Handouts Grade 11 University Level

Unit 1Nomenclature

Handout 1 – Grade 11 NomenclatureHandout 2 – Recognizing Patterns 1 and 2Handout 3 – Combining Patterns 1 and 2

Types of ReactionsHandout 4 – Types of ReactionsHandout 5 – Net Ionic Equations

Atomic StructureHandout 6 – History of the Atom (Powerpoint)Handout 7 – The Electromagnetic SpectrumHandout 8 – The Emission Spectrum of HydrogenHandout 9 – Theories of the AtomHandout 10 – Atomic OrbitalsHandout 11 – Aufbau Diagram

Unit 2Periodicity

Handout 12 – Atomic RadiiHandout 13 – 1st Ionization EnergyHandout 14 – Successive Ionization Energies Handout 15 – Electron Affinity/Eletronegativty

BondingHandout 16 – Covalent CompoundsHandout 17 – Ionic to Covalent to Metallic Bonding

Handout 1 – Grade 11 Nomenclature

Hydrated Salts

Some solids are crystals that regularly associate with water- SiO2 placed in shoes to absorb water to protect the leather

- when these compounds are associated with H2O we call them hydrated - when water is removed we call them anhydrous

- to name hydrate compounds we use a prefix (same Greek prefixes from molecular

compound naming) followed by hydrate to indicate the number of H2O molecules associated to each formula unit.

- i.e. MgSO4∙7H2O is the chemical formula for magnesium sulfate heptahydrate- could also be called hydrated magnesium sulfate but the above name is better because indicates number of water molecules.

- CuSO4∙5H2O could be copper (II) sulfate pentahydrate or cupric sulfate pentahydrate

Acids

There are two kinds of acids that you will learn to name Binary acids and Ternary Acids

Binary Acids

- the term binary indicates that a compound has only two types of atoms- all acids must contain H. Therefore, there is only one other atom that can vary

for each binary acid- Binary acids commonly contain elements from the halogen group, but other

examples include sulfur and selenium.- HCl(aq) is a very common binary acid you have used. Remember all acids are

aqueous because they dissolve in water- HCl(aq) is called aqueous hydrogen chloride by the more modern IUPAC

system, or more commonly, hydrochloric acid by the classical system.- Two naming systems for acids

IUPAC (modern) very simple just put aqueous in front of regular chemical name.

Classical system follows this general formula hydro ________ic acidThe blank is filled in with the associated anion name after

removing the “ide”

- an unusual acid is HCN(aq), which is named using binary acid rules

Chemical IUPAC Nomenclature Classical NomenclatureFormula

HCl(aq) __________________________________________________HF(aq) __________________________________________________H2S(aq) __________________________________________________H2Se(aq) __________________________________________________HCN(aq) __________________________________________________

Ternary Acids

- Ternary means 3, therefore this is an acid containing three types of atoms- After the H the rest of the atoms in the ternary acid are polyatomic ions that

contain oxygen, or also called oxyanions. Therefore, ternary acids can also be called oxyacids.

- HNO2(aq) is called - aqueous hydrogen nitrite by the IUPAC system -nitrous acid by the Classical system

Conversion from IUPAC to Classical for ternary acids

1) Replace the words “aqueous hydrogen” with the word “acid” at the end

2) Change the ending: - “ate” to “ic” or “ite” to “ous”

IUPAC Classicalaqueous hydrogen hypo______ite hypo_______ous acidaqueous hydrogen ______ite _______ous acidaqueous hydrogen ______ate _______ic acidaqueous hydrogen per ______ate per _______ic acid

Chemical IUPAC Nomenclature Classical Nomenclature Formula

HNO2(aq) _________________________________________________HBrO2(aq) _________________________________________________H3PO4(aq) _________________________________________________H2CO3(aq) _________________________________________________H2SO4(aq) _________________________________________________HClO4(aq) _________________________________________________

- remember only do classical system rules for binary and ternary acids when you know it is an acid. For example “(aq)” symbols are a good indication the compound is an acid if it also contains an H at the beginning. There are some exceptions such as acetic acid CH3COOH, the last H being the acidic proton.

- example HBrO2(g) would be called hydrogen bromite gas

Handout 2 - Recognizing Patterns #1

Oxyanions (polyatomic ions containing oxygen) have a pattern in their names to indicate amount of oxygens. Look at the list of oxyanions that have chlorine in them and see if you notice the pattern.

