Atomic Structure
Atom
Definitions
Mass Number (A): The relative mass of an atom of an element
Atomic number (P): The amount of protons per atom of that element
Subatomic Particles
Valence is the amount of electrons in the outer shell of the atom
Atomic Models
Isotopes
Definition
An atom that has more or less neutrons in its nucleus than normal, and therefore has a change in atomic mass but not atomic number.
Examples and Uses
Heavy Water: Water made up of oxygen and isotopes of hydrogen \((\mathrm{H}-1, \mathrm{H}-2\), and \(\mathrm{H}-3)\) is used to slow down neutrons in order to increase the likelihood of a nuclear reaction.
Uranium 235: Used as an energy source in a nuclear power plant.
Relative Atomic Mass based on Abundance
Average Relative Atomic Mass \(=(\) (Mass of Isotope 1 * Percentage Abundance \()+(\) Mass of Isotope 1 * Percentage Abundance)) / 100
Balancing Equations
Law of Conservation of Mass
In any chemical reaction, mass cannot be created or destroyed.
Rules of Balancing Equations
There should be the same proportion of each element on each side of the equation.
Ionic Bonding
Ions
- Ions are atoms that are positively or negatively charged. When electron transfer happens, atoms have more or less electrons than protons, making them ions.
- AnIons: Negatively charged Ions
- CatIons: Positively charged Ions
The Process
All atoms want to have a full outer shell. Ionic bonding occurs when atoms exchange electrons with each other to fulfill this. Because one atom loses an electron, making it positively charged, and vice versa for the other atom, they are attracted to each other, and therefore they bond. This happens between metals and non-metals.
Diagram
Covalent Bonding
The Process
Covalent bonding is the sharing of electrons for atoms to fill each other’s outer shells. The positive nucleuses are attracted to the shared electrons, thus they become a bond.
Single, Double and Triple Bonds
Single bonds occur when there is a single pair of electrons shared (2 electrons)
Double bonds occur when there is a double pair of electrons shared (4 electrons)
Triple bonds occur when there is a triple pair of electrons shared (6 electrons)
Diagram
Simple and Giant Covalent Structures
Simple covalent structures are made up of individual molecules. Giant covalent structures consist of rigid 3D lattices where atoms are held in place
Metallic Bonding
The Process
Atoms share delocalized electrons which float around in a ‘sea of electrons.’ Since the atoms have lost electrons, they become Cations. The positively charged atoms are attracted to the negatively charged delocalized electrons. The atoms form a grid.
Diagram
Properties of Metals
Conductive, as the delocalized electrons are free to move and have a charge Malleable, as the metals form layers, which are easy to bend Ductile, as the metal forms layers, which can be stripped off
Skills
Properties of Substances
The properties of a substance can be linked to what kind of compound it is, for example, since oxygen is a covalent bond, it cannot conduct electricity, as it has no free-to-move charged particles.
Types of Molecular Forces
Intermolecular Forces that take place between multiple molecules
Hydrogen Bonding – Is an electrostatic attraction created between covalently bonded hydrogen atom to an electronegative atom (Oxygen, Fluorine, and Nitrogen). This creates strong dipoles that can then interact.
Dipole-Dipolele Action – Different atoms have different electronegativity values hence dipoles are created as the shared electron are more attracted to one side.
London Dispersion Forces- These are temporary dipoles created in a molecule through the movement of electrons. Often large molecules have very strong diples created by LDF’s; this is cause they have many electrons.
Intramolecular Forces: Forces that take place within a molecule
Mole
- Much like the word dozen, a mole represent a certain amount of a substance. Mole is basically \(6{ }^* 10^{\wedge} 26\) of anything. This is often used to convert between amu (atomic mass units) and grams. Additionally, it is a convention and used in experiments.
- A formal definition is that mass of substance containing the same number of fundamental units as there units as there are atoms in exactly \(12.00 \mathrm{~g}\) of carbon – 12
- A mole ratio is the ratio between the amounts in moles of any two compounds involved in a chemical reaction. Mole ratios are used as conversion factors between products and reactants in many problems.
- A few definitions
- Mass Number: The total number of protons and neutrons in a nucleus
- Relative Atomic Mass: The ratio of the average mass of one atom of an element to one twelfth of the mass of an atom of carbon – 12
- Relative molecular mass: the ratio of the average mass of one molecule of an element or compound to one twelfth of the mass of an atom of carbon-12.
- Solute: The minor component in a solution, dissolved in the solvent. Or in other words; it is the smaller part of the solution.
