Unit 1 - Chemical Changes And Structures

Question 1

(Need data booklet)
What elements have:
1. Metallic lattice bonding and structure
2. Covalent molecular bonding and structure
3. Covalent network bonding and structure
4. Monatomic

Answer 1

1. Lithium, beryllium, sodium, magnesium, potassium, calcium, aluminium (all the metals (duh))
2. Boron, carbon, silicon
3. Hydrogen, carbon, nitrogen, phosphorus, oxygen, sulfur, fluorine, chlorine
4. Helium, neon, argon

Question 2

How do metals bond?

Answer 2

They have 1, 2 or 3 electrons in outer shell. The atoms readily lose these, forming “positive ions”. The “positive ions” are arranged in a giant 3D pattern called a lattice. The outer electrons can move freely within the lattice - they are delocalised.

Question 3

What is metallic bonding?

Answer 3

An electrostatic attraction between positive ions and negative delocalised electrons.
The presence of delocalised electrons explains some of the physical properties of metals

Question 4

(Nat 5)
Why are metals electrical conductors

Answer 4

The delocalised electrons are free to move, therefore can carry electrical current.

Question 5

(Nat 5)
Why do metals tend to be strong and have fairly high melting points and boiling points?

Answer 5

Metallic bonding is strong due to the electrostatic force of attraction between the positive ions and
negative electrons.

Question 6

(Nat 5)
Why are metals are malleable (hammered into shapes) and ductile (can be drawn into wire).

Answer 6

Metallic bonding is non-directional so the bonds can be re-arranged different ways. Delocalised electrons are attracted by any of the positive ions around them, allowing metals to be stretched into wires and hammered into shape.

Question 7

(Nat 5)
Explain why group 2 metals have higher melting points than group 1 metals.

Answer 7

Group 2 metals have two outer electrons and Group 1 metals have one outer electron. Group 2 metals
have stronger bonding because they have more delocalised electrons. More energy is needed to break
the bonds.

Question 8

Why do covalent network elements have very high melting and boiling points?

Answer 8

Covalent bonds are very strong. Melting a covalent network means breaking all the covalent bonds. A huge amount of energy is needed.

Question 9

What two covalent structures does carbon have?

Answer 9

Graphite and diamond

Question 10

Explain bonding in diamond

Answer 10

In diamond, each carbon is covalently bonded to four other carbon atoms in a tetrahedral network. All of the four outer electrons are used to make strong covalent bonds. There are no individual molecules, no weak forces and no delocalised electrons.

Question 11

Explain bonding in graphite

Answer 11

In graphite, each carbon is bonded to three other carbon atoms in a hexagonal layer. Each carbon atom uses three of its outer electrons to make strong covalent bonds in the carbon layer. Layers of carbon atoms are stacked and held together by London dispersion forces. The fourth electron is delocalised and can move freely through the structure.

Question 12

What forces hold monatomic atoms together?

Answer 12

London dispersion forces

Question 13

What is an instantaneous temporary dipole?

Answer 13

When one side of an atom is slightly more negative (δ-) and the other side slightly more positive (δ+) due to the random movements of electrons (at any instant there may be more electron on one side than the other).

Question 14

What is an induced dipole?

Answer 14

When the negative side of an atom (due to instantaneous temporary dipoles) repels electrons in a neighbouring atom.

Question 15

What are London dispersion forces (LDFs)?

Answer 15

Electrostatic attraction between instantaneous and induces temporary dipoles, caused by the random movement of electrons inside atoms.

Question 16

What causes different melting and boiling points in terms of LDFs?

Answer 16

The more electrons an atom or molecule has, the stronger the LDFs, therefore more energy is needed to break them and so the higher the melting and boiling points.

Question 17

What is a moleule?

Answer 17

A discrete (unbonded) group of atoms joined by strong covalent bond.

Question 18

What type of bonding occurs inside molecules?

Answer 18

Intramolecular.

Question 19

What bonding occurs between molecules?

Answer 19

Intermolecular (LDFs)

Question 20

What is sublimation?

Answer 20

When a substance transitions from a solid straight to the gas state, without passing through the liquid state.

Question 21

What is the molecular formula for hydrogen, nitrogen, oxygen, fluorine, chlorine, phosphorus, sulphur and fullerene ?

Answer 21

H₂ , N₂ , O₂ , F₂ , Cl₂ , P₄ , S₈ , C₆₀

Question 22

What is covalent radius?

Answer 22

The measure of the size of an atom.