- the base ion is the one with “ate” and no prefix- when the suffix “ite” is used, subtract 1 oxygen atom from the base ion- when the prefix “hypo” and the suffix “ite” is used, subtract 2 oxygen atoms- when the prefix “per” and the suffix “ate” is used, add 1 oxygen atom to the base ion

Use your polyatomic ions list to find the formulas for the following “base” ions: carbonate, nitrate, phosphate, sulfate, iodate, bromate.

Use the 6 base polyatomic ions above and fill in the boxes below with the 3 other polyatomic ions that can be known from the base ion.

Name Formula Name Formula Name Formula

Name Formula Name Formula Name Formula

ClO- hypochloriteClO2

- chloriteClO3

- chlorateClO4

- perchlorate

Ion charge is (-1), but oxygens are increasing by 1

Recognizing Patterns #2

Acid Anions

- an acid anion is created when one or more H+ ions covalently bond with an oxyanion (i.e. HCO3

-, HPO42-)

- when acid anions combine with cations, acid salts are created (i.e. CaHPO4)

- using the base polyatomic ions from Recognizing Patterns #1 (carbonate, phosphate, sulfate) and the pattern below you can create the acid anions

Base Ion Acid Anion

+H+

carbonate hydrogen carbonate CO3

2- -2 + 1 = -1 HCO3-

+2H+

phosphate dihydrogen phosphate PO4

3- -3 + 2 = -1 ______

______sulfate hydrogen sulfate SO4

2- ________ ______

______phosphate hydrogen phosphate PO4

3- ________ ______

Handout 3 - Combining Patterns 1 and 2

phosphate PO4

3-

(Pattern 1)

+2H+

Phosphite (Pattern 2) dihydrogen phosphite PO3

3- ________ ________

Sulfate SO4

2-

(Pattern 1)

+H+

Sulfite (Pattern 2) ___________________ SO3

2- ________ ________

Practicse Questions

1. Oxyanion Pattern: Fill in the table below using the patterns for oxyanions

Chemical Formula Chemical NameCalcium hypochlorite

Zn(BrO4)2

Barium phosphateAurous nitrite

Mg(IO)2

Lithium persulfateIron (III) percarbonate

SnSO3 or

2. Oxyanion Pattern: Fill in the table below using the patterns for oxyanions

Chemical Formula Chemical NameSr(HCO3)2

Sodium hydrogen sulfateCu(H2PO3)2 or

Aluminum dihydrogen phosphateRb2HPO4

Gold (I) hydrogen sulfite

Handout 4 - Types of Reactions

1) Synthesis Reactions (A + B AB)

I) Simple Binary Ionic Compounds i.e. solid aluminum reacts with chlorine gas

Al(s) + Cl2(g) AlCl3(s)

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___________________________________________

II) Slightly More complicated Synthesis Reactions

- non-metal oxides such as CO2, SO3, N2O5 react with H2O to form acids

i.e. CO2(g) + H2O(l) H2CO3(aq)

___________________________________________

___________________________________________

- metal oxides such as Li2O, CaO react with H2O to form bases

i.e. CaO(s) + H2O(l) Ca(OH)2(aq)

___________________________________________

___________________________________________

Salts containing oxyanions i.e. Li2CO3(s)

Basei.e.LiOH(aq), Ca(OH)2(aq)

Acid i.e.H2CO3(aq) , H2SO4(aq)

Metal oxides(basic oxides)i.e. Li2O(s) , CaO(s)

Non-metal oxides (acidic oxides)i.e. CO2(g) , SO3(g)

H2O

- non-metal oxides and metal oxides can react to form salts containing oxyanions

i.e. CaO(s) + CO2(g) CaCO3(s)

___________________________________________

___________________________________________

2) Decomposition Reactions (AB A + B)

-reverse of the above reactions-often heat is needed; this is called “thermal decomposition”

I) Simple Binary Ionic Compounds i.e. aluminum chloride is heated

II) Slightly More complicated Reactions

- acids will decompose into non-metal oxide and water

i.e. H2CO3(aq) CO2(g) + H2O(l)

___________________________________________

___________________________________________

- bases will decompose into metal oxides and water

i.e. Ca(OH)2(aq) CaO(s) + H2O(l)

___________________________________________

___________________________________________

- salts containing oxyanions decompose into non-metal oxides and metal oxides

i.e. CaCO3(s) CaO(s) + CO2(g)

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___________________________________________

3) Single Displacement Reactions (AX + B A + BX)

Create your own activity series mini-lab

Hypothesis: Predict the order of reactivity of the 5 metals in the lab from most to least reactive. (Hint: use their position on the periodic table)

Most _____ _____ _____ _____ _____ Least

When complete show your teacher the order and obtain the materials for the lab.

