Chemistry Regents Reviw

What is a mole in chemistry?

A mole is a unit that describes the quantity of 6.02 x 10^23, similar to how a dozen refers to 12.

What does the dot in a hydrate formula, such as CuSO4 • 5H2O, signify?

The dot indicates that for every copper sulfate molecule, there are 5 water molecules, not multiplication.

How do you calculate the percent composition of a component in a hydrate?

Percent composition = (mass of component / total mass of hydrate) x 100.

What is the percent composition of water in CuSO4 • 5H2O?

36% (calculated as (90g/250g) x 100).

When are the empirical and molecular formulas the same?

They are the same when the largest common multiple in the molecular formula is 1, meaning it cannot be further reduced.

Where should special conditions like catalysts be placed in a chemical equation?

They should be placed above the yield arrow in the chemical equation.

What is the first step in balancing a chemical equation?

Always write out the skeleton equation first.

How do you balance the equation NH3(g) + O2(g) → NO(g) + H2O(g)?

Start with the skeleton equation, then adjust coefficients to balance atoms on both sides.

What is the stoichiometric relationship in the reaction 2H2 + O2 → 2H2O?

For every 1 mole of O2, 2 moles of H2O are produced.

How do you convert grams to moles in a chemical reaction?

Divide the mass in grams by the molar mass of the substance.

What is ΔH in thermodynamics?

ΔH represents the heat of reaction, calculated as the difference between the potential energy of products and reactants.

What does a positive ΔH indicate about a reaction?

It indicates an endothermic reaction, where heat is absorbed.

What does a negative ΔH indicate about a reaction?

It indicates an exothermic reaction, where heat is released.

How is heat represented in a chemical equation for an exothermic reaction?

Heat is shown on the right side of the equation, e.g., A + B → C + heat.

How is heat represented in a chemical equation for an endothermic reaction?

Heat is shown on the left side of the equation, e.g., A + B + heat → C.

What is the activation energy in a chemical reaction?

The minimum amount of energy required for reactants to collide and produce products.

What is collision theory?

It states that reactants must collide with the correct orientation and sufficient energy to form products.

What is the difference between exothermic and endothermic potential energy diagrams?

Exothermic diagrams show products with less potential energy than reactants, while endothermic diagrams show products with more potential energy.

What is Le Chatelier's Principle?

It states that if a system at equilibrium is disturbed, the system will shift to counteract the disturbance and restore equilibrium.

What are the phases typically included in chemical equations?

Phases are indicated as (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solutions.

What is the purpose of balancing chemical equations?

To ensure that the number of atoms of each element is conserved in the reaction.

How do you calculate total heat released or absorbed in a reaction?

Set up a proportion based on the stoichiometry of the reaction and the heat values from the balanced equation.

What is the significance of the heat of reaction in a chemical equation?

It indicates the energy change associated with the reaction, affecting the reaction's spontaneity and direction.

What is the role of a catalyst in a chemical reaction?

A catalyst lowers the activation energy, increasing the rate of the reaction without being consumed.

What is the potential energy of the activation complex in a reaction?

It is the highest potential energy state during the transition from reactants to products.

Molarity (M)

The concentration of a solution expressed as moles of solute per liter of solution.

ppm

Parts per million; calculated as grams of solute divided by grams of solution, multiplied by 1,000,000.

Dilution

The process of reducing the concentration of a solute in a solution, calculated using the formula M1V1 = M2V2.

Homogeneous Mixtures

Mixtures that are uniform in composition, such as solutions.

Solvent

The substance that dissolves another substance in a solution.

Solute

The substance that is dissolved in a solvent.

Solubility

The quantity of a solute that may be added to a given quantity of solvent at a specific temperature and pressure.

Concentration

The amount of solute dissolved in a given amount of solvent.

Unsaturated Solutions

Solutions that can still dissolve more solute.

Saturated Solutions

Solutions that have dissolved the maximum amount of solute possible.

Supersaturated Solutions

Solutions that contain more dissolved solute than saturated solutions, often forming crystals when disturbed.

Precipitate

An insoluble solid that forms in a liquid when a substance is added beyond its solubility limit.

Solvation

The process by which solvent particles surround solute particles, separating them.

Dissolving Rate

The speed at which a solute dissolves in a solvent, influenced by agitation, surface area, pressure, and temperature.

Molarity Calculation

Calculated using the formula: moles of solute divided by liters of solution.

Example of Molarity

0.5M NaCl (aq) indicates a concentration of 0.5 moles of NaCl dissolved in 1 liter of solution.

Volume Conversion

To use the molarity formula, convert all volumes to liters (e.g., 250ml = 0.250 liters).

Parts per Million Calculation

Calculated using the formula: grams of solute divided by grams of solution, multiplied by 1,000,000.

Dilution Problem Example

To find the concentration of a 300ml solution of 0.25M HCl diluted to 500ml, use M1V1 = M2V2.

Types of Solutions

Includes solid in liquid, gas in liquid, liquid in liquid, gas in gas, and solid in solid solutions.

