Module 2.2 - Electrons, Bonding and Structure Flashcards
How are quantum numbers used to describe the electrons in atoms?
- Principal quantum number, n, indicates the shell the electron is in
- Different shells have different principal quantum numbers
- Larger the value of n, further the shell is from the nucleus + highest energy level
What phrase is ‘shell’ equivalent to?
Energy level
How do you work out the number of electrons the first 4 shells hold?
2n^2
E.g. 2nd shell = 2 x 2^2 = 8
What is the quantum number of the first shell?
1
How many electrons does the first shell hold?
2
What is the quantum number of the second shell?
2
How many electrons does the second shell hold?
8
What is the quantum number of the third shell?
3
How many electrons does the third shell hold?
18
What is the quantum number of the fourth shell?
4
How many electrons does the fourth shell hold?
32
How many electrons can an orbital contain?
2
How many s-orbitals are there in one shell?
1
What is the shape of an s orbital?
Spherical
How many p-orbitals are there in a shell, and therefore how many electrons are there in p-orbitals per electron shell (/quantum number)?
- 3; px, py, pz
- 6
What is the shape of the p-orbital?
Dumbbell shaped, 8/∞
What shell do d-blocks start in?
n=3 (3rd shell)
How many d-orbitals are there in a shell, and therefore how many electrons can it hold?
- 5
- 10
What shell does the f-blocks start in?
n=4
How many f-orbitals are there in each shell?
7 (therefore 14 electrons)
What method is used to show the electrons in orbitals?
‘Electrons in orbitals’
Up and down arrows in boxes
Why do the two electrons in an orbital not repel one another?
- Opposite spins
- Represent the opposite spins by an up and down arrow
What are the 4 sub-shells?
s, p, d, f
How do the types of sub-shell change as the electron shell increases?
One more type is added
Shell 1: s
Shell 2: s and p etc
How do the energy levels of each type of orbital change?
s = lowest
p
d
f
From 1s to 4f, what is the order energy levels of each orbital from lowest to highest?
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f
What are the rules for the arrangement of electrons in an atom?
- Electrons are added, one at a time, to ‘build up’ the atom
- Lowest available energy level is filled first (can consider this level as being closest to the nucleus)
- Each energy level must be filled before the next higher energy level starts to fill
What are the rules of filling up orbitals in the same energy level?
- Each orbital in a sub-shell is filled singly before pairing starts
- 4s orbital is at a slightly lower energy level than the 3d orbital, so fills before it
In an orbital, how do paired electrons move?
With opposite spins
What form is electron configuration written in?
nx^y
n=shell number
x=type of orbital
y=number of electrons in orbitals making up the sub-shell
What are the orbitals occupied and the electron configuration of:
a) Boron
b) Carbon
c) Nitrogen
d) Oxygen
Boron: 1s2 2s2 2px1 /// 1s2 2s2 2p1
Carbon: 1s2 2s2 2px1 2py1 /// 1s2 2s2 2p2
Nitrogen: 1s2 2s2 2px1 2py1 2pz1 /// 1s2 2s2 2p3
Oxygen: 1s2 2s2 2px2 2py1 2pz1 /// 1s2 2s2 2p4
Which electrons are lost to form positive ions?
Electrons in the highest energy level are lost first
What are the most stable and unreactive elements?
Noble gases, group 18. Already have a full outer shell of electrons (other elements react to try to get the same electron configuration as a noble gas)
What are the 3 main types of chemical bonding?
- Metallic
- Ionic
- Covalent
What kind of materials are involved in an ionic bond?
Metal and a non metal (electrons transferred to metal to non metal forming oppositely charged ions)
Using magnesium oxide as an example, show how an ionic bond can be done to give the elements involved the electron configuration of a noble gas.
- Mg forms Mg2+ (1s2 2s2 2p6, same as neon)
- O forms O2- (1s2 2s2 2s6, same as neon)
What materials do covalent bonds form between?