Observation Chart:

Metals

SolutionsIron Magnesium Copper Zinc Calcium

Iron nitrate

Magnesium nitrate

Copper nitrate

Zinc nitrate

Calcium nitrate

Indicate a reaction with a checkmark and no reaction with an X.

Conclusion: Using your observation chart order the metals from most to least reactive

Most _____ _____ _____ _____ _____ Least

From your observations write the products of the following reactions that would react and if there is no reaction indicate no reaction.

Ca(s) + Mg(NO3)2(aq)

Cu(s) + Zn(NO3)2(aq)

Mg(s) + Cu(NO3)2(aq)

Activity Series: is an arrangement of metals in order of their relative reactivities. Knowing the order allows you to predict if a single displacement reaction will take place or not. Any metal higher on the list can displace any metal lower on the list.

Metal Displaces hydrogen from acids

Displaces hydrogen from cold water

Lithium Most ReactivePotassium

BariumCalciumSodium

MagnesiumAluminum

ZincChromium

IronCobaltNickel

TinLead

HydrogenCopperMercury

SilverPlatinum

Gold Least Reactive

Using the activity series which of the following reactions will occur?

a) Au(s) + CuSO4(aq)

Co(s) + HgClO2(aq)

Na(s) + Sn(IO3)2(aq)

b) How could the metal activity series be used to predict reactions with acids? Give 2 examples of reactions between acids and metals that would occur. What gas is produced?

_________________________________________________

_________________________________________________

c) How is the metal activity series used to predict reactions with water? Give 2 examples of reactions with metal and water. What gas is produced?

_________________________________________________

_________________________________________________

d) There is also a Halogen Activity Series. Give an example of 2 reactions that could be predicted by the Halogen Activity Series below.

Halogen Series: Most _____ _____ _____ _____ Least _________________________________________________

_________________________________________________

4) Double Displacement (AX + BY AY + BX)

- always occur between two soluble ionic compounds- there are three possible outcomes a) precipitate forms

b) gas is produced c) water is produced

a) Precipitate Forms

- know how to use the solubility chart below to identify if a solid is produced

Rule ExceptionNitrates (NO3

-) are soluble NoneHalides (Cl-, Br-, I-) are soluble Ag+, Hg2

2+, Pb2+

Sulfates (SO42-) are soluble Ca2+, Ba2+, Pb2+, Hg2

2+, Ag+

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Sulfides (S2-) are insoluble NH4+ and ions of groups 1 and 2 elements

Carbonates (CO32-) are insoluble NH4

+ and ions of group 1 elementsPhosphates (PO4

3-) are insoluble NH4+ and ions of group 1 elements

Hydroxides (OH-) are insoluble Ba2+, Sr2+, Ca2+ and ions of group 1 elements

Use the solubility chart above to identify if the following reactions will produce a precipitate or not. If they do write the products and indicate the precipitate by using a subscript (s).

K2CO3(aq) + CaCl2(aq)

Pb(NO3)2(aq) + KI(aq) b) Gas is Produced

- there are four cases in which a gas is formed. The first 3 cases are a double displacement reaction followed by a decomposition.

i) acids + carbonates ii) acids + sulfites iii) bases + ammonium

- the double displacement reactions produce products such as carbonic acid, sulfurous acid and ammonium hydroxide, which then decompose into gas and water. Try completing the following reactions.

i) CaCO3(s) + HCl(aq)

ii) Na2SO3(aq) + HCl(aq)

iii) NH4Cl(aq) + NaOH(aq)

Come up with your own examples of the three kinds of reactions above

i)_____________________________________________________________

ii)_____________________________________________________________

iii)____________________________________________________________

- the 4th case of a reaction that produces a gas involves acids and sulfides. This case only requires the double displacement, the gas is produced immediately, H2S(g).

iv) Na2S(aq) + HCl(aq)

iv)____________________________________________________________

c) Water is Produced

- these double displacement reactions are more specifically named neutralization reactions. They occur when acids are combined with bases and the products are water and salt. Most times the salt can be labeled aqueous, but make sure by checking the solubility table. Look at the following reactions and write the products and their states.