- Solvent: Is the part of the solution in which the solute is dissolved. Or is the part of the solution present in greater amounts
- Solution: A mixture of solvent and solute.
Common Equations
Percentage Composition/ Limiting Reactant
- Percentage Composition
- Percentage Composition is a technique in chemistry in which a certain elements mass is calculated from the complete element.
- The formula for percentage composition is Mass/ Total mass * 100
- Limiting Reactant
- Limiting reactants are important to calculate as they are used when forming mole ratios.
- There are a few steps when calculating limiting reactants
- Balance the equation for the chemical reaction.
- Convert the given information into moles.
- Use stoichiometry for each individual reactant to find the mass of product produced.
- The reactant that produces a lesser amount of product is the limiting reagent.
- The reactant that produces a larger amount of product is the excess reagent.
- To find the amount of remaining excess reactant, subtract the mass of excess reagent consumed from the total mass of excess reagent given.
- Empirical Formula is the smallest Ratio while molecular formula is when the equation is not simplified.
States of Matter & Kinetic Theory
● There are many different states of matter each have different properties
● Kinetic Molecular Theory states that gas particles are in constant motion and exhibit perfectly elastic collisions. This can be used to explain Charles’ and Boyle’s Law. The average energy of a collection of gas particles is directly proportional to absolute temperature.
Collision Theory
● Collision theory is normally used to predict rates of chemical reaction, particularly for gases. The theory is based on the assumption that for a reaction to occur it is necessary for the reaction species to come together.
● There are three main points listed in collision theory
○ Molecules must collide to react
○ Collision must have the correct orientation
○ Collision must have enough energy
Equilibrium
● Definitions
○ Thermal Dissociation: The breaking apart of a molecule’s bond due to the introduction of heat. Or it is the breaking down of a large substance into smaller substance.
○ Reversible Reaction: It a chemical reaction where the reactants from product in turn can be reversed and give back reactants.
○ Thermal Decomposition: Is a simple single step reaction where a molecule splits into two products. It normally takes place due to ionization of a substance of heat.
● Chemical equilibrium is a state in which the rate of the forward reaction equals the rate of the backward reaction. In other words there is no net change in concentration. Otherwise this is known as dynamic equilibrium.
● A physical equilibrium is a system whose physical state does not change when dynamic equilibrium is reached in a system
● A Catalyst is used to find an alternative pathway to reaction with a lower activation energy.
● Le Chatelier Principle is used to predict the behavior of a system due to changes in temperature, concentration and pressure.
○ If the temperature in a system changes the behavior will change. If the system is exothermic then an addition of heat will favor the front direction. In endothermic reverse is applicable.
○ If pressure is increased then it depends where the most gas molecules are present.
○ If the concentration of the products or reactants are increased respectively you will get a change in the rate of reaction for that side.
● The Haber Process
Rate of Reaction
- Rate of Reaction – Is the speed at which reactants are converted into products
- There are multiple different factors that impact the rate of a reaction however the most common include temprature, pressure/ Concentration, and catalyst.
- Temperature affects the rate of reactions as it increases the speed at which particles collide; otherwise known as the kinetic energy. An increase in KE means a higher percentage of particles have the minimum activation energy. Another way in which temp can impact a reaction is that it increases the random motion of particles. This means collision can happen more often; hence being successful more often
- An increase in concentration and pressure means that there are more particles in a given volume. More particles in the same volume mean that there is a higher chance of a collision to take place. It is more likely for a reaction to take place.
- Catalyst even impacts the rate of reaction by using an alternative pathway that has lower activation energy. A lower activation energy means more particles have the ability to pass the activation energy barrier. A catalyst increases the percentage of particles with suitable activation energy by introducing a new pathway with less activation energy.
- Surface Area is an important factor. A higher surface area means there is more area for the reaction to take place on.
- Common Experimental Procedures
Temperature –
Measure out \(50 \mathrm{~cm} 3\) of sodium thiosulfate and pour into the conical flask. Draw a cross on the piece of paper and then place the conical flask on top of it.
Use the thermometer to measure the temperature of the sodium thiosulfate. Record this value in the Results table below.
Measure out \(5 \mathrm{~cm} 3\) of hydrochloric acid and pour into the conical flask. Start the stop clock straightaway.
Looking from above, time how long it takes for the cross to ‘disappear’. Record this time in seconds in the Results table
Pour the solution away as quickly as possible and rinse out and the flask.
Repeat steps 1 to 6 using sodium thiosulfate from one of the water (or ice) baths.
Continue until you have done the experiment with sodium thiosulfate from all of the different water baths.
