Lecture notes: CHEM103 November 20, 2008

more stoichiometry examples:

Aluminum sulfate can be made by the following UNBALANCED reaction:

___ AlCl_3(aq) + ___ H₂SO_4(aq) à ___ Al₂(SO₄)_3(aq) + ___ HCl_(aq)

In an experiment, 25.0 g of AlCl₃ was mixed with 612 mL of 0.5 M H₂SO₄. The products actually contained 28.46 g of pure Al₂(SO₄)₃.

Which is limiting?

Which is in excess?

What amount (in g) of Al₂(SO₄)₃ should have formed (the theoretical yield)?

Considering the amount of aluminum sulfate actually formed, what is the percent yield?

Finally, what is the molar concentration of the excess reactant at the end of the reaction?

(You may assume the volume remained constant at 612 mL.)

When coal is burned to make electricity, the common contaminant mineral pyrite (FeS₂) reacts with O₂ in a single displacement reaction.

FeS₂ (s) + 2 O₂ (g) à ?

From the combustion of 1.00 kg of pyrite with a volume of 655 L of oxygen,

…how many moles of SO₂ is formed? (Make sure to check for limiting reactant!)

(Assume a combustion temperature of 1500 C and a pressure of 11 atm.)

A propane (C₃H₈)/oxygen (O₂) mixture (fuel is 15.1% by mass) is ignited.

Which of these two reactants is limiting?

If 65 grams of THIS mixture is burned, how much CO₂ (in mass) will be formed?

What mass of the excess reactant will remain?

What is the percentage (by mass) of the excess reactant in the new mixture (after the reaction)?

(Assume that the reaction occurs in a sealed container, and that NO MASS IS LOST.

That is: the total mass remains constant, even while most of the reactants are used

up when they are converted to products.)

Another raw iron ore (hematite) may be refined to make iron by the following process:

Fe₂O₃ (s) + 3 CO (g) à 2 Fe (s) + 3 CO₂ (g)

An impure sample of iron ore (containing more than just Fe and O!)

weighing 1805 kg is mixed with 274 L of carbon monoxide (12.2 atm at 657 K)

If 438 kg of pure iron is refined from the reaction, what is the % hematite (by mass) in the original ore sample?

CREATING MODELS OF COVALENT COMPOUNDS

As with everything so far, we generate models with predictive power. We will see 2 or 3 different models of varying complexity and sophistication used to describe covalently bonded molecules.

First model of covalent bonding: LEWIS DOT STRUCTURES or THE OCTET RULE

What can this do for us?

1. Generate molecular formulas (how do different elements combine to form compounds?)

2. First step toward understanding molecular structure of non-ionic (covalent) molecules.

3. BUT – this is a deeply flawed model (overly simple)

4. (Don’t worry - we’ll get on to more detailed, better models soon!)

SIMPLE RULES: (for diatomic molecules)

1) Place each atom, together with dots corresponding to its valence electrons: Ar, He

2) position atoms adjacent to each other so that the "octet rule" is satisfied for each atom (a pair of electrons between atoms forms a single bond; remaining "non-bonding" electrons are lone pairs): F₂, H₂, HCl

3) if necessary, rearrange electron pairs to create multiple bonds - satisfy the "octet rule": O₂, SO, N₂

NOTE: these simple rules have predictive power: Why are covalently bonded gases (H, O, Cl, etc.) diatomic instead of monatomic?

COMPLETE RULES: (for polyatomic molecules)

BUT First you need to know about "Pauling Electronegativities"

A periodic trend related to electron affinity (same trend w/o the exceptions: MORE E.A. => higher E.N.) and also to ionization energy (lower I.E. => lower E.N.)

http://acswebcontent.acs.org/periodic/tools/PT.html

What do electronegativities tell us?

· Tendency of an atom to draw electron density toward itself (but in the context of covalent compounds, not becoming an ion)

{translation: shared, but not shared equally!}

· This holds implications for arrangement of atoms within a covalent molecule

· Affects the polarity of covalent molecules - discussed later