2 non metals
Using hydrogen as an example, show how a covalent bond can be used to give elements the same electron configuration as a noble gas.
-2 H atoms covalently bond + share electrons, giving 1s2 configuration, the same as helium
What kind of materials do metallic bonds form between?
Metals (e.g. Zinc, iron, aluminium, and their alloys such as brass)
In a metallic structure, how many cations are the delocalised electrons shared between?
All of them
How does an ionic bond form?
- Electrons are transferred from the metal to the non metal
- Oppositely charged ions form, bonded together by electrostatic attraction
- Metal ion is positive (cation)
- Non metal ion is negative (anion)
How does an ionic bond form between sodium and oxygen, to form sodium oxide?
- 2 sodium atoms (each with one electron in outer shell) give an electron to an oxygen (6 electrons in its outer shell)
- Forms 2 Na+ ions and an O2- ion
- 2Na –> 2Na + + 2e- (1s2 2s2 2p6 3s1 –> 1s2 2s2 2p6 [Ne])
- O + 2e- –> O2- (1s2 2s2 2p4 –> 1s2 2s2 2p6 [Ne])
Describe the structure of a giant ionic lattice.
- Each ion is surrounded by oppositely charged ions
- Ions attract to each other from all directions, forming a 3D giant ionic lattice
- All ionic compounds exist as giant ionic lattices in the solid state
Describe the giant ionic lattice of sodium chloride.
- Sodium transfers one electron to chlorine, forming Na+ and Cl-
- Forms a giant ionic lattice
- Each Na+ is surrounded by 6 Cl- ions
- Each Cl- is surrounded by 6 Na+ ions
Describe the ionic bonding of calcium oxide.
- Calcium transfers 2 electrons to oxygen, forming Ca2+ and O2-
- Calcium gets the same electron configuration of argon, and oxygen the same as neon
- Ca –> Ca2+ + 2e- (oxidised)
- O + 2e- –> O2- (reduced)
Describe the ionic bond of aluminium fluoride.
- Aluminium loses 3 electrons forming Al3+ (same electron configuration of neon)
- 3 fluorine atoms each gain an electron each, forming 3 x F- (same electron configuration of neon)
- Al –> Al3+ + 3e- (oxidised)
- 3F + 3e- –> 3F- (reduced)
What are the properties of ionic compounds?
- High melting and boiling points
- Conduct electricity (aqueous or molten)
- Soluble
Why do ionic compounds have high melting and boiling points?
- Ionic compounds are solid at room temperature (in a giant ionic lattice)
- Lots of energy is required to break the strong electrostatic forces of attraction between the oppositely charged ions
Explain why sodium chloride’s melting point (801ºC) is lower than that of magnesium oxide (2852ºC).
- Both are ionic compounds
- Charges on the ions in MgO (Mg2+ and O2-) are greater than the charge of the ions in NaCl (Na+ and Cl-)
- Greater charge = stronger electrostatic forces of attraction between the ions so more energy is required to break up the ionic lattice
Explain the electrical conductivity of a solid ionic lattice.
- Ions are held in a fixed position and no ions can move
- Therefore can not conduct electricity
Explain the electrical conductivity of molten or aqueous ionic compounds.
- Solid lattice breaks down and the ions are free to move
- Therefore can conduct elecricity
What kind of solvents do ionic compounds dissolve in?
Polar solvents
Explain how ionic lattices are soluble in polar solvents.
- Polar molecules break down an ionic lattice by surrounding each ion to form a solution
- Slight charges on the polar substance are able to attract charged ions from the giant ionic lattice
- Lattice is disrupted and ions are pulled out of it
Explain sodium chloride’s solubility in water.
- Water molecules attract the Na+ and Cl- ions
- Ionic lattice breaks down as it dissolves and water molecules surround the ions
- Na+ attracts delta- charges on the O atoms of the water molecules
- Cl- attracts delta+ charges on the H atoms of the water molecules
Describe a covalent bond.
