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Higher Chemistry > Unit 1 (H) > Flashcards

Unit 1 (H) Flashcards

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1
Q

What is a lattice

A

A lattice is a regular repeating pattern

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2
Q

How are atoms arranged in a solid metal

A

In a solid metal, the atoms are arranged neatly in a simple repeating pattern known as a lattice.

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3
Q

how do atoms in a lattice stay together?

A

Atoms in a metal contribiute their outer most electrons to a pool of free or delocalised electrons.

Each positively charged ion is attracted to the pool of delocalised electrons.

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4
Q

Describe the metallic bond

A

The metallic bond is the electrostatic force of attraction between delocalised electrons and the nuclei of neighbouring atoms

The metallic bond is quite strong and acts in all directions around each metal atom.

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5
Q

What does the term delocalised electron mean

A

In a metallic lattice the outermost electrons are said to be delocalised because they are not fixed to one atom - they can move freely from atom to atom

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6
Q

What is the significance of delocalised electrons in a metallic bond

A

Each free moving electron is attracted to the nucleus of more than one neighbouring atom.

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7
Q

What are the physical properties of a substance

A

the physical properties of a substance are those properties which do not involve the substance in chemical reactions.

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8
Q

What are common examples of physical properties

A

Colour, State, Melting point, conductivity, malleability and ductility

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9
Q

What are the physical properties of metals and what causes them

A

The metallic bonding attractions are not directional and this results in the metal being malleable and ductile. They are also conductive due to the free moving electrons in the structure which allow electricity to flow.

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10
Q

Define malleable

A

Able to be rolled into sheets. (or hammered into a specific shape)

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11
Q

Define ductile

A

Able to be stretched out.

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12
Q

Why are metals malleable

A

Since delocalised electrons move freely. Bonds can change direction without breaking.

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13
Q

Why are metals ductile

A

Since delocalised electrons move freely. Bonds can change direction without breaking.

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14
Q

What must happen for a metal to boil

A

When a metal is boiled, all the forces of attraction between the atoms must be overcome.

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15
Q

What is the relationship between the BPs of metals and their attractive forces

A

The boiling point of metals must be a guide to the strengths of the attractive forces between the atoms.

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16
Q

What happens to the boiling point as you move down group 1 metals

A

the boiling point decreases

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17
Q

Why does the boiling point of group 1 metals go down as you move down the group

A

The metallic bond strength is decreasing. This is because the outer electrons are further from the nucleus and so the electrostatic attraction becomes less.

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18
Q

What happens to the boiling point of metals as you move along period 4

A

The boiling point increases.

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19
Q

Why does the boiling point change as you move along period 4 metals

A

The metallic bond strength increases. This is because of an increased nuclear charge and increased number of electrons which makes the electrostatic attractions more frequent and stronger.

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20
Q

What are monatomic elements

A

The monatomic elements exist as single atoms in the gas state.

At extremely cool temperatures they will turn from a gas to a liquid and then to a solid

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21
Q

What Atractions cause monatomic elements to turn into liquid/solid

A

The attractions between the atoms which cause this are known as London dispersion forces.

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22
Q

What are the monatomic elements

A

The noble gases helium, neon, argon, krypton, xenon, and radon are the only monatomic elements.

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23
Q

Why do the noble gases (monatomic elements) have such low boiling/melting points

A

The London dispersion forces which keep the atoms together are very weak which means little energy is needed to break the attractions.

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24
Q

Describe London dispersion forces and what causes them

A

London dispersion forces are caused by a vibration or wobbling in the electron clouds. This results in very small and very tempory charges on the atom.

These charges are written as delta negative and delta positive.

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25
Q

What are dipoles

A

Dipoles are points in the atom where the charge is negative or positive

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26
Q

How do London dispersion forces in one atom effect neighbouring atoms

A

The dipole moment will induce a temporary dipole in a neighbouring atom by attracting/repelling its electron charge cloud.

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27
Q

Why do the boiling points of noble gases increase as you go down a group

A

The atoms get larger which causes more and greater wobble in the electron cloud which causes more and greater dipoles which means more and stronger London dispersion forces. (The more electrons the greater the attractive force.)

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28
Q

What are molecular elements

A

Molecular elements are elements that are made of individual molecules. The bonding within these molecules is covalent.

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29
Q

What is a covalent bond?

A

A covalent bond is the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the two atoms which are sharing these electrons

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30
Q

Define intermolecular forces

A

Attractive forces between molecules. Responsibile for keeping matter in solid or liquid phase. These are most commonly London dispersion forces as they can occur in any atom.

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30
Q

What is a diatomic element

A

A diatomic element is an element which exists as molecule made up of just two atoms.

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31
Q

What are the 7 diatomic elements

A

Hydrogen, oxygen, iodine, bromine, Flourine, Chlorine, Nitrogen

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32
Q

Define intramolecular forces

A

Forces which keep molecules together. ie bonds

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33
Q

What elements other than the noble gases can exist as molecular elements

A

Phosphorous, Sulphur

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34
Q

Why are sulfur and phosphorous solids?

