Chapter 2 and 3 QUIZ

Chapter 2

Dalton’s Model

• All matter is made up of indivisible particles called atoms

• Atoms cannot be created or destroyed

• Atoms of the same element are identical

• Atoms of different elements are different

• Atoms can combine to form molecules

J.J. Thomson Model

• Stated that atoms had negatively charged electrons scattered in a positively-charged area (“plum pudding”)

Ernest Rutherford Model

• Stated that there is a small positively charged nucleus in the centre

Bohr Model

• Positive nucleus with negative electrons surrounding it

• Electrostatic force of attraction between the oppositely charged subatomic particles prevents electrons from leaving the atom

• Evidence:

  • Electron moves into an orbit (higher energy level) further away from the nucleus when an atom absorbs energy

  • Excited state is unstable and electron falls back to the lowest level; ground state

  • Energy the electron gives out when falling is in the form of EM radiation

  • When an atom absorbs a photon, it reaches this excited state and jumps to a higher energy level

  • One photon is released for each electron transitions

    • Energy of the photon is proportional to the frequency of the radiation

      • Planck’s equation: E_photon = hf (h = Planck’s constant)

      • Shows that line spectra allows us to look inside of the atom

    • Energy of the photon is equal to the energy change in the atom

      • Atoms emit photons of certain energies which give lines of certain frequencies as the electron can only occupy certain orbits

      • Electron can change its energy by discrete amounts

      • Energy of the atom is quantized (concept that a system cannot have a possible energy value but is limited to certain energy values)

        • Quantization: process of mapping continuous infinite values to a smaller set of discrete finite values

        • If energy were not quantized, emission spectra would be continuous

Atomic Number & Atomic Mass Number

• Atomic Number: number of protons (Z)

• Atomic Mass Number: number of protons and neutrons (A)

Isotopes

• Atoms have the same number of protons but different numbers of neutrons

  • Same chemical properties and reactions

  • Different physical properties, ie, boiling and melting points (heavy isotopes move slowly at a given temperature)

  • Positions can be monitored by detecting radiation levels

• Atoms of the same element with different mass numbers exist so there can be relative mass numbers where the average atomic mass is found

Radioisotopes

• Graph where x = # of protons and y = # of neutrons, the stable nuclei within these curved lines the band of stability

• Nucleus’ stability depends on the balance between protons and neutrons

  • Too many/few neutrons, it is radioactive and changes to a more stable nucleus by giving out radiation

  • May be of different forms which differ in ionization and penetration abilities

    • Alpha particles: emitted by nuclei with too many protons composed of 2 protons and 2 neutrons

    • Beta particles: emitted by nuclei with too many neutrons composed of electrons ejected from the nucleus due to neutron decay 

• Carbon-14 dating:

  • Carbon-14 has 8 neutrons and is unstable (too many)

    • Reduces when neutron changes to a proton and electron

    • Proton stays in the nucleus but electron is ejected as a beta particle

  • Abundant in plants as carbon atoms are constantly refreshed from CO₂ around us

    • When organisms die, no more carbon-14 is absorbed and carbon-14 levels drop due to nucleus decay

  • Can be used to date carbon-containing compounds

    • Rate of decay measured by its half-life; time taken for ½ the atoms to decay

• Cobalt-60

  • Radiotherapy damages DNA by knocking electrons and not allowing the cell to grow

  • Eg, can treat localized solid tumours (ie, cancer of skin, tongue) and blood cancer

  • Commonly used as it emits penetrating gamma radiation when its protons and neutrons change their relative positions in the nucleus

• Iodine-131

  • Emitter of beta and gamma rays

  • Can be used as sodium iodide to investigate thyroid gland and diagnose/treat thyroid cancer

  • Short half-life of 8 days, thus quickly eliminated from the body

• Iodine-125

  • Used for prostate cancer treatment

  • Isotopes’ pellets are implanted into the gland

  • Half-life of 80 days, allowing low levels of beta radiation to be emitted over an extended period

Ions

• # of protons and neutrons never change during a chemical reaction = electrons responsible for chemical change