- Negatively charged shared pair of electrons are attracted to the positive nuclear charge of both the nuclei
- This attraction overcomes the repulsion between the 2 positively charged nuclei
- Resulting attraction is the covalent bond that holds the atoms together
- The 2 electrons are shared
What is the name for a covalent bond with only one shared pair of electrons?
Single covalent bond
Describe the single covalent bond in hydrogen.
-Each hydrogen atom has 1 electron in its outer shell
-Each hydrogen atom contributes 1 electron to the covalent bond
H-H
-Each hydrogen fills its 1s sub-shell, achieving the electron configuration of helium
How are single covalent bonds sometimes drawn?
By a single line e.g. H-H
Explain the covalent bonding in an oxygen molecule.
-Share 2 pairs of electrons in a double covalent bond
O=O
Explain the covalent bond of nitrogen.
-Share 3 pairs of electrons to form a triple bond
_
N=N
Explain the covalent bond of carbon dioxide.
-2 double bonds, each between the carbon atom and an oxygen atom
O=C=O
What is average bond enthalpy measured in?
kJ mol^-1
What is a dative covalent bond also known as?
Coordinate bond
What is a dative covalent bond?
A covalent bond in which one of the atoms supplies both of the shared electrons to the covalent bond
How is a dative covalent bond written?
A–>B
Arrow shows direction in which the pairs are being donated (e.g. A is donating a pair to B)
Describe the bonding of an ammonium ion.
- Has 3 covalent bonds and one dative covalent bond
- Formed from ammonia, NH3, and H+
- One of the electron pairs around the nitrogen in ammonia is a lone pair
In an ammonium ion, how can you tell which bond is the dative covalent bond?
You can’t, once formed a dative covalent bond is equivalent to all other covalent bonds
Explain the bonding in an oxonium ion., H3O+.
- Forms when an acid is added to water. Responsible for reactions of acids (in equations simplified to H+)
- One of the lone pairs around the oxygen in H2O provides both electrons to form a dative covalent bond with a H+ ion
Why can covalent bonds result in atoms that don’t follow the Octet Rule (noble gas electron configuration)?
- May not be enough electrons to reach an octet
- More than 4 electrons may pair up in the bonding (expansion of the octet)
Give one example of a covalent bond where there aren’t enough electrons to reach an octet.
- Boron trifluoride (BF3)
- Boron has 3 electrons in outer shell + each fluorine atom has 7
- 3 covalent bonds can form
- Each of boron’s 3 outer electrons is paired so there’s now 6 in its outer shell
- Each of the 3 fluorine atoms now have 8 outer shell electrons, attaining the octet
- Central boron atom does not achieve the octet
In what elements does expansion of the octet usually occur?
Groups 15-17 from period 3
(As you go down the period table, there’s more outer shell electrons able to take part in bonding due to distance)
One of the atoms may have more than 8 electrons in the outer shell, breaking the Octet Rule
Describe how expansion of the octet may occur in groups 15, 16 and 17.
- 15: can form 3 or 5 covalent bonds, depending on how many electrons used in bonding
- 16: can form 2, 4 or 6 covalent bonds, depending on how many electrons used in bonding
- 17: can form 1, 3, 5 or 7 covalent bonds, depending on how many electrons are used in bonding
Describe how expansion of the octet occurs in sulphur hexafluroide, SF6.
- S has 6 electrons in its outer shell
- 6 covalent bonds can form
- Each of S’s 6 electrons is paired, so sulphur has 12 outer shell electrons (expanded the octet)
- Each of the 6 fluorine atoms has eight electrons in its outer shell, attaining the octet
What are the 2 types of covalent structure?
- Simple molecular structure
- Giant molecular structure
Describe the bonding in a simple molecular covalent structure.
