Untitled Flashcards Set
- Across each period, there is a repeating trend in properties of the elements, called periodicity. The most obvious periodicity in properties is the trend from metals to non-metals. There is a trend in these 4 properties:
o electron configuration
o ionisation energy
o structure
o melting points
Periodic trend in electron configuration:
- The chemistry of each element is determined by its electron configuration, particular the outer, highest energy electron shell.
Trend across a period:
- Each period starts with an electron in a new highest energy shell.
Trend down a group -> Elements in each group also have atoms with the same number of electrons in each sub-shell. This similarity in electron configuration gives elements in the same group their similar chemistry.
Blocks:
- The elements in the periodic table can be divided into blocks. Blocks correspond to their highest energy sub-shell. This gives 4 distinct sub shells.
- The order of sub-shells filling up is the order of increased number of protons.
Names and numbers for groups:
- 1 = alkali metals
- 2 = alkaline earth metals
- 3-12 = transition metals
- 15 = pnictogens
- 16 = chalcogens
- 17 = halogens
- 18/0 = noble gases
What is ionisation energy? Ionisation energy measures how easily an atom loses electrons to form positive ions. The first ionisation energy is required to remove one electron from each atom in one mole of gaseous atoms of an element, to form one mole of gaseous +1 ions.
Na(g) -> Na+(g) + e- first ionisation energy = +496 kJ mol-1
Factors affecting ionisation energies:
Electrons are held in their shells by attraction from the nucleus. The first electron lost will be in the highest energy level and will experience the least amount of energy from the nucleus. For example, the first electron lost from a sodium atom (1s22s22p63s1) is from the 3s sub-shell.
1. Atomic radius
a. The greater the distance between the outer electron and the nucleus, the less the nuclear attraction. The force of attraction falls off sharply with increasing distance, so atomic radius has a large effect.
2. Nuclear Charge
a. The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons.
3. Electron shielding
a. Electrons are negatively charged, and so inner-shell electrons repel outer-shell electrons. This repulsion, called the shielding effect, reduces the attraction between the nucleus and the outer electrons.
Successive ionisation energies:
An element has as many ionisation energies as there are electrons. This is because each electron has an ionisation energy if it were to be taken away.
Example: Helium has 2 electrons, so it has 2 ionisation energies.
He(g) -> He+ + e- first ionisation energy
He+(g) -> He2+ + e- second ionisation energy
The second ionisation energy of Helium is greater than the first ionisation energy. In a helium atom, there are two protons attracting two electrons in the 1s sub-shell.
After the first electron is lost, the single electron is pulled closer to the nucleus. The nuclear attraction on the remaining electron increases and more ionisation energy will be needed to remove this second electron.
2nd ionisation energy definition: the energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.
The number of the ionisation energy = the charge produced!!
- Changes in shell numbers can cause big jumps in the ionisation energy.
- Electrons with the largest ionisation energy require to remove it are the closest to the shells!!
Making predictions from successive ionisation energies:
1. The number of electrons in outer shells
2. The group of the element in the periodic table
3. The identity of an element.
Trends in ionisation energies:
1. A general increase in first ionisation energies across the period.
2. A sharp decrease in first ionisation energy between the end of one period and the start of the next period.
These trends can be explained in terms of atomic radius, electron shielding, and nuclear charge.
Trend in first ionisation energy going down a group ⬇:
- First ionisation energies decrease going down a group.
1. The atomic radius increases.
2. More inner shells so shielding increases.
3. Nuclear attraction to outer shell decreases.
4. First ionisation energy decreases.
Trend in first ionisation energy across a period ➡:
1. Nuclear charge increases
2. Same shell: similar shielding
3. Nuclear attraction increases
4. Atomic radius decreases
5. First ionisation energy increases.
Across the period, the increased nuclear charge is the most important factor for the general increase in first ionisation energy.
Sub-shell trends in first ionisation energy:
- Although there is a general increase about both period 2 and period 3.
- There are 2 drops – both in same place in both periods 2 and 3.
- This suggests that the drop is due to something to do with the groups.
Period 2 drops:
Beryllium to Boron -> this marks the filling of the 2p sub-shell.
The 2p sub-shell in Boron has a higher energy than the 2s sub-shell in beryllium.
We know that higher energy subshells are emptied first.
Therefore, in boron the 2p electron is easier to remove that one of the 2s electrons in beryllium. The first ionisation energy of boron is less that the first ionisation energy of beryllium.
💡Essentially, there is a drop in the ionisation energy because the electron that Beryllium has that Lithium doesn’t is the first electron in the 2p sub-shell.
Comparing nitrogen and oxygen:
- The fall in the first ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p-orbitals of the 2p sub-shell.
o In Nitrogen and Oxygen, highest energy electrons are in the 2p sub-shell.
o In oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom.
o The 2p electrons start to pair
o Paired electrons repel each other -> less energy is required to remove an electron.
