abg Flashcards
Why does ionisation energy decrease down a group?
- shielding of outer electrons increases as does atomic radius as the no. of shells increases
- this is despite the increased nuclear charge down the group
Explain the increase in ionisation energy across a period?
- nuclear charge increases
- no change in shielding as electrons are removed from the same shell
- atomic radius decreases slightly due to increasing nuclear charge attracting outer electrons with more force & pulling them in closer to the nucleus
Equation representing the 2nd ionisation energy of oxygen - with state symbols
O+(g) –> O2+(g) + e-
How to calculate the number of particles in a given sample?
multiply the number of moles of substance in the sample and multiply this by avogadro’s constant (6.02 x10^23)
acid + base –>
salt + water
acid + metal carbonate –>
salt + water + carbon dioxide
acid + metal –>
salt + hydrogen
what can oxidation be defined as?
loss of electrons, gain of oxygen or increase in oxidation number
what can reduction be defined as?
gain of electrons, loss of oxygen or decrease in oxidation number
What is the oxidation number of an element e.g. H2?
0
what is the sum of the oxidation numbers in a neutral compound?
0
What determines which element in a compound will be given the negative number? e.g. F2O
The most electronegative element, e.g. F is more electronegative than O so O will take on an oxidation number of 2+
What is disproportionation?
when one species is both oxidised and reduced in a reaction
What is an acid?
A proton donor - releases H+ ions in aqueous solution
when hydrocarbonic acids (eg methanoic acid) dissociate and release H+ ions which H+ ion always splits off and why? HCOOH –> ?
the one on the end, that isn’t bonded to carbon eg HCOOH –> H+ + HCOO-
This is because the hydrogen carbon bonds cannot be broken by the polarised charge of water- they are very strong.
What are bases?
proton acceptors
what are three different TYPES of bases?
ionic - a metal and a non-metal bonding, usually opposite sides of periodic table
covalent - non-metals bonding, usually closer on periodic table
what is an alkali?
a soluble base that releases OH- ions in water
describe the bonding in metals?
metallic bonding: giant metal structure, lattice held together by the electrostatic attraction between positive metal ions in a sea of delocalised electrons (lost from the outer shells of each atom)
Why do metals conduct electricity?
Delocalised electrons free to move and carry charge
Describe the bonding in ionic compounds?
giant ionic structure, attic held together by the electrostatic attraction between cations and ions
Describe the conductivity of ionic compounds?
Conduct electricity when molten/in solution as ions are then free to move and carry charge, however do not conduct electricity when solid as ions are fixed in place in the lattice
What are the three possible structures of covalent substances?
- giant covalent
- polymer structure
- simple molecular structure
what does the bond strength (and therefore melting point) of metals depend on?
- size and charge of cation
- number of electrons delocalised
what does the bond strength (and therefore melting point) of ionic compounds depend on?
size and charge of cation and anion
what does the bond strength (and therefore melting point) of giant covalent substances depend on?
the length of the bond: larger atom = longer and weaker covalent bond
What does the strength of intermolecular forces in simple molecular structures depend on?
- number of electrons: more electrons = stronger Van der Waals’
- polarity: polar molecules have stronger Van der Waals’ than non-polar substances with a similar Mr due to permanent dipole-dipole interactions between polar molecules
- hydrogen bonding: strong dipole-dipole attraction - strongest type of intermolecular force
What are 2 anomalous properties of water and what causes them?
- relatively high melting and boiling point thanks to strong hydrogen bonds that require a lot of energy to overcome
- ice is less dense than water as molecules are held further apart in an open lattice structure as 4 hydrogen bonds form with other molecules
Describe the trend in melting point across period 3
Increases from sodium to aluminium as metallic bond strength increases (increase in charge density and no. of delocalised outer electrons).
Silicon has the highest of the period thanks to strong covalent bonds.
phosphorus, sulphur and chlorine have simple molecular structures do m.p depends on strength of VdWs, so S8 with the most electrons has the highest and Cl with the least has the lowest of the 3. Argon has the weakest VdWs as fewest electrons.
what is covalent bonding?
a shared pair of electrons
what is the attraction between in covalent bonding?
between the negatively charged electrons and the positive chaarges of both nuclei - this attraction overcomes the repulsion between the two positively charged nuclei
what shape and bond angle will a molecule with 2 bonded pairs have?
