AOS1: The Periodic TableAtomic Theory RevisionMass number: Total number of protons and neutrons in the nucleus of a give isotopeIsotope: Atoms of a given element that have different numbers of neutrons

History

0222222220200!"#$%&'()*"+'*,"-."/$&&*%

01

2.4

In the current model of the atom: electrons are thought to occupy regions of space or orbitals

around the nucleus the maximum number of electrons in an orbital is two orbitals of similar energy are grouped in subshells that are

labelled s, p, d and f subshells are energy levels within the major shells

the number of subshells in a shell is the same as the shell number, and the lowest energy subshell is always an s-subshell

electrons generally fi ll shells and subshells of lowest energy fi rst, the order of fi lling being 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p <… This is called the ground state electronic confi guration. If one or more electrons are present in a higher energy subshell when one of lower energy is not fi lled, then the atom is in an excited state.

summary

14 Copy and complete the following table by using the examples given to write the electronic confi guration of each the atoms in its electronic ground state.

Element (atomic number)

Electronic confi guration (using the shell model)

Electronic confi guration (using the subshell model)

Boron (5) 2,3 1s22s22p1

Lithium (3)

Chlorine (17)

Sodium (11)

Neon (10)

Potassium (19)

Scandium (21)

Iron (26)

Bromine (35)

15 In terms of energy levels, what is the essential difference between the shell model and the subshell model of the atom?

key questions

Aristotle:no atoms:matter iscontinuous

Dalton:atoms

1804

Thomson:electrons

Rutherford:nuclear atom

1897 1911

Bohr:shell model

1913

Schrödinger:quantummechanics

1926

Chadwick:neutrons

1932

Figure 2.22Our model of an atom underwent substantial change during the 20th century. Will it change as much during the 21st century?

DaltonModel: A ʻbilliard shapeʼ ballExperiment: After observing different chemical reactions and looking elements

ThomsonModel: A ʻplum puddingʼ i.e. Positively charged sphere with electrons embedded in surfaceExperiment: Cathode ray experiment i.e. Different metals and gases in a tube creating a ray. When an electric field was applied, ray would move meaning that there were negatively charged particles: electrons

RutherfordModel: Electrons revolve around nucleus Experiment: Alpha particle experiment i.e. alpha particles at gold foil; some passed through and some didnʼt therefore most of the volume must be empty space

BohrModel: Electrons is shells surrounding nucleus Experiment: Emission spectra i.e. Coloured lines emitted by electrons as they move energy states indicated that there were shells

SchrodingerModel: Electrons as waves i.e. electrons move in regions of space: orbitals Experiment: Quantum mechanics, further exploring emission spectra

ChadwickModel: Discovery of nucleus Experiment: Observing that atoms were heavier than expected i.e. alpha particle experiment with beryllium: scattered particles had a neutral charge

Mendeleev arranged elements by • Increasing atomic mass• Similar properties

AOS1 & AOS2!

Electron Configuration• Electrons are grouped in shells

Shell Electrons

Ist 2

2nd 8

3rd 18

4th 2n^2

ShellsWithin an atom there are energy levels called shellsOrbitals:• ʻsʼ = 1 orbital

• Spherical shape• ʻpʼ = 3 orbitals

• Dumbbell shaped • ʻdʼ = 5 orbitals

• Complex• ʻfʼ = 7 orbitals

• Lots of balloons *Each orbital can hold a maximum of 2 electrons Electrons in the same orbital have opposing spin

Groups Group 1 - Alkali Metals • One electron in ʻsʼ subshellGroup 2 - Alkali Earth Metals • Full ʻsʼ subshellGroup 17 - HalogensGroup 18 - Noble gases

TrendsAtomic Size• Decreases across periods

• Greater pull makes it smaller• Increase down groups

• Due to more shells Ionisation EnergyThe energy required to lose an electron• Decreases down a group

• Further away means that there is less pull• Increases across a period

• More electrons

Electronegativity Ability to attract electrons• Decreases down a group

• Further away from pull

AOS1 & AOS2!

• Increases across periods• More electrons

Core Charge • Number of protons - Number of non-valence electrons

Relative Isotopic MassMass Spectrometer determines relative isotopic mass number:1. Ionisation i.e. to make a positive ion2. Acceleration i.e. same kinetic energy3. Deflection i.e. lighter ions deflected more than heavier ones4. Detection i.e. relative abundance vs. relative atomic mass

Relative Atomic Mass• The average relative isotopic mass of one atom

AR = isotope 1 (relative mass x abundance) + isotope 2 (relative mass x abundance)

100

Relative Molecular Mass• The sum of the Relative Atomic Masses of each atoms within the molecule

Relative Formula Mass • Sum of the relative atomic masses of atoms given in the chemical formula

Avogadroʼs Number 6.02 x 1023 = 1 mol

Molar Mass• Mass of one mol of any substance• Measured in grams/mol

RelationshipsM = m/n# # # n=m/M

Percentage Composition• Percentage by mass of each element within a compoundM(element1)/M(total compound) x 100/1

AOS1 & AOS2!

Empirical Formula• Simplest whole number ratio of atoms

Calculating Empirical Formula1. Write down symbols2. Write down mass3. Convert masses to moles i.e. put over molar mass4. Find the simplest whole number ratio of all atoms by dividing them by the smallest number of moles5. If necessary convert to whole numbers

Molecular Formula• Actual number of each atom present in compound

Calculating Molecular Formula1. Divide actual compound by empirical formula to get n2. Multiply empirical formula by n

AOS2: Materials Types of Bonding Metal + Non-metal = Ionic Bonding • Gives away electrons Non-metal + Non-metal = Covalent Bonding • Shares electrons Metal + Metal = Metallic • Shares i.e. Lattice with SODE

Properties of Metals All metals are:• Lustrous

• reflect light• Opaque

• do not allow light to pass through • Good conductors of heat and electricity• Malleable

• Can be beaten into a sheet• Ductile

• Can be drawn into wire

Most metals are:• Hard• Donʼt snap• High melting and boiling points • High density

AOS1 & AOS2!

