Flashcards
covalent bonding
Flashcards do conjunto (61)
- 1A covalent bondis a shared pair of electrons one from each donor atom
- 2Covalent bondingis the electrostatic attraction between the nuclei and the bonding pair of electrons which holds the 2 nuclei together.
- 3The Bond lengthis the distance between the nuclei of the two atoms that are covalently bonded together.
- 4The electrostatic forces of attraction and repulsion within a hydrogen molecule are depicted:-The electrostatic forces of attraction and repulsion within a hydrogen molecule are depicted:- NOTE: Electron density anywhere between these 2 lines
will hold the 2 nuclei together making a covalent bond. - 5As a general rule the shorter the bond length thestronger the bond
- 6What advantage has each chlorine atom gained from forming covalent bond?Each chlorine atom gains a full outer shell of electrons
- 7A double covalent bond isone in which two pairs of electrons are shared between two atoms.
- 8A triplecovalent bond is one in which three pairs of electrons are shared between two atoms.
- 9Look again at the dot and cross diagram of an ammonia molecule, NH3.There are 5 electrons in the outer shell of the nitrogen atom.
Three unpaired electrons form normal covalent bonds, one with each hydrogen atom.
There is a pair of non-bonding (un-bonded) electrons remaining known as a lone pair.
A hydrogen ion, H+, has an empty electron shell. It will bond with the lone pair of the NH3 molecule to form a dative covalent bond, represented by an arrow from the donating pair of electrons to the atom which accepts them (see diagram below). - 10A dative covalent bondis a shared pair of electrons both from one donor atom
- 11The dative covalent bond will be indistinguishable from the normal covalent bonds. It will:have the same strength (bond energy)
be the same length. - 12MOLECULAR LATTICESMost covalent structure are molecules which are held together with strong bonds within the molecule but weak forces between the molecules examples include CO2, CH4, H2O, at room temperature most of these are liquids or gases
However compounds can also form molecular solids, especially if the temperature is low enough . The two examples are ice and iodine - 13Iodine Iodine is a diatomic molecule held together by weak Intermolecular forces. In the solid form the molecules are held in a regular lattice. What happens to the structure as it is heated?As the temperature increases the molecules gain
kinetic energy and vibrate more, enough to overcome the weak intermolecular forces between the molecules (the strong covalent bonds remain unchanged). - 14IceAnother example of a molecular solid is ice, again held together by weak intermolecular forces. (you will study these in a later topic). Other molecular solids are Sulphur (S8), phosphorus (P4) C60 etc. Note that they have a specified number of atoms in the molecule. This differentiates them from Giant covalent lattices.
- 15When melting a molecular lattice it is theweak intermolecular forces that are overcome NOT the strong covalent bonds within the molecule.
- 16Allotropes –two or more forms of the same element in which the atoms or molecules are arranged in different ways
- 17DIAMOND -a giant, covalently bonded structure.
In a diamond, each carbon atom is bonded to four other carbon atoms oriented around it in a tetrahedral manner to form a giant molecule. The bonds are covalent and because the bonding electrons are localised close to the nucleus, they are exceptionally strong. Diamond is the hardest natural substance known. - 18In terms of structure and bonding, explain the differences between diamond and iodine.
- 19Why is there such a difference in the M.Pt. of iodine and diamond?Diamond : giant covalent structure need to break many covalent bonds in 3D
Covalent bonds need to be broken
I2 molecule: only has weak attractions (London forces) between molecules
Little energy is needed to overcome these weak forces of attraction. - 20Silicon is also in group 4 and the element shows the same structure as diamond. Why does silicon have a very high melting point?Si has a giant covalent structure and a lot of heat energy is needed to break the 4 covalent bonds between each Si atom
It has a lower M.Pt. than diamond as the atoms have a larger radius and so the covalent bonds will be longer and weaker - 21Silicon dioxide (sand) If silicon forms four covalent bonds and oxygen forms two bonds how can a giant structure be formed?Each Si atom makes single covalent bonds to four O atoms
Each O atom makes a single covalent bond to two Si atoms - 22Diagram of SiO2 structure
- 23SiO2 structure Physical propertieshard
High M Pt.
Does not conduct electricity
Insoluble - 24Giant covalent and simple covalent structures- similarities and differencesAs we have seen covalent compounds can be described as either simple molecular or giant atomic/giant covalent.
Examples of simple covalent are water, ammonia, methane, carbon dioxide etc.
Examples of giant atomic (or giant covalent) are sand (SiO2), graphite and diamond
They have some quite different physical properties but one similarity.
Fill in the table below to illustrate the differences and similarities between giant atomic and simple
covalent structures using carbon dioxide and silicon dioxide as typical examples. - 25covalent structures using carbon dioxide and silicon dioxide as typical examples.
- 26GRAPHITEGraphite is one of three crystalline forms (allotropes) of carbon. It is an unusual non-metal as it is able to conduct both heat and electricity. The reason lies in its structure. It is composed of carbon atoms in rings of six, bonded together and formed into sheets. (The sheets are then stacked in a layer structure).
- 27A carbon atom has four electrons in its outer shell. Looking at the diagram, to how many other carbon atoms is each carbon atom bonded?3 others
- 28Each of these bonds is astrong covalent bond.
- 29This means that each carbon atom has a valence. electron free which is not localised. These electrons are mobile and move freely across the plane of the sheet. They are said to be delocalised as they form adelocalised π molecular orbital.
Between each layer are weak intermolecular forces.
(called London forces) - 30Between each layer are weakintermolecular forces.
(called London forces) - 31Explain why graphite can conduct heat and electricity.Delocalised electrons are free to move ALONG a layer (not between layers)
When a p.d.is applied they move to the +ve terminal conducting electricity.
