Unit 2 Atomic Structure and Periodicity

1.2.1 1.2.2 1.2.3 1.3.1 1.3.2 1.3.3 1.3.4 1.3.5 1.3.6 1.3.7 3.1.1 3.1.2 3.1.3 3.1.4 3.1.5 3.1.6 3.1.7

The atomic model

Atoms contain a positively charged, dense nucleus composed of protons and neutrons (nucleons). Negatively charged electrons occupy the space outside the nucleus.

Ions

An atom or molecule with a net electrical charge. Occurs when the number of protons is no longer balanced by the number of electrons.

Cation

An atom loses an electron to form a positive ion as the number of protons is now greater than the number of electrons.

Anions

Formed when atoms gain electrons.

Isotopes

Atoms of the same element with different types of neutrons and so they have different mass numbers. Have the same chemical properties as the difference in the number of neutrons makes no difference to how atoms react.

Mass sepctra

Mass spectra are used to determine the relative atomic masses of elements from their isotopic composition.

Mass spectrometer

A mass spectrometer can measure the mass and abundance of isotopes. This information can then be presented in the form of a mass spectrum. Multiply the mass/charge by the % abundance and divide the total sum by 100.

Emission spectrum

Produced by atoms emitting photos when electrons in excited states (higher energy level) return to lower energy levels.

Color

Atoms of different elements give out light of distinctive colors.

Relationship between wavelength and the emission spectrum

Different colors of visible light have different wavelengths. Towards the red end there is less energy, thus the wavelengths will be longer.

Relationship between frequency and the emission spectrum

The shorter the wavelength, the higher the frequency.

Energy

The energy of a photon is proportional to the frequency of the radiation. Energy increases as you go down the emission spectrum.

Difference between a continuous and line spectrum

A continuous spectrum contains all wavelengths of light in a certain range. A line spectrum contains light of only a certain wavelength that is emitted or absorbed.

Emission spectrum of hydrogen

When a hydrogen atom is excited its electrons jump to a higher energy level. As the electron returns to lower energy levels it emits photons with energies specific to the energy level transitions. These emitted photons correspond to specific wavelengths, which appear as distinct coloured lines in the spectrum.

Series of lines. Electrons transfer from a higher energy level to a lower energy level. Energy levels converge at higher frequency and energy and shorter wavelength. Transition to 3rd IR, transition to 2nd visible light, transitions to 1 UV.

Line emission spectrum of hydrogen

Provides evidence for discrete energy levels. The transitions to the first energy level (n=1) correspond to the highest energy change and are in the ultraviolet region of the spectrum. The lines converege at higher energies because the energy levels inside the atoms are closer together.

Ionization energy

Ionization energy is the energy needed to remove one electron from the ground state of one mole of gaseous atoms, ions, or molecules. Transition to from n = 1 to n = infinity. It is a positive enthalpy change (endothermic)

Emission spectra

• electrons release energy

• electrons transition down from higher to lower energy levels

• energy released is equal to the difference in energy of the 2 energy levels the electron is transitioning between.

Absorption spectra

• electrons absorb energy

• electrons transition up from lower to higher energy levels

• energy absorbed is equal to the difference in energy of the 2 energy levels the electron is transitioning between.

Energy levels in the atom

• converge at higher energies, as they get farther from the nucleus

• energy levels are farther apart in energy closer to the nucelus

Describe the appearance of the hydrogen emission spectrum

Series of colored lines which converge at higher energies.

Explain how the hydrogen emission spectrum is created

• when electrons transition from higher to lower energy levels, they release electromagnetic radiation

• transitioning to the n = 2 energy level releases visible light

• the energy of light is the difference in energy of the energy levels

Explain the model of electron configurations (the Bohr Model) that was determined by the hydrogen emission spectrum

• electrons can only exist at specific energy levels

• Energy levels converge at higher energies (farther from the nucleus)

Limit of convergence

The limit of convergence is a theoretical value for ionisation energy which can be experimentally measured by physically removing electrons from atoms.

