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75 Cards in this Set
- Front
- Back
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What is an ion and what are the two types of ions? |
An ion is a charged particle, formed when an atom or group of atoms has lost or gained on or more electrons. A positively charged ion is a cation A negatively charged ion is a anion |
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When ions in the same period have the same electron configuration what are they known as? |
Isoelectronic |
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What process can be used as evidence for ions? |
Electrolysis |
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What is ionic bonding? |
The strong electrostatic force of attraction between oppositely charged ions |
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Describe a giant ionic lattice |
The structure of solid ionic compounds comprising oppositely charge ions arranged in a highly ordered way |
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Where does ionic attraction occur? |
Between oppositely charged ions |
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Where does ionic repulsion occur? |
Between ions with the same charge (may also be between electrons, or between nuclei) |
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Describe the trend in ionic radius down a group |
The ionic radius decreases down the group since the ions have more occupied shells (more shielding/screening). Therefore there is greater repulsion between electrons and the ionic radius is larger. |
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What happens the ionic radius as the atomic number increases in isoelectronic ions? |
As the atomic number increases, the nuclear charge increases. The attraction between the nucleus and electrons increases, so the ionic radius decreases |
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What is charge density |
A measure of the charge/size ratio. In general, ions with a high charge density form stronger bonds than those with a low charge density |
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Why is the ionic bond between F- ions and Mg2+ ions stronger than between F- ions and Li+ ions? |
The sizes of lithium ions and magnesium ions are similar, but magnesium ions have a higher charge. This means that magnesium ions have a higher charge density and are more strongly attracted to fluoride ions. |
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How do ionic charge and ionic radius effect the strength of ionic bonding? |
An increase in ionic charge and a decrease in ionic radius cause an increase in bond strength. |
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Which of the ions S2-, Cl-, K+, Ca2+ has the largest ionic radius and why is this. |
S2-. The ions are isoelectronic and S2- has the lowest atomic number meaning it has the smallest nuclear charge. Therefore, S2- has the weakest attraction between its nucleus and its electrons to has the largest ionic radius |
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What is a covalent bond? |
The strong electrostatic attraction between two nuclei and the shared pair of electrons between them. |
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Describe three ways in which a covalent bond can be displayed |
1. As a straight line between two atoms e.g. H-H 2. A dot and cross in overlapping circles 3. A dot and cross between two symbols without the circles |
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How can a sigma bond be formed by overlapping different types of orbitals? |
A sigma bond can be formed by: 1. Overlapping two s orbitals 2. Overlapping two p orbitals horizontally 3. Overlapping an s and a p orbital horizontally |
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How can a pi bond be formed by overlapping different types of orbitals? |
By overlapping two p orbitals in the vertical direction so that they overlap in two places |
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What type of bond is formed when an orbital with a lone pair of electrons in one atom overlaps with a vacant orbital in the other atom? |
A dative covalent bond |
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How can a dative covalent bond be represented? |
1. By an arrow pointing from the atom that provides the electron pair towards the atom with the vacant orbital 2. As a dot and cross diagram, where two dots/crosses from the donating atom are in the place between the donating atom and the receiving atom |
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Define bond length |
The distance between the nuclei of two atoms that are covalently bonded together. Bond lengths are measure in nanometres, nm (1nm = 10-9m) |
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Define bond strength |
Bond strength is given by the bond enthalpy for a particular covalent bond |
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Define bond enthalpy |
The enthalpy change when one mole of a bond in the gaseous state is broken. Bond enthalpy is measured in kJmol-1 |
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How does a multiple bond arise? |
When two electron pairs are shared (double bonds) or when three electron pairs are shared (triple bond) |
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How does the bond length effect the bond strength? |
For Cl2, Br2 and I2, as the bond length increases the halogen-halogen bond becomes weaker. This is because the lone pairs of electrons on atoms with shorter bond lengths are close enough to produce a lot of repulsion |
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How does the number of bonds effect the bond length and strength? |
For a carbon-carbon bonds, as the number of bonds decreases, the bond length increases and the bond strength decreases. The electrostatic attraction between the two nuclei and the shared electrons is decreased. |
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What is the name of the theory which allows us to predict shapes of molecules and ions? |
The Valence shell electron pair repulsion (VSEPR) theory |
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By how many degrees, does each lone pair of electrons decrease bond angle? |
Approx. 2.5 degrees |
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Describe the bonding pairs, lone pairs and shape name of bonding in BCl2 |
Bonding pairs: 2 Lone pairs: 0 Shape: Linear |
