Unit 5: Bonding Flashcards
Ionic Bond Materials
Metals with Non-metals
Covalent Bond Materials
Non-metals with Non-metals
Metallic Bond Materials
Metals
Explain Metallic Bonding
Solid metals exist as a lattice. The atoms let go of their electrons, creating Delocalised Electrons. Letting go of e- turns all the atoms into + ions.
Now we have + ions and e- in a lattice. Of course there will be a lot of electrostatic forces.
Properties of a Metallic Bond
Very Strong (High boiling point)
Very Conductive (because of delocalised electrons)
Pure Metals - Malleable and Ductile (because the lattice slides easily without alloy structure)
Explain Ionic Compounds
Metals like to donate electrons and non-metals like to accept electrons, because metals have a lower valency and non-metals have a higher valency. Thus they lean toward the least energy needed.
eg. Mg needs to lose 2 electrons and O needs to gain 2 electrons. Thus, it will donate.
This adds a +2 charge on Mg and -2 charge on O. They get attracted to the charges due to electrostatic forces.
Properties of Ionic Compounds
Very Strong (High melting point) - due to electroststic forces
Can Conduct
Brittle - due to structure
Why are Ionic Compounds Brittle?
Ionic Compounds exists as crystalline structures/ionic solids, where ions are located in specific positions.
A hammer blow distorts the arrangement, brining like-charged ions cost together due t the sliding nature of the structure. Causing the compound to shatter.
Covalent Bonds
When 2 unstable non metals share electrons to complete their outer shells, bonding them.
Properties of Covalent Bonds
Do not conduct electricity - no free electrons
Not soluble in water.
Low melting and boiling points.
Intermolecular Forces
The forces between a molecule (like h20) and another molecule. Often Occur in Covalent Bonds.
Intramolecular Forces
The forces between atoms (like Hydrogen and Oxygen)
Polar and Non-Polar molecules
Polar: When the electrons in an atom/molecule are unequal. Example, Hydrogen and Fluorine. Fluorine is very electronegative. So, in their covalent bond, it ends up pulling the electrons. However, not fully. So instead of a negative charge toward F, we use a partial-. At the same time making H partially +.
So a polarised covalent bond would look like: partial+H partial-F
This also establishes + and - poles as dipoles.
Non Polar: When electrons are equally distributed. Like, have the same electronegativity or not a large enough different. Producing so partial charges.
Electronegativity
More electronegativity = More attractive the element is to electrons
Electron Cloud
Just the general electrons
Polarised
A state where the atom/molecule has unequal electron distribution through to partial charges.
Name the intermolecular forces
Dipole-Dipole interactions, Vander Waals and Hydrogen Bonds
Permanent Dipole-Dipole
The the partial - from one molecule attracts the partial - from another through electrostatic forces.
Temporary Dipole
Because electrons are constantly moving, at some point they end up imbalanced with more electrons on 1 side. This makes the electron heavy side the partial- pole and vice versa. Thus making it a temporary Dipole, because once the electrons go back to normal, the poles will balance out.
Induced Dipole
When a Temporary Dipole uses its partial- or + charge to induce a temporary dipole. Like by attracting the electrons to 1 side with a partial +.
This makes that atom polarised as well.
Vander Walls Force
The Temporary attraction between the partial+ of a temporary dipole and the partial- of an induced/temporary dipole.
Strength of Dipole Dipole Interactions/Vander Walls Force
Very weak
Hydrogen Bonds
Hydrogen + An electronegative atom like F, O or N (usually)
(Example: H20)
- Strongest intermolecular force, but weaker than intramolecular forces
Name the 3 possible ion and dipole interactions
ion-ion (formal charge)
ion-dipole
dipole-dipole
Highest Boiling Point in Intermolecular Forces
Hydrogen Bonds
Why is graphite less conductive than the rest of the metals?
Because graphite’s structure relies on pushing it down into a sheet. It uses covalent bonds. Thus, it is only a semi-conductor, because it does not have any delocalised electrons.
Allotrope
An alternative structural arrangement of atoms in the same element. Similar to isomers.
Allotropes of Carbon: Diamond, Graphite
Ionic Compounds and Structure
Compounds formed through ionic bonding
Form regular lattice structures (use a Ball and Stick Diagram to show)
Properties of Ionic Compounds
Conduct electricity when melted or dissolved in water (in a regular lattice structure the atoms are fixed. But when its in these forms, it allows them to move because)
High melting and boiling point (because there are a LOT of strong ionic bonds in a regular lattice)
Formula for number of Moles
Mass of substance * Molar Mass
How do you find Molar Mass of a substance
Assuming there are two substances
A has a relative atomic mass of 2 or 2 g/mol
B has a relative atomic mass of 4 or 4 g/mol
When riding the molar mass of AB
We add 2+4 = 6 g/mol
What is Avagardo’s Constant
6.02 * 10^23
Formula for number of Atoms or Molecules in a Mole
(Number of Moles)(6.0221023)
Formula for Concentration (mol/m3)
Moles/Density (kgm^3)
Instantaneous Rate
The rate of a reaction at any given moment on the graph. Measured in Tangents.
What Factors affect the Rate of Reaction
Main, Frequency of Collisions
- Concentration
- Surface area
- Catalyst
- Temperate
Explain Collision Theory
Particles must collide to react, but its not necessary for every collision to lead to a reaction.
What is a Fruitful Collision and How does it happen?
A successful collision that leads to a reaction
Happens when the colliding particles have enough activation energy and proper orientation
Activation Energy
The min amount of energy needed for particles to form a product in a reaction.
Explain a why a Rate of Reaction graph rises and falls
Rises - the energy starts the increase in the beginning of the reaction, leading to more collisions and a faster reaction rate
Peak - then the energy peaks, this is the activation energy. Officially commencing the reaction.
Falls - from here it falls because as the reaction starts making the product, the concentration of the reactant starts to reduce, leading to less collisions.
The Basic Idea of Identifying Exothermic and Endothermic Reactions
- It takes energy to break bonds
- Energy is released with bonds are formed
Define Exothermic and Endothermic reactions
Exo release heat and light (eg. combustion reaction)
End take heat and light (eg. decomposition)
What are positive and negative enthalpies and how are they realised to Exothermic/Endothermic reactions
Negative enthalpy: when the energy required to break the bonds is less than the energy released. Meaning the reaction is released excess energy as light. - Exothermic Reactions
It is called a negative enthalpy because when the energy needed and released are subtracted, the answer is negative. Same for Positive enthalpy.
Positive enthalpy: when the energy required to break the bonds is more than the energy released or if a reaction uses more energy than it released. - Endothermic Reactions
How to calculate Change in enthalopy
energy used (activation) + energy released - be vary of signs
Different in Edo and Exo graph
An exo reaction’s activation energy is less that what it releases , thats what makes it a negative enthalpy. An Edo graph shows this, where the Reacantant’s energy is higher than the products. With the difference between the two showing more much energy the product has lost/released.
Opposite for Edo - R below P
What are Reversible Reactions
Reactions that go both ways. Always done in a closed system. One always has to be exothermic and the other endothermic.
What is equilibrium in a Reversible Reaction
A middle point, where both the reactions are taking place but they’re cancelling each other out.
Eg, we have A and it turns into B, but it is a reversible reaction. So at first, it will go forward. But once it researches B, it will go backwards, till the two reach equilibrium.
Combustion Reaction + Example
Fuel + Oxygen = Energy + __oxide
Eg. a car engine combines petrol or diesel with air, igniting it with a spark plug. This leads to an “explosion”. However, the same fuel can be used on a stove top, at a slower rate. This can lead to “burning”.