Electron affinity

The amount of energy released when a mole of a gaseous atom gains a mole of electron. It is a negative enthalpy change (exothermic)

The Bohr model of the atom

• electrons exist in discrete energy levels

• lower energy levels are closer to the nucleus and are more stable

• energy levels converge at higher energies

• electrons can only exist at distinct energy levels

Valence electrons

Outer shell electrons determined by the group of the periodic table an element is in.

Electron orbitals

3-dimensional shapes produced by the mathematical wave equations used to describe the probability of electron motion. They tell us where electrons are most likely to exist around an atom. An orbital can hold two electrons of opposite spin. There are 4 types of orbitals: s, p, d and f.

Energy sublevels

Within the main energy levels of electrons, electrons can occupy different orbitals. These orbitals have slightly different energies from one another, so they can be referred to as energy sublevels.

Order of increasing energy of orbitals within a given energy level

s < p < d < f

Mathematical patterns of orbitals

• Number of sublevels (types of orbitals) = n

• Total number of orbitals = n²

• Total number of e^- = 2n²

where n = principle energy level

Pauli exclusion principle

An orbital can hold only two electrons of opposite spin. Spin is an important factor in electron-electron interactions because electrons can occupy the same region of space despite their mutual repulsion if they spin in opposite directions.

First energy level

• The first energy level has one 1s orbital.

• This energy level can hold a maximum of two electrons.

Second energy level

• The second energy level is split into two sublevels

• The 2s sublevel is one 2s orbital and can hold a maximum of two electrons

• The 2p sublevel is three 2p orbitals and can hold six electrons

• The 2s orbital has the same symmetry as a 1s orbital but extends over a larger volume. The electron in a 2s orbital is also on average further from the nucleus and has higher energy.

• The three 2p atomic orbitals have equal energy and are said to be degenerate. They all have the same dumbbell shape with only different orientation (arranged at right angels to each other with the nucleus at the center)

Third energy level

• made up of three sublevels: 3s, 3p and 3d

• the d sublevel is made up of five d atomic orbitals (10 electrons)

Degenerate orbitals

Degenerate orbitals of the same energy form a sublevel: three p orbitals form a p sublevel, five d orbitals form a d sublevel, a single s orbital makes up an s sublevel.

Max amount of electrons per sublevel

• s: 2 electrons

• p: 6 electrons

• d: 10 electrons

• f: 14 electrons

Aufbau principle

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

Hund’s third rule

Electrons are placed into separate orbitals first because this configuration minimizes the mutual repulsion between them.

Relative energy 3d and 4s

• the 3d and 4s levels are very close in energy and their relative seperation is sensitive to inter-electron repulsion

• Electrons first fill 4s before 3d

• However, electrons are first lost from 4s sublevel before the 3d sublevel when transition metals form their ions, as once the 3d sublevel is occupied the 3d electrons push the 4s electrons to higher energy.

Exceptions for orbitals

• Chromium has the electron configuration [Ar]3d⁵4s¹

• Copper has the electron configuration [Ar]3d¹⁰4s¹

Electron configuration of ions

• Positive ions are formed by the loss of electrons, these electrons are lost from the outer sublevels.

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

• The electron configurations of negative ions are determined by adding the electrons into the next available electron orbital.

Electron configuration and the periodic table

• the position of an element in the periodic table is based on the occupied sublevel of highest energy in the ground-state atom.

Successive ionization energies

On a succesive ionization energy graph, there are jumps between the ionization energies as electrons are removed from lower energy levels, closer to the nucleus. Since the electrons in those levels are more exposed to the positive charge of the nucleus it needs significantly more energ to be removed.

Trend in first ionization energy across periods

• ionization energies increase from left to right across a period as the nuclear charge increases. There is an increase in the force of electrostatic attraction between the nucleus and outer electrons.

• Ionization energy decreases down a group as a new energy level, which is further from the nucleus is occupied. Less energy is required to remove outer electrons tht are further from the attractive pull of the nucleus.