- Atoms within each molecule are held together by strong covalent bonds
- Different molecules are held together by weak intermolecular forces (London forces)
Describe the simple molecular covalent structure of iodine.
- Within each I2 molecule, I atoms are held together by strong covalent bonds
- When solid I2 (simple molecular lattice) changes state, weak intermolecular forces between the I2 molecules break
What are the properties of a simple molecular covalent structure.
- Low melting and boiling points
- Don’t conduct electricity
- Soluble in polar substances
Why do simple molecular covalent structures have low melting and boiling points?
Weak intermolecular forces so only a relatively small amount of energy needed to break them
Why can’t simple molecular covalent structures conduct electricity?
There are no charged particles that can move
Why are simple molecular covalent structures soluble in polar molecules?
- Weak London forces can form between covalent molecules and these solvents
- This helps the lattice break down and the substance to dissolve
Give some examples of a giant covalent structure.
- Diamond
- Graphite
- SiO2
What are the properties of giant covalent structures?
- High melting and boiling points
- Don’t conduct electricity
- Insoluble in polar and non polar solvents
Why do giant covalent structures have high melting and boiling points?
High temperatures are needed to break the strong covalent bonds between the atoms
Why don’t giant covalent structures conduct electricity?
There are no free charged particles, except for graphite
Why are giant covalent structures insoluble in polar and non polar substances?
Covalent bonds in the lattice are too strong to be broken by either polar and non polar solvents
Describe the electron repulsion theory.
- All electrons have a negative charge, so electron pairs repel other electron pairs
- Shape of the molecule is the one that allows the electrons to be as far away from each other as possible
What is the bonding region?
Where a bond, e.g. double covalent bond, forms
Give the number of bonded electron pairs around the central atom of a linear structure, and explain the bond angle using an example.
- 1 or 2
- H2, H-H (no bond angle)
- CO2, O=C=O 180º
Give the number of bonded electron pairs around the central atom of a trigonal planar molecule, and explain the bond angle using and example.
- 3
- BF3, 120º
Give the number of bonded electron pairs around the central atom of a tetrahedral shape, and explain the bond angle using an example.
- 4
- CH4, 109.5º
Give the number of bonded electron pairs around the central atom of a trigonal bipyramid molecule, and explain the bond angle using an example.
- 5
- PCl5, 90º and 120º**
Give the number of bonded electron pairs around the central atom of an octahedral molecule, and explain the bond angle using an example.
- 6
- SF6, 90º
Describe the order of repulsiveness between 2 different types of paired electrons, from most repulsive to least repulsive.
Lone pair/lone pair –> lone pair/bonded pair –> bonded pair/bonded pair
How much does a lone pair reduce the bond angle by?
Around 2.5º
Describe the shape and bond angle of a methane molecule.
- CH4
- Tetrahedral
- 109.5º
Describe the shape and bond angle of an ammonia molecule.
- NH3
- Pyramidal
- 107º (one lone pair)
Describe the shape and bond angle of water.
- H2O
- Non-linear
- 104.5º (2 lone pairs)
Describe the shape and bond angle of an ammonia ion.
- NH4+
- 4 electron pairs around the central N atom
- Tetrahedral
- -Charge will be distributed across the whole molecule (non-polar as symmetrical) shown by the square brackets with the + sign
- 109.5º
Where does electronegativity increase towards of the periodic table?
Top right (fluorine has the most electronegative atoms)
Why is a hydrogen molecule, H2, non-polar?
- The 2 bonding atoms have the same electronegativities
- Nucleus of each atom is equally attracted to the bonding electron
- Electrons are evenly distributed between the atoms
Why is hydrogen chloride, HCl, a polar molecule?
- Non symmetrical
- Cl is more electronegative than H
- Cl has a greater attraction on the bonding pair of electrons than the H atom
- Bonding atoms are held closer to the Cl atom than to the H atom
- H has a δ+ charge (small positive charge)
- Cl has a δ- charge (small positive charge)
Why is tetrachloromethane, CCl4, non-polar?