Question 23

What happens to the covalent radii across a period and why?

Answer 23

It decreases. This is because across the period the number of protons increases. This results in an increased nuclear charge and so the electrons are held more tightly.

Question 24

What do outer electrons experience in atoms?

Answer 24

A shielding effect as the inner electron shells block the outer electrons from the full attractive force of the nucleus.

Question 25

What happens to the covalent radii down a group and why?

Answer 25

It increases. This is because as you go down the group the number of electron shells increases. Therefore there I increases shielding and the attraction of the nucleus for the outer electrons decreases.

Question 26

What is the ionisation energy?

Answer 26

A measure of how tightly an atom holds on to its outer electrons.

Question 27

What is the first ionisation energy?

Answer 27

The energy needed to remove one mole of electrons from one mole of gaseous atoms, measured in kJmol⁻¹.

Question 28

What is the second ionisation energy?

Answer 28

The energy needed to remove one mole of electrons from one mole of gaseous 1+ (already lost one electron) ions.

Question 29

What happens to the ionisation energy across a period and why?

Answer 29

It increases. As you go across the period the number of protons increases. This results in an increased nuclear charge; the electrons are held more tightly and so more energy is required to remove them.

Question 30

What happens to the ionisation energy down groups and why?

Answer 30

It decreases. As you go down the group, there is increased shielding and the attraction of the nucleus for the outer electrons decreases.

Question 31

Is the second ionisation energy higher or lower than the first ionisation energy for elements in group 1 and why?

Answer 31

It is much higher than the first ionisation energy. The second ionisation energy involves removing an electron from a full shell which is closer to the nucleus. The second electron is less shielded, therefore more strongly attracted to the nucleus.

Question 32

How do you calculate total ionisation energy?

Answer 32

Add both the quantitites. together.

Question 33

What is electronegativity?

Answer 33

Electronegativity is a measure of the attraction that a bonded atom has for the electrons in the bond.

Question 34

What happens to electronegativity across a period and why?

Answer 34

It increases. As you go across the period the number of protons increases. This results in an increased nuclear charge and so the electrons are held more tightly.

Question 35

What happens to electronegativity down the groups and why?

Answer 35

It decreases. As you go down the group the number of electron shells increases. Therefore, as you go down the group, there is increased shielding and the attraction of the nucleus for the outer electrons decreases.

Question 36

What is a covalent molecule?

Answer 36

A small group of non-metal atoms held together by strong covalent bonds.

Question 37

How are covalent bonds formed?

Answer 37

By the electrostatic force of attraction between two positive nuclei and a shared pair of electrons.

Question 38

Why do covalent molecules have low melting and boiling points?

Answer 38

The forces of attraction between molecules (LDFs) are very weak and very little energy is needed to separate the molecules.

Question 39

What is a covalent network?

Answer 39

A giant 3D structure in which atoms are joined by strong covalent bonds.

Question 40

Why do covalent networks have have high melting and boiling points?

Answer 40

The atoms are tightly held together by strong covalent bonds and a lot of energy is needed to break these.

Question 41

What is non-polar covalent bonding?

Answer 41

Bonding in which the electrons are equally shared (both atoms have same electronegativity from a bond).

Question 42

What type of bonds do covalent elements always have and why?

Answer 42

Non-polar, because two atoms of the same element must have the same electronegativity. They must share their electrons equally.

Question 43

How do you know if a bonding is non-polar (typically)?

Answer 43

The electronegativity between the atoms is almost zero.

Question 44

What is polar covalent bonding?

Answer 44

Bonding in which there is a permanent dipole.

Question 45

How are permanent dipoles formed?

Answer 45

Most covalent compounds are formed between elements with different electronegativity. This means that one atom has a greater attraction for bonding electrons than the other. The atom with the higher electronegativity value has a greater attraction for bonding electrons and has a slight negative charge (δ-). This leaves the other atom with a slight positive charge (δ +).

Question 46

What makes a bond more polar?

Answer 46

A bigger difference in electronegativity.

Question 47

How do you know if bonding is polar covalent (typically)?

Answer 47

If the difference in electronegativity between the atoms is small (less than 2).

Question 48

What is an ionic bond?

Answer 48

The electrostatic attraction between a positive ion and a negative ion.

Question 49

Explain what happens to the metal and non-metal atom to from an ionic bond.

Answer 49

The metal atom loses an electron, forming a positive ion, and the non-metal atom gains an electron, forming a negative ion.

Question 50

How do you now if bonding is ionic (typically)?