• Cation: positive ion, lost electrons

• Anion: negative ion, gains electrons

  • Magnitude of charge depends on how many electrons lost/gained

Mass Spectrometer

• Masses of different isotopes and their relative abundance can be measured using a mass spectrometer

• 5 operations:

  • Vaporization: vaporized sample is injected, allowing the individual atoms to be analyzed

  • Ionization: atoms are hit with high-energy electrons which knock out electrons, producing positively charged ions

  • Acceleration: positive ions are attracted to negatively-charged plates, and are accelerated by an electric field and pass through a hole in the plate

  • Deflection: accelerated cations are deflected by a magnetic field at right angles to their path

    • Amount of deflection is proportional to the charge/mass ratio

      • Ions with smaller mass deflected more than heavier ions

      • Ions with higher charges deflected more as they interact more efficiently with the magnetic field

  • Detection: cations of a particular mass/charge ratio are detection and signal sent to recorder

    • Strength of signal is a measure of # of ions with that charge/mass ratio that are detected

• Can measure the mass of individual atoms

  • Since values are so small, relative values are used (must be agreed upon)

Mass Spectrum

• Results of analysis by mass spectrometer are presented in a mass spectrum

• Horizontal axis shows the mass/charge ratio of the ions

• Vertical axis shows the relative abundance of the ions

Electromagnetic Spectrum

• All electromagnetic waves travel at the same speed (c) but can be distinguished by their different wavelengths (λ)

• Frequency: (f) number of waves that pass a given point per second

  • Shorter wavelength = higher frequency

• Relation: c = (f)(λ)

• White light is a mixture of light waves of different wavelengths

• Continuous spectrum: containing all of the wavelengths of a given range

  • Sunlight through a prism

• Line spectrum: containing some colours of the continuous spectrum missing (not all)

• Absorption spectrum: contains dark lines/gaps in the spectrum corresponding to wavelengths absorbed by the gas

• Emission line spectrum: produced if a high voltage is applied to the gas

  • Colours present are missing from the absorption spectra

Hydrogen Spectrum

• Gives out energy when an electron falls from a higher to a lower energy level

• Produces visible light when the electron falls to the second energy level (n = 2)

• Transitions to the first energy level produces ultraviolet radiation

• Infrared radiation is produced when an electron falls to the third or higher energy levels

• Lines in diagram converge at higher energies as the energy levels inside the atoms are closer together

• When an electron is at the highest energy (n = ∞), it is no longer in the atom and it has been ionized

• Ionization energy: energy needed to remove an electron from the ground state of each atom in a mole of gaseous atoms, ions, or molecules

  • Patterns in successive ionization energies:

    • Decreases down a group

    • Increases from left to right

    • Increase in successive ionization charges

    • Jumps when electrons are removed from levels closer to the nucleus

    • First 3 ionization energies involve removing electrons from the 3rd level

    • Electron is removed from the second level for the fourth ionization energy

      • Electron closer to nucleus and is more exposed to the nucleus’ positive charge and needs significantly more energy to be removed

Electron Sublevels

• Understood if electron is treated as a wave and not as a particle

• Each level can hold a max of 2n² electrons

  • n² atomic orbitals available at the nth level

Atomic Orbitals

• Can give a picture of where the electron may be, thus using orbitals; highlights distinction between Bohr’s orbits and wave descriptions

• Each orbital holds a max of 2 electrons

• Electrons in an orbital spin – 2 electrons can occupy an orbital if they have opposite spins

• Denser dots = higher chance electron is there

• s sublevels hold a max of 2 electrons

  • Spherical

• p sublevels hold a max of 6 electrons

  • 3 p atomic orbitals arranged at right angles with the nucleus at the centre

• d sublevels hold a max of 10 electrons

  • 5 d atomic orbitals

• f sublevels hold a max of 14 electrons

  • 7 f atomic orbitals

Pauli Exclusion Principle

• No more than 2 electrons can occupy any one orbital, and if two electrons are in the same orbital they must spin in opposite directions

