a measure of how much surface area of an object is exposed. - the more surface contact between reactants the more collisions so the faster the Rx.

3) ___________________________ – is the number of particles of each reactant that is involved in the reaction. - the greater the concentration, the more collisions there will be so the faster the Rx rate.

4) _________________________________________– which is a substance that reduces the energy needed to break molecular bonds (ex: enzymes in the body) - the catalyst helps the Rx go faster but is still present in the same amount at the end of the Rx -less energy is needed for a collision to be effective at creating products

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Section 3.2 Review: Rates of Reactions Different rates of reactions 1. Indicate whether each of the following would increase or decrease the rate of reaction. a) adding heat -

b) removing heat -

c) adding a catalyst -

d) diluting a solution -

e) removing an enzyme -

f) lowering the temperature -

g) decreasing the surface area -

h) increasing the temperature -

i) increasing the concentration of a solution -

j) breaking a reactant down into smaller pieces -

2. Identify which situation would have a higher reaction rate. Then state the factor that affected the rate of the reaction in each situation.

Situation X Situation Y Situation with a higher Rx rate

Factor affecting the Rate of Rx

a) 1 g of sugar (cubes) 1 g of sugar (powder)

b) 50oC 0oC

c) low number of particles = few collisions

high number of particles = more collisions

d) enzyme added no enzyme added

e) twigs logs

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Four Factors affecting the rate of reactions

Use the following graph to answer question 1.

1. The graph above shows the difference in the rate of reaction at different temperatures, concentrations, surface area, and the presence or absence of a catalyst. A steeper line represents a greater rate of reaction. Indicate which line (X or Y) each of the following are associated with. a) lower temperature b) higher temperature

c) lower concentration d) higher concentration

e) absence of a catalyst f) presence of a catalyst

g) larger pieces (small surface area)

h) smaller pieces (large surface area)

2. Which of the following four factors affecting reaction rate is most important in each of the following examples? Choose from concentration, temperature, surface area, and catalyst. a) Raw carrots are cut into thin slices for cooking.

b) Protein is broken down in the stomach by the enzyme pepsin.

c) A woolly mammoth is found, perfectly preserved, near the Arctic.

d) More bubbles appear when a concentrated solution of hydrochloric acid is added to a magnesium strip than when a dilute solution of the acid is added.