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Higher Chemistry > Unit 1 - Chemical Changes And Structures > Flashcards

Unit 1 - Chemical Changes And Structures Flashcards

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1
Q

(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

A
  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
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2
Q

How do metals bond?

A

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.

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3
Q

What is metallic bonding?

A

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

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4
Q

(Nat 5)
Why are metals electrical conductors

A

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

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5
Q

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

A

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

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6
Q

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

A

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.

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7
Q

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

A

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.

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8
Q

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

A

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

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9
Q

What two covalent structures does carbon have?

A

Graphite and diamond

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10
Q

Explain bonding in diamond

A

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.

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11
Q

Explain bonding in graphite

A

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.

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12
Q

What forces hold monatomic atoms together?

A

London dispersion forces

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13
Q

What is an instantaneous temporary dipole?

A

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).

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14
Q

What is an induced dipole?

A

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

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15
Q

What are London dispersion forces (LDFs)?

A

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

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16
Q

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

A

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.

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17
Q

What is a moleule?

A

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

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18
Q

What type of bonding occurs inside molecules?

A

Intramolecular.

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19
Q

What bonding occurs between molecules?

A

Intermolecular (LDFs)

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20
Q

What is sublimation?

A

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

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21
Q

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

A

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

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22
Q

What is covalent radius?

A

The measure of the size of an atom.

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23
Q

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

A

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.

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24
Q

What do outer electrons experience in atoms?