Waves & Particle Models

• 2 models to explain scientific phenomena: wave model & particle model

  • Light can be described by its frequency (wave characteristic) or by its photons (particle characteristic)

    • Both properties are related by Planck’s equation (E_photon = hf)

    • Neither model gives a complete explanation of light’s properties; both models are needed

• Quantum theory suggests it is preferable to think of an electron to have wave properties

Uncertainty Principle

• Problem with the Bohr model is that it assumes the electron’s trajectory can be precisely described

• Now known to be impossible as any attempt to measure its position will disturb its motion

  • To focus radiation to locate an electron gives it a random kick which causes it to scramble to another direction

• According to Heisenberg’s Uncertainty Principle, we cannot know where an electron is at any given moment; best to hope for is a probability picture of where it may be

Aufbau Principle

• Electron configuration of the ground state can be determined by this principle

• States that electrons are placed into orbitals of lowest energy first

• Boxes are used to represent atomic orbitals with single-headed arrows representing the spinning electrons

• Hund’s Rule: every orbital in a subshell is singly occupied with one electron with the same spin before any one orbital is doubly occupied

Electron Configuration of Ions

• In cations, electrons are lost from the outer sublevels

• Electron in a doubly occupied orbital is repelled by its partner with the same negative charge, thus it is easier to remove here than remove electrons in a singly occupied orbital

• When positive ions are formed for transition metals, the outer 4s electrons are removed before the 3d electrons

Electron Configuration & Periodic Table

• Position is based on the occupied sublevel of highest energy in the ground-state atom

• Periodic arrangement also reflected by patterns in first ionization energies:

  • Increase from left to right as the nuclear charge increases

  • As electrons are removed from the main energy level, there is an increase in the force of electrostatic attraction between the nucleus and outer electrons

  • Decrease to a lower level at the start of the next period (new energy level) that is farther away from the nucleus 

Chapter 3

• Periodic table demonstrates trends/patterns of elements’ periodicity

• Proposed in 1869 by Dmitri Mendeleyev who made predictions about the elements and left gaps for elements that were not discovered yet

• Periods: left to right, increasing in atomic mass

  • How many energy levels

• Groups: up to down, increasing in atomic radius

  • How many valence electrons

Periodicity in Physical Properties

• Effective nuclear charge

  • Nuclear charge given by atomic number, increasing by one (+ proton)

  • Valence electrons do not experience the full attraction of this charge as they are shielded from the nucleus and are repelled by inner electrons

    • Presence of inner electrons reduces the attraction of the nucleus and the valence electrons

    • Effective charge experienced by the valence electrons is less than the full nuclear charge

  • Ie, in sodium, nuclear charge = 10, but the valence electron is shielded from this charge due to the other 10 electrons in the inner energy levels 

  • Effective charge increases with the nuclear charge as there is no change in the # of inner electrons

  • Effective nuclear charge by the outer electrons remains the same in a group

• Atomic Radius

  • Atomic radius (r) is measured as half the distance between neighbouring nuclei

    • Can be considered the distance from the nuclear to the outer electrons

  • Increase down a group as # of occupied electron shells increases

  • Decrease across a period

    • Attraction between the nucleus + outer electrons increases as the nuclear charge increases – leads to general decrease in atomic radii 

• Ionic Radius

  • Cations are smaller than parent atoms

    • Formation of cations involves the loss of electrons from the outer shell

  • Anions are larger than parent atoms

    • Formation of anions involves the addition of electrons to the outer shell

    • Increased electron repulsion between valence electrons cause them to move further apart – increases outer shell’s radius

  • Ionic radii decrease from groups 1-4 for cations

    • Due to increase in nuclear charge with atomic number across the period

    • Increased attraction between the nucleus and electrons pulls the outer shell closer to the nucleus

  • Ionic radii decrease from groups 4-7 for anions

    • Due to increase in nuclear charge across the period

  • Ionic radii increase down a group as the # of electron shells increase

• Ionization Energies

  • First ionization energies: measure of the attraction between the nucleus + valence electrons

