Module 3.1 - The Periodic Table Flashcards
Describe the relationship in elements in the same period.
-Trends are repeated for each (periodicity)
-Atomic number increases
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What is the relationship between elements in the same group?
Same physical and chemical properties (as same outer electrons + type of orbitals). Repeating pattern of similarity caused by underlying repeating pattern in electron configuration
Put each orbital in order from lowest energy to highest.
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Put each orbital in order of ascending energy up to 4d.
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d
The 4s orbital has a lower energy than the 3d orbital. What does this mean?
- 4s fills before 3d
- Empties before 3d during ionisation
How can you shorten electron configuration?
Use the noble gas before the atom and add any extra orbitals
What is ionisation?
When atoms gain or lose electrons to form ions
What is the equation for the first ionisation of sodium?
Na(g) –> Na+(g) + e-
What factors affect ionisation energy?
- Atomic radius
- Nuclear charge
- Electron shielding or screening
How does atomic radium affect the ionisation energy?
- Larger atomic radius = smaller nuclear attraction to outer electrons
- Positive nuclear charge is further away from negative electrons
How does nuclear charge affect the ionisation energy?
-Higher nuclear charge = stronger attraction
How does the electron shielding affect ionisation energy?
- Inner shells of electrons repel outer electrons as all negative so easier to lose
- Repelling effect is called electron shielding/screening
- More inner shells there are, larger the shielding effect + smaller nuclear attraction experiences by outer electrons
What is successive ionisation energy?
Measure of the amount of energy required to remove each electron in turn
Why is each successive ionisation energy higher than the one before?
- As each electron is removed, there’s less repulsion between remaining electrons + each shell will be drawn slightly closer to nucleus
- +ve nuclear charge outweighs -ve charge every time an electron in removed
- As distance of each electron from nucleus decreases slightly, nuclear attraction increases. More energy required to remove each successive electron
Which element in each period has the highest ionisation energies?
-Noble gas as they have a full outer shell of electrons + high positive attraction from nucleus
What trends are there across each period?
- Number of protons in nucleus increases so higher attraction to electrons
- Atomic radius decreases (as more electrons so more attraction)
- Shielding is around the same as same inner shells
Why is there a slight decrease in first ionisation energy between groups 2 and 13?
- Group 2’s highest energy electron: s orbital
- Group 13’s highest energy electron: p orbital
- P orbitals have a little more energy so are slightly further away from nucleus so electrons are slightly easier to remove
Why is there a slight decrease in first ionisation energy between groups 15 and 16?
-Between groups 13 and 15 each electron gets its own p-orbital but from 16 they start pairing in each orbital so slightly easier to remove (as slight repulsion) hence slightly lower first ionisation energy
Why is there a sharp drop in ionisation energy between the noble gas of a period and the group 1 element of the next?
- Increased atomic radius as another shell so further away from nucleus
- Increase in electron shielding of outermost shell by inner electrons
Why does the first ionisation energy decrease as you go down a group?
- More shells so greater distance between outer electron + nucleus so weaker attraction to outer electrons
- More inner shells so shielding effect on outer electrons increases so weaker attraction
- Increase in shielding outweighs the increase in nuclear charge
What are the trends as you go down a group?
- Number of shells increases
- Shielding increases
- Atomic radius increases
- First ionisation energy decreases
Describe the structure of metallic lattices.
- Delocalised electrons spread through whole structure
- Electrons can move within the structure
- Impossible to tell which electron originated from which cation
- Charges must balance over the whole structure
- Cations in fixed positions in the lattice
- Outer shell electrons are delocalised
What are the properties of giant metallic lattices?
- High melting + boiling points
- Good electrical conductor
- Malleability and ductility
Why do giant metallic lattices have high melting and boiling points?
- Electrons free to move throughout structure but cations remain where they are
- Attraction between cations + delocalised electrons is very strong
- High temperature needed to overcome metallic bonds to dislodge cations from rigid positions in lattice
Why do metallic lattices conduct electricity?
- Delocalised electrons can move anywhere in the metallic lattice
- Allows solid to conduct electricity, even when solid
Metallic lattices are malleable and ductile. What does this mean and why do they behave this way?
- Ductile: can be drawn out or stretched, permitting metals to be drawn into wires
- Malleable: can be hammered in different shapes. Metals can be hammered into shapes or hammered in thin sheets
- Delocalised electron can move giving structure a degree of give, so layers can slide over each other
Is silicon considered a metal or a non-metal?