Metallic elements

• Found on the left hand side of the table in the s block, in the central d block and in the island of the f block.

• A small number of metals such as Al and Pb are also found on the left of the p block.

Periods and groups

• period number shows the outer energy level that is occupied by electrons

• elements in a group have a common number of valence electrons

Alkali metals

Group 1. Silvery metals and are too reactive to be found in nature.

physical properties:

• good conductors of electricity and heat

• low densities

• shiny grey surfaces when freshly cut with a knife

chemical properties:

• very reactive

• form ionic compounds with non-metals

Halogens

Group 17, exist as diatomic molecules X₂

physical properties:

• they are coloured

• show a gradual change from gases, to liquid and solids

Chemical properties:

• very reactive non-metlas

• reactivity decreases down the group

• form ionic compounds with metals and covalent compounds with other non-metals

Noble gases

Group 18

Transition metals

Have an incomplete d-sublevel in one or more oxidation state

What properties compose the periodicity of elements?

• atomic and ionic radii

• electronegativity

• ionization energy

• electron affinity

• Effective nuclear charge

Atomic radius trends

• atomic radius increases down a group as the number of occupied electron levels increases

• atomic radius decreases across a period as the attraction between the nucleus and the outer electrons increases as the nuclear charge increases

Ionic radius trends

• increases down a group and decreases across a period

• positive ions are smaller than their parent atoms as the formation of positive ions involves the loss of the outer energy level.

• negative ions are larger than their parent atoms as the formation of negative ions involves the addition of electrons into the outer energy level and the increased electron repulsion causes them to move further apart increasing the radius.

• ionic radius decreases from groups 1 to 14 for positive ions due to increase in nuclear charge pulling the outer energy level closer to the nucleus.

• ionic radii decreases from groups 14 to 17 for negative ions due to the increase in nuclear charge across the period. The positive ions are smaller than the negative ions as they former have only two occupied electron principal energy levels and the latter has 3.

• ionic radii increases down a group as number of electron energy levels increases.

Ionization energy trends

• ionization energies increase across a period as the increase in effective nuclear charge increases the attraction between the outer electrons and the nucleus making the electrons more difficult to remove

• ionization energies decrease down a group as the electron removed is from the energy level furthest from the nucleus. Although the nuclear charges increase, the effective nuclear charge is about the same due to the shielding by the inner electrons so the increased distance between the electron and the nucelus reduces the attraction betwen them.

Electron affinity trends

• electron affinity decreases down a group and increases across a period

• Group 17 elements have incomplete outer energy levels and a high effective nuclear charge so they attract electrons the most

• Group 2 elements have an electron configuration ns², so the added electron must be placed into a 2p orbital which is further from the nucleus and so experiences reduced electrostati

Electronegativity trends

• electronegativity increases across a period due to the increase in nuclear charge resulting in increased attraction between the nucleus and the bond electrons

• electronegativity decreases down a group as the bonding electrons are furthest from the nucleus so there is reduced attraction.

Electronegativity

A measure of the ability of its atoms to attract electrons in a covalent bond. Also a measure of the attraction between the nucleus and its outer electrons (bonding electrons). An element with high EN has strong electron puling power and an element with low EN has weak pulling power.

Bonding of period 3 oxides

• oxides of Na to Al have giant ionic structures

• oxides of P, S and Cl are molecular covalent

• Oxide of Si (metalloid), has a giant covalent structure

• Ionic character of the period 3 oxides decrease from left to right as the electronegativity values approach this value.

