Unit 6 - Stoichiometry & Chemical Formulas

Question 1

What are the three types of Chemical Formulas?

Answer 1

1. Empirical Formula
2. Molecular Formula
3. Structural Formula

Question 2

Definition of Empirical Formula (3)

Definition & Useful

Answer 2

1. Simplest Ratio of Atoms
2. All Ionic Compounds Formulas are Empirical
3. Covalent Formula can be Empirical too

Question 3

Why are all ionic compound formulas empirical?

Answer 3

They do not form a discrete molecule, but rather a crystal lattice, so there’s no point in have unsimplified ratios

Question 4

Definition of Structural Formula (2)

Definition & Useful

Answer 4

1. Provides info on HOW elements are bonded
2. Only useful for covalent compounds

Question 4

Definition of Molecular Formula (3)

Answer 4

1. Actual number of atoms per element
2. May be multiple of the empirical formula
3. May be the same as the empirical formula

Question 5

Definition of Molecular Weight/Formula Mass

Answer 5

Sum of masses of atoms in the molecule

Question 6

What is “u” short for?

Answer 6

Atomic Mass Units

Question 7

How do you find the molecular formula when given percent compositions? (5)

Answer 7

1. Find Empirical Formula
2. Find Molecular Weight of E. Form.
3. Molecular Weight/Given Molar Mass
4. Round Quotient (nearest ones place)
5. Multiply all subscripts by quotient

Question 7

How do you determine the Empirical Formula when given percent compositions? (5)

Answer 7

1. Change all % to grams (Ex: 22.1% to 22.1 g)
2. Identify atomic mass of elements in PT (Ex: Al - 27)
3. Divide grams by the molar mass (Ex: 22.1/27 = 0.82)
4. Divide the quotients of each element by the smallest quotient (Ex: 3.3/0.82 & 0.82/0.82)
5. A Ratio is Formed (if a decimal, take smallest multiple) (Ex: 1:1.5 = 2:3)

Ex:

Element - Percent to Gram - Molar Mass - Quotient - Ratio
Al - 22.1 - 27 - 0.82 - 1
P - 25.4 - 31 - 0.82 - 1
O - 52.5 - 16 - 3.3 - 4
AlPO4

Question 8

What are the four types of Chemical Reactions?

Answer 8

1. Synthesis
2. Decomposition
3. Single Replacement
4. Double Replacement

Question 9

Definition of Synthesis Rxn

Answer 9

Elements/Compounds Combine for a more Complex Compound
Ex: A + X = AX

Question 10

Definition of Decomposition Rxn

Answer 10

Complex Compounds Breaks Down to Simpler Compounds
Ex: AX = A + X

Question 11

Definition of Single Replacement Rxn

Answer 11

An element replaces another element in a compound
Ex: A + XY = X + AY

Question 11

Definition of Double Replacement Rxn

Answer 11

One of the elements in each of the compounds are interchangeable
Ex: AB + XY = AY + XB

Question 12

What are the 4 Cardinal Rules when Balancing Equations? (3)

Answer 12

1. Count each type of atom on both sides of the equation
2. If a polyatomic ion remains intact on both sides, leave it alone
3. NEVER CHANGE THE SUBSCRIPTS
4. Automatically Convert H2O as HOH

Question 13

What is the difference between the moleular weight, formula mass, molar mass and gram formula mass?

Answer 13

Molecular Weight - Covalently Bonded Molecules
Formula Mass - Ionic Compounds
Molar Mass - Mass of one mole (g/mol)
Gram Formula Mass - Same as Molar Mass, but specific to ionic compounds

ALL HAS THE SAME VALUE
Molecular Weight of H20 - 18 g
Formula Mass of H20 - 18 g
Molar Mass of H20 - 18 g/mol
Gram Formula Mass of H20 - 18 g/mol

Question 14

Definition of Mole (3)

Answer 14

1. A way to count a large quantity of very small particles
2. 1 mole = 6.022 × 10²³ (Avogadro’s number)
3. A “dozen” counts 12 items, a “mole” counts 6.022 × 10²³ particles

Question 15

How do you write the Molar Mass?

Answer 15

Formula Mass & add units g/mol

Question 16

What do coefficients in a balanced equation actually indicate?

Answer 16

Mole - Mole Ratios between Reactants & Products

Ex: 2H₂ + O₂ – 2H₂O
- 2 moles of H reacts with 1 mole of O to produce 2 moles of water
- For every mole of O, 2 moles of water is created
- 5 moles of O will react with 10 moles of H to produce 10 moles of water

Question 17

Definition of Percent Composition

Answer 17

Def: Expresses relative amount of each element within a compound (%)
Formula: PercentComposition = (MassofElementinCompound/TotalMassofCompound) x 100