Chapter 3/ 13 Flashcards
Periodicity
Elements in periodic table
Arranged in order of increasing atomic number (Z)
- each successive element has one more proton than the last
Periodicity
Repeating patterns of chemical and physical properties
Importance of atomic number
- defining property of an element
- came from Moseley’s studies of X-rays released when atoms of different metallic elements were bombarded by electrons
- discovered that frequency of X-rays released showed a direct relationship to atomic number of element
Groups
Vertical columns
Periods
Horizontal rows
- period no. = no. of occupied main energy levels in the atom
Group 1 elements
Alkali metals
- most reactive group of metals
- react strongly with cold water
Group 2 elements
Alkaline earth metals
- reactive metals
- less reactive than group 1
Group 17 elements
Halogens
- most reactive group of non-metals
Group 18 elements
Noble gases: so called because of their lack of reactivity
- group of very unreactive monoatomic gases
Metalloids
Located in a ‘diagonal’ staircase that forms a boundary between metals and non-metals
- have properties of both metals and non-metals
Covalent radius and metallic radius
Measured as half the distance between two neighbouring nuclei
- valent radii and metallic radii can be obtained for most elements (except those that don’t from compounds)
- used for comparison across a period in the periodic table
van der Waal’s radius
- atomic radii values quoted for the noble gases (group 18)
- they don’t form compounds
- larger than covalent radii
Atomic radius trends
Across a period:
- atomic radii of atoms decreases across the period
- no. of protons increases by one with each successive element, hence, no. of electrons also increases by one
- extra electron occupies the same main energy level, so, electron shielding remains roughly constant across a period
- results in stronger attraction between nucleus and valence electrons, pulls them closer to nucleus
Down a group:
- atomic radii increases down a group
- each additional period means valence electrons occupy a main energy level further from nucleus
Electron shielding
Occurs when outer electrons are shielded from attraction of nucleus by inner electrons (shielding electrons)
Ionic radius
Metallic elements tend to lose electron to form cations
Non-metal elements gain electrons to form anions
Isoelectronic species
Ions that have different numbers of protons, as they are different elements, but the same number of electrons
Ionic radius trends (positive ions)
Positive ions:
- Lose electrons from valence shell
- each ion still has same no. of protons in nucleus as parent atom, but are pulling on fewer electrons
- increases attraction between nucleus and valence electrons
- hence, positive ions are smaller than their parent atoms
- So, trend across a period is that ionic radii decrease as nuclear charge increases
Ionic radius trends (negative ions)
Negative ions:
- gain electrons; have more electrons than protons
- decreases attraction between nucleus and valence electrons; weaker attraction causes ionic radius to increase
- these extra electrons also increase repulsion between electrons, contributes to increase in ionic radii
- across period, ionic radii of negative ions decrease as nuclear charge increases
Electrostatic attraction
The force of attraction between the positively charged nucleus and negatively charged electrons within an atom
Strength of electrostatic attraction depends on:
- Atomic radius (distance between nucleus and electrons)
2. Number of shielding electrons within the atom
Shielding electrons
Inner electrons that tend to ‘shield’ the outer electrons from the full attraction of the nucleus
Hence, valence electrons don’t feel full attraction from protons in the nucleus
Effective nuclear charge
Attraction felt by valence electrons
- is less than the actual nuclear charge of the atom
Trends in effective nuclear charge
Across a period:
- no. of protons increases by one for each successive element, but no. of electrons in inner energy levels doesn’t change
- effective nuclear charge increases by one until group 18
Down a group:
- effective nuclear charge remains approx. constant moving down a group
- attraction from increasing no. of protons in nucleus is offset by increase in occupied energy levels down group
First ionisation energy
Energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions
- measure of attraction between positively charged nucleus and negatively charged outer valence electrons
First ionisation energy trends
Across a period:
- increase across a period
- due to increase in nuclear charge and decrease in atomic radii
- results in an increased attraction between nucleus and valence electrons
Down a group:
- decrease down a group
- due to increase in atomic radius
- increased distance results in a weaker attraction between nucleus and valence electrons of an atom
Exceptions to the ionisation energy trend across a period
IE decreases from:
- Beryllium to boron
electrons in p orbitals are higher in energy and further from nucleus than electrons in s orbitals, hence less energy is required to remove
- Magnesium to aluminum
electrons in p orbitals are higher in energy and further from nucleus than electrons in s orbitals, hence less energy is required to remove - Nitrogen to oxygen
electron is removed from a doubly occupied p-orbital. An electron in a doubly occupied orbital is repelled by other electron and requires less energy to remove than an electron in a half-filled orbital - Phosphorus to sulphur