- Si has shiny appearance of metal but is very brittle
- Si conducts electricity but very poorly
- ‘In between’ element, usually classified as a metalloid or semi-metal
Describe the trend in melting point across a period.
- 1-14: increases steadily. If metallic: nuclear charge increases + so does no. of outer electrons so stronger attraction. If giant covalent: more electrons in covalent bond as go across
- 14-15: sharp decrease in melting point as simple covalent so weak London forces
- 15-18: melting remains low due to simple structure
As you go across the period of elements forming metallic lattices, what are the trends?
- Ionic charge increases
- Ionic size decreases
- No. of outer shell electrons increases
- Attraction increases: melting + boiling point increases
What are the physical properties of group 2 elements?
- Reasonably high melting + boiling points
- Light with low densities
- Form colourless (white) compounds
Describe generally how group 2 elements react with other elements.
- They’re reactive metals
- Strong reducing agents
- Oxidised in their reactions to form 2+ ions
- M –> M+ + e-
- M+ –> M2+ + e-
Describe the reaction between group 2 elements and oxygen.
-React vigorously with oxygen
-Redox reaction
-Product: ionic oxygen with the formula MO
-e.g. calcium + oxygen –> calcium oxide
2Ca(g) + O2(g) –> 2CaO(s)
-calcium oxidised as oxidation number increased from 0 to +2
-oxygen reduced as oxidation number decreased from 0 to -2
Describe the reaction between group 2 elements and water.
-All (except beryllium) react to form hydroxide M(OH)2 and H2(g)
-Mg reacts slowly with water. Moving down group = more vigorous reaction
-e.g. calcium + water –> calcium oxide + hydrogen
Ca(s) + 2H2O(l) –> Ca(OH)2(aq) + H2(g)
-Calcium oxidised as oxidation number increased from 0 to +2
-Hydrogen reduced as oxidation number decreased from +1 to 0
-Redox reaction
Describe the reaction between group 2 elements and dilute acids.
-All (except Be) react to form a salt + hydrogen
-Becomes more vigorous as you go down the group
-e.g. calcium + hydrochloric acid –> calcium chloride + hydrogen
Ca(s) + 2HCl(aq) –> CaCl2 + H2
-Calcium oxidised as oxidation number increases from 0 to +2
-Hydrogen reduced as oxidation number decreases from +1 to 0
-Redox
Describe the reactions between group 2 oxides and water.
- React with water to form metal hydroxides
- General equation: MO(s) + H2O(l) –> M(OH)2(aq)
- Metal hydroxides soluble in water forming alkaline solutions with water as they release OH- ions
- Typical ph: 10-12
Describe the solubility of group 2 metal hydroxides.
- Increases down the group. More soluble hydroxide = release more OH- ions making more alkaline solution so higher pH
- Be: BeO is insoluble in water
- Mg: forms Mg(OH)2(s), only slightly soluble in water, dilute, comparatively low OH-(aq) concentration
- Ba(OH)2(s) more soluble so higher OH-(aq) conc so more alkaline
What are some uses of group 2 compounds?
- Neutralising acidic soils: Ca(OH)2 used by farmers + gardeners as lime to reduce acidity of soil
- Indigestion remedies: indigestion is build up of too much HCl. E.g. milk of magnesia contains Mg(OH)2 neutralising excess acids Mg(OH)2 + HCl –> MgCl2 + 2H2O
- Building + constructing: CaCO3 useful in building material (in limestone + marble). Used to manufacture glass + steel. Bad as it readily reacts with acids e.g. CaCO3(s) + HCl(l) –> CaCl2(aq) + H2O(l) + CO2(g). Most rainwater has acidic pH so gradual erosion of limestone or marble
What are the properties of the halogens?
- Low melting + boiling point
- Exists as diatomic molecules
Explain the trend in boiling points down group 17 (the halogens).
- Boiling point increases as you go down the group + physical state changes from gas to liquid to solid
- Due to successive elements having an extra shell leading to higher level of London forces (more electrons = larger instantaneous dipole so stronger London dispersion forces)
Explain the reactivity of the halogens.
-Halogens are very reactive + highly electronegative so good oxidising agents (attracting electrons)
-Form halide 1- ions
-As you go down reactivity/oxidising power decreases as:
> Atomic radius increases
> More shielding
> Ability to gain an electron in p sub-shell to form 1- ions decreases
Describe and explain the redox reactions of halogens.
