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Chemistry > AoS1 - Ionic Compounds and Metals > Flashcards

AoS1 - Ionic Compounds and Metals Flashcards

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1
Q

What is an Ionic Compound?

A

A compound made up of cations and anions in which produce a neutral substance.
- Metal transfers valence electrons to the Non-metal.

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2
Q

Electrostatic Forces in Ionic Bonding

A

The forces between the anion and cation that hold the compound together, creating ionic bonding.

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3
Q

Neutral Substance

A

The Ion charges cancel out

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4
Q

Arrangement of Ionic compounds

A

Crystal lattice structure
- Ions are in fixed positions in a lattice.
- based on size and ratio of the ions

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5
Q

How are ions held together in Ionic Bonding?

A

The attraction between the anions and cations. This is called the ionic bond.

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6
Q

The most stable arrangement of ions in Ionic Bonding

A

Positively charged ions are packed to the nucleus as closely as possible to the negatively charged particles. Ions with the same charge are as far away as possible.

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7
Q

Co-ordination number

A

Used to describe the ratio of ions.

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8
Q

Properties of Ionic Compounds

A
  • High melting and
    boiling points
  • hard
  • brittle
  • conductivity
  • solubility
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9
Q

Explain the high melting and boiling point of Ionic Compounds.

A

Attractive forces between the ions are strong and therefore a large amount of energy is required to break these bonds.

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10
Q

Explain the hard but brittle property of Ionic Compounds.

A

Forces between ions are strong, hence it is very hard to break these forces. However, when these forces are broken and the lattice structure shifts and repulsions occur between like charges, structure is shattered.

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11
Q

Conductivity of Solid Ionic Compounds.

A

Do not conduct electricity as there is no free moving ions

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12
Q

What is a molten compound?

A

Are compounds that are heated into a liquid

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13
Q

Conductivity of Molten and Aqueous Compounds in Ionic Bonding

A

There are free moving ions that can therefore conduct electricity

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14
Q

Solubility of Ionic Compounds

A

Solubility varies depending on the ionic bond strength. Compounds that form really strong bonds will be insoluble in solution

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15
Q

Monoatomic Ion

A

Are ions containing one atom only

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16
Q

Polyatomic Ion

A

Are ions made up of more than one atom

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17
Q

If an ion forms more than one ion

A

Must specify the charge when naming

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18
Q

Transition metals (in terms of forming ions)

A
  • Tend to loose electrons to form cations
  • some elements will form more than one stable ion
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19
Q

Why can transition metals form more than one ion?

A
  • similar energy levels in 3d and 4s orbitals
  • valence electrons can easily jump to both
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20
Q

Naming anions

A

Suffix “ide” is added + ion

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21
Q

Empirical Formula

A

The simplest whole number ratio of atoms in a compound

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22
Q

Chemical Formulae Rules

A
  1. When naming and writing ionic compounds, cations written first
  2. Positive and negative charges must be balanced
  3. Subscripts indicate number of ions
  4. Electrovalency must be specified
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23
Q

Solubility Rules

A

Can be used to predict whether an ionic compound will be soluble or insoluble under standard Laboratory Conditions

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24
Q

Standard Laboratory Conditions

A

25 degrees and 1 atmosphere pressure

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25
Q

Determining Solubility

A

Identify anion and determine if soluble then identify if cation is soluble or an insoluble exception.

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26
Q

Precipitation reaction

A

A reaction in which a precipitate is formed

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27
Q

Precipitate

A

An insoluble solid that forms when two or more solutions are mixed. Solid seperates from the solution.

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28
Q

Why does a solid form in a precipitate reaction?

A

The attraction between cations and anions is greater than the attraction between the individual ions and the surrounding water molecules.

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29
Q

Balanced formula Equation

A

Shows all reactants and products (including spectator ions)

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30
Q

Spectator Ions

A

Ions which remain in solution and are unchanged (not involved).

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31
Q

Ionic Equation

A

Only shows reacting species which produce the precipitate.

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32
Q

What percent of atoms in the earths crust are metals?

A

25%

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33
Q

Metals exist in

A

Element form or as compounds

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34
Q

Element form

A

very unreactive eg. Gold and Silver

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35
Q

Compound form

A

combined with non-metal elements called ores

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36
Q

Properties of Metal compounds

A
  • Strong
  • Can be light
  • Ductile
  • Conducts Electricity
  • Malleable
  • Lustrous
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37
Q

Metallic bonding model

A

Arranged in a metal lattice.
Lattices differ in charge on the metal ion and the size of the metal ion.