electron is removed from a doubly occupied p-orbital. An electron in a doubly occupied orbital is repelled by other electron and requires less energy to remove than an electron in a half-filled orbital (electrons are being removed from 3p sub-level)
Electronegativity
Attraction of an atom for a bonding pair of electrons
- measured on Pauling scale
Electronegativity trends
Across a period:
- increases across a period duet to increase in nuclear charge and decrease in atomic radius
- results in a stronger attraction between nucleus and bonding electrons
Down a group:
- electronegativity decreases down a group
- due to increase in atomic radius, increasing distance between nucleus and shared pair of electrons, hence, attraction for these electrons decreases
- increase in nuclear charge down a group is counteracted by increased shielding caused by extra occupied main energy levels within the atom
Electronegativity and bonding
Difference in electronegativity between atoms in a compound determines type of bonding that occurs
- difference up to 1.7 units = forms covalent bond
- difference of 1.8 units or more = forms ionic bond
Electron affinity
First electron affinity: energy released when one mole of electrons are added to one mole of gaseous atoms to form one mole of gaseous 1- ions
- exothermic
Group 17 elements and electron affinities
Group 17 elements have highest electron affinities (most exothermic)
- relatively small atoms can accommodate an electron in their unfilled outer shell
- hence, it’s strongly attracted to the nucleus of the atom
Electron affinity trends
Across a period:
- values increase, becoming more exothermic
- progressing to the filling of the outer valence shell of the atom
- group 17 atom releases more energy than a group 1 atom on gaining an electron
- on gaining an electron it achieves a filled valence shell, is more stable
Down a group:
- values decrease
- additional electron gained enters an energy level further from nucleus
- added electron has a weaker attraction to nucleus, releases less energy when added
Metals and ionisation energies
- have low IE
- lose their valence electrons relatively easily to form positive ions
Non-metals and ionisation energies
- high ionisation energies
- gain electrons to form negative ions
Metallic character trends
Across a period:
- metallic character decreases
- due to increase in IE across a period
Down a group:
- metallic character of elements increases as IE decreases
Metallic character
- strongly related to its ionisation energy
- IE depends on nuclear charge and atomic radius of an atom
- higher the nuclear charge, smaller the atomic radius of an atom, lower its metallic character and vice versa
Bonding trends
Across a period, bonding and structure change from:
Metallic –> Giant covalent –> Molecular covalent
Period 3 oxides
Show gradual trend of decreasing metallic character across the period
- type of bonding changes from ionic to covalent across period, determined by difference in electronegativity between bonding atoms
Metal oxides
- giant ionic structures
- solids under standard conditions due to strong electrostatic attractions between ions
Non-metal oxides
- molecular covalent structures
- gases of liquids under standard conditions
Acid-base properties of period 3 oxides
- change from basic through amphoteric to acidic across a period
- metal oxides form metal hydroxides when reacted with water
- non-metal oxides form acidic solutions when reacted with water
Amphoteric
Can act as both an acid and a base
Basic oxides
Dissolve in water to produce basic solutions
Acidic oxides
Dissolve in water to produce acidic solutions
Properties of alkali metals
- very reactive
- soft (can be easily cut with a knife; reveals a shiny surface that quickly tarnishes as a layer of oxide is formed)
- MP and BP are relatively low, decrease down the group (due to metallic bonding getting weaker as ionic radii of metals increase)
- stored in oil to prevent reaction with oxygen in the air
- one electron in valence shell, lose relatively easily (low IE) to form 1+ ions with halogens
- undergo vigorous reactions with water to form a metal hydroxide and hydrogen gas; resulting solution has high pH
- reactivity increases down the group
Properties of halogens
- very reactive group of non-metals
- characteristic colours
- physical state changes down the group (related to increasing molar mass of halogens, results in stronger intermolecular forces between molecules)
- diatomic; they exist as two atoms bonded together, rather than separate atoms
- MP and BP increases down the group, related to increasing strength of intermolecular forces
- reactivity of halogens decreases down the group
- react with group 1 elements to form salts
- undergo displacement reactions with each other
Properties of chlorine
- dense pale-green gas
- smelly and poisonous
- occurs as chloride in the sea
Properties of bromine
- deep-red liquid with red-brown vapour
- smelly and poisonous
- occurs as bromide in the sea
Properties of iodine
- grey solid with purple vapour
- smelly and poisonous
- occurs as iodides and iodates in some rocks and seaweed
d-block elements
- occupy central region of periodic table
- so called because they have valence electrons in d sub-level
Transition elements
An element that has an incomplete d sub-level in its atom or one or more of its ions
- found in the d-block region (excludes zinc)
- show and oxidation state of +2 when the s-electrons are removed
Scandium as a transition metal
- transition element because of electron configuration of atom [Ar] 3d1 4s2
- Sc3+ ion has no 3d electrons at all, as atom loses both 4s electrons and one 3d electron when forming the ion
