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How are the elements in the periodic table arranged?
in increasing atomic number, periods are rows, groups and columns and table is split into s,p,d,f blocks according to the highest occupied sub-level
What group make up the alkali metals?
group 1
What groups make up the transition metals?
group 3-12
What group makes up the alkaline earth metals?
group 2
What groups make up the S block?
groups 1 and 2
What section of the periodic table makes up the D block?
transition metals
What group makes up the noble gases?
group 18
What elements are metalliods?
B, Si, Ge, As, Sb, Te
What group makes up the halogens?
Group 17
What groups make up the P block?
groups 13-18
What groups make up the f block?
lathanoids and actinoids
Define Electronegativity
the ability of an atom to attract a pair of electrons towards itself in a shared covalent bond.
How does electronegativity change down a group?
electronegativity decreases down a group as shielding increases and nucleus’ power of attraction gets weaker.
How does electronegativity change across a period (left to right)?
electronegativity increases across a period as shielding stays the same but proton number increases, making nucleus’ power of attraction greater.
Define Atomic Radius
the radius of an a atom
How does atomic radius change down a group?
atomic radius increases down a group as shielding increases
How does atomic radius change across a period?
atomic radius decreases across a period as shielding stays the same but atomic number increases, causing electrons in atom’s energy levels to be attracted more closely to nucleus, making the atom denser.
Define ionisation energy
the energy required to remove one mole electron from one mole of a gaseous atom M(g) → M+(g) + e-
How does ionisation energy change going down a group?
ionisation energy decreases down a group
How does ionisation energy change across a period?
ionisation energy generally increases across a period but there are exceptions
Define electron affinity
The amount of energy released when one mole of electrons is gained by one mole of atoms of an element in the gaseous state to form one mole of gaseous ions. M(g) + e- → M-(g)
How does electron affinity change going down a group?
generally electron affinity decreases going down a group and becomes less exothermic - but beware of exceptions. The size of the atom increases, weakening the attraction between the nucleus and the added electron, making the addition of an electron require more energy.
How does electron affinity change across a period?
the general trend is that electron affinity increases across a period and becomes more exothermic - there are exceptions
Why are radii of positive ions smaller than their atomic radii?
M and M+ have the same number of protons and therefore the same amount of nuclear attraction, however M+ has on fewer electron so less electron-electron repulsion, making the electrons pulled closer together and therefore creating a smaller atomic radii.
Why are the radii of negative ions larger than their atomic radii?
X and X^- have the same number of protons so their nuclear attraction is the same, however X- has an extra electron so there is greater electron-electron repulsion, making the electrons spread further apart, increasing the ionic radii.
Why is there a sudden decrease in electron affinity from F to Cl?
From F to Cl, a huge decrease in electron-electron repulsion occurs as F is a very small atom with a lot of electron electron repulsion whereas Cl is a larger atom that has less electron electron repulsion.
Define a metallic structure
a structure consisting of a regular lattice of positive ions in a sea of delocalised electrons
What features do most elements capable of forming metallic structures have?
a low ionisation energy - meaning they can form positive ions easily
Define a Basic Oxide (e.g Na2O)
a basic oxide is an oxide that will react with an acid to form a salt and, if soluble in water, will produce an alkaline solution
Define an amphoteric oxide (e.g. Al2O3)
an amphoteric oxide is an oxide that reacts with both acids and bases
Define an acidic oxide (e.g. SO2)
an acidic oxide is an oxide that reacts with bases/alkalis to form a salt and, if soluble in water, will produce an acidic solution
Describe how oxides change across a period
oxides change from basic to amphoteric to acidic across a period
Give the formula of the oxide, nature of element and nature of oxide for Sodium
Na2O, metal, basic
Give the formula of the oxide, nature of element and nature of oxide for Magnesium
MgO, metal, basic
Give the formula of the oxide, nature of element and nature of oxide for Aluminium
Al2O3, metal, amphoteric
Give the formula of the oxide, nature of element and nature of oxide for Silicon
SiO2, non metal, acidic
Give the formula of the oxide, nature of element and nature of oxide for phosphorus
P4O10, non metal, acidic
Give the formula of the oxide, nature of element and nature of oxide for Sulfur
SO2, SO3, non metal, acidic
Write the chemical equation for the reaction between sodium oxide and water
Na2O(s) + H2O(l) → 2NaOH(aq)
Write the chemical equation for the reaction between Magnesium Oxide and water
MgO(s) + H2O(l) → Mg(OH)2(aq)
Write the chemical equation for the reaction between Phosphorus oxide and water
P4O10(s) + 6H20(l) → 4H3PO4(aq)
Write the chemical equation for the reaction between Nitrogen dioxide and water
2NO2(g) + H2O(l) → HNO2(aq) + HNO3(aq)
Write the chemical equation for the reaction between sulfur (2) oxide and water
SO2(g) + H2O(l) → H2SO3(aq)
Write the chemical equation for the reaction between sulfur (3) oxide and water
SO3(g) + H2O (l) → H2SO4 (aq)
What two oxides contribute to acid deposition in the form of acid rain?
nitrogen and sulfur oxides
Describe the reactions of group 1 (alkali metals)
all group 1 elements have one outer shell electron so virtually all reactions involve the loss of this outer shell electron to form the positive M+ ion. Group 1 metals rapidly react with water and react with halogens to form salts. Reactivity increases down the group.