H3PO4(aq) + NaOH(aq) H2CO3(aq) + CaOH(aq)

Handout 5 – Net Ionic Equations

Writing net ionic equations

1. Write balanced equation first

2. To do your total ionic equationa) Break up all ionic compounds EXCEPT precipitatesb) DO NOT break up molecular compounds EXCEPT acids

3. To get net ionic equation cross out all spectator ions. Because they do not really take part in the reaction.

Example 1Write the balanced net ionic equation for the reaction of aqueous sodium carbonate with aqueous calcium nitrate.

Balanced Chemical Equation

Na2CO3(aq) + Ca(NO3)2(aq) 2NaNO3(aq) + CaCO3(s)

Total Ionic Equation (break into ions)

2Na+(aq) + CO3

2-(aq) +Ca2+

(aq) + 2NO3-(aq) 2Na+

(aq) +2NO3-(aq) + CaCO3(s)

- total charges should be the same on both sides of equation

Cross out spectator ions

2Na+(aq) + CO3

2-(aq) +Ca2+

(aq) + 2NO3-(aq) 2Na+

(aq) + 2NO3-(aq) + CaCO3(s)

Net Ionic Equations

CO32-

(aq) + Ca2+(aq) CaCO3(s)

Balanced Chemical Equation

2NaI(aq) + Br2(aq) 2NaBr(aq) + I2(g)

Total Ionic Equation (break into ions)

2Na+(aq) + 2I-

(aq) + Br2(aq) 2Na+(aq) + 2Br-

(aq) + I2(g)

- total charges should be the same on both sides of equation

Cross out spectator ions

2Na+(aq) + 2I-

(aq) + Br2(aq) 2Na+(aq) + 2Br-

(aq) + I2(g)

Net Ionic Equations

2I-(aq) + Br2(aq) 2Br-

(aq) + I2(g)

Handout 6 – History of the Atom

Handout 7 – The Electromagnetic Spectrum

Handout 8 – The Emission Spectrum of Hydrogen

Handout 9 – Theories of the Atom

Handout 10 – Atomic Orbitals

The first 3 S - orbitals

Shapes of the s, p, and d orbitals

The s and p orbitals around a single atom

Handout 11 – Aufbau Diagram

Handout 12 – Atomic Radii

Handout 13 – 1st Ionization Energy

Handout 14 – Successive Ionization Energies

Handout 15 – Electron Affinity

Handout 16 - Covalent Compounds

Remember: - an ionic compound has a metal + non-metal- a molecular compound only has non-metals

Hydrogen - neither a metal nor a non-metal - when we draw the Lewis/Electron dot diagrams for

compounds containing H, we consider the bond covalent. - when we name the compound containing H, we use ionic

nomenclature rules

- a compound only needs one ionic bond to be classified as ionic (even if there are many covalent bonds)

Calculating Polarity of Covalent Bonds

To this point we have described bonds as either ionic or covalent. Due to electronegativity we can now classify the covalent bonds as non-polar, slightly polar, polar and very polar.

Step #1 – We need to consider the dipoles of each bond in the molecule by examining the differences in electronegativities.

ElectronegativityDifference

Polarity

0 Non-polar, ie/ diatomic molecules: H2, O2, F2

0.5 or less Slightly polar0.6 – 1 Polar

1.1 – 1.6 Very polarGreater then 1.7 Usually indicates an ionic bond

Step #2 – Dipoles act as forces that pull electrons toward the more electronegative atom - Dipoles that are equal in magnitude but opposite in direction will cancel out

Step #3 –Then add labels to your electron dot/structural diagrams to show that the atom with a greater electronegativity pulls on the electrons better and thus has a partial negative charge, δ-, the atom that is least electronegativity has a partial positive charge, δ+.

NOTE: The general rule for writing a chemical formula is to write the most electronegative atom last. Therefore water’s chemical formula is H2O and not OH2.

Handout 17 – Ionic to Covalent to Metallic Bonding