AoS1 - Ionic Compounds and Metals Flashcards
(110 cards)
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.
2
Q
Electrostatic Forces in Ionic Bonding
A
The forces between the anion and cation that hold the compound together, creating ionic bonding.
3
Q
Neutral Substance
A
The Ion charges cancel out
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
5
Q
How are ions held together in Ionic Bonding?
A
The attraction between the anions and cations. This is called the ionic bond.
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.
7
Q
Co-ordination number
A
Used to describe the ratio of ions.
8
Q
Properties of Ionic Compounds
A
- High melting and
boiling points - hard
- brittle
- conductivity
- solubility
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.
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.
11
Q
Conductivity of Solid Ionic Compounds.
A
Do not conduct electricity as there is no free moving ions
12
Q
What is a molten compound?
A
Are compounds that are heated into a liquid
13
Q
Conductivity of Molten and Aqueous Compounds in Ionic Bonding
A
There are free moving ions that can therefore conduct electricity
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
15
Q
Monoatomic Ion
A
Are ions containing one atom only
16
Q
Polyatomic Ion
A
Are ions made up of more than one atom
17
Q
If an ion forms more than one ion
A
Must specify the charge when naming
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
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
20
Q
Naming anions
A
Suffix “ide” is added + ion
21
Q
Empirical Formula
A
The simplest whole number ratio of atoms in a compound
22
Q
Chemical Formulae Rules
A
- When naming and writing ionic compounds, cations written first
- Positive and negative charges must be balanced
- Subscripts indicate number of ions
- Electrovalency must be specified
23
Q
Solubility Rules
A
Can be used to predict whether an ionic compound will be soluble or insoluble under standard Laboratory Conditions
24
Q
Standard Laboratory Conditions
A
25 degrees and 1 atmosphere pressure
25
Determining Solubility
Identify anion and determine if soluble then identify if cation is soluble or an insoluble exception.
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Precipitation reaction
A reaction in which a precipitate is formed
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Precipitate
An insoluble solid that forms when two or more solutions are mixed. Solid seperates from the solution.
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Why does a solid form in a precipitate reaction?
The attraction between cations and anions is greater than the attraction between the individual ions and the surrounding water molecules.
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Balanced formula Equation
Shows all reactants and products (including spectator ions)
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Spectator Ions
Ions which remain in solution and are unchanged (not involved).
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Ionic Equation
Only shows reacting species which produce the precipitate.
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What percent of atoms in the earths crust are metals?
25%
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Metals exist in
Element form or as compounds
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Element form
very unreactive eg. Gold and Silver
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Compound form
combined with non-metal elements called ores
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Properties of Metal compounds
- Strong
- Can be light
- Ductile
- Conducts Electricity
- Malleable
- Lustrous
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Metallic bonding model
Arranged in a metal lattice.
Lattices differ in charge on the metal ion and the size of the metal ion.
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What is a lattice?
3-d regular arrangement of particles
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Positive Ions in Metal Lattice
Form regular 3d lattice, fixed positions and closely packed.
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Valence Electrons in Metal Lattice
Free to move, delocalised electrons that belong to the lattice as a whole, sea of electrons
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Electrons in Inner shells of the Metal Lattice
Localised
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Attraction in metal compound
Attraction between positive cations and delocalized electrons creates metallic bonds.
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Conductivity of Electricity - Metal
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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Conductivity of Heat - Metal
When delocalised electrons bump into each other and the positive ions, they transfer electricity
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Lustrous - Metal
Light can be reflected due to the presence of free moving and delocalised electrons
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Malleable and Ductile - Metal
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.
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Malleable definition
Can be hammered into sheets.
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Ductile definition
Can be drawn into a wire
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High melting and boiling temperatures - Metal
Strong forces of attraction between the positive ions and the negative delocalised electrons aka Electrostatic forces.
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The higher the charge (referring to melting and boiling points)
The stronger the attraction and thus the higher melting/boiling point.
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Density
Mass per unit volume
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Density - Metals
Positive ions are closely packed. Depends on the mass of the metal ions, their radius and the way they are packed in the lattice.
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Metal reactions
React with a variety of substances to form ionic compounds.
Involves losing electrons to become a cation (oxidation)
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Predictable reactions of Metal
Those reacting with water, oxygen and acids.
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Single displacement Reactions
Metal + Water, Metal + Acid
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Three reactions
Metal + water, Metal + Oxygen, Metal + acid
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Metal + Water
Metal Hydroxide + Hydrogen Gas
Single Displacement
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Metal + Oxygen
Metal Oxide
Synthesis
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Metal + Acid
Salt (IC) + Hydrogen Gas
Single Displacement
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Reactivity depends on
Ionization energy (atomic radius + core charge)
- The lower the Ionization energy, the easier it is to lose electrons.