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Describe the bonding pairs, lone pairs and shape name of bonding in BCl3 |
Bonding pairs: 3 Lone pairs: 0 Shape: Trigonal planar |
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Describe the bonding pairs, lone pairs and shape name of bonding in NH3 |
Bonding pairs: 3 Lone pairs: 1 Shape: Trigonal pyramidal |
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Describe the bonding pairs, lone pairs and shape name of bonding in CH4 |
Bonding pairs: 4 Lone pairs: 0 Shape: Tetrahedral |
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Describe the bonding pairs, lone pairs and shape name of bonding in H2O |
Bonding pairs: 2 Lone pairs: 2 Shape: V-Shaped/Bent |
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Describe the bonding pairs, lone pairs and shape name of bonding in PCl5 |
Bonding pairs: 5 Lone pairs: 0 Shape: Trigonal bipyramidal |
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Describe the bonding pairs, lone pairs and shape name of bonding in SF6 |
Bonding pairs: 6 Lone pairs: 0 Shape: Octahedral |
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Describe the effect of lone pairs on the arrangement of electron pairs around the central atom |
Lone pair- lone pair repulsion > lone pair-bonded pair repulsion > Bonded pair-bonded pair repulsion |
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What is electronegativity? |
The ability of an atom to attract the bonding electrons in a covalent bond |
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What happens to the electronegativity of atoms across a period? |
The electronegativity increases. Fluorine is the most electronegative element |
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What happens to the electronegativity of atoms down a group? |
The electronegativity decreases down the group. |
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How can a polar covalent bond arise? |
A covalent bond is polar if the two bonded atoms have different electronegativities. |
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Why are some molecules non-polar? |
The two atoms in the covalent bond may be the same, in which case there is no difference in electronegativity e.g. H-H or Cl-Cl Or, the atoms in the molecule have an equal spread of electronegativity across the molecule, so the differences in electronegativity balance each other out e.g. CO2 |
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How are polar covalent bonds displayed in a diagram? |
The least electronegative atom has a delta+ charge and the most electronegative atom has a delta- charge. |
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What must polar molecules have? |
1. One or more polar bonds 2. Unbalanced dipoles because of their partial charges or the shape of the molecule (a polar bond has a permanent dipole) |
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How can polar bonds be describe in terms of covalent and ionic bonding? |
They can be said to have a degree of ionic character. Polar molecules with a greater difference in electronegativity have a greater % ionic character. |
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Define intermolecular force |
Any force which acts between molecules |
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Define instantaneous dipole |
Instantaneous dipoles occur between all simple molecules, even non-polar molecules. They occur due to fluctuation of electron density in a molecule. |
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Define instantaneously induced (London) forces |
An induced dipole is caused when a molecule comes close to a molecule with a permanent dipole or an instantaneous dipole. Electrons are repelled by delta-charges and attracted to delta+ charges. London forces are attractive forces between instantaneous dipoles and induced dipoles |
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Describe three features of London forces |
1. They exist between all simple molecules 2. Their strength depends upon the number of electrons in the molecule 3. Their strength also depends upon the number of points of contact between molecules |
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Explain in terms of the intermolecular forces present, why chlorine is a gas at room temperature, bromine is a liquid and iodine is a solid. |
The number of electrons / Mr increases going from chlorine through bromine to iodine. The London forces increase Cl2 < Br2 < I2 Chlorine is a gas because it has the wekest London forces / Iodine is a solid because it has the strongest London forces since it has the most electrons |
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What are permanent dipoles? |
Dipoles which occur between polar molecules i.e. the atoms in the molecule have a fixed difference in electronegativity |
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What are hydrogen bonds? |
Intermolecular forces formed if the compound contains a hydrogen atom covalently bonded to an atom of N, O or F, or if there is a lone pair of electrons on an atom of N, O or F. The bond angle between the three atoms involved in a hydrogen bond is often 180 degrees |
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Explain the differences in boiling points of the hydrogen halides (HF, HCl, HBr, HI) |
The intermolecular forces between HCl, HBr and HI are London forces. Their boiling temperatures increase as the number of electrons per molecule increases. HF has the highest boiling temperature, even though its London forces are the weakest, since it also has hydrogen bonding between molecules. |
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Describe the anomalous properties of water in terms of its intermolecular bonding |
Water has a relatively high melting and boiling point for a molecule with its number of electrons. Its solid (ice) is less dense that its liquid. Each water molecule can form two hydrogen bonds on average, increasing the amount of energy needed to melt or boil water. Ice formed open rings of six water molecules, joined by hydrogen bonds. |
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Ethanol is a liquid at room temperature and propane is a gas, but both have similar numbers of electrons. Explain the difference. |
Both ethanol and propane have London forces, but ethanol contains a hydroxyl group -OH. This allows it to form hydrogen bonds (to other ethanol molecules) More energy is needed to separate ethanol molecules than to separate propane molecules. |