A

They are large molecules (P4, S8) which means greater intermolecular forces making the boiling/melting points higher. As the number of similar sized atoms whithin the molecule increases the number of london dispersion forces increases leading to an increase in melting point.

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35
Q

What are molecular forms of carbon known as

A

Fullerenes, these exist as large molecules containing 60 or more carbon atoms.

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36
Q

Describe fullerene

A

Fullerene has a low density due to its hollow sphere shape with moderate melting point due to stronger london dispersion forces between molecules. THEY ARE NOT COVALENT NETWORKS.

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37
Q

Give an example of a fullerene

A

Buckminister fullerene (C60)

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38
Q

What two groups can covalent substances be divided into

A

Covalent network.
Discrete covalent molecules.

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39
Q

Name common examples of covalent network substances.

A

Covalent network substances include Boron, Silicon, Carbon (graphite, diamond)

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40
Q

Describe covalent network structures

A

Substances are made up of covalently bonded atoms in a three dimensional network or array.

40
Q

Describe the structure of diamond

A

In diamond, each carbon atom is joined to four others by strong covalent bonds. The covalent network structure is very strong and hard.

41
Q

What is the name of the structure of diamond

A

tetrahedral.

42
Q

Why are diamonds not malleable or conducting

A

The bonding electrons between the carbon atoms are not free moving.

43
Q

What makes graphite (carbon) different from diamond (carbon)

A

In graphite, London dispersion forces hold together layers of covalently bonded carbon atoms

44
Q

Why is graphite a conductor

A

One electron per carbon atom is not involved in covalent bonding and is free to move allowing it to conduct.

45
Q

Why do layers of graphite slide over each other easily

A

Long weak london dispersion forces break readily (between the layers) so the layers slide over each other

46
Q

What is diamond used for

A

cutting tools as it is very hard

47
Q

What are common uses of graphite

A

Electrodes (as it can conduct electricity), Solid lubricant and pencil “lead” (as layers slide each other)

48
Q

Why do covalent networks have such high melting points

A

They are a large network of strong covalent bonds which take a large amount of energy to break.

49
Q

How can we tell boron and silicon are covalent networks

A

their high melting points

50
Q

What is the covalent radius of an atom

A

half the distance between the nuclei of two identical atoms when they are joined by a single covalent bond. The covalent radius is measured in picometres and gives us a measure of atomic size.

(you can use covalent radius to find the bond length by adding 2 covalent radii)

51
Q

What happens to atomic size going across a period

A

the atomic size decreases as the nuclear charge increases and attracts the outer electrons towards the nucleus.

52
Q

What happens to atomic size going down a group

A

the atomic size increases as an extra electron shell is added (but the number of outer electrons stay the same)

53
Q

Why is there no known covalent radii for the noble gases

A

they do not form bonds

54
Q

What is the first ionisation energy of an element.

A

The first ionisation energy of an element is the energy required to remove one mole of elecrtrons from one mole of atoms in the gaseous state.

55
Q

What is the second ionisation energy of an element

A

The second ionisation energy of an element is the energy required to remove a second mole of electrons from a mole of free atoms once a first mole of electrons has already been removed.

56
Q

Why is the second ionisation energy greater than the first ionisation energy

A

In the second ionisation energy negative electrons are being removed from positive ions rather than neutral atoms. The positive ion has a greater attraction for the electron so more energy is needed.

57
Q

What happens to ionisation energy as you go across periods and why

A

Ionisation energy increases due to the nuclear charge increasing which holds the electrons more tightly requiring more energy to be used to remove a mole of electrons.

58
Q

What happens to ionisation energy going down a group and why

A

Ionisation energy decreases, this is due to more electron shells, which means the outer electrons are further from the nucleus making them harder to attract additionally the screening effect of the inner electrons means the attraction for the outer electrons decreases despite the increases nuclear charge.

59
Q

Why is the jump between ionisation energies sometimes particullarly large

A

This is when the electron to be removed is in a closer electron shell meaning more attraction.

60
Q

Why is it harder to remove electrons from certain atoms (like noble gases or the second ionisation energy of pottasium)

A

If the electron to be removed comes from a full outer shell a greater amount of energy is needed.

61
Q

What is an element

A

An element is a chemical substance which cannot be broken down into anything simpler. Each element has its own type of atom.

62
Q

What are the 8 groups in the periodic table called

A

Group 1 -Alkali metals
Group 2 - Alkali-earth metals
Group 3 - [N/A]
Group 4 - Carbon group
Group 5 - [N/A]
Group 6 - [N/A]
Group 7 - Halogens
Group 8/0 - Noble gases

63
Q

What are transition metals

A

The 38 elements placed between groups 2 and 3.

64
Q

Which of the first 20 elements are monatomic

A

Helium, Neon and Argon

65
Q

Which of the first 20 elements are covalent molecular

A

Phosphorous, Oxygen, Sulfur, Nitrogen, Chlorine, Flourine, hydrogen

66
Q

Which of the first 20 elements are covalent networks

A

Boron, Carbon, Silicon

67
Q

Which of the first 20 elements are metallic laticces

A

Lithium, Magnesium, Potassium, aluminium. berylium, Sodium, calcium

68
Q

What are the first 20 elements

A

Hydrogen, Helium, Lithium, Beryllium, Boron, Carbon, Nitrogen, Oxygen, Fluorine, Neon, Sodium, Magnesium, Aluminium, Silicon, Phosphorus, Sulphur, Chlorine, Argon, Potassium, and Calcium.