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Most reactive group
Group 1 Metals
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Qualitative relative reactivity
amount of bubbles produced or the extent to which metals react.
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Molecule Definition
- 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
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Types of forces in Covalent Compounds
Intramolecular and Intermolecular Forces
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Intramolecular Forces
Strong forces of attraction within the molecule
- Also known as covalent bonds
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Intermolecular Forces
Weak forces of attraction between the molecules
- the forces much weaker than both covalent or ionic bonds
- determines properties.
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Covalent Bonds
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.
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Octet Rule
8 electrons in the outer shell
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Single Covalent Bonds
One pair of electrons being shared
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Double covalent bond
2 pairs of electrons being shared to create a double bond
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Triple covalent bond
3 pairs of electrons being shared to form a triple bond
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Name of the electrons not involved in covalent bonding
Lone pairs or non-bonding pairs
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Valence shell electron pair repulsion theory
assumes electron pairs are located as far away as possible to minimise repulsion
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Molecular formula
provides the actual number of each type of element present in a compound
no information about structure or bonding
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Lewis structure
Illustrated number of valence electrons an atom has
Valence electrons represented by dots
Electrons involved are between the atoms
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Electronegativity
a measure of an atoms ability to attract electrons.
The higher the electronegativity, the greater the attraction.
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Difference in electronegativity
- Valence electrons are not always equally shared between two elements
- If the difference is greater than 0.4, they will not share equally.
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Two charge types of covlent bonds
Polar and non polar.
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Molecule shape
Describes the way in which the atoms are arranged around a central atom in 3d space
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Central Atom
always has the smallest electronegativity and hence smallest number of valence electrons.
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Electron domains
Total number of pairs, repulsion results in them being as far away as possible.
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Linear
Diatomic molecules or 2 atoms coming off a central atom. Central atom has no lone pairs of electrons.
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Diatomic molecule
made of two atoms
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Tetrahedral
four atoms surrounding a central atom.
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Pyramidal
Three atoms surrounding a central atom + one lone pair of electrons coming off the central atom.
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Bent
Two atoms surrounding a central atom + 2 lone pairs of electrons coming off the central atom.
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Polar Molecule
A polar bond has one or more polar covalent bonds that are arranged asymmetrically.
- result is called a dipole
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Dipole
Two oppositely charged ends or poles.
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Polar bond
- electrons pulled towards more electronegative atom
- creates partially +ve and partially -ve.
- The greater the difference the more polar the bond.
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Non-polar molecule
- No polar covalent bonds
- Has polar bonds but arranged symmetrically
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Dipole- Dipole bond
Permanent dipole-dipole bond forces are an intermolecular, electrostatic attraction between dipoles of polar molecules.
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Types of intermolecular bonds
Hydrogen bonds, dipole-dipole bonds, dispersion bonds
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Hydrogen bonding
- Stronger form of dipole-dipole bonding
- occurs when Hydrogen attracted to Fluorine, Oxygen, Nitrogen
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- Dispersion forces
- 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)
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Properties of covalent molecules
1. Do not conduct electricity in solid or molten state
2. Low melting and boiling points
3. Soft as a solid
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Conductivity of covalent molecules
No charged particles as the molecules are neutral.
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Low melting and boiling points of covalent compounds
- weak intermolecular forces means only a small amount of energy is needed to change the state of a molecular substance.
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Soft as a solid - Covalent compound
Weak intermolecular forces of attraction
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Melting point and boiling point of Non-Polar Covalent compounds
- only have weak dispersion forces
- stronger dispersion forces increase points
- greater number of electrons increase the dispersion forces.
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Melting point definition
Temperature where solid becomes a liquid.
98
Boiling point definition
Temperature in which liquid becomes a gas.
99
Melting and boiling points of Polar Molecules.
- Hydrogen bonds have highest
- Dipole- dipole second highest.
100
Allotrope
Different physical form of the same element
101
Allotropes of Carbon
Diamond
Graphite
Fullerenes - bucky balls, nanotubes and graphemes
102
Diamond
- 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
Other network Lattices
Silicon, silicon carbide, silicon dioxide
104
Graphite
- 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
Why is diamond the hardest substance?
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
Why does silicon dioxide have a high melting point?
- 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
Why can graphite be used as a lubricant?
- Layers held together by weak dispersion forces so they can readily slide past one another.
