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.