When do oxides conduct electricity

Oxides only conduct electricity in the liquid or aqueous state when the ions are free to move

Acid-base character of the period 3 oxides

• Na and Mg basic

• Al amphoteric

• Si, P, S, Cl acidic

Basic oxides

• Sodium oxide and magnesium oxide dissolve in water to form alkaline solutions owing to the presence of hydroxide ions

• A basic oxide reacts with an acid to form a salt and water

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

Non-metallic oxides are acidic

• react readily with water to produce acidic solutions

• carbon dioxide dissolves in water to form the weak acid, carbonic acid

• sulfur trioxide reacts with water to produce sulfuric(VI)acid

• sulfur dioxide reacts with water to produce sulfuric (IV) acid

Amphoteric oxides

• aluminium oxide does not affect te pH when it is added to water as it is essentially insoluble

• It shows both acid and base behavior

  • behaves as a base when reacts with sulfuric acid

  • behaves as an acid when it reacts with alkalis such as sodium hydroxide

Acid rain and acid deposition

• produced by non-metal oxides

• all rainwater is naturally acidic due to the presence of dissolved carbon dioxide

• sulfur dioxide dissolves in rainwater to form sulfuric(IV) or sulfurous acid

• Sulfur dioxide can be oxidized to sulfur trioxide, in the atmosphere, which then dissolves in rainwater to form sulfuric (VI) acid.

• Nitrogen monoxide is produced from internal combustion engines where the burning of the fuel releases heat energy that causes nitrogen and oxygen from the air to combine

  • a similar reaction gives rise directly to the brown gas, nitrogen dioxide which also forms from the oxidation of nitrogen monoxide

  • nitrogen dioxide dissolves in rainwater to form a mixture of nitric (III) or nitrous acid (HNO₂) and nitric (V) acid (HNO₃)

Ocean acidification

• occurs as carbon dioxide dissolves in the oceans

• reacts with the water to form carbonic acid

• the carbonic acid then reacts to form hydrogen carbonates or carbonates

• this leads to higher ocean acidity, mainly near the surface which inhibits shell growth in marine animals and is the cause of reproductive disorders in some fish

Oxidation state

A number assigned to an atom to show the number of electrons transferred in forming a bond. It is the charge that the atom would have if the compound were composed of ions. A measure of how electrons are distributed.

Rules to assign oxidation states

• atoms in the free element have an oxidation state of zero

• in simple ions, the oxidation state is the same as the charge on the ion

• the oxidation states of all atoms in a neutral compound must add up to zero

• the oxidation states of all the atoms in a polyatomic ion must add up to the charge on the ion

• the usual oxidation state for an element in a compound is the same as the charge on its most common ion

• F has the oxidation state of -1 in all of its compounds as it is the most electronegative element

• O has the oxidation state of -2 except in H₂O₂ where it is -1 or bonded to Flourine where it is +2

• Cl has the oxidation state of -1 except when bonded to F or O

• H has the oxidation state of +1 except when bonded to the group 1 or 2 metals when it forms ionic hydrides (-1)

• the oxidation state of a transition metal in a complex ion can be worked out from the charge on the ligands.

Explain how the discontinuities that occur in the trend of increasing first ionization energy across a period provide evidence for the existence of energy sublevels

Group 13 elements have lower first ionization energies than group 2 elements because p orbitals have higher energy than s orbitals. The drop between groups 15 and 16 occurs as the electron removed from the group 16 element is taken from a doubly occupied p orbital. The group 16 electron is easier to remove because it is repelled by its partner.

Rutherford gold foil experiment conclusions

Most He²⁺ passing straight through:

• most of the atom is empty space

Very few He²⁺ deviating largely from their path:

• the positive charge is concentrated in the nucleus

Why transition metals form coloured compounds

Why is the ionic radius of Nitrogen much larger than its atomic radius?

It has more electrons than protons and the electrons are held less tightly together.

Effective nuclear charge

The net positive charge pulling these electrons towards the nucleus

State based on successive ionization energies, why transition metals can form ions with multiple charges

Successive ionization energies change very gradually in transition metals since s and d electrons are very close in energy. Multiple electrons can be removed from the d-sublevel.

Why do most magnesium compounds tested in a school laboratory show traces of yellow win the flame even tho they dont produce emission or absorption lines in the visible region.

Contamination with other compounds