Answer 50

The difference in electronegativity between the atoms is large (more than 2).

Question 51

Explain ionic bonding.

Answer 51

The metal and non-metal atoms have a large difference in electronegativity values so that bonding electrons are not shared but are completely transferred from the atom with the lower electronegativity value to the atom with the greater electronegativity value.

Question 52

What structure do ionic compounds exist as and why?

Answer 52

Lattice structures. Each positive ion attracts all the negative ions around it, and each negative ion attracts all the positive ions around it. A giant structure of repeating units is formed, which stretch in all directions.

Question 53

What are the melting and boiling points like for ionic compounds and why?

Answer 53

They are high and exist as solids at room temperature. The ions are tightly held together by strong ionic bonds and a lot of energy is needed to break the bonds.

Question 54

What are Van Der Waals’ forces?

Answer 54

The forces of attraction acting between the molecules which hold them together.

Question 55

What are three main types of Van Der Waals’ forces?

Answer 55

Temporary attractions:
* London Dispersion Forces

Permanent attractions:
* Permanent dipole-permanent dipole attractions
* Hydrogen bonding

Question 56

Where do permanent dipole- permanent dipole (pd-pd) attractions occur?

Answer 56

Between polar molecules.

Question 57

What do all polar molecules have?

Answer 57

A permanent dipole: one side of the molecule is always slightly negative (-δ) and the other side is always positive (+δ).

Question 58

How do pd-pd forces occur?

Answer 58

One side of a molecule will be slightly negative and the other slightly positive due to the difference in electronegativity (the atom with smaller electronegativity will be slightly positive). The negative side will be attracted to the positive side of other molecules.

Question 59

What determines if a molecule is polar or non-polar?

Answer 59

The electronegativity of its atoms and the shape of the molecule.

Question 60

What makes a polar molecule?

Answer 60

Must contain polar covalent bonds and have a non-symmetrical shape.

Question 61

Is water a polar or non-polar molecule and why?

Answer 61

Polar. It has 2 polar bonds and has no symmetry as it has a bent bond making it polar.

Question 62

What makes a molecule non-polar?

Answer 62

They have a symmetrical shape, even though they might contain polar bonds.

Question 63

Why can non-polar molecules contain polar bonds?

Answer 63

Because the symmetrical shape causes the outer atoms to “block” the charge from the central atoms.

Question 64

What is hydrogen bonding?

Answer 64

In some polar molecules, a hydrogen atom is bonded to a very electronegative atom. The intermolecular force of attraction between these very polar molecules is called hydrogen bonding.

Question 65

What are the three bonds that result in hydrogen bonding between very polar molecules?

Answer 65

Nitrogen, N - H
Oxygen, O-H
Fluorine, F - H
NOF

Question 66

Rank the order of strength for Van der Waals’ forces from 1 (strongest) - 3 (weakest).

Answer 66

1. Hydrogen Bonding
2. Pd-pd
3. London dispersion forces

Question 67

What happens to the boiling/melting point of molecules the stronger the intermolecular forces and why?

Answer 67

The stronger the intermolecular force the higher the boiling/melting point. This is due to more energy being required to break the stronger force.

Question 68

What similarities are needed when comparing intermolecular forces?

Answer 68

The same number of electrons, therefore they will have the same strength of London dispersion forces. This allows us to guarantee any difference in melting/boiling point is due to permanent dipole-permanent dipole attractions or hydrogen bonding.

Question 69

What is viscosity and what does it mean if something has a high or low viscosity?

Answer 69

Viscosity is a measure of the thickness of a liquid. A thick, syrupy liquid has a high viscosity – it is viscous. A thin, runny liquid has a low viscosity.

Question 70

What does the viscosity of a liquid depend on?

Answer 70

The strength of intermolecular forces present. The stronger the intermolecular forces, the more slowly the liquid will move when poured or stirred and so the greater the viscosity.

Question 71

What is water good at dissolving and why?

Answer 71

Water is very efficient at dissolving polar substances and ionic substances. This is because water molecules are very polar.

Question 72

What is miscibility?

Answer 72

The ability of liquids to mix in all proportions, forming a solution.

Question 73

Why is water a good solvent for polar covalent compounds and ionic compounds?

Answer 73

The positive δ+ and negative δ- charges of the water molecules are attracted to other positive and negative charges.

Question 74

Why are non-polar substances insoluble in polar substances like water?

Answer 74

Non-polar substances only have London dispersion forces and so cannot form intermolecular forces with polar substances.