- Symmetrical so the dipoles cancel out
- Each CCl bond is polar
Why is carbon dioxide non polar?
Although the C=O bonds are polar bonds (as oxygen is more electronegative than carbon), the molecule is symmetrical (due to 180º bond angle) so the dipoles cancel out
Why is water a polar molecule?
The H-O bonds are polar bonds (as oxygen is more electronegative than hydrogen) and the molecule is not symmetrical therefore the dipoles do not cancel each other out
What are the 2 main types of intermolecular forces?
- Hydrogen bonding
- London forces
Order all of the types of bonds by their relative strength, from strongest to weakest.
- Ionic and covalent (relative strength 1000)
- Hydrogen bonds (relative strength 50)
- Permanent dipole-dipole forces (relative strength 10)
- London dispersion forces (relative strength 1)
What is a permanent dipole-dipole force?
Any intermolecular force containing permanent dipoles that aren’t hydrogen bonds
What is a permanent dipole-induced dipole interaction and how do they form?
- Molecules with a polar bond (permanent dipoles) bond to induced dipoles from a non polar bond
- When a polar molecule goes towards a non polar molecule, the δ charge causes the electrons on the non polar molecule to move towards the polar molecule (for δ+) or away from the polar molecule (for δ-)
- This causes the non-polar molecule to become slightly polar
What is a permanent dipole-permanent dipole interaction and how do they form?
-Molecules with permanent dipoles are attracted to other molecules with permanent dipoles
-Oppositely charged dipoles of different molecules attract to one another
E.g. HCl: H-Cl——–H-Cl (as Cl is δ- and H is δ+)
What are the 2 types of permanent dipole-dipole interactions?
- Permanent dipole-permanent dipole interactions
- Permanent dipole-induced dipole interactions
What is the intermolecular force between two non polar molecules?
London dispersion forces
How do London dispersion forces form?
- Electrons move constantly and randomly in atom’s shells. This movement unbalances the distribution of charge within the electron shells (like the density in the shells ‘wobbling’ from side to side)
- At any moment, there’ll be an instantaneous dipole across the molecule
- Instantaneous dipole induces a dipole in a neighbouring molecule, in turn inducing further dipoles on their neighbouring molecules
- Small induced dipoles attract one another, causing a weak intermolecular force (London dispersion forces or instantaneous dipole-dipole forces)
What affects the strength of the London dispersion forces?
Number of electrons of the atoms (as creates a larger induced dipole so grater attractive forces between molecules)
What is the boiling points of non polar molecules?
Low as the weak London forces are all that need to be broken which does not require much energy
What atoms does a hydrogen bond form between?
Hydrogen and the lone pair of a fluorine, nitrogen or oxygen
What anomalous properties of water arise as a result of the hydrogen bond?
- Ice floats on water
- Higher melting and boiling point than expected
- High surface tension of water
Why does ice float on water?
- When ice forms, water molecules arrange into an orderly pattern and hydrogen bonds form between the molecules (hydrogen bonds occur in a liquid but not as often as molecules move past each other and therefore overcome these bonds)
- Ice has an open lattice with hydrogen bonds holding water molecules apart
- When ice melts, the rigid hydrogen bonds collapse, allowing the H2O molecules to move closer together as hydrogen bonds are long
- So ice is less dense than water, hence why it floats
Why does water have a higher melting and boiling point than expected?
-The hydrogen bonds are much stronger than other intermolecular forces
-Extra strength of these bonds has to be overcome to melt/boil H2O so has a higher melting/boiling point than if there weren’t any hydrogen bonds
(-Other group 16 elements have the same structure as water but don’t have hydrogen bonds and therefore have a lower melting/boiling point)
Why does water have a higher surface tension and viscosity than expected?
- Hydrogen bonds
- Allows insects to walk along water (walking across a raft of hydrogen bonds)