A

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

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25
What happens to the covalent radii down a group and why?
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.
26
What is the ionisation energy?
A measure of how tightly an atom holds on to its outer electrons.
27
What is the first ionisation energy?
The energy needed to remove one mole of electrons from one mole of gaseous atoms, measured in kJmol⁻¹.
28
What is the second ionisation energy?
The energy needed to remove one mole of electrons from one mole of gaseous 1+ (already lost one electron) ions.
29
What happens to the ionisation energy across a period and why?
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.
30
What happens to the ionisation energy down groups and why?
It decreases. As you go down the group, there is increased shielding and the attraction of the nucleus for the outer electrons decreases.
31
Is the second ionisation energy higher or lower than the first ionisation energy for elements in group 1 and why?
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.
32
How do you calculate total ionisation energy?
Add both the quantitites. together.
33
What is electronegativity?
Electronegativity is a measure of the attraction that a bonded atom has for the electrons in the bond.
34
What happens to electronegativity across a period and why?
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.
35
What happens to electronegativity down the groups and why?
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.
36
What is a covalent molecule?
A small group of non-metal atoms held together by strong covalent bonds.
37
How are covalent bonds formed?
By the electrostatic force of attraction between two positive nuclei and a shared pair of electrons.
38
Why do covalent molecules have low melting and boiling points?
The forces of attraction between molecules (LDFs) are very weak and very little energy is needed to separate the molecules.
39
What is a covalent network?
A giant 3D structure in which atoms are joined by strong covalent bonds.
40
Why do covalent networks have have high melting and boiling points?
The atoms are tightly held together by strong covalent bonds and a lot of energy is needed to break these.
41
What is non-polar covalent bonding?
Bonding in which the electrons are equally shared (both atoms have same electronegativity from a bond).
42
What type of bonds do covalent elements always have and why?
Non-polar, because two atoms of the same element must have the same electronegativity. They must share their electrons equally.
43
How do you know if a bonding is non-polar (typically)?
The electronegativity between the atoms is almost zero.
44
What is polar covalent bonding?
Bonding in which there is a permanent dipole.
45
How are permanent dipoles formed?
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 +).
46
What makes a bond more polar?
A bigger difference in electronegativity.
47
How do you know if bonding is polar covalent (typically)?
If the difference in electronegativity between the atoms is small (less than 2).
48
What is an ionic bond?
The electrostatic attraction between a positive ion and a negative ion.
49
Explain what happens to the metal and non-metal atom to from an ionic bond.
The metal atom loses an electron, forming a positive ion, and the non-metal atom gains an electron, forming a negative ion.
50
How do you now if bonding is ionic (typically)?
The difference in electronegativity between the atoms is large (more than 2).
51
Explain ionic bonding.
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.
52
What structure do ionic compounds exist as and why?
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.
53
What are the melting and boiling points like for ionic compounds and why?
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.
54
What are Van Der Waals' forces?
The forces of attraction acting between the molecules which hold them together.
55
What are three main types of Van Der Waals' forces?
**Temporary attractions:** * London Dispersion Forces **Permanent attractions:** * Permanent dipole-permanent dipole attractions * Hydrogen bonding
56
Where do permanent dipole- permanent dipole (pd-pd) attractions occur?
Between polar molecules.
57
What do all polar molecules have?
A permanent dipole: one side of the molecule is always slightly negative (-δ) and the other side is always positive (+δ).
58
How do pd-pd forces occur?
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.
59
What determines if a molecule is polar or non-polar?
The electronegativity of its atoms and the shape of the molecule.
60
What makes a polar molecule?
Must contain polar covalent bonds and have a non-symmetrical shape.
61
Is water a polar or non-polar molecule and why?
Polar. It has 2 polar bonds and has no symmetry as it has a bent bond making it polar.
62
What makes a molecule non-polar?
They have a symmetrical shape, even though they might contain polar bonds.
63
Why can non-polar molecules contain polar bonds?
Because the symmetrical shape causes the outer atoms to "block" the charge from the central atoms.
64
What is hydrogen bonding?
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.
65
What are the three bonds that result in hydrogen bonding between very polar molecules?
Nitrogen, N - H Oxygen, O-H Fluorine, F - H NOF
66
Rank the order of strength for Van der Waals' forces from 1 (strongest) - 3 (weakest).
1. Hydrogen Bonding 2. Pd-pd 3. London dispersion forces
67
What happens to the boiling/melting point of molecules the stronger the intermolecular forces and why?
The stronger the intermolecular force the higher the boiling/melting point. This is due to more energy being required to break the stronger force.
68
What similarities are needed when comparing intermolecular forces?
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.
69
What is viscosity and what does it mean if something has a high or low viscosity?
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.
70
What does the viscosity of a liquid depend on?
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.
71
What is water good at dissolving and why?
Water is very efficient at dissolving polar substances and ionic substances. This is because water molecules are very polar.
72
What is miscibility?
The ability of liquids to mix in all proportions, forming a solution.
73
Why is water a good solvent for polar covalent compounds and ionic compounds?
The positive δ+ and negative δ- charges of the water molecules are attracted to other positive and negative charges.
74
Why are non-polar substances insoluble in polar substances like water?
Non-polar substances only have London dispersion forces and so cannot form intermolecular forces with polar substances.
75
What is the rule for solubility?
Non-Polar substances are soluble in non-polar solvents. Polar substances are soluble in polar solvents. **Like dissolves like.**
76
In what way is water's density different?
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**
77
Why is ice less dense than liquid water?
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.
78
How do you predict the solubility of a compound?
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
79
What is oxidation?
The loss of elections
80
What is an oxidising agent?
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.
81
What is reduction?
The gain of electrons
82
What is a reducing agent?
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.
83
Are metals good oxidising or reducing agents and why?
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.**
84
Are non-metals good oxidising or reducing agents and why?
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.**
85
What does a high electronegativity mean?
The **higher** the electronegativity of an element, the more likely the element is to gain electrons and act as an **oxidising agent**.
86
What does a lower electronegativity mean?
The **lower** the electronegativity of an element, the more likely the element is to lose electrons and act as a **reducing agent**.
87
What is carbon monoxide (CO)?
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)
88
What is hydrogen peroxide (H₂O₂)?
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)
89
What is dichromate ions (Cr₂O₇²⁻) and permanganate ions (MnO₄⁻)?
**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.)
90
What is hydrogen peroxide used as?
**Bleach**. It can oxidise coloured compounds, producing colourless products. It is used to bleach wool, cotton, and paper, lighten har colour and whiten teeth.
91
What is potassium permanganate used for?
As an antiseptic and disinfectant. Dilute potassium permanganate solution can kill fungi and bacteria and inactivate some viruses.
92
How do you balance ion-electron equations?
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.
93
What is the electrochemical series rule?
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

Higher Chemistry (1 decks)

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