| 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 |
