Topic 3: The periodic table Flashcards
Which way are groups and which was are periods?
Groups are the vertical columns
Periods are the horizontal rows
Roughly where are metals, non metals and metalloids on the periodic table? (imagine)
Yellow: Metals
Pink: Non metals
Green: Metaloids
What are the elements called when they have both properties of both metals and non metals?
Metalloids
What are metalloids?
Elements when they have both properties of both metals and non metals
Which is the most abundant element in the Universe: about 90% of the atoms?
Hydrogen
What are noble gases also known as?
Inert gases
What are the elements that have been omitted called?
Lanthanoid elements and actinoid elements
What does a metallic structure consist of?
consists of a regular lattice of positive ions in a sea of delocalised electrons
What can the metallic and non-metallic properties of elements be related to? What energy?
Ionisation energies
What must the element do to form a metallic structure?
They must be able to lose electrons fairly readily to form positive ions
How does the ionisation energy change across a period and down a group?
Across: ionisation energy increases → so elements lose electrons less easily
- so metallic structures are formed by elements on the left hand side of the periodic table, which have lower ionisation energies
Down: Ionisation energy decrease → therefore elements are much more likely to exhibit metallic behaviour lower down a group
What are the properties a metallic element would generally have?
- Large atomic radii
- Low ionisation energies
- Less exothermic electron affinity values
- Low electronegativity
How are valence electrons presented on the periodic table?
As groups
For groups 13-18, how are valence electrons shown?
group number - 10
ex. elements in group 13 have 3 valence electrons
What does the period number indicate?
The number of shells (main energy levels) in the atoms - or which is the outer shell (main energy level)
What is the periodic table divided into blocks according to?
According to the highest energy sub-shell (sub-level) occupied by electrons.
→ in the s block, all elements have atoms in which the outer shell electron config. is ns1 or ns2 (n is the shell number) and in the p block it is the p subshell that is being filled
What are the physical properties that can be deduced from the periodic table?
- Atomic radius
- Ionic radius
- First ionisation energy
- Electron affinity
- Electronegativity
What does the atomic radius describe?
The size of an atom
The larger the atomic radius, the larger the atom
How does atomic radius change across and down a period?
Across: Decreases across a period, because nuclear charge increases across the period with no significant change in shielding → Meaning that the outer electrons are pulled in more strongly and the radius is smaller
Down: Increase because atoms have increasingly more shells
What does the ionic radius measure?
The size of an ion
In general, are positive/negative ions bigger or smaller than their atomic radii?
In general, the ionic radii of positive ions are smaller than their atomic radii, and the ionic radii of negative ions are grater than their atomic radii
Why does Na have a bigger ionic radius than Na+?
Because Na has one extra shell whereas Na+ lost one.
There is also a greater amount of electron-electron repulsion in Na because there are 11 electrons compared to 10. The electron cloud is therefore larger in Na than in Na+ because there are more electrons repelling for the same nuclear charge pulling the electrons in.
How does ionic radius change across a period?
It is not a clear-cut trend because the type of ion changes going from one side to the other
Positive ions are formed on the left hand side and the negative ion on the right hand side.
Even though on the same period (have same shells), there are just more valence electrons on the right hand side (negative) for there to be a stronger election-electron repulsion
How does the ionic radius change for negative and positive ions?
For positive ions, there is a decrease in ionic radius as the charge on the ion increases
For negative ions, the size increases as the charge increases
Which one has a larger ionic radius? Na+ or Mg2+?
Both ions have the same electron configuration, same number of electrons → same electron-electron repulsion
However there is one more proton in Mg2+ and the higher nuclear charge in Mg2+ means that the electrons are pulled in more strongly and so the ionic radius is smaller
What is the definition of first ionisation energy?
The energy required to remove the outermost electron from a gaseous atom
How does the first ionisation energy change down and across a period?
Down: the first ionisation energy decreases
Across: the general trend is that it increases from left to right due to increase in nuclear charge across the period
Why is the first ionisation energy decreasing down a group?
Because there are more shells, size of the atom increase so the outer electron is further from the nucleus and therefore less strongly attracted by the nucleus
Although the nuclear charge also increases, it is largely balanced out by the increase in shielding down the group.
What governs the change in first ionisation energy?
the increase in size
What is the general trend for first ionisation energy from left to right across a period?
The general trend is that first ionisation energy increases from left to right across a period. This is because of an increase in nuclear charge across the period (protons are added)
Why does the first ionisation energy increase across a period?
the nuclear charge increases as protons are added to the nucleus.
Therefore the attractive force on the outer electrons due to the nucleus increases from left to right across the period.
