Chemistry - Unit 9

Synthesis (reaction)

A synthesis reaction is a chemical reaction in which two or more simple substances combine to form a single, more complex compound.

Multiple reactants → ONE product

Decomposition (reaction)

A single compound breaks down into two or more simpler substances.

ONE reactant → multiple products

What are the two types of decomposition?

Thermal Decomposition and Electrolytic Decomposition

Thermal Decomposition

• The compound breaks down when heated.

• Heat is required to split the substance.

A triangle would be drawn on top of the arrow.

Electrolytic decomposition

• The compound breaks down using electricity.

• This usually happens with ionic compounds dissolved in water (electrolysis).

Single-replacement or single displacement

• One element replaces another

• Usually happens with metals replacing metals or halogens replacing halogens

• Only occurs if the replacing element is more reactive than the one it replaces

Spontaneous - will not occur instantly (More reactive element will kick out the less reactive one)

Double Displacement

It is a chemical reaction in which the positive ions (cations) of two compounds switch places to form two new compounds.

• Usually occurs in aqueous solutions

• Often forms a precipitate, gas, or water as a driving force

• No single element is “replaced”; it’s an ion swap

There are two types: Preceptitation (ppt) and Acid-Base Reaction

Precipitation (ppt)

One of the products is insoluble in water → forms a solid called a precipitate

Ex. AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Acid-base (neutralization) reactions

An acid reacts with a base → produces water and a salt

Activity Series

It lists the relative reactivities of elements within groups, such as metal and halogens. Only used for single-replacement reactions. (Depends of relativity)

The more easily a metal loses electrons → more reactive it is

• A single replacement reaction depends on whether one element can “kick out” another from a compound.

• The reactivity series ranks metals (or halogens) from most to least reactive.

• If the replacing element is above the element in the compound → reaction occurs spontaneously

• If it’s below → no reaction happens

When it comes to metals reacting with acids, if a metal is more active than H in the activity series, it will react with an acid and form a hydroge n gas.

Combustion

is a chemical reaction in which a substance reacts with oxygen (O₂) and releases energy as heat and light.

• Always involves oxygen

• Always produces energy (flame, heat, or light)

• Usually produces CO₂ and H₂O if it’s a hydrocarbon

Hydrocarbon

A hydrocarbon is a chemical compound made only of hydrogen (H) and carbon (C) atoms.

Mole Ratios (molecules)

The mole ratio is the ratio of moles of reactants and products in a balanced chemical equation.

• It tells you how many moles of one substance react with or produce how many moles of another.

• Comes directly from the coefficients in the balanced equation.

Law of Definite Proportions

Law of Definite Proportions says the elements always combine in the same fixed ratio by mass

Mole Tunnel

(1) Convert grams of reactants to moles of reactants

moles = grams/molar mass

(2) relate moles of reactants to moles of products using mole ratios

(3) Convert moles of products to grams of products.

gram= (molar mass) x moles

d

d

Key Concept: Element in relation to a compound

In a given compound, the ratio of atoms of each element is fixed. For example, CO2 the ratio of C:) is 1mol:2mol

Percent Composition using mass

Percent by mass of element = (mass of element/mass of compound) x 100

Percent Composition using chemical formula

% by mass of element = (mass of element in 1 mol of compound)/(molar mass of compound) x 100

Percent Composition

tells you the percentage by mass of each element in a compound.

Molecular Formula

Tells you the exact number of each type of atom in a molecular compound

Ex. Hydrogen Peroxide H2O2

Empirical Formula

Shows the simplest whole-number ratio of elements in a compound
Ex. Hydrogen Peroxide HO

How to find empirical formula?

(1) Determine the mass (g) of each component element

(2) Convert the mass (g) of each element to moles and write a preliminary formula

(3) Convert the moles mathematically to whole number subscripts

a. Divide each subscript by the smallest subscript

b. If necessary, multiply through by the smallest integer that turns all subscripts into whole numbers

1.Determine the mass (g) of each element

• Often, the problem gives percent composition (e.g., 40% C, 6.7% H, 53.3% O).

• Or you have the total mass of the compound, so you calculate the grams of each element from the percentages.

Use 100g as it's pretty useful here

2.Convert the mass (g) of each element to moles and write a preliminary formula

• Use the molar mass of each element to figure out how many moles you have.

mass x 1 mol /molar mass (average atomic mass)

Ex. Preliminary Formula: Mg 0.370 I 0.717

3. Convert the moles mathematically to whole-number subscripts

3. Divide each subscript by the smallest subscript (found in the preliminary formula)

   b. If necessary, multiply through by the smallest integer that turns all subscripts into whole numbers

Common Multipliers: Empirical Formulas

0.33 Multiply by 3 1: 1.33 3:4

0.67 Multiply by 3. 1: 0.67. 3:2

0.5 Multiply by 2 1: 0.5 2:1

Empirical to Molecular Formula

Molecular formula=(Empirical formula)×n

n= molar mass of molecule/molar mass of imperial formula

Writing formulas:

For organic or molecular compounds, the order of elements is usually standardized:

1. C first (carbon) → if present

2. H second (hydrogen) → if present

3. Then all other elements in alphabetical order

   • Example: N (nitrogen), O (oxygen), P (phosphorus), S (sulfur), halogens, etc.