Concentration
Method
A Set up a gas syringe as shown in the diagram so that it is at the right height for the bung to go into the boiling tube in the beaker.
B Pour \(20 \mathrm{~cm}^3\) of \(2.0 \mathrm{~mol} / \mathrm{dm}^3\) hydrochloric acid into a boiling tube.
C Put the boiling tube in the water bath and have the delivery tube and bung ready to place in the boiling tube.
D Drop a \(5 \mathrm{~cm}\) strip of magnesium ribbon into the acid and immediately put the bung in the boiling tube. Your partner should start the stopclock at the same time.
E Measure the volume of hydrogen gas formed in 20 seconds and record the result in the table. During this time, shake the reaction mixture.
F Repeat the experiment using different concentrations of acid. These can be made by mixing \(2.0 \mathrm{~mol} / \mathrm{dm}^3\) hydrochloric acid with delivery tube water, using the quantities shown in the table.
Catalyst
• The minimum quantity of energy that the reacting species must possess in order to undergo a specified reaction.
• Catalyst does impact the rate of a reaction by finding an alternative pathway of energy for the reactants
• Catalyst does impact the rate of a reaction by finding an alternative pathway of energy for the reactants
Combustion
Definitions
Flash Point: The lowest temperature at which the vapors of that material will ignite Ignition Temperature: The lowest temperature at which a combustible substance when heated in air takes fire and continues to burn
Complete and Incomplete Combustion
Combustion, otherwise known as burning, involves the reaction of a hydrocarbon and oxygen to produce carbon dioxide and water.
If there is sufficient oxygen, carbon dioxide is produced. This is known as complete combustion.
If there is not enough oxygen, carbon monoxide is produced. This is known as incomplete combustion
Chemical Equations
Complete: \(\mathrm{C}_{\mathrm{x}} \mathrm{H}_{\mathrm{y}}+\mathrm{O}_2 \rightarrow \mathrm{CO}_2+\mathrm{H}_2 \mathrm{O}\)
Incomplete: \(\mathrm{C}_{\mathrm{x}} \mathrm{H}_{\mathrm{y}}+\mathrm{O}_2 \rightarrow \mathrm{CO}+\mathrm{H}_2 \mathrm{O}\)
Enthalpy
Definitions
Standard Average Bond Enthalpy: The amount of energy required to break a specific type of bond per mole of the substance.
Standard Enthalpy Change of a Reaction: The enthalpy change that will occur in the system when matter is transformed by a chemical reaction.
Standard Enthalpy of Formation: Enthalpy during the formation of 1 mole of the substance from its constituent elements
Hess’s Law
Regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes.
Source of Energy
Energy can come from an external heat source, or by formation of chemical bonds, which releases energy.
Calculating Enthalpy Change
\(\Delta H=\) Energy given in formation of Products – Energy used in Breaking Apart Reactant Thermochemical Equations
(Balanced equation) \(\Delta H=x y z . k J\)
Using Experimental Data
$
\Delta H=m c \Delta T
$
\(\Delta H\) : Enthalpy change
M: Mass
C: Specific Heat Capacity (Heat it takes to increase the temperature of 1 gram of the substance by \(1^{\circ} \mathrm{C}\) )
\(\Delta T\) : Change in temperature
Exothermic and Endothermic
Definitions and Examples
An exothermic reaction is a reaction that releases heat energy as the reaction happens An endothermic reaction is a reaction that absorbs heat energy as the reaction happens Examples:
Identifying Type of Reaction
If \(\triangle H H \geq 0\) then the reaction is endothermic and vice versa
In terms of bond enthalpies, when bonds are broken, energy is required, but when they are made they release energy. If the new bonds release more energy than the previous bonds broken, then it is exothermic; vice versa.
If \(\Delta H<0\) then the reaction is exothermic and vice versa
Use in Industry
- Exothermic reactions are used to heat up steam in order to move turbines and generate electricity
- Endothermic reactions are used in cold packs to treat bruises.
Heat
Calorimetry
Calorimetry is the process of measuring the amount of heat released or absorbed during a chemical reaction. By knowing the change in heat, it can be determined whether or not a reaction is exothermic or endothermic.
Assumptions of Calorimetry
The substance is pure
No heat is absorbed by the calorimeter
A concentration of \(1 \mathrm{~mol} / \mathrm{dm}^{\wedge} 3\) is used
Calorimeter Experiments
Entropy
Definition
The measure of a system’s thermal energy per unit temperature that is unavailable for doing useful work. Because work is obtained from ordered molecular motion, the amount of entropy is also a measure of the molecular disorder, or randomness, of a system.