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CHEMISTRY: TOPIC 3 PERIODICITY


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CHEMISTRY: TOPIC 3 PERIODICITY


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[Front]


What is relative atomic mass?
[Back]


the weighted average mass of an atom compared to the isotope carbon-12

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35 questions
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What is relative atomic mass?
The weighted average mass of an atom compared to the isotope carbon-12
What is a group on the periodic table?
Vertical column (1-18) no of electrons in the outer shell
What is a period on the periodic table?
Horizontal row = number of shells
What is electron shielding? (4)
When the inner electrons shield the outer valence electrons from the the full attraction from the nucleus so valence electrons needs less energy to remove than the inner electrons (lower ionization energy) across period: electron shielding is constant bec number of shells stays the same = same number of shielding electrons down group = electron shielding increases bec number of shells increase = number of shielding electrons increase
What is effective nuclear charge? (4)
Net positive charge on valence electrons atomic number - shielding electrons across period = increases bec atomic number increases but no of valence electrons are the same (more positive charge from the nucleus) Down group = stays the same
What are the trends in atomic radius? and why? (3)
Across period: decreases bec nuclear charge increases so more attraction to nucleus + constant shielding bec same no of shells Down group: increases bec of more shells being occupied = so the outer electrons are further away from the attraction of the nucleus (more distance) + more shielding = more repulsion = bigger radius
What are the trends in ionic radius? and why? (2)
Across period: first 4 decreases then it increases down group: increases bec of more occupied shells
What does isoelectronic mean?
Same electron configuration
Why are positive ions smaller than the original?
Bec they lose electrons to get the full outer shell so they have more protons than electrons = stronger bond + fewer occupied shells
Why are negative ions bigger than the original?
Bec they gain electrons to get a full shell so more electrons than protons = weaker attraction bw nucleus and electrons
What is first ionisation energy? (2)
Energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to produce 1 mole of gaseous 1+ ions Positive values (endothermic) as energy is coming in to overcome the attraction
What is the trend of ionisation energy? (3)
Across period: increases bec of increasing nuclear charge = decrease in atomic radius = more energy is needed to overcome bec of stronger attraction down group =decreases bec no of shells increase and so there is more electron shielding of the attraction from the nucleus Group 1 has the lowest ionisation energy whereas group 18 has the highest
Why is beryllium to boron an exception to ionisation energy?
Beryllium to Boron = decreases instead of increasing (across period) Be: 1s2 2s2 B: 1s2 2s2 2p1 electron lost in boron is in the p orbital which is slightly higher in energy = further from the nucleus than the s orbital = lower energy to remove Same for Magnesium to Aluminium
Why is nitrogen to oxygen an exception to ionisation energy?
The electron removed from oxygen is removed from a double occupied p orbital so it is repelled by the other electron = less energy to remove than a half filled orbital
What is electronegativity?
Measure of attraction of an atom in a molecule for a bonding pair of electrons as the 2 electrons don't bond equally. Depends on the atomic size and nuclear charge Fluorine = most Francium = least
What are the trends in electronegativity?
Across period: increases bec of increasing nuclear charge (protons) = stronger attraction bw nucleus + bonding pair of electrons + constant shielding Down group: decreases bec the atom size increases so bonding electrons are further from the nucleus attraction
What are the trends in melting point for period 2?
Lithium => carbon increase (Li, Be is metallic bonding and carbon is a giant covalent structure) Big decrease from nitrogen to neon increase again from Na to Si (Na, Mg, Al metallic and Si is a giant covalent) decrease from P => Ar
What does melting point depend upon?
Bonding type (covalent, metallic, ionic) structure (ionic lattice, molecular covalent, giant covalent and metallic) Metals: MP increases as bond gets stronger Giant covalent: highest peaks Molecular covalent: MP decreases Noble gases = lowest
What is metallic character?
How easily an atom can lose electrons metallic e = low IE = lose electrons = + ions non metallic e = high IE = gains electrons = - ions
What is the trend for metallic character?
Across period : decreases bec increasing nuclear charge = decreasing atomic radius = stronger attraction = harder to lose down group: increases bec of increasing atomic radius = weaker attraction bw nucleus and e = easier to lose e
What is First electron affinity? (4)
Energy released when one mole of electrons is added to one mole of gaseous atoms to form one mole of negative ions energy released = exothermic (negative) greater atomic radius + shielding = less energy released non metals are more exothermic than metals
Why is second electron affinity positive?
Bec of extra repulsion when adding a electrons to a negative ion
What are the characteristics of group 1 metals? (11)
Soft shiny = easily cut MP decreases as bonds gets weaker Atomic radii increases Li, Na, K float on water bec of low densities reacts with air reactivity increases down group IE : decreases bec radii increases = weaker attraction Metallic character = increases electronegativity = decreases bec of bigger radii forms ionic bonds with group 17 2NA + Cl2 = 2NaCl
How does lithium react with water and oxygen?
2Li + 2H2O = 2LiOH (Aq) + H2 (forms alkaline solution and hydrogen gas) 4Li + O2 = 2Li2O (s)
What are the colours and states for group 17 elements?
F = pale yellow gas Cl = greenish yellow gas Br = red liquid, Orange vapour I = grey solid, purple vapour, red-brown liquid
What are the characteristics of group 17 elements?
Soluble in H2O = becomes colourless MP increases bec of increasing molar mass (london forces get stronger) Atomic radii increases Electronegativity decreases as atomic radii increase IE decreases Electron affinity: less exothermic down reactivity decreases = F is strongest oxidising agent more reactive halogen can displace the less reactive halogen
What happens when a halogen gets displaced?
Colour changes depending on the halogen being displaced
What is the bonding for period 3 oxides?
Na2O , MgO, Al2O3 = ionic SiO2 (s) = giant covalent P4O10, P4O6, SO3 (l), SO2, Cl207(l), Cl2O = molecular covalent
Why does the bonding change from ionic to covalent in period 3 oxides?
Bec decreasing diff bw electronegativity = smaller diff = covalent Na+ o = 2.5 (ionic) S+ O = 0.8 (covalent)
What are the acid base properties of period 3? (3)
Na2o , MgO = Basic (base) Al2O3 = amphoteric (Acid/base) SiO2, P4O10, P4O6, SO3, SO2, Cl2O7, Cl2O = acidic (Acids)
Chemical formula for basic acids reacting with H2O in period 3?
Na2O + H2O = 2NaOH (aq) MgO + H2O = Mg (OH)2 (aq)
Chemical formula for acidic oxides reacting with H2O? (SO3, P4O10)
SO3 (g) + H2O = H2SO4 P4010 (s)+ 6H2O = 4H3PO4
How does nitrogen react with water and air?
2NO2 + H2O = HNO2 + HNO3 N2 + O2 = 2NO neutral oxide
How is the metallic structure related to ionisation energy?
A metallic structure has a regular lattice of positive ions in a sea of delocalised electrons they lose electrons to form positive ions across a period , IE increases = harder to lose electrons so metalic structures are formed on the left side = lower IEs down a group IE decreases = elements are more likely to show metallic behaviour lower down
What are the general characteristics of metallic elements?
Large atomic radii low IE less exothermic electron affinity and low electronegativity