This is often remembered as CHON (Carbon → Hydrogen → Oxygen → Nitrogen) for common organic/biochemical compounds.

NOT IONS

Two polyatomics combine….

MOLECULAR FORMULA (If they can be simplified, then new empirical formula)

Representative Particles (Atoms, molecules, formula units) to Moles

x 1 mol/avogadro’s #

Moles to Representative particles (Atoms, molecules, formula units)

x Avogadro’s #/ 1 mol

Volume of gas (STP) to Moles

x 1 mol/22.4 L

Moles to Volume of gas

x 22.4 L/1 mol

Mass to Moles

x 1 mol/molar mass

Moles to Mass

x molar mass/ 1 mol

“How many ___ are in ___ moles of ___?”

Step 1: What does the formula tell me?

Look at the subscript of the thing they’re asking about.

Ask yourself:

• How many of those particles are in ONE formula unit?

Step 2: Convert moles of compound → moles of particle

Multiply by the ratio from the formula.

0.75 mol Al₂O₃ (Part of the question)

× 3 mol O (how many of the particles in one formula unit) / 1 mol Al₂O₃

= 2.25 mol O

Step 3: Convert moles → particles

Now think:

“They want actual atoms/ions, not moles.”

So multiply by Avogadro’s number:

6.022 × 10²³ particles per mole

That’s ALWAYS the last step.

What does a subscript mean in a chemical formula?

The subscript tells how many of that atom are in ONE formula unit.

If give moles of a compound and asked for moles of one element inside it what do it do?

Multiply by the subscript

Ex. Al2O3 has 3 Oxygens, 0.5 × 3 = 1.5 mol O

If given molecules (or formula units) and asked for atoms, what do I do?

Multiply by subscript

Ex. CO2 has 2 Oxygens

4.0 × 10²2 × 2

What type of question is this? (MS1-2)

Asking about mass? Molar Mass (4 sig figs)

Converting moles into particles? (Avogadro)

Counting atoms inside a compound?(Multiply by subscript)

Mole (mol)

A mole is the SI unit for amount of substance and is defined as the number of atoms in exactly 12 grams of Carbon-12. This number is called the Avogadro constant and equals 6.02 × 10²3 particles.

Mass of a single element =atomic weight of an element expressed in amu. The number comes directly from the periodic table. And…

Coincidentally, the mass of 1 mole (in grams) is numerically the same.

How do you determine the molecular or molar mass of a compound or molecule? (On test when it says gram formula mass)

You add the atomic weights of all the atoms that make up that substance.

REMEMBER: 4 sig figs for atomic weight

Ex. Methane (CH4)

= 1 (atomic mass of carbon) + 4(atomic mass of hydrogen)

= 1 (12.01 g/mol) + 4(1.008 g/mol)

= 16.04 g/mol'

DIATOMIC MOLECULES: Remember to double it’s atomic mass

Standard Mole Conversion:

1 amu = 1gram/6.02 × 10²3

Diatomic Molecules

(1) Hydrogen (2) Nitrogen (3) Fluorien (4) Oxygen (5) Iodine (6) Chlorine (7) Bromine

Atomic ___ = 1 atom

Important Diatomic Calculation Rule:

Double the atomic mass only when the element is diatomic and alone (like O₂). In a compound, count each atom as written in the formula.

Representative Particle

Element (metal/non-metal) - atom

Molecular element (diatomic and polyatomic) - molecule

Molecular Compound (Covelant) - molecule

Ionic Compound - Formula Unit

Representative particles in 1.00 mol = 6.02 × 10²3

(s), (l), (g), (aq)

Stands for the state of matter; solid, gas, liquid, and aqueous (dissolved in solution)

+ (On reactant side)

Added to

➡ (In chemical equation)

yields too

+ (product side)

and

Splint test (Hydrogen)

• Use a lit splint

• Bring it near the gas
  👉 If you hear a “pop” sound, the gas is hydrogen (H₂)

Why? Hydrogen combusts quickly with oxygen, making that popping sound.

How to balance chemical equations?

• Keep polyatomic ions together

• Start with metals; then carbon and hydrogen

• Leave oxygen for last

• If you get 7/2 or a decimal, multiply everything by 2

• Balance elements that appear in only ONE compound on each side first

• Never change subscripts

• Reduce coeffecients (by the same factor on reactant and product sides)