Unit 3: Bonding and Energy



Molecular dipole: The sum of all bond dipoles in a molecule. A molecule only has a molecular dipole if they contain polar bonds and they are arranged asymmetrically, showing that they don’t balance each other out.  


  • Bond dipole: The unequal sharing of electrons between two atoms. It is shown with an arrow that points towards the more electronegative atom, which will carry the partial negative charge.


  • Electronegativity: The tendency of an atom to attract electrons towards itself when it is a part of a covalent bond 


  • Polyatomic ions: Charged particles made up of two or more elements that are bonded covalently but act as a single unit with an overall positive or negative charge (e.g. NH4+ or NO3-)


  • Asymmetric structures: Molecules which are not symmetrical (if you flip vertically, it wouldn’t be the same) 


  • Ionic bonds: Bond between metal and non-metal 


  • Covalent bonds: Bond between two non-metals  


  • Transition state: Highest point of energy on a reaction coordinate diagram. Breaking down bonds are always endothermic (requires energy), while forming bonds are exothermic (releases energy).  


  • Bond Enthalpy: Total amount of energy taken to break down and reform 


  • Chemical Potential Energy: energy stored within the chemical bonds of substances 


  • Thermal Energy: energy associated with the motion of particles 


  • Lewis Structures: Models to show how valence electrons are arranged in molecules and polyatomic ions which can help you identify single, double, or triple bonds.


  • System vs. Surroundings: The system is a specific chemical process you are experimenting with. The surroundings are everything else, including the container or the person doing the experiment. 


  • Exothermic: A system (chemical process) releases energy into the surroundings (often feeling hot to the touch), because more energy is released forming the bonds than it was to break them. 


  • Endothermic: A change in which the system absorbs energy from the surroundings (often feeling cold to the touch), because more energy is required to break the bonds than form them. 


  • Activation energy: The minimal amount of energy needed required to initiate a chemical reaction


  • Enthalpy Change(▵H): The difference in potential energy between the reactants and products 


  • Lattice Energy: A measure of the strength of an ionic bond. Higher lattice energy indicates stronger attractions between the cations and anions in a crystal lattice. 

  • Dissociation: process in which ionic bond breaks apart into individual ions when dissolved in water


  • Physical change: Something that alters the form or state of the substance but does not change its chemical properties (e.g. ice melting) 


  • Chemical change: A change in which a new substance is formed through breaking and forming new chemical bonds (e.g. wood burning) 

Enthalpy

Change in enthalpy (H) = a change in chemical potential energy between the reactants and products


    H = Chemical Potential Energy of Products - Chemical Potential Energy of Reactants




Topics I have trouble with:


Part 1: Understanding Bond Strength


Covalent Bond Strength


Electronegativity Difference: Generally, a greater electronegativity difference between two atoms in a covalent bond leads to a stronger bond (higher bond enthalpy). This is because the unequal sharing creates partial charges that increase the attraction between the atoms.

Atomic Radius: There is an inverse relationship between atomic radius and bond strength. Smaller atoms can get their nuclei closer together, resulting in a stronger force of attraction for the shared electrons.


Periodic Table Trends:

Down a group: Atomic radius increases, so bond strength decreases (e.g., H−F>H−Cl>H−Br>H−I).

Across a period: Atomic radius generally decreases, which typically leads to stronger bonds between smaller atoms.

Bond Type: Triple bonds are the strongest and shortest, followed by double bonds, with single bonds being the weakest and longest.


Ionic Bond Strength


Magnitude of Charge: This is the most important factor in determining ionic bond strength. The greater the charges on the ions, the stronger the bond (e.g., ScN with 3+/3− ions has a much higher lattice energy than LiF with 1+/1− ions).


Ionic Radius: Smaller ions form stronger ionic bonds. Because smaller ions can get closer to one another, the electric force of attraction between the centers of charge is greater.

Periodic Table Trends: Ionic radius increases down a group, so the strength of ionic bonds decreases as you move down (e.g., LiCl>NaCl>KCl).


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Part 2: Energy Calculations and Stoichiometry


Calculating Heat (q=mcΔT)


To calculate the heat energy (q) involved in a process, use the formula: q=mcΔT.

m: Mass of the substance (usually the water or solution in a calorimeter).

c: Specific heat capacity (for water, this is 4.18 J/g⋅

ΔT: The change in temperature (Final Temp−Initial Temp).



Stoichiometry and Enthalpy (ΔH)


Enthalpy changes are proportional to the amount of substance reacting.

Using Balanced Equations: The ΔH value given for an equation refers to the moles specified by the coefficients. For example, if ΔH=−405.8 kJ for an equation with 4 moles of reactant, then reacting only 2 moles would release half that energy (202.9 kJ).


Experimental Data (Molar Enthalpy): To find the ΔH in kJ/mol from experiment:

Calculate q using mcΔT.

Convert the mass of the limiting reactant used into moles.

Divide the energy (q) by the number of moles to get the molar enthalpy.


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Part 3: Reaction Coordinate Diagrams


When sketching or interpreting these diagrams, you must include and label these four key components:

Reactants and Products: Plotted as horizontal lines.

In Exothermic reactions, products are lower than reactants (energy is released).

In Endothermic reactions, products are higher than reactants (energy is absorbed).

Enthalpy Change (ΔH): The vertical distance between the reactant line and the product line.

Activation Energy (Ea): The energy "hill" the reactants must climb. It is the distance from the reactant line to the peak of the graph.

Transition State: The very peak of the curve, representing the highest energy point where bonds are breaking/forming.