Example of Solid in Liquid Solution

Salt dissolving in water, where salt is the solute and water is the solvent.

Example of Gas in Liquid Solution

Carbon dioxide dissolved in soda, where carbon dioxide is the solute.

Example of Liquid in Liquid Solution

Ethanol dissolved in water.

Example of Gas in Gas Solution

Air, which includes oxygen dissolved in nitrogen.

Example of Solid in Solid Solution

Alloys such as steel, brass, and bronze.

M1

0.25M

V1

300ml

M2

unknown or X

V2

500ml

Equation for M2

0.25M x 300ml = M2 x 500ml

Solved M2

0.15M

Electrolytes

Substances that form ions in solution (conduct electricity)

Solubility of ionic compounds

Refer to Table F for solubility rules.

CaCl2

Soluble since Cl is a halide and Ca is not in the exception column in Table F.

SrSO4

Insoluble since sulfates are soluble but Sr is in the exception column in Table F.

BaCO3

Not soluble and would form a precipitate (insoluble solid) since CO3 is a carbonate and Ba is not in the exception column in Table F.

Ba(OH)2

Soluble because even if hydroxides are not soluble in Table F, Ba is in the exception column.

Precipitation Reaction Example

CaCl2 (aq) + Na2SO4 (aq) → 2NaCl(aq) + CaSO4(s)

Double Replacement Reaction

A chemical reaction where two compounds exchange ions.

Saturated Solution

On the line of the solubility curve at a given temperature.

Supersaturated Solution

Above the line of the solubility curve at a given temperature.

Unsaturated Solution

Under the line of the solubility curve at a given temperature.

Maximum NH4Cl in 100ml of water at 40 degrees Celsius

47 grams, representing a saturated solution.

Colligative Properties

Properties that depend on the number of solute particles in a solution.

Boiling Point Elevation

Solute particles in solution increase the boiling point.

Melting Point Depression

Solute particles in solution decrease the melting point.

Dissociation of NaCl

NaCl (s) → Na+ (aq) + Cl- (aq) produces 2 ions.

Dissociation of CaCl2

CaCl2 (s) → Ca2+ (aq) + 2Cl- (aq) produces 3 ions.

Dissociation of Na3PO4

Na3PO4 (s) → 3Na+ (aq) + PO4 3- (aq) produces 4 ions.

Ionic Compound Increasing Boiling Point

Na3PO4 increases the boiling point of 100 mL water the most.

Oxidation States

Oxidation States (or Oxidation Numbers) are used to keep track of electrons in redox reactions and identify which species is oxidized or reduced.

Rule 1 for Oxidation Numbers

Elements not combined with a different element have oxidation number of zero.

Example of Rule 1

Examples: Na, Fe, Cl2, O2

Rule 2 for Oxidation Numbers

Ions have an oxidation number equal to the charge of the ion.

Example of Ions Alone

Examples of Ions alone: Na1+, Fe+2, Cl1-

Example of Ions in Ionic Compounds

Na in NaCl has a +1 oxidation state and Cl in NaCl has a -1 oxidation state.

Rule 3 for Oxidation Numbers

All group 1 elements have a +1 oxidation number as seen on the periodic table (except H).

Example of Group 1 Oxidation Number

In KCl, K must be +1 and Cl is -1.

Example of Group 2 Oxidation Number

In MgCl2, Mg must be +2 and each Cl is -1.

Rule 4 for Oxidation Numbers

Oxygen has an oxidation number of -2, unless it is in a peroxide (e.g., H2O2), in which case it has an oxidation number of -1.

Rule 5 for Oxidation Numbers

Hydrogen has an oxidation number of +1, unless it is bonded to a metal.

Example of Hydrogen Oxidation Number

In HF, H=+1; In LiH, H=-1.

Rule 6 for Oxidation Numbers

The sum of oxidation numbers for a polyatomic ion must equal the ion's charge.

Example of Polyatomic Ion

In CO3^2-, O is -2 but there are 3, so we have a total of (-6). C is +4.

Example of NH4+ Oxidation Numbers

In NH4+, H is +1 and there are 4, so we have (+4). N is -3.

Law of Conservation of Charge

The charge before a chemical or physical change equals the total charge after the change.

Oxidation Half Reaction Example

Zn0 → Zn+2 + 2e- (show electrons on right side).

Reduction Half Reaction Example

Cu+2 + 2e- → Cu0 (show electrons on left side).

Conservation of Charge

Charge on both sides of the equation have to equal each other.

Oxidation State Increase

Oxidation State or Oxidation Number increases when electrons are lost (Oxidation).

Reduction State Decrease

Oxidation State or Oxidation Number decreases when electrons are gained (Reduction).

Diatomic Elements in Reactions

When diatomic elements are in a chemical equation, keep the diatomic formula in the half reaction.

Diatomic Elements

Common diatomic elements include O₂, Cl₂, F₂, N₂, Br₂, H₂, I₂.

Oxidation Half Reaction

Fe⁰ → Fe²⁺ + 2e⁻.

Reduction Half Reaction

Cl₂⁰ + 2e⁻ → 2Cl¹⁻.