  • Measured directly and are properties of gaseous atoms

  • Ionization energies increase across a period

    • Increase in nuclear charge causes an increase in attraction between outer electrons and nucleus, making the electrons more difficult to remove

  • Ionization energies decrease down a group

    • Nuclear energies increase so effective nuclear charge is about the same due to the shielding of inner electrons

    • Increased distance between valence electrons and nucleus reduces the attraction

  • Group 3 elements (ns² np¹) have lower first ionization energies than group 2 elements (ns²), since p orbitals have higher energy than s orbitals

• Electronegativity

  • Measure of an atom’s ability to attract electrons in a covalent bond

  • Measure of the attraction between the nucleus and outer electrons (bonding electrons)

  • Property of an atom of a molecule derived indirectly from experimental bond energy data

  • High electronegativity = strong electron pulling power 

  • Low electronegativity = weak electron pulling power 

  • Increases from left to right due to nuclear charge increase – increased attraction between nucleus and bonding electrons

  • Decreases going down a group as bond electrons are farthest from the nucleus – reduced attraction

  • Top right = most electronegative, bottom left = least electronegative

• Melting Point

  • Melting points decrease down group 1

    • Metallic structures held together by attractive forces between delocalized outer electrons and cations – attraction decreases with distance

  • Melting points increase down group 7

    • Elements have structures held together by van der Waals’ intermolecular forces

      • Increase with the number of electrons in the molecule

  • Rise across a period and reach a maximum at group 4

  • Fall to a minimum at group 0

Periodicity in Chemical Properties

• Group 0: Noble Gases

  • Colourless gases, monatomic (occur as single atoms), unreactive due to inability to form an ion

  • Highest ionization energies 

  • No anions formed as extra electrons would have to be added to an empty outer shell where they would experience a negligible effective nuclear force, with protons shielded by an equal number of inner electrons

• Group 1: Alkali Metals

  • Silvery metals, too reactive to be found in nature, stored in oil to prevent contact with water/air

• Reactivity increases down the group as elements with higher atomic number have the lowest ionization energies

• Ability to conduct electricity is due to mobility of outer electron

• Reaction with water:

  • Produces hydrogen and the metal hydroxide

    • Lithium floats and reacts slowly; releases hydrogen but same shape

    • Sodium reacts with vigorous release of hydrogen; heat produced can melt the unreacted metal; forms a small ball on the surface

    • Potassium reacts more vigorously to produce enough heat to ignite the produced hydrogen

      • Produces lilac coloured flame and moves excitedly to water’s surface

• Group 7: Halogens

  • Exist as diatomic molecules

• Reactivity decreases down the group as the atomic radii increases and attraction for other electrons decreases

• Reactivity due to readiness to accept electrons

• Nuclei have high effective charge and exert a strong pull on any electron from other atoms

  • Electron then occupies the outer energy level to form a stable octet

• Attraction greatest for smallest atom; fluorine, most reactive nonmetal

• Reaction with alkali metals:

  • Form ionic halides

    • Halogen gains one electron from group 1 element to form a halide ion X^–

    • Resulting ions both have stable octet 

    • Most vigorous reaction is between furthest apart elements (francium + fluorine)

• Displacement reactions:

  • Relative reactivity can be seen by placing them in direct competition for an extra electron

  • Chlorine nucleus has stronger attraction to electron than a bromine nucleus due to chlorine’s smaller atomic radius and takes bromine’s electron

    • Chlorine gains an electron to form the chloride ion, while bromide ion loses an electron to form bromine

      • Distinguish between bromine + iodine by shaking with a hydrocarbon solvent

        • Iodine = violet solution

        • Bromine = dark orange solvent

• Halides

  • Halogens form insoluble salts with silver

  • Adding a solution with a halide to a solution with silver ions produces a precipitate 

Bonding of Period 3 Oxides

• Transition from metallic to non-metallic is displayed by the bonding of period 3 oxides

• Ionic compounds are made of metal + non-metal, so the oxides of these elements (ie, Na + Al) have giant ionic structures