Explain the exception to the general increase in first ionisation energy across a period
The major difference is that the electron to be removed from the boron atom is 2p sub-level whereas it is in a 2s sub-level in beryllium. The 2p sub-level in B is higher in energy than 2s sub-level in beryllium, therefore less energy is required to remove an electron from boron
What are the two exceptions in first ionisation energies?
Boron vs beryllium
- Main difference is that the electron to be removed in born is in the 2p sub-level, whereas it is in 2s sub-level in beryllium. The sp sub-level in B is higher in energy than the 2s sub-level in beryllium, therefore requiring less energy to remove an electron from boron
Oxygen vs nitrogen
- the first ionisation energy of oxygen is lower than that of nitrogen
- The major difference is that oxygen has two electrons paired up in the same p orbital, but nitrogen does not. An electron in the same p orbital as another electron is easier to remove than one in an orbital by itself because of the repulsion from the other electron.
- When two electrons are in the same p orbital they are closer together than if there is one in each p orbital. If the electrons are closer together, they repel each other more strongly. If there is a greater repulsion, an electron is easier to move
What is electron affinity defined as?
It involves the energy change when one electron is added to a gaseous atom
the enthalpy change when one electron is added to each atom in one mole of gaseous atoms under standard conditions
Why is the first electron affinity is exothermic for virtually all elements?
It is a favourable process to bring an electron from far away to the outer shell of an atom, where it feels the attractive force of the nucleus
How does the electron affinity change down a group and across a period?
Down: electron affinity decrease, becomes less exothermic as the size of the atom increases.
There is a weaker attraction between the added electron and the nucleus if the atom is big (brought to a position which is further from the nucleus)
Across: the electron affinity becomes more exothermic.
This is because of an increase in nuclear charge and a decrease in atomic radius from left to right across the period. → electrons will be more strongly attracted when brought into the outer shell of the atom
How does group 1 elements react with oxygen and water? How does the reaction with water change as it goes down a group
Oxygen: The alkali metals react vigorously and all tarnish rapidly in air
Water: React rapidly. Alkaline solution is formed. The alkali metal hydroxides are strong bases and ionise completely in aqueous solution
Reaction with water becomes more vigorous going down the group
Properties of halogens (F, Cl, Br, I)
M/B point/ colour/ physical state at room temp. and pressure
What do halogens from when reacting with alkali metals?
Salts
- All white/colourless
- fairly typical ionic compounds
What will the results of reactions between halogen solutions and solutions containing halide ions be?
(Cl2, Br2, I2 vs KCl, KBr, KI)
Why would a solution turn orange?
Bromine solution - Due to the production of bromine
Why would a solution turn red-brown
Iodine solution - Due to the production of iodine
What is the colour of a chlorine solution?
Pale yellow-green
In a reaction, what determines which halogen displaces the halide ion of another halogen?
The more reactive halogen displaces the halide ion of the less reactive halogen from the solution.
e.g. chlorine displaces bromide ions and iodide ions from solution, and bromine displaces iodide ions from solution
→ Cl is a stronger oxidising agent than bromine and iodine so it will oxidise bromide ions to bromine and iodide ions to iodine
→ Bromine is a stronger oxidising agent than iodine and will oxidise iodide ions to iodine
→ in terms of electrons, Cl has the strongest affinity for electrons and will remove electrons from bromide ions and iodide ions
What does sodium oxide reacted with water form?
Sodium hydroxide
What do sodium oxides react with to form salts?
Acids such as sulfuric acid
Why is magnesium oxide not very soluble in water?
Because of the relatively high charges on the ions, but it does react to a small extent to form a solution of magnesium hydroxide, which is alkaline
How does aluminium oxide display amphoteric behaviour?
It does not react with water but it does display amphoteric behaviour in that is reacts with both acids and bases to form salts
What do amphoteric oxides react with?
Both acids and bases
What does P4O6 (phosphorus (III) oxide) and P4O10 (phosphorus (V) oxide) form when reacting with water?
Phosphoric (III) and Phosphoric (V) acid
What does SO2 (sulfuric (IV) oxide) and SO3 (sulfur (VI) oxide) form when reacting with water?
Sulfuric (IV) and Sulfuric (VI) acid
What is formed when nitrogen reacts with oxygen at very high temperatures?
to form NO (Nitrogen monoxide, nitric oxide or nitrogen (II) oxide)
Where does this reaction occur?
In the internal combustion engine
What can NO2 form when reacting with water?
forms nitric (V) acid (HNO3), which is one of the acids responsible for acid deposition.
What can nitric oxide be classified as?
An acidic oxide