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38
Q

What is a lattice?

A

3-d regular arrangement of particles

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39
Q

Positive Ions in Metal Lattice

A

Form regular 3d lattice, fixed positions and closely packed.

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40
Q

Valence Electrons in Metal Lattice

A

Free to move, delocalised electrons that belong to the lattice as a whole, sea of electrons

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41
Q

Electrons in Inner shells of the Metal Lattice

A

Localised

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42
Q

Attraction in metal compound

A

Attraction between positive cations and delocalized electrons creates metallic bonds.

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43
Q

Conductivity of Electricity - Metal

A

Contains charged particles that are free to move (delocalised electrons)
If a current is applied, the electrons are forced in at one end and an equal number flow out the other, hence producing a current.

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44
Q

Conductivity of Heat - Metal

A

When delocalised electrons bump into each other and the positive ions, they transfer electricity

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45
Q

Lustrous - Metal

A

Light can be reflected due to the presence of free moving and delocalised electrons

46
Q

Malleable and Ductile - Metal

A

Forces between the particles can adjust when the particles move, without breaking. When rearranging a metal, the positive ions are forced across each other. The delocalised electrons also move to compensate.

47
Q

Malleable definition

A

Can be hammered into sheets.

48
Q

Ductile definition

A

Can be drawn into a wire

49
Q

High melting and boiling temperatures - Metal

A

Strong forces of attraction between the positive ions and the negative delocalised electrons aka Electrostatic forces.

50
Q

The higher the charge (referring to melting and boiling points)

A

The stronger the attraction and thus the higher melting/boiling point.

51
Q

Density

A

Mass per unit volume

52
Q

Density - Metals

A

Positive ions are closely packed. Depends on the mass of the metal ions, their radius and the way they are packed in the lattice.

53
Q

Metal reactions

A

React with a variety of substances to form ionic compounds.
Involves losing electrons to become a cation (oxidation)

54
Q

Predictable reactions of Metal

A

Those reacting with water, oxygen and acids.

55
Q

Single displacement Reactions

A

Metal + Water, Metal + Acid

56
Q

Three reactions

A

Metal + water, Metal + Oxygen, Metal + acid

57
Q

Metal + Water

A

Metal Hydroxide + Hydrogen Gas
Single Displacement

58
Q

Metal + Oxygen

A

Metal Oxide
Synthesis

59
Q

Metal + Acid

A

Salt (IC) + Hydrogen Gas
Single Displacement

60
Q

Reactivity depends on

A

Ionization energy (atomic radius + core charge)
- The lower the Ionization energy, the easier it is to lose electrons.

61
Q

Most reactive group

A

Group 1 Metals

61
Q

Qualitative relative reactivity

A

amount of bubbles produced or the extent to which metals react.

62
Q

Molecule Definition

A
  • A discrete group of two or more non-metal atoms covalently bonded to one another.
  • The overall charge is neutral
  • Can contain more than one type of atom
63
Q

Types of forces in Covalent Compounds

A

Intramolecular and Intermolecular Forces

64
Q

Intramolecular Forces

A

Strong forces of attraction within the molecule
- Also known as covalent bonds

65
Q

Intermolecular Forces

A

Weak forces of attraction between the molecules
- the forces much weaker than both covalent or ionic bonds
- determines properties.

66
Q

Covalent Bonds

A

Bonds that hold atoms within molecules
Sharing of valence electrons between 2 or more atoms to produce a stable outer shell configuration
- Bonding pairs are localised.

67
Q

Octet Rule

A

8 electrons in the outer shell

68
Q

Single Covalent Bonds

A

One pair of electrons being shared

69
Q

Double covalent bond

A

2 pairs of electrons being shared to create a double bond

70
Q

Triple covalent bond

A

3 pairs of electrons being shared to form a triple bond

71
Q

Name of the electrons not involved in covalent bonding

A

Lone pairs or non-bonding pairs

72
Q

Valence shell electron pair repulsion theory

A

assumes electron pairs are located as far away as possible to minimise repulsion

73
Q

Molecular formula

A

provides the actual number of each type of element present in a compound
no information about structure or bonding

74
Q

Lewis structure

A

Illustrated number of valence electrons an atom has
Valence electrons represented by dots
Electrons involved are between the atoms

75
Q

Electronegativity

A

a measure of an atoms ability to attract electrons.
The higher the electronegativity, the greater the attraction.

76
Q

Difference in electronegativity

A
  • Valence electrons are not always equally shared between two elements
  • If the difference is greater than 0.4, they will not share equally.
77
Q

Two charge types of covlent bonds

A

Polar and non polar.

