⏱️ August 5, 2020
Pure substances are generally defined by the three common states of matter:
Solids are composed of densely packed atoms which have slight movements known as vibrations. Solids have a fixed volume and fixed shape.
Atoms in liquids flow to assume the shape of the container that they are occupying. Liquids have a fixed volume, but no fixed shape.
Atoms in gases flow freely and expand to fill their container. Gases have no fixed volume or shape.
Here’s where things can get a little tough, but yay more calculations!
You can look at individual particles in a compound, but you could also look at formula units. A formula unit is the smallest collection of atoms from which the formula of the compound can be established. Understanding formula units is useful as it can be used as another means of conversion.
How many formula units of KCl would be needed if 1.76mol KCl were used in a chemical reaction?
The law of definite proportions states that the ratio of mass of the constituent elements in any pure sample of that compound is always the same. Constituent elements are the elements that the compound is made up of. For example, no matter what sample of water you obtain, you will always have the same ratio of hydrogen to oxygen. Whether it be a 100 g sample, a 200 g sample, or a 5000 g sample, the ratio will always reduce to be the same. The empirical formula uses this concept.
The empirical formula of a compound is the simplest whole number combination of elements in a compound. In other words, it is a way to simplify the chemical formula of a compound.
Image courtesy of scienceabc
No matter how absurd the numbers in the molecular formula are, like C90H180O90, the ratio will always lead to the empirical formula of CH2O.
Empirical Formula - Practice Question
Sometimes, you could be asked to find the empirical formula from a percent composition. Let's try this:
(1) A carbohydrate, which contains C, H, and O, has a % composition of 33.3% C and 7.4% H. Find the empirical formula of this carbohydrate.
My teacher always had this poem for us that lined out the steps:
Percent to grams, Grams to Moles, Divide by Small, Times to whole
Step 1 and 4 are optional but in this problem, we are going to have to do both. Step 1 is only necessary if the composition is given in percents. If it is already given in grams, you don't have to do it.
Step 1) Drop the % sign. 33.3% C --> 33.3g C and 7.4% H --> 7.4 g H.
But what about oxygen? Let's use the 100% concept: all percents have to add up to 100. Therefore, we can simply subtract 100%-33.3%-7.4% to obtain the percent composition that is oxygen. This would get us 59.3% or 59.3 g O.
Step 2) Now we have 33.3g C, 7.4g H, and 59.3g O and we have to convert all of these to moles. This is done simply by dividing each of these gram amounts by the element's molar mass.
C: 33.3g / 12g = 2.775 moles C
H: 7.4g / 1g = 7.4 moles H
O: 59.3g / 16g = 3.70625 moles O
Step 3) Now, we have to divide each mole amount by the smallest mole amount that we have. 2.775 is the smallest amount so...
C: 2.775 / 2.775 = 1
H: 7.4 / 2.775 = 2.66
O: 3.70625 / 2.775 = 1.33
Step 4) Since the ratios for hydrogen and oxygen aren't full numbers, we have to "times the numbers to whole numbers." Whenever the numbers end in 0.5, 0.33/0.66, or 0.25/0.75, multiply all by 2, 3, and 4, respectively. Here, since the values end in .33 or .66, we have to multiply all of them by 3 to get whole numbers.
C: 1 x 3 = 3 C
H: 2.66 x 3 = 8 H
O: 1.33 x 3 = 4 O
Therefore, the empirical formula for this carbohydrate is C3H8O4. This question didn't ask for the molecular formula but an example would be C6H16O8. There are millions of possibilities for a molecular formula.
✍️ Free Response Questions
AP Chemistry Free Response Questions
⚛️ Unit 1: Atomic Structure and Properties
1.1Moles and Molar Mass
1.2Mass Spectroscopy of Elements
1.3Elemental Composition of Pure Substances
1.4Composition of Mixtures
1.5Atomic Structure and Electron Configurations
1.6Photoelectron Spectroscopy & Graph Interp.
🤓 Unit 2: Molecular and Ionic Compound Structures and Properties
2.0Unit 2 Overview: Molecular and Ionic Bonding
2.1Types of Chemical Bonds
2.2Intramolecular Force and Potential Energy
2.3Ionic Bonding and Ionic Solids
2.4Metallic Bonding and Alloys
2.5Lewis Dot Diagrams
2.6Resonance and Formal Charge
🌀 Unit 3: Intermolecular Forces and Properties
3.0Unit 3 Overview: Intermolecular Forces and Properties
3.2Properties of Solids
3.3Solids, Liquids, and Gases
3.4The Ideal Gas Law
3.5The Kinetic Molecular Theory of Gases
3.6Deviations from the Ideal Gas Law
3.7Mixtures and Solutions
3.8Representations of Solutions
3.9Separation of Solids/Mixtures
3.10Solubility and Solubility Rules
3.11Spectroscopy and the Electromagnetic Spectrum
3.12Quantum Mechanics and the Photoelectric Effect
🧪 Unit 4: Chemical Reactions
4.0Unit 4 Overview: Chemical Reactions
4.1Recognizing Chemical Reactions
4.2Net Ionic Equations
4.4Physical vs. Chemical Changes
4.5Stoichiometry & Calculations
4.6Titrations - Intro and Calculations
4.8Intro to Acid-Base Neutralization Reactions
👟 Unit 5: Kinetics
5.0Unit 5 Overview: Kinetics
5.1Defining Rate of Reaction
5.2Introduction to Rate Laws
5.3Rate and Concentration Change
5.4Writing Rate Laws
5.5Collision Model of Kinetics
5.6Reaction Energy and Graphs w/ Energy
5.7Reaction Mechanisms and Elementary Steps
5.8Writing Rate Laws Using Mechanisms
🔥 Unit 6: Thermodynamics
6.0 Unit 6 Overview: Thermochemistry and Reaction Thermodynamics
6.1Endothermic Processes vs. Exothermic Processes
6.2Energy Diagrams of Reactions
6.3Kinetic Energy, Heat Transfer, and Thermal Equilibrium
6.4Heat Capacity and Coffee-Cup Calorimetry
6.5Phase Changes and Energy
6.6Introduction to Enthalpy of Reaction
6.7Bond Enthalpy and Bond Dissociation Energy
6.8Enthalpies of Formation
⚖️ Unit 7: Equilibrium
🍊 Unit 8: Acids and Bases
8.0Unit 8 Overview: Acids and Bases
8.1Introduction to Acids and Bases
Unit 9: Applications of Thermodynamics
🤺 AP Chemistry Essentials
🧐 Multiple Choice Questions
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