Give the symbol equation for the reaction between an alkali metal and water
2M(s) + 2H20(l) → 2MOH(aq) + H2(g)
Give the equation for the reaction between an alkali metal and a halogen
2M(s) + X2(g) → 2MX(s)
Describe the reactions that occur with elements in Group 17
all elements in group 17 have 7 outer shell electrons so usually react by gaining one electron (X-) or forming a covalent compound. Reactivity decreases down the group and displacement reactions are very common.
What is a displacement reaction? (involving halogens)
a displacement reaction involves a reaction between a halogen solution and a solution containing halide ions. The more reactive halogen displaces the halide ion of the less reactive halogen from the solution.
What occurs during this reaction? Cl2 + 2KBr
→2KCl + Br2, chlorine displaces bromide ions and solution turns orange
What occurs during this reaction? Cl2 + 2KI
→ 2KCl + I2, chlorine displaces Iodide ions and solution turns brown
What occurs during this reaction? Br2 + 2KI
→ 2KBr + I2, Bromine displaces iodide ions and solution turns brown
Define Transition Metal
an element which forms at least one stable ion with a partially-filled d sub-shell
Give the electron configuration of Sc (21)
[Ar]4s2 3d1
Is Zinc a transition metal according to the definition?
No Zinc is not a transition metal as it has a complete 3d sub shell
Is Scandium a transition metal according to the definition?
No Sc is not a transition metal according to the definition as it only forms the ion Sc3+ which doesn’t have a 3d sub orbital
Give 5 characteristics of transition metals
can exhibit more than one oxidation state 2. form complex ions 3. form coloured compounds 4. act as catalysts in many reactions 5. can exhibit magnetic properties
What oxidation state do all transition metal elements exhibit when the 4s electrons are removed
+2
How do you calculate the maximum oxidation state exhibited by a transition metal?
add the number of electrons in their 4s and 3d sub orbitals
Why can transition metals have more than one oxidation state?
the 4s and 3d sub orbitals are close in energy and there are no big jumps in the successive ionisation energies when the 4s and 3d electrons are removed
Define Complex Ion
a central transition metal ion surrounded by ligands
Define Ligand
negative ions or neutral molecules which use lone pairs of electrons to bond to a transition metal ion to form a complex ion
What bond is formed between the ligand and the transition metal ion?
a coordinate bond as the ligand donates a pair of electrons to the transition metal ion
Give the formula of neutral ligands
H2O, NH3, CO
Give the formula of -1 ligands
Cl-, CN-, OH-
Define Paramagnetic substances
substances that are attracted by a magnetic field, caused by the presence of unpaired electrons
Define Diamagnetic substances
substances that are repelled slightly by a magnetic field, caused by the presence of paired electrons
Why in transition metals do we remove electrons from the 4s sub orbital before the 3d sub orbital when an ion is formed?
in transition metals the 4s sub shell is at a lower energy than the 3d sub shell so electrons are removed their first as it requires less energy.
True or False Cr is [Ar] 4s1 3d5 not [Ar] 4s2 3d4
True - [Ar] 4s1 3d5 is more energetically stable
What two transition metals have 4s1 instead of 4s2?
Cr and Cu
How are coloured complexes formed?
Complex ions e.g. [Ni(H2O)6]2+ have an incomplete d -sub-shell which splits into two groups with different energies. Certain frequencies of visible light can be absorbed to promote an electron from the lower set of d orbitals to the higher set. Since [Ni(H2O)6]2+ is green, it must absorb the complementary colour red and transmit green.
What complexes are colourless?
any complexes containing Sc3+ or Ti4+ are colourless as they have no electrons in the 3d sub-shell. any complexes containing Cu+ or Zn2+ are also colourless as they have a full 3d subshell.
What can affect the colour of a complex ion?
the amount of splitting of the d orbitals - the bigger the difference in energy between the lower and higher set of d-orbitals, the higher the frequency of light absorbed. 2. the identity of the metal (different electron configurations) 3. the oxidation state of the metal (higher charge in ion causes greater splitting) 4. the nature of the ligand (ligands have different splitting powers)
What is the spectrochemical series?
the arrangement of ligands according to how much they split d orbitals.
CN- is after F- in the spectrochemical series. What does this mean?
CN- causes greater splitting of the d orbitals than F- so therefore a higher frequency of light will be absorbed by complexes containing CN- ions.
What factor causes some ligands to split d orbitals more than others?
the charge density of the ligand can explain its splitting power, the higher the charge density the greater the splitting power.
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