- More reactive halogens oxidise + displace halides of a less reactive halogen: a displacement reaction
- Halogens change colour to indicate a redox reaction has taken place
- Mixture usually shaken with an organic solvent i.e. cylohexane to distinguish between iodine + bromine
- Cl2; in water: pale green, in cylohexane: pale green
- Br2; in water: orange, in cylohexane: orange
- I2; in water: brown, in cylohexane: violet
Describe the displacement reaction between chlorine and bromide ions.
- Chlorine more reactive so oxidises Br-
- Cl2(aq) + 2Br-(aq) –> 2Cl-(aq) + Br2(aq)
- Br2 orange in water + cylohexane
- Chlorine reduced (0 to -1)
- Bromine oxidised (-1 to 0)
Describe the displacement reaction between chlorine and iodide ions.
- Chlorine more reactive so oxidises I-
- Cl2(aq) + 2I-(aq) –> 2Cl-(aq) + I2(aq)
- I2 brown in water + violet in cylohexane
- Chlorine reduced (0 to -1)
- Iodine oxidised (-1 to 0)
Describe the displacement reaction between bromine and iodide ions.
-Bromine more reactive so oxidises I-
-Br2(aq) + 2I-(aq) –> 2Br-(aq) + I2(aq)
-I2 brown in water + violet in cylohexane
-Bromine reduced (0 to -1)
Iodine oxidised (-1 to 0)
What is a disproportionation reaction?
A reaction in which the same element is both oxidised and reduced
Describe the reaction of chlorine with water and what this is used for.
-Chlorine kills bacteria making water safe to drink
-Reacts with water forming HCl + chloric(I) acid HClO
Cl2(aq) + H2O(l) –> HClO(aq) + HCl(aq)
-Chlorine reduced (0 to -1 in HCl)
-Chlorine oxidised (0 to +1 in HClO)
-Disproportionation reaction
Describe the reaction between chlorine and dilute aqueous sodium hydroxide.
- Chlorine only slightly soluble in water + has mild bleaching action
- Household bleach formed from chlorine + dilute aq NaOH reacting at room temp
- Disproportionation reaction
- Cl2(aq) + 2NaOH –> NaCl(aq) + NaClO(aq) + H2O(l)
- Chlorine reduced (0 to -1 in NaCl)
- Chlorine oxidised (0 to +1 in NaClO)
Describe how household bleach is formed.
- Chlorine only slightly soluble in water + has mild bleaching action
- Household bleach formed from chlorine + dilute aq NaOH reacting at room temp
- Disproportionation reaction
- Cl2(aq) + 2NaOH –> NaCl(aq) + NaClO(aq) + H2O(l)
- Chlorine reduced (0 to -1 in NaCl)
- Chlorine oxidised (0 to +1 in NaClO)
How do you test for carbonate ions (CO3 2-)?
- Add a dilute strong acid to suspected carbonate
- Collect any gas for me + pass through limewater
- Fizzing/colourless gas is produced that turns limewater cloudy (if positive result)
- CO3 2-(aq) + 2H+(aq) –> H2O(aq) + CO2(g)
How do you test for sulfate ions (SO4 2-)?
- Add dilute HCl and BaCl2 to suspected sulfate
- White precipitate of barium sulfate is produced (if positive result)
- Ba2+(aq) + SO4 2-(aq) –> BaSO4(aq)
How do you test for halides?
- Dissolve suspected halide in water
- Add an aqueous solution of silver nitrate
- Note colour of precipitate formed
- If colour is hard to distinguish add aqueous ammonia (first dilute then concentrated)
- Note solubility of precipitate in aqueous ammonia
- Positive test; silver chlorine: white ppt, soluble in dilute ammonia; silver bromide: cream ppt, soluble in concentrated ammonia only; silver iodide: yellow ppt, insoluble in dilute + concentrated ammonia
- Ag+(aq) + X-(aq) –> AgX(s)
How do you test for ammonium ions (NH4+)?
- Add NaOH solution to suspected ammonium compound + warm v gently
- Test any gas evolved with red litmus paper
- Ammonia gas turns red litmus paper blue + has a distinct smell (ammonia gas hazardous - done with care)
- NH4+(aq) + OH-(aq) –> NH3(aq) + H2O(aq)