Question 75

What is the rule for solubility?

Answer 75

Non-Polar substances are soluble in non-polar solvents. Polar substances are soluble in polar solvents.
Like dissolves like.

Question 76

In what way is water’s density different?

Answer 76

Due to the presence of hydrogen bonding. When water reaches 4°C it begins to expand again and become less dense until it freezes at 0°C. Solid ice is less dense than liquid

Question 77

Why is ice less dense than liquid water?

Answer 77

As water cool the molecules move more slowly. This allows more hydrogen bonds to form between the molecules, until every molecule has six hydrogen bonds, the maximum possible. This forms a very regular “open network” structure. This cause water to expand as it freezes to form ice.

Question 78

How do you predict the solubility of a compound?

Answer 78

1. Presence in molecules of O-H or N-H bonds, which implies hydrogen bonding
2. Spatial arrangement of polar covalent bonds, which could result in a molecule possessing a permanent dipole

Question 79

What is oxidation?

Answer 79

The loss of elections

Question 80

What is an oxidising agent?

Answer 80

A substance which causes oxidation of another substance. The oxidising agent must be an electron acceptor, so that the other substance is able to lose electrons. This means that the oxidising agent is itself reduced as it works during a redox reaction.

Question 81

What is reduction?

Answer 81

The gain of electrons

Question 82

What is a reducing agent?

Answer 82

A substance which causes reduction of another substance. The reducing agent must be an electron donor, so that the other substance is able to gain electrons. This means that the reducing agent is itself oxidised as it works during a redox reaction.

Question 83

Are metals good oxidising or reducing agents and why?

Answer 83

Metal elements have one, two or three electrons in their outer energy levels. They have low electronegativity values and tend to lose their outer electrons easily - they undergo oxidation.
∴They are good reducing agents.

Question 84

Are non-metals good oxidising or reducing agents and why?

Answer 84

Non-metal elements have electron arrangements that need only one, two or three electrons to give a full outer energy level. They have high electronegativity values and tend to gain electrons easily – they undergo reduction.
∴They are good oxidising agents.

Question 85

What does a high electronegativity mean?

Answer 85

The higher the electronegativity of an element, the more likely the element is to gain electrons and act as an oxidising agent.

Question 86

What does a lower electronegativity mean?

Answer 86

The lower the electronegativity of an element, the more likely the element is to lose electrons and act as a reducing agent.

Question 87

What is carbon monoxide (CO)?

Answer 87

Carbon monoxide (CO) is the reducing agent in the industrial extraction of iron from iron or as carried out in a blast furnace.
Fe2O3 (s) + 3CO (g) → 2Fe (s) + 3CO2 (g)

Question 88

What is hydrogen peroxide (H₂O₂)?

Answer 88

Hydrogen peroxide (H₂O₂) is one of the most powerful oxidising agents. As it acts, it is reduced to water.
H2O2 (l) + 2H+ (aq) + 2eˉ → 2H2O (l)

Question 89

What is dichromate ions (Cr₂O₇²⁻) and permanganate ions (MnO₄⁻)?

Answer 89

Dichromate ions (Cr₂O₇²⁻) and permanganate ions (MnO₄⁻) are powerful oxidising agents. They are both reduced. (I can’t be bothered typing out the reactions, fuck you, the school and the SQA if for some reason you actually need them.)

Question 90

What is hydrogen peroxide used as?

Answer 90

Bleach. It can oxidise coloured compounds, producing colourless products. It is used to bleach wool, cotton, and paper, lighten har colour and whiten teeth.

Question 91

What is potassium permanganate used for?

Answer 91

As an antiseptic and disinfectant. Dilute potassium permanganate solution can kill fungi and bacteria and inactivate some viruses.

Question 92

How do you balance ion-electron equations?

Answer 92

1. Write the main reactant and product for the reduction or oxidation. Balance the metal or non-
   metal element present (but not the oxygen – see the next step).
2. Add water to one side to balance the oxygen.
3. Add hydrogen ions to the other side to balance the hydrogen.
4. Add electrons to the same side as the hydrogen ions so that both sides of the equation have the
   same total charge.

Question 93

What is the electrochemical series rule?

Answer 93

Applying the rule that “electrons flow from high to low”, this means that:
* the oxidation ion-electron equation (electrons lost) must be higher in the electrochemical series
than the reduction ion-electron equation (electrons gained)
* the oxidising agent must be lower down on the left and the reducing agent must be higher up
on the right