69
Q

What is electronegativity

A

Electronegativity is a measure of the attraction an atom, which is involved in a covalent bond, has for the shared pair of electrons.

70
Q

What is the ionic bond

A

The ionic bond is the electrostatic attraction between oppositely charged ions in a compound.

71
Q

Describe an ionic compound

A

Metals lose electrons to form positive ions and non metals gain electrons to form negative ions creating an attraction between the two.

72
Q

The formula of an ionic substance tells us ____

A

the ratio of the ions present in the compound

73
Q

Describe the melting point of ionic compounds

A

the melting points of ionic compounds tend to be quite high due to the large amount of attractions that need to be broken in order to seperate the compound.

74
Q

What are molecular ions

A

This where both covalent bonding and ionic bonding are both present.

75
Q

Define molecular

A

Known number of atoms in a substance, atoms share electrons to achieve a full outer electron shell and so are non metals.

76
Q

Define compound

A

Made up of atleast 2 different elements.

77
Q

What is a pure covalent bond (non polar covalent bond)?

A

A covalent bond in which the electrons are shared equaly between the two atoms, this is the case for all non metals which have covalent bonds (phosphorous, Silicon, Carbon) This is because their attractions for the electrons are exactly the same.

78
Q

Describe polar covalent bonds.

A

In covalent compounds there might not be an equal sharing of the bonding electrons. One of the atoms might have a greater attractive force on the electrons and so end up with a bigger share of it. This attractive force is called electronegativity. Polar bonds are unevenly shared electrons within a covalent bond leading to a slight dipole.

79
Q

Why are there no electronegativity values for noble gases.

A

They do not form covalent bonds.

80
Q

What does the electronegativity value depend on?

A

The electronegativity value depends on the number of protons within the nuclei and the number of electron shells (extra layer of electrons results in a screening effect)

81
Q

What happens to electronegativity as you move left to right across a period.

A

the electronegativity value increases due to the increased number of protons.

82
Q

What happens to electronegativity moving down a group

A

the electronegativity value decreases due to the increasing influence of the screening effect and greater distance to the bonding electrons.

83
Q

What happens to the charge of atoms participating in a polar covalent bond

A

The atom with a stronger electronegativity will pull the electrons closer resulting in a permanent negative dipole, where as the other bonding electron will have a permanent positive dipole.

84
Q

What must to be true for a polar molecule to exist

A

There must be a dipole across the whole molecule (they can be split by a straight line into a positive and negative half)

85
Q

What are permanent dipole - permanent dipole attractions

A

Permanent dipole - Permanent dipole attractions are intermolecular forces between molecules which have permanent overall dipole (one exposed side of the molecule is slightly postive and one slightly negative.) such molecules are polar molecules

86
Q

What does the term like dissolves like mean

A

The term like dissolves like is based on the polarity of the molecules involved ie polar molecules will dissolve (mix completely) in other polar molecules and non-polar will dissolve in non-polar.

87
Q

Why do polar molecules dissolve in polar molecules

A

If two polar molecules are added to together permanent dipole - permanent dipole attractions between the different substances can form (as well as between different molecules of the same substance) this means the substances are miscible

88
Q

What happens to BPs of molecules as number of atoms increases

A

Increases due to more LDFs being present.

88
Q

Why can two non polar molecules mix

A

LDFs can be formed between the different molecules.

88
Q

Why can polar molecules and non polar molecules not mix

A

If a polar molecule and non polar molecule are added together permanent dipole attractions are not possible. The polar molecules clump together due to permanent dipole - permanenet dipole attractions the non polar molecules can only form LDFs which are very weak making the substance immsicible.

89
Q

What happens to the BPs of molecues as the mass of the molecules increases

A

generally increases as the size of LDFs increases (due to larger atoms)

89
Q

What are the three types of intermolecular forces (in order of strength)

A

LDFs, Polar polar attractions and hydrogen bonding.

90
Q

What 3 things are needed for hydrogen bonding to occur

A

A large dipole between hydrogen and another atom.
A lone pair of electrons in the other atom
The other atom must be a small atom.

91
Q

What 3 elements meet the conditions for hydrogen bonding

A

Nitrogen oxygen and flourine.

92
Q

Why can hydrogen bonding not occur with chlorine

A

The chlorine atom is too large, the two neighbouring molecules are attracted to each other but the large size of the chlorine atoms with negative dipoles will cause them to repell

93
Q

What are the 3 intermolecular forces known as

A

Van der Waals forces

94
Q

What type of bonding does water have

A

Hydrogen bonding given that oxygen is directly bonded to hydrogen

95
Q

Why does water have a high boiling point

A

Caused by hydrogen bonding.

96
Q

Why does ice float in water

A

Ice is less dense than water (Ice is an open network structure.)

Higher Chemistry (1 decks)

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