• Covalent compounds are made of nonmetal + nonmetal, so oxides of ie, phosphorus, sulfur, and chlorine are molecular covalent

  • Oxide of silicon (metalloid) has a giant covalent structure

• Ionic character of a compounds depends on the difference in electronegativity between its elements

• Ie, oxygen has electronegativity of 3.5, so the ionic character of the oxides decrease from left to right

• Oxides become more ionic down a group as electronegativity decreases

• Conductivity of molten oxides gives an experimental measure of their ionic character

  • Only conduct electricity in the liquid state when ions are free to move

    • Maximum oxidation # of a period 3 element corresponds to group #

Acid-base Character of Period 3 Oxides

• Linked to their bonding

• Metallic elements (ionic oxides) are basic

• Nonmetal oxides (covalent) are acidic

• Aluminum oxide (can be considered an ionic oxide with some covalent character) shows amphoteric properties reacting with both acids and bases

  • Amphoteric: substance with the ability to act as a base or acid

• Basic oxides:

  • Sodium and magnesium oxides dissolve in water to form alkaline solutions (neutral in pH) due to the presence of hydroxide ions

  • Reacts with an acid to form a salt and water

  • Oxide ion combines with two H⁺ ions to form water

• Acidic oxides:

  • Nonmetallic oxides react with water to form acidic solutions

  • Phosphorus (v) oxide reacts with water to produce phosphoric (v) acid

  • The further right you go on the table, the more acidic compounds are

• Amphoteric oxides:

  • Aluminum oxide does not affect the pH when it is added to water since it is insoluble

  • With amphoteric properties, it shows both acid and base behaviour 

    • Behaves as a base when it reacts with acids

    • Behaves as an acid when it reacts with bases

Trends Across Period 3

• Bonding of period 3 chlorides

  • Demonstrate similar periodic pattern in their chemical and physical properties

  • Transition from ionic to covalent occurs earlier than with the corresponding oxides as chlorine is less electronegative than oxygen

  • Aluminum oxide is ionic but aluminum chloride is considered covalent as shown by its low conductivity in liquid form

• Reaction of period 3 chlorides with water

  • Reaction of chlorine with water

    • Chlorine reacts slowly with water in a reversible reaction to produce a mixture of hydrochloric and chloric (I) acids = disproportionation reaction as chlorine is simultaneously oxidized and reduced

  • Hydration of ionic chlorides

    • Metallic elements form ionic chlorides, which break up their lattice structure when they dissolve

    • Cations are attracted by the partially charged negative oxygen atom in the water molecules

    • Anions are attracted by the partially charged positive hydrogen atoms in the water molecules

    • Ions separated from the lattice in this way become surrounded by water molecules are said to be hydrated

    • Since resulting solution contains free ions, it can conduct electricity

  • Hydrolysis of covalent chlorides

    • Hydrolysis: chemical breakdown of a compound due to reaction with water

    • Covalent chlorides are broken up or hydrolysed when added to water

  • Hydrolysis of aluminum chloride

    • Aluminum chloride dissociates into ions when added to water

    • Aluminum ion has a high charge density due to its relatively high charge and small ionic radius, attracting water molecules

    • Water molecules form a dative covalent bond with the ion to form an octahedral complex ion

      • Complex is formed when a central ion is surrounded by molecules or ions with a lone pair of electrons

      • Has an independent existence, as the surrounding species or ligands are attached via a dative covalent bond

      • Identified by the use of square brackets

      • Dative covalent bond: both electrons in a covalent bond come from the same atom

      • Ligand: ion or molecule attached to a metal atom by a dative covalent bond

  • Hydrolysis of silicon and phosphorus chlorides

    • Silicon chloride reacts with water to produce hydrochloric acid and insoluble silicon dioxide

    • Both phosphorus chlorides produce acidic solutions due to the formation of hydrochloric acid and the corresponding phosphoric acid

      • Liquid phosphorus (III) chloride produces phosphoric (III) acid

      • Solid phosphorus (v) chloride produces phosphoric (v) acid

First-Row D-Block Elements

• Characteristics properties:

  • Electron configuration

    • Relatively small range in atomic radii due to corresponding small increase in effective nuclear charge by the outer 4s electrons

      • Small increase in effective nuclear charge accounts for small range in first ionization energies 

    • Increase in nuclear charge due to the added proton is opposed by the addition of an element in an inner 3d sublevel

    • Explains ability of transition metals to form alloys

      • Atoms here can be replaced by atoms of another without too much disruption of the solid structure

    • Unusual electron configuration of chromium and copper are due to stability of the half-filled and filled 3d sublevel

  • Physical properties:

    • High electrical and thermal conductivity

    • High melting point

    • Malleable (easily beaten into shape)

    • High tensile strength (can hold large loads without breaking)

    • Ductile (can be easily drawn into wires)

      • All properties due to strong metallic bonding 

      • Since 3d and 4s electrons are close in energy, they are all involved in bonding and form part of the delocalized sea of electrons which holds the metal lattice together

      • These account for the strength of the metallic bond and high electrical conductivity

      • Smaller atomic radii account for their higher densities

  • Chemical properties:

    • Form compounds with more than one oxidation number

    • Form a variety of complex ions

    • Form coloured compounds

    • Act as catalysts when either elements or compounds

Scandium and Zinc - NOT TRANSITION METALS

• Do not form coloured solutions

• Are part of the d-block elements but ARE NOT transition metals

• Do not display above characteristics:

  • Metals show one oxidation state in their compounds

• Reason for exception is due to the electronic configuration of their ions and lack of a partially filled d orbital

Explanation of Variable Oxidation Number of Transition Elements

• Transition metals have a wide range of oxidation numbers in their compounds

• Should be contrasted with the s-block metals which show only the oxidation state corresponding to their group number in their compounds

• Difference in behaviour can be related to patterns in successive ionization energies

• Ca³⁺ ion is energetically unstable due to large jump in ionization as third electron is removed from a 3p orbital

• Increase in successive energies for titanium is gradual as 3d and 4s orbitals are close

• All transition metals show both the +2 and +3 oxidation states

  • M³⁺ ion is the stable state for elements from scandium to chromium, but the M²⁺ state is more common for the later elements

  • Increased nuclear charge of later elements make it difficult to remove a 3rd electron

• Maximum oxidation state increases by +1 and reaches a maximum at manganese

  • States correspond to the use of both the 4s and 3d electrons in bonding

  • Maximum oxidation state decreases by -1

• Oxidation states above +3 generally shows covalent behaviour

• Compounds with higher oxidation states tend to be oxidizing agents

Complexes

• Ligands

  • Have at least one atom with a lone pair of electrons to form a dative covalent bond with a central metal atom

  • Number of dative covalent (coordinate bonds) from the ligands to the central ion is called the coordination number

  • In aqueous solution, water molecules act as ligands but can be replaced through ligand exchange

• Ability to form complex ions is a characteristic of transition metal ions

  • Relatively high charge and small size of transition metal allows them to attract the ligand’s lone pair of electrons

Colour of Transition Metal Ion Complexes

• Since complexes often have distinctive colours, they can be used in qualitative analysis

• Colour can be related to the presence of partially filled d orbitals 

• Ion Sc³⁺ is colourless since the 3d sublevel is empty

• Ion Zn²⁺ is colourless since the 3d sublevel is full

• Appear coloured because they absorb visible light

  • Ions absorb some of these colours

• Absorb light as the d orbitals split into two sublevels

  • D orbitals in an isolated transition metal atom are said to be degenerate as they all have the same energy

  • In the electric field produced by the ligand’s lone pair of electrons, they split into 2 sublevels

• Colour of complex depends on:

  • Nuclear charge and identity of the central metal atom

  • Charge density of the ligand

  • Number of d electrons present and oxidation number of the central ion

  • Shape of complex ion

    • Electric field created by the ligand’s lone pair depends on the geometry of the complex ion 

• Resonance structures are a set of two or more Lewis Structures that collectively describe the electronic bonding of a single polyatomic species including fractional bonds and fractional charges