Periodic trends in bonding and structure:
- One of the key changes in the periodic table is from metals to non-metals.
- It takes place on a diagonal line from the top of group 13 to the bottom of group 17.
- There are far more metallic elements that non-metallic.
- At room temperature, all metals are solids except mercury. They also have a wide range of properties.
- One constant property of metals is their ability to conduct electricity. This is a remarkable property for a solid – a charge must be able to move within a rigid structure for conduction to take place.
Metallic bonding and structure:
- We already know about ionic and covalent bonding. Metals have a third:
- Metallic bonding:
o In a solid metal structure, each atom donated its negative-outer shell electrons to a shared pool of electrons, which are delocalised (spread out) throughout the whole structure.
o The positive ions (cations) left behind consist of the nucleus and the inner electron shells of the metal atoms.
- Metallic bonding is the strong electrostatic attraction between cations and the delocalised electrons.
- The cations are fixed in position, maintaining the structure and shape of the metal.
- The delocalised electrons are mobile and are able to move throughout the structure – only the electrons move.
Charged metals:
- In non-charged metals, there are an equal number of cations and anions.
- In a 2+ charged metals, there are twice the number of cations compared to electrons.
Properties of metals:
- Strong metallic bonds – attraction between positive ions and delocalised electrons.
- High electrical conductivity
- High melting and boiling points
Electrical conductivity:
- Metals conduct electricity in solid and liquid states – when a voltage is applied through a metal, the delocalised electrons can move through the structure, carrying charge. Contrast this ability with the conductivity of ionic compounds, which have no mobile carriers in their solid state.
Melting and Boiling points:
- Tungsten has the highest melting point of the transition metals.
- Mercury has the lowest
The melting point depends on the strength of the metallic bonds holding together the atoms in the giant metallic structure.
- For most metals, the high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons
- This strong attraction results in most metals having high melting and boiling points.
Solubility:
- Metals do not dissolve. It might be expected that there would be some interaction between polar solvents and the charges in a metallic lattice (as there is with ionic compounds) but any interactions would lead to a reaction rather than dissolving, as with sodium and water.
Giant Covalent Structures:
- Many non-metallic elements exist as simple covalently bonded molecules – F2, H2
- In the solid state, these molecules form a simple molecular lattice structure held together by weak intermolecular forces. These structures therefore have low melting and boiling points.
- The non-metals: boron, carbon, and silicon have very different lattice structures. Instead of small molecules and intermolecular forces, many billions of atoms are held together by a network of strong covalent bonds to form a giant covalent lattice.
o Carbon and Silicon are in group 4 (14) of the periodic table and their atoms have four electrons in the outer shell.
o Carbon (in its diamond form) and silicon use these four electrons to form covalent bonds to other carbon and silicon atoms. The result is a tetrahedral structure
o This diagram shows a tetrahedral arrangement of atoms in the diamond form of carbon – all bond angles are 109.5o because of electron-pair repulsion. The dot-and-cross diagram shows part of the covalently bonded network of carbon atoms.
- Properties:
o High melting and boiling points – large amounts of energy are required to overcome the strong covalent bonds.
o Insoluble – insoluble in almost all solvents – covalent bonds holding the atoms together are far too strong to be broken by interaction with solvents.
o Electrical conductivity – non-conductors of electricity.
§ Only exceptions are graphene and graphite.
§ In carbon (diamond) and silicon, all four outer-shell electrons are involved in covalent bonding, so none are available for conducting electricity.
§ Carbon is special in forming several structures in which one of the electrons is available for conductivity. Graphene and graphite are able to conduct electricity.
Graphene + Graphite:
- Apart from diamond, carbon forms giant covalent structures based on planar hexagonal layers.
- Only 3 out of the 4 electrons on the outer shell are used – the remaining electron is released into a pool of delocalised electrons, shared by all atoms in the structure. Structures of carbon containing planar hexagonal layers are therefore good electrical conductors.
- Graphene and Graphite are both covalent structures of carbon based on planar hexagonal layers with bond angles of 120o (electron pair repulsion).
Graphene:
- Single layer of graphite, composed of hexagonally arranged carbon atoms linked by strong covalent bonds. Graphene has the same electrical conductivity as copper – is the thinnest and strongest material made ever.
Graphite:
- Composed of parallel layers of hexagonally arranged carbon atoms, like a stack of graphene layers.
- The layers are bonded by weak London forces.
- The bonding in hexagonal layers only uses three of the four electrons in the carbon’s outer shells – electricity can be conducted.
Periodic trends in melting points:
- Across Periods 2 & 3.
- The melting point increases between Group 1 – Group 4 (14).
- There is a sharp decrease between Groups 4 – 5.
- The melting points are comparatively low from Group 15 – 18.
- The sharp decrease in melting point marks a change from giant to simple molecular structures.
- You can also start to see the metal -> non-metal transition.
- Giant covalent = strong forces to overcome (high mp)
- Simple molecular = weak forces to overcome (low mp)
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