Unit 5: Bonding

Question 1

Ionic Bond Materials

Answer 1

Metals with Non-metals

Question 2

Covalent Bond Materials

Answer 2

Non-metals with Non-metals

Question 3

Metallic Bond Materials

Answer 3

Metals

Question 4

Explain Metallic Bonding

Answer 4

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.

Question 5

Properties of a Metallic Bond

Answer 5

Very Strong (High boiling point)

Very Conductive (because of delocalised electrons)

Pure Metals - Malleable and Ductile (because the lattice slides easily without alloy structure)

Question 6

Explain Ionic Compounds

Answer 6

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.

Question 7

Properties of Ionic Compounds

Answer 7

Very Strong (High melting point) - due to electroststic forces

Can Conduct

Brittle - due to structure

Question 8

Why are Ionic Compounds Brittle?

Answer 8

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.

Question 9

Covalent Bonds

Answer 9

When 2 unstable non metals share electrons to complete their outer shells, bonding them.

Question 10

Properties of Covalent Bonds

Answer 10

Do not conduct electricity - no free electrons

Not soluble in water.

Low melting and boiling points.

Question 11

Intermolecular Forces

Answer 11

The forces between a molecule (like h20) and another molecule. Often Occur in Covalent Bonds.

Question 12

Intramolecular Forces

Answer 12

The forces between atoms (like Hydrogen and Oxygen)

Question 13

Polar and Non-Polar molecules

Answer 13

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.

Question 14

Electronegativity

Answer 14

More electronegativity = More attractive the element is to electrons

Question 15

Electron Cloud

Answer 15

Just the general electrons

Question 16

Polarised

Answer 16

A state where the atom/molecule has unequal electron distribution through to partial charges.

Question 17

Name the intermolecular forces

Answer 17

Dipole-Dipole interactions, Vander Waals and Hydrogen Bonds

Question 18

Permanent Dipole-Dipole

Answer 18

The the partial - from one molecule attracts the partial - from another through electrostatic forces.

Question 19

Temporary Dipole

Answer 19

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.

Question 20

Induced Dipole

Answer 20

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.

Question 21

Vander Walls Force

Answer 21

The Temporary attraction between the partial+ of a temporary dipole and the partial- of an induced/temporary dipole.

Question 22

Strength of Dipole Dipole Interactions/Vander Walls Force

Answer 22

Very weak

Question 23

Hydrogen Bonds

Answer 23

Hydrogen + An electronegative atom like F, O or N (usually)

(Example: H20)

• Strongest intermolecular force, but weaker than intramolecular forces

Question 24

Name the 3 possible ion and dipole interactions

Answer 24

ion-ion (formal charge)
ion-dipole
dipole-dipole

Question 25

Highest Boiling Point in Intermolecular Forces

Answer 25

Hydrogen Bonds

Question 26

Why is graphite less conductive than the rest of the metals?

Answer 26

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.

Question 27

Allotrope

Answer 27

An alternative structural arrangement of atoms in the same element. Similar to isomers.

Allotropes of Carbon: Diamond, Graphite

Question 28

Ionic Compounds and Structure

Answer 28

Compounds formed through ionic bonding

Form regular lattice structures (use a Ball and Stick Diagram to show)

Question 29

Properties of Ionic Compounds

Answer 29

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)

Question 30

Formula for number of Moles

Answer 30

Mass of substance * Molar Mass

Question 31

How do you find Molar Mass of a substance

Answer 31

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

Question 32

What is Avagardo’s Constant

Answer 32

6.02 * 10^23

Question 33

Formula for number of Atoms or Molecules in a Mole

Answer 33

(Number of Moles)(6.0221023)

Question 34

Formula for Concentration (mol/m3)

Answer 34

Moles/Density (kgm^3)

Question 35

Instantaneous Rate

Answer 35

The rate of a reaction at any given moment on the graph. Measured in Tangents.

Question 36

What Factors affect the Rate of Reaction

Answer 36

Main, Frequency of Collisions

1. Concentration
2. Surface area
3. Catalyst
4. Temperate

Question 37

Explain Collision Theory

Answer 37

Particles must collide to react, but its not necessary for every collision to lead to a reaction.

Question 38

What is a Fruitful Collision and How does it happen?

Answer 38

A successful collision that leads to a reaction

Happens when the colliding particles have enough activation energy and proper orientation

Question 39

Activation Energy

Answer 39

The min amount of energy needed for particles to form a product in a reaction.

Question 40

Explain a why a Rate of Reaction graph rises and falls

Answer 40

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.

Question 41

The Basic Idea of Identifying Exothermic and Endothermic Reactions

Answer 41

1. It takes energy to break bonds
2. Energy is released with bonds are formed

Question 42

Define Exothermic and Endothermic reactions

Answer 42

Exo release heat and light (eg. combustion reaction)

End take heat and light (eg. decomposition)

Question 43

What are positive and negative enthalpies and how are they realised to Exothermic/Endothermic reactions

Answer 43

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

Question 44

How to calculate Change in enthalopy

Answer 44

energy used (activation) + energy released - be vary of signs

Question 45

Different in Edo and Exo graph

Answer 45

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

Question 46

What are Reversible Reactions

Answer 46

Reactions that go both ways. Always done in a closed system. One always has to be exothermic and the other endothermic.

Question 47

What is equilibrium in a Reversible Reaction

Answer 47

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.

Question 48

Combustion Reaction + Example

Answer 48

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”.