In a hot region they move faster colliding with other electrons so transferring KE - 32In which direction can graphite conduct and why?Graphite only conducts ALONG sheets, electrons are only moving within sheets,
adjacent sheets/ individual planes are too far apart. - 33Explain why graphite is a good lubricant or useful in lead pencils?Weak London forces allow layers to slip leaving a thin layer of graphite on the paper.
- 34GRAPHENEThe 2010 Nobel prize for physics was awarded for the discovery of graphene another allotrope of carbon, similar in structure to graphite it can be thought of as a single sheet, one atom thick of graphite.
It can then be rolled into a ball to form a
Fullerene or into a cylinder to form a carbon
Nanotubes. Like graphite it forms 3 covalent
bonds with the last electron delocalised over the surface. Hence it can conduct electricity.
Like graphite and diamond graphene is - 35Why is Graphene a giant covalent lattice and C60 a molecular solid.C60 has a fixed number of atoms in the molecule, whereas graphene (as well as graphite and diamond) have a variable number of atoms determined by the size of the crystal. Similar to giant ionic structures e.g.NaCl in ionic bonding.
- 36Diamond -a giant, covalently bonded structure.
In a diamond, each carbon atom is bonded to four other carbon atoms oriented around it in a tetrahedral manner to form a giant structure. The bonds are covalent and because the bonding electrons are localised close to the nucleus, they are exceptionally strong. Diamond is the hardest natural substance known. - 37Diamond -
- 38Graphite
- 39diamond vs graphite.
- 40METALLIC BONDINGenergies. This means that metal atoms have a high tendency to lose electrons and form positive ions.
The properties of metals depend on the way the atoms are packed (structure) and the bonding of the atoms making up the metal structure.
Most metals have fairly high melting points suggesting that the attractive forces between cations and delocalised electrons are strong. Metallic atoms readily lose their outer shell electrons which move randomly throughout the crystal structure, are shared by all the neighbouring positive ions and are said to be delocalised. - 41Metals of group 1 loseone electron.
- 42Metals of group 2 losetwo electrons
- 43Metals of group 3 losethree electrons
- 44Metallic bonding isthe electrostatic attractions between positive metal cations and the delocalised electrons.
- 45The strength of the metallic bond depends on:the number of electrons which can become delocalised
the size of the ions - 46The …........... the number of electrons delocalised the stronger the bonding. The …....... the size of the ions the stronger the bondinggreater
smaller - 47State and explain which of the following will have the strongest metallic bonding. Sodium or potassium?Na – smaller ionic radius, same no of delocalised electrons therefore greater electrostatic attraction and more energy required
- 48State and explain which of the following will have the strongest metallic bonding. Sodium or magnesium?Magnesium ion/ Mg2+ is smaller (than sodium ions) (1)
The magnesium ion contributes two electrons (to the sea of electrons) compared to sodium’s one. (1)
More energy/heat required to overcome (attractive) forces/bonds (between cations and “sea” of electrons) in magnesium (compared to sodium) - 49State and explain which of the following will have the strongest metallic bonding. Magnesium or aluminiumAl – smaller ion, 3e- delocalised.
- 50Explain why the delocalised electrons in metals make them:- Good electrical conductors:delocalised electrons are free to move in any direction
when a potential difference is applied the electrons move to the positive terminal
a flow of electrons is an electric current. - 51Explain why the delocalised electrons in metals make them:- Good conductors of heat:delocalised electrons are free to move in any direction
in hot areas electrons have more K.E., move faster
electrons collide more, passing on K E to other electrons. - 52Explain why the delocalised electrons in metals make them:- Explain why the electrical conductivity increases Na → Mg →AlMore electrons are delocalised Na(1) → Mg(2) → Al(3) are free to move in any direction.
when a p.d. is applied the electrons move to the positive terminal
larger flow of electrons results in a larger current and so greater conductivity. - 53Explain why the delocalised electrons in metals make them: Explain why the thermal conductivity increases Na → Mg →AlMore electrons are delocalised Na(1) → Mg(2) → Al(3)
in hot areas electrons have more K.E.
more electrons collide more passing on K E to other electrons more efficiently. - 54The stronger the metallic bond the higher the melting point. Most metals are solids. Which one is liquid at room temperature?mercury
- 55Which other two have the lowest melting points?Caesium, francium
- 56Sort the following into metallic, covalent or ionic Sodium chloride, Magnesium, Magnesium oxide, Methane, Silicon dioxide, Bromine, Aluminium, Oxygen, Graphite, Barium iodide, Water, Ammonia, (now add some of your own)IONIC (giant ionic lattices)
Sodium chloride - conducts electricity when molten or (aq)
Magnesium oxide - conducts electricity when molten or (aq)
Barium Iodide - conducts electricity when molten or (aq)
METALLIC (giant metallic lattices)
Magnesium – conducts when molten or solid
Aluminium - conducts when molten or solid
COVALENT
Giant Covalent
Silicon Dioxide - does not conduct
Graphite - conducts due to delocalised electrons in layers between hexagonal sheets
Simple Covalent (molecular)
Methane
Bromine
Oxygen
Water
Ammonia - 57noteNone of the simple covalent compounds are said to conduct electricity. Pure water will not conduct electricity, however most water contains dissolved salts which make it highly conductive.
Ammonium chloride: Simple covalent within the Ammonium ion (NH4+) ionic between the chloride and ammonium
Potassium dichromate: Simple covalent within the dichromate ion, ionic between the potassium and the dichromate - 58metal properties
- 59covalent bonds may be polarised by differences in electronegativity
- 60use the pauling scale to work out the percentage ionic character
- 61—
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- A level Chemistry FlashCards sobre covalent bonding , criado por megan trent em 19-03-2019.
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