Periodicity and Group 2 Flashcards
-What do groups and periods tell us in the periodic table?
Groups = number of outer electrons
Period = the number of electron shells an element has
Draw a periodic table in terms of S, P, D and F blocks
Draw the trend in melting/boiling points in period 3 elements.
Explain the trend in melting point for sodium to aluminium
From sodium to aluminium it increases:
- They are metallic structures, so they consist of cations surrounded by delocalised electrons.
- More protons in the nucleus and more electrons per cation = cation are smaller and increased attraction of cation for the delocalised electrons due to the “sea” being more negatively charged (as you go across there are more outer shell electrons, so therefore more delocalised electrons) and the electrons being closer to the nuclei due to reduced atomic radius= metallic bond gets stronger = melting point increases.
Explain the melting point for Silicon (trend in period 3 elements)
the highest melting point out of period 3 elements:
- macromolecular with a tetrahedral structure
- strong covalent bonds links each of the silicon atoms together = lots of energy needed to break these bonds.
Explain the melting point for phosphorus, sulfur, chlorine and Argon (trend in period 3 elements)
A general dramatic decrease in melting point due to being simple molecular substances with only Van der Waals attractions between the molecules = little energy needed to overcome this
Phosphorus contains P4 molecules. To melt phosphorus you don’t have to break any covalent bonds - just the much weaker van der Waals forces between the molecules.
Sulphur consists of S8 rings of atoms. The molecules are bigger than phosphorus molecules, and so the van der Waals attractions will be stronger, leading to a higher melting and boiling point.
Chlorine, Cl2, is a much smaller molecule with comparatively weak van der Waals attractions, and so chlorine will have a lower melting and boiling point than sulphur or phosphorus.
Argon molecules are just single argon atoms, with electrons closer to nucleus Electrons are not easily polarised since argon is very unreactive
Why does atomic radius decrease across a period?
more protons in the nucleus = no additional shielding because electrons are added to the same shell= attraction for the electrons is greater so the atoms shrink
What is the first ionisation energy trend across a period? Explain
more protons in the nucleus as you go across = no additional shielding, as the electrons are added into the same shell = atoms are slightly smaller = attraction for the outer electron is stronger = outer electron is harder to lose.
Why is there a drop in ionisation energy from one period to the next?
the next period has a new main energy level = increase in atomic radius = outer electron is further from the nucleus= less strongly attracted and easier to remove
What are group 2 metals sometimes called?
the alkaline earth metals
What are the physical properties and trends as you go down group 2 elements (magnesium to barium) in terms of atomic arrangement? Explain them
- all have two electrons in an outer s-orbital and lose them when they react to become 2+ ions
- lower charge density as you go down the group (strong nuclear force from protons not as powerful)
- increases in atomic radius as you go down the group due to an increase in atomic number = more electron shells due to more electrons
- Decreasing first ionisation energy = more shielding so shielded from attraction to nucleus by the repulsive forces of inner-shell electrons. This means less effective nuclear charge do outer electrons are lost more easily.
- Melting point decreases from Ca to Ba (magnesium is an anomaly with no explanation)= bigger ionic radius for positive ions = increased distance between positive ions and delocalised electrons= weaker electrostatic attraction between delocalised electrons and positive ions, therefore the metallic bonds are weaker
- reactivity increases due to increased atomic radius = easier to lose outer electrons
Describe the reactions of group 2 metals with water. Write the general equation
X(s) + H2O → X(OH)2 (s) + H2 g)
metal + water = metal hydroxide + hydrogen
Mg = reacts very slowly with cold water
Ca = react with cold water more readily than Mg , effervescence and a white precipitate forms, GRANULES SINK AND RISE
Ba = vigorous reaction with effervescence
- As you go down the group: fizzing becomes more vigorous, the metal dissolves faster, the solution heats up more, less precipitate forms
Describe the reaction of Magnesium with steam
- reacts rapidly to produce MgO and H2(g)
- burns with a bright white light
- MgO in the form of a white powder
Describe the solubility of group 2 metal hydroxides
Mg(OH)2 = sparingly soluble
Ca(OH)2 = soluble (forms limewater)
Sr(OH)2 = soluble
Ba(OH)2 = soluble
Why is magnesium hydroxide solution the least alkaline? Which group 2 metal hydroxide solution is the most alkaline?
It has the least amount of hydroxide ions dissolved.
Barium hydroxide
Write the chemical equation for the solubility of magnesium hydroxides and the solubility of barium hydroxides
Write the ionic equation for the test for sulfate ions
What is the full equation and the ionic equation for the formation of a precipitate when group 2 metals react with, for example, sodium sulfate
What is the equation for the extraction of titanium using magnesium (including the chlorine and carbon bit)
What are the chemical equations for when calcium oxide or calcium carbonate neutralises sulfur dioxide