Copper as a transition metal
- transition element
- although Cu and Cu+ have complete 3d sub-levels, copper (II) ion (Cu2+) has electron configuration [Ar]3d9, signifies an incomplete d sub-level
Zinc as a transition metal
- not a transition metal because it doesn’t form ions w/ incomplete d-orbitals
- both atom and ion formed (Zn2+) have a complete 3d sub-level
Characteristic properties of transition elements
Presence of unpaired electrons in d sub-level gives rise to characteristic properties of transition elements
- variable oxidation states
- have coloured compounds
- elements display catalytic and magnetic properties
- form complex ions w/ ligands
Variable oxidation states
Reason:
- closeness in energy of 3d and 4s sub-levels
- there’s little change in atomic radii from left to right, as electrons are added to inner sub-level (3d)
- so, effective nuclear charge experienced by outer valence electrons in 4s sub-level remains constant
- atoms of d-block elements are of similar size, and effective nuclear charge on outer 4s electrons is also similar
- results in there being only a small range in values of first IE across first row of d-block
Transition metals display a wide range of oxidation states in their compounds
- contrast to s-block metals, only show oxidation states of +1 (for Group 1) and +2 (for Group 2)
Copper (II) sulphate solution
- solution is blue due to presence of hydrated copper (II) ions
- copper (II) ions in solution are surrounded by 6 water molecules (ligands), results in formation of complex ions
Complex ion
Made up of a central metal ion surrounded by ligands
Ligands
Species with lone pairs of electrons that act as Lewis bases, they donate a lone pair of electrons to the central metal ion
- can be neutral molecules, w/ lone pairs of electrons or charged ions
- all ligands are lone pair donors (all ligands act as Lewis bases)
Central metal ion (in a complex ion)
Central metal ion = Lewis acid
- accept lone pair of electron from ligand, forming a coordinate covalent bond
Coordinate covalent bond
Bond formed between a ligand and the central metal ion
Coordination number
Number of coordinate covalent bonds formed between the ligands and central metal ion
- no. of coordinate bonds to one central ion
Monodentate ligands
Ligands that form just one coordinate covalent bonds
Bidentate ligands
Ligands that form two coordinate bonds from each ion or molecule to the central metal ion
Overall charge on a complex ion
Sum of the charges of the central metal ion and the charges contributed by the ligands
Shapes of complex ions
Depend on no. and size of ligands involved
look at notes for shapes
Catalytic behaviour
Presence of unpaired electrons in 3d sub-level of transition metal atoms or ions gives rise to two characteristic properties of these elements
- catalysts
- magnetic properties
Catalyst
Increases rate of a chemical reaction by providing an alternative reaction path w/ a lower activation energy (Ea)
Transition elements as catalysts
Play an important role as catalysts in chemical industry
- allow chemical reactions to take place at lower temp. and pressures than they would otherwise need
Heterogenous catalysis
Where catalyst is in a different state to that of the reactants
- Transition elements are effective heterogeneous catalysts
- eg. use of Iron in Haber process
Haber process
Reactants nitrogen and hydrogen are in gaseous state w/ iron being in solid state
- iron provides a surface on which reactant molecules can adsorb
- hence, come together w/ correct orientation to react
Homogeneous catalyst
A catalyst that is in the same physical state as the reactants
- transition metals are also effective as homogeneous catalysts in redox reactions due to their ability to have variable oxidation states
Examples of transition metals as homogeneous catalysts
- enzymes (biological states) enable chemical reactions within human body cells to take place at lower temp. than would be otherwise necessary (iron and cobalt are important components of these enzymes)
- iron (II) ion is central to haeme group in haemoglobin, responsible for carrying oxygen around the body
- cobalt (III) ions are found in vitamin B12, essential to maintain good health
Electrons and orbitals
Two electrons in same atomic orbital have opposite spins- represented by single-headed arrows
- electron spins give rise to magnetic effects
- classified as: dimagnetism, paramagnetism or ferromagnetism
- if an atomic orbital contains two electrons, opposite spins cancel each other out
Dimagnetism
Substances with any no. of paired electrons
- diamagnetic materials show a weak repulsion in an external magnetic field (these are normally considered as non-magnetic)
Pyrolytic carbon
- similar to graphite
- is diamagnetic
Paramagnetism
Substances w/ half-filled atomic orbitals (contain only one electron)
- paramagnetic materials are attracted to an external magnetic field
- greater no. of unpaired electrons, stronger the attraction
Ferromagnetism
- the largest effect
- unpaired electrons become aligned with an external magnetic field
- this alignment persists even after the external field is removed- object itself becomes magnetised
- only certain metals, some rare earth elements and their alloys are ferromagnetic
Transition metal and paramagnetism
- transition metal complexes have unpaired electrons
- attracted to a magnetic field, showing paramagnetic properties
Dimagnetism vs. Paramagnetism
Dimagnetic = substances with only paired electrons Paramagnetic = substances with unpaired electrons
Complex ions and d-orbitals