78
Q

Molecule shape

A

Describes the way in which the atoms are arranged around a central atom in 3d space

79
Q

Central Atom

A

always has the smallest electronegativity and hence smallest number of valence electrons.

80
Q

Electron domains

A

Total number of pairs, repulsion results in them being as far away as possible.

81
Q

Linear

A

Diatomic molecules or 2 atoms coming off a central atom. Central atom has no lone pairs of electrons.

82
Q

Diatomic molecule

A

made of two atoms

83
Q

Tetrahedral

A

four atoms surrounding a central atom.

84
Q

Pyramidal

A

Three atoms surrounding a central atom + one lone pair of electrons coming off the central atom.

85
Q

Bent

A

Two atoms surrounding a central atom + 2 lone pairs of electrons coming off the central atom.

86
Q

Polar Molecule

A

A polar bond has one or more polar covalent bonds that are arranged asymmetrically.
- result is called a dipole

87
Q

Dipole

A

Two oppositely charged ends or poles.

88
Q

Polar bond

A
  • electrons pulled towards more electronegative atom
  • creates partially +ve and partially -ve.
  • The greater the difference the more polar the bond.
89
Q

Non-polar molecule

A
  • No polar covalent bonds
  • Has polar bonds but arranged symmetrically
90
Q

Dipole- Dipole bond

A

Permanent dipole-dipole bond forces are an intermolecular, electrostatic attraction between dipoles of polar molecules.

90
Q

Types of intermolecular bonds

A

Hydrogen bonds, dipole-dipole bonds, dispersion bonds

90
Q

Hydrogen bonding

A
  • Stronger form of dipole-dipole bonding
  • occurs when Hydrogen attracted to Fluorine, Oxygen, Nitrogen
91
Q
  • Dispersion forces
A
  • induced non-permanent dipole-dipole forces
  • constant movement within atoms and ions produce temporary positive and negative regions in the molecule (instantaneous dipole)
  • OCCUR IN ALL SUBSTANCES (NON-POLAR)
  • the larger the molecule, the greater the dispersion forces (as there are more electrons)
92
Q

Properties of covalent molecules

A
  1. Do not conduct electricity in solid or molten state
  2. Low melting and boiling points
  3. Soft as a solid
93
Q

Conductivity of covalent molecules

A

No charged particles as the molecules are neutral.

94
Q

Low melting and boiling points of covalent compounds

A
  • weak intermolecular forces means only a small amount of energy is needed to change the state of a molecular substance.
95
Q

Soft as a solid - Covalent compound

A

Weak intermolecular forces of attraction

96
Q

Melting point and boiling point of Non-Polar Covalent compounds

A
  • only have weak dispersion forces
  • stronger dispersion forces increase points
  • greater number of electrons increase the dispersion forces.
97
Q

Melting point definition

A

Temperature where solid becomes a liquid.

98
Q

Boiling point definition

A

Temperature in which liquid becomes a gas.

99
Q

Melting and boiling points of Polar Molecules.

A
  • Hydrogen bonds have highest
  • Dipole- dipole second highest.
100
Q

Allotrope

A

Different physical form of the same element

101
Q

Allotropes of Carbon

A

Diamond
Graphite
Fullerenes - bucky balls, nanotubes and graphemes

102
Q

Diamond

A
  • covalent network lattice
  • each C atom is covalently bonded to 4 other C atoms in 3 dimensions
  • High melting point
  • Chemically Inert
  • Hardest Known substance
  • Non conductive
  • Insoluble
103
Q

Other network Lattices

A

Silicon, silicon carbide, silicon dioxide

104
Q

Graphite

A
  • Covalent layer lattice
  • Each C atom bonded to 3 other C atoms
  • Fourth electron delocalised and free to move within layers
  • weak dispersion forces between layers
  • conductor of electricity
  • soft, weak, greasy
  • high melting temp.
105
Q

Why is diamond the hardest substance?

A

Has the strongest form of intramolecular forces due to the network lattice in which atoms are arranged.
Each carbon atom is covalently bonded to four other carbons.

106
Q

Why does silicon dioxide have a high melting point?

A
  • carbon atoms arranged in a network lattice
  • network lattice significantly stronger than the intermolecular forces of carbon dioxide
  • strong covalent bonds in all three dimensions require great energy to break.
107
Q

Why can graphite be used as a lubricant?

A
  • Layers held together by weak dispersion forces so they can readily slide past one another.

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