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Define solvent |
A substance in which another substance can dissolve, forming a solution |
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Describe the three conditions necessary for a substance to dissolve |
1. The solute particles must be able to separate from one another 2. The separated solute particles must become surrounded by solvent particles 3. The solute-solvent forces are greater than the solute-solute and solvent-solvent forces |
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Define solute |
The substance that dissolves in the solvent. For example in salty water, salt is the solute and water is the solvent. |
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Define soluble |
A solute is soluble if it can dissolve in a given solvent. |
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Define solution |
The mixture formed between a solute and its solvent. Solutions are clear and can be coloured or colourless. |
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What is hydration |
Hydration occurs when water molecules surround individual ions and form strong electrostatic attractions with them |
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What are ion-dipole interactions? |
Interactions between oppositely charged molecules of ions and the solvent dipole: Na+ ions and the delta- of water molecules, or Cl- ions and delta+ of water molecules |
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What happens to the solubility of an alcohol as the number of carbon atoms in the alcohol molecule increases? |
As the number of carbon atoms in an alcohol molecule increases, the London forces become more important and the alcohol becomes less soluble |
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In terms of hydrogen bonding, suggest why chlorobutane C4H9Cl is less soluble in water than it is in ethanol |
Water forms more hydrogen bonds per molecule than ethanol does. More energy would be needed to break the hydrogen bonds in water than in ethanol, so chlorobutane is less soluble in water than it is in ethanol |
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What kind of lattice is present in solid metal, ionic solids and covalently bonded solids like diamond? |
Giant lattice |
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Define metallic bonding |
The strong electrostatic force of attraction between metal ions and the delocalised electrons |
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How does iodine exist in terms of covalent bonding? |
As diatomic molecules In a solid, iodine molecules are arranged in a simple molecular lattice, attracted to each other by London forces |
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What are the names of the two giant covalent lattices that carbon can form? |
Diamond - a tetrahedral structure where each C atom is joined to 4 others by sigma covalent bonds Graphite - an interlocked structure consisting of hexagonal rings of atoms in layers, where each C atom is joined to 3 others by sigma covalent bonds |
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Graphene is a form of pure carbon. It exists in sheets just one atom thick. What type of crystal structure does graphene have? |
Graphene has a giant covalent lattice structure. |
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Describe the properties of a giant metallic lattice |
Metal lattices consist of metal ions attracted to delocalised electrons (electrons that are not fixed/associated with any one atom). They do not have intermolecular forces. The melting and boiling temperatures are fairly high to high. They are good electrical conductors in solid and liquid form. They are insoluble in water. Metals react with water to produce soluble products. |
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Describe the properties of a giant ionic lattice |
Ionic lattices consist of oppositely charged ions. The melting and boiling temperatures are fairly high to high. They are good electrical conductors in the liquid and aqueous states, but not as solids. They are soluble in water (some are insoluble e.g. AgCl, AgBr, AgI and BaSO4) |
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Describe the properties of a giant covalent lattice |
Giant covalent lattices consist of atoms in a regular arrangement held together by covalent bonds. They do not have intermolecular forces. Their melting and boiling temperatures are high to very high. They generally do not conduct electricity (exceptions include graphite and graphene which have delocalised electrons) and are insoluble in water. |
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Describe the properties of simple molecular substances |
Simple molecules have covalent bonding. They do have intermolecular bonding and have melting and boiling temperatures which are generally low. They are poor electrical conductors and are generally insoluble in water. They may be soluble if they can form hydrogen bonds (e.g. glucose) or if they can react with water (e.g. chlorine) |
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The structure and bonding in the hexagonal form of boron nitride is similar to that of graphite. Predict, with reasons, its melting temperature and electrical conductivity. |
It should have a high melting temperature because of its giant covalent lattice structure with many strong covalent bonds that must be broken. It should conduct electricity when solid because it will have delocalised electrons. |
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Explain why ionic compounds have relatively high melting temperatures |
They have strong electrostatic forces of attraction between their ions. The ions are held in giant ionic lattices so there are many ionic bonds |
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Explain why the boiling temperature of aluminium is higher than the boiling temperature of magnesium |
Aluminium ions have a greater charge than magnesium ions. Aluminium ions are also smaller than magnesium ions. This means that aluminium ions have a greater attraction for the "sea" of delocalised electrons, so more energy is needed to overcome these forces |
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Explain why ammonia has a higher boiling temperature than methane |
Both substances exist as simple molecules with instantaneous dipole-induced dipole forces between them. however, ammonia molecules also have hydrogen bonds. These are stronger and need more energy to overcome them. |