d-orbitals have same energy in an isolated atom
- but split into two sub-levels in a complex ion
- electric field of ligands may cause d-orbitals in complex ions to split
- so energy of an electronic transition between them corresponds to a photon of visible light
Coloured compounds
Formation of coloured compounds in solution is a characteristic property of transition elements
- formation of coloured compounds is due to presence of partially-filled orbitals in 3d sub-level
Electromagnetic spectrum
- made up of a range of wavelengths/frequencies of electromagnetic radiation
- very little is ‘seen’ by the human eye
- small portion of spectrum that we can see, visible spectrum; 400-700nm
White light in solution
- white light passes through a solution containing transition metal ions
- certain wavelengths of visible spectrum are absorbed, certain wavelengths are transmitted
Colour wheel and complementary colours
- use a colour wheel to explain relationship between absorbed and transmitted light
- each colour in wheel has a complementary colour
- for coloured solutions, light that we observe is the transmitted light- this is the complementary colour of the light that is absorbed
NB/ colour of complex ion seen is the complementary colour to that light which is absorbed
Exceptions to forming coloured compounds
- Cu+ and Sc3+ don’t form coloured compounds
- in these ions, d sub-level is either completely full or completely empty
Splitting d sub-level
- 5 d-orbitals in an isolated transition metal atom/ion are degenerate
- ligands bond to central metal ion
- repulsion between electrons in ligands and those in d orbitals in transition metals causes 5 d orbitals to split into 2 sets of different energy (non-degenerate orbitals)
Delta E (change in E)
Difference in energy between two sets of d orbitals
- exact energy difference between non-degenerate d orbitals in a transition metal is determined by several factors
- one being the identity of ligands that surround the transition metal ion
Transition metals and d orbitals
Transition metal ions have at least one partially filled d orbital
- hence, electrons can transition from lower set to higher set of d orbitals by absorbing energy
- energy absorbed corresponds to wavelength of visible light, w/ complementary colour of colour that is absorbed being transmitted
- amount of energy absorbed during the electron transition (delta E) is calculated using E=hv
Why are compounds of transition elements coloured?
Colour doesn’t occur due to transitions between principal energy levels but because some colour is absorbed when electrons are excited between split d-orbitals
- complexes of d-block elements are coloured because light is absorbed when an electron is excited between the d-orbitals
Spectrochemical series
Effect of different ligands on degree of d orbital splitting in an octahedral complex, hence size of the change in energy is given in this series
Presence of ammonia ligands
- causes a difference in colour, due to presence of ammonia ligands, cause greater splitting of d orbitals
- means that size of deltaE increases, resulting in a greater amount of energy being absorbed when electrons transition to higher set of d orbitals
- hence, a shorter wavelength of light is absorbed, and solution appears a darker blue colour
Factors that determine colour of complex ions
- identity of metal ion at centre of complex ion
- this affects no. of electrons involved in d orbitals, oxidation state of metal and nuclear charge at centre of ion
- oxidation states of the metal ion
- higher the oxidation state, higher the charge, lower the no. of e-
- higher the electron repulsion between the ligand and d electrons
- hence, higher the energy
- nuclear charge
- higher the no. of protons, higher the nuclear charge
- this increases the electrostatic attraction between donated pairs of e- and nucleus
- thus, change in energy is greater, resulting in a higher wavelength of light being emitted
- ligand identity
- higher the ligand on the spectrochemical series, higher the charge density
- higher the charge density, higher the split in d orbitals, due tot increased repulsion within orbitals
- higher the split, higher the energy
Crystal field theory
Used to explain colour in transition metal complexes
- theory accounts for colour of complex ions as arising from d-d electron transitions caused by splitting of d sub-level orbitals by repulsion effect of ligands present
- type of splitting that takes place depends on no. of ligands attached as they interact w/ electron distributions in space of 5 d-orbitals
- theory is based on assumption that ligands are point charges
Anion ligands
Should have greatest splitting effect
- but anion ligands are found at low end of spectrochemical series, causing the least splitting
Weakness of crystal field theory
- doesn’t take account of covalent bonding between ligand and central metal cation
- these are explained by alternative ligand field theory, model based on molecular orbital (MO) theory
Transition metal ions in solution
- have a high charge density
- thus, act as Lewis acids and attract species rich in e- (ligands)
Ligands
neutral molecules or anions that contain one or more non-bonding pairs of electrons
- ligands form covalent bonds w/ central transition metal ion to form a complex ion
- if a ligand is higher on the spectrochemical series than another, it can displace that ligand
Complex ion
has a central metal ion at its centre w/ a no. of other molecules surrounding it
Polydentate ligands
ligands that can utilise 2 or more pairs of lone electrons to form a coordinate bond