Linear, 180o bond angle
what shape and bond angle will a molecule with 3 bonded pairs have?
trigonal planar, 120o bond angle
what shape and bond angle will a molecule with 4 bonded pairs have?
tetrahedral, 109.5o bond angle
what shape and bond angle will a molecule with 6 bonded pairs have?
octahedral, 90o bond angle
what shape and bond angle will a molecule with 3 bonded pairs and one lone pair of electrons have?
pyramidal, 107o bond angle. This is because lone pairs repel more strongly
how do lone pairs of electrons affect the bond angle?
lone pairs reduce bond angles as they repel bonded pairs more than bonded pairs repel each other
What are two factors bond polarity depends on?
- electronegativity: different electronegativities of the bonded atoms result in a partial positive/partial negative charge on either part of the molecule known as a dipole
- shape of molecule: if dipoles act in opposite directions e.g. in linear, symmetrical CO2 dipoles cancel, but if dipoles act in the same direction e.g. in asymmetrical water dipoles do not cancel out and the molecule is polar.
Describe the trend in ionisation energy, and therefore reactivity, down group 2
Ionisation energy decreases down the group as atomic radius increases so outer electrons are further from the nucleus and more shielded from the nucleus by inner shells. This outweighs the increasing nuclear charge.
Reactivity thus increases down the group due to the decreasing ionisation energy.
Describe the reactions of group 2 elements with water?
The Group 2 metals become more reactive towards water as you go down the Group
Only Ca and below will react with COLD water, and Mg will react with steam.
Ca sinks, fizzes and forms a white precipitate of Ca(OH)2
Describe the reaction od group 2 elements with oxygen?
Form white, ionic oxides. Mg will burn with a bright white flame.
What is the trend in ease of thermal decomposition as you go down group 2 carbonates?
It becomes progressively harder to decompose group 2 carbonates as you go down the group
What is the industrial importance of the reaction between CaO (lime) and water?
Formed slaked lime, Ca(OH)2 which forms a strongly alkaline solution (pH 14) used in agriculture to neutralise acid soils
state and appearance of chlorine, bromine and iodine at room temperature? why is the change is states the case
chlorine - green gas
bromine - orange liquid
iodine - grey shiny solid
Increase in melting and boiling point down the group, due to increasing strengths of Van der Waals’ forces due to an increasing number of electron within the molecule
trend in electronegativity down the halogens? why?
electronegativity decreases as the nucleus attracts the shared pair less as the atom gets larger (more shells). Fluorine is the most reactive
Describe the trend in oxidising strength (reactivity) down the halogens?
oxidising strength (e.g. the attraction they have towards an electron to accept it) decreases down the group as the atomic radius increases and the electron accepted is less strongly attracted to the nucleus e.g as ionisation energy decreases
which ionic equations from displacement reactions show the trend in reactivity down the halogens?
Cl2 + 2Br- –> 2Cl- + Br2
Br2 + 2I- –> I2 + 2Br-
What colour is the hexane layer for the following ionic equations? and if Cl2 was present?
- Cl2 + 2Br- –> 2Cl- + Br2
- Br2 + 2I- –> I2 + 2Br-
- = orange
- = purple
With Cl2 = pale green
what is the name given to a reaction in which the same species is both oxidised and reduced?
disproportionation
How is chlorine used in public health? & reaction.
What is the risk in doing this?
to sterilise drinking water as HOCl kills bacteria
Cl2 + H2O –> HOCl + HCl
Risk -can react instead with organic compounds in water to form carcinogenic compounds such as CH3Cl
What is the reaction of chlorine with cold dilute sodium hydroxide compared with hot sodium hydroxide?
Cold:
Cl2 + 2NaOH –> NaCl + NaClO + H2O
Hot - further disproportionation:
3Cl2 + 6NaOH –> 5NaCl + NaClO3 + 3H2O
how are elements arranged of elements in the Periodic Table? (3)
- by increasing atomic (proton) number,
- in periods showing repeating trends in physical and chemical properties,
- in group shaving similar physical and chemical properties;
what is used as the standard measurement of relative masses?
carbon-12 isotope
what are common bases?
metal oxides, metal hydroxides and ammonia
when is a salt produced?
when the H+ ion of an acid is replaced by a metal ion or NH4+
describe the term electronegativity
the ability of an atom to attract the bonding electrons in a covalent bond
describe periodicity
a repeating pattern across different periods
why do atoms of elements in the same group have similar properties?
because they have similar outer shell electron configurations