!" 05000000555000#$%&'($)*!"##

5.1

Although there are exceptions, compared with non-metals, metals: have high melting temperatures conduct electricity and heat well in the solid state have a lustre

have high density are malleable and ductile are often hard.

summary

1 Where would you draw a line in Table 5.1 to divide the metals from the non-metals?

2 a Potassium is classed as a metal. In what properties is potassium similar to the metal gold? In what ways is it different?

b Identify another element in Table 5.1 that has similar properties to potassium.

c Identify another metal in Table 5.1 that has similar properties to gold.

d Where are these four metals in the periodic table?

3 a Which metals would you select if you wanted a good electrical conductor?

b What other factors not listed here might infl uence your choice?

4 Suggest some properties not included in Table 5.1 that you would need to consider before choosing between aluminium and iron for building a bridge.

key questions

5.2

Properties and structureA satisfactory model for the structure of metals will explain the properties we have discussed. It will also help chemists and materials engineers to understand why metals behave as they do and how their behaviour can be modifi ed. Each one of the properties in Table 5.2 gives some information about a metal’s structure.

Using information such as that in Table 5.2, chemists have developed a model to explain the properties of metals. Such a model needs to describe the particles in a metal and how they are arranged. We can deduce from the properties and inferences in Table 5.2, that the model must be one in which: some of the particles are charged and free to move there are strong forces of attraction between particles throughout the

metal structure.

Property What this tells us about structureMetals conduct electricity in the solid state.

Metals have charged particles that are free to move.

Metals are malleable and ductile. The forces between the particles must be able to adjust when the particles are moved.

Metals generally have high densities. The particles are closely packed in a metal.

Metals tend to have high boiling temperatures.

The forces between the particles must be strong.

Metals are lustrous or refl ective. Metals can refl ect light.

TABLE 5.2 The physical properties and inferred structural features of metals

!An electric current is a fl ow of charged particles. In solids, the charged particles are electrons. In liquids or solutions, the charged particles are ions.

Structure of Metals • Small number of electrons in their valence shell

• Low electronegativity; electrons can be released • Metal cations arrange into a closely packed lattice

• Attraction between SODE and cations

Qualities of Metals • High melting and boiling points

• Due to strong electrostatic attraction• Electrical conductivity

• Free electrons • Good conductors of heat

• SODE bump into each other and transfer energy• Malleable and ductile

• Layers held together by SODE• Lustrous

• Free electrons reflect light• Dense

• Closely packed cations

AlloysInvolves adding small amounts of other substances to metals to change the properties Substitutional Alloys• Some of the host metals atoms are replaced by other metal atoms of similar sizeInterstitial Alloys• A small proportion of a significantly smaller atom is added

• Makes the layers harder to slide past each other

Work Harding and Heat Treatment Metal solids are compsed of many small crystals• Each individual crystal is a regular lattice of ions surrounded by SODEWork Hardening • Hammering or working cold metals causes crystals to re-arrange

• Becomes more brittle due to the smaller crystals Heat Treatment • Annealing

• Heating a metal to high temp then cooling slowly • RESULT: larger crystals = softer, but less brittle

• Quenching • Heating to a moderate temp then cooling quickly • RESULT: smaller crystals = harder, but more brittle

• Tempering • Heating the quenched metal again, but to a lower temperature• RESULT: reduces brittleness

LimitationsDoes not explain• Range of melting points and densities• Magnetic nature of some metals

AOS1 & AOS2!

05555550500 !"#$%&

'(

The metallic bonding modelWhat are the charged particles that enable metals to conduct electricity? How are they arranged? The only particles that are small enough to move through a solid lattice are electrons. If a metal atom loses one or more electrons from its outer shell, it forms a positive ion or cation.

Chemists now believe that, in a solid sample of a metal: Positive ions are arranged in a closely packed structure. This structure is

described as a regular, three-dimensional lattice of positive ions. The ions occupy fi xed positions in the lattice.

The much smaller negatively charged electrons that have been released from the outer shell of the metal atoms are free to move throughout the lattice. These electrons are called delocalised electrons because they belong to the lattice as a whole. The delocalised electrons come from the outer or valence shell. Electrons that are not free to move throughout the lattice are said to be localised. Electrons in the inner shells are localised.

The ions are held in the lattice by the electrostatic force of attraction between these positive ions and the delocalised electrons. This attraction extends throughout the lattice and is called metallic bonding.

Explaining the properties of metalsFigure 5.6 shows how the metallic bonding model is consistent with the relatively high boiling temperature, the electrical conductivity and malleability and ductility of metals.

!The electronic confi guration of magnesium is 1s22s22p63s2. When a magnesium atom loses its two valence electrons, it will still have its 12 protons but only 10 electrons with an electronic confi guration of 1s22s22p6. The resultant magnesium cation, Mg2+, has a charge of +2. Other examples of cations are Na+, K+ and Ca2+.

!A lattice is a regular arrangement of particles, similar to the way that oranges are stacked together in a greengrocer’s shop.

delocalised‘sea’ ofelectrons

positivelychargedmetal ions

+–

+ + +

+ + + +

+ + + +

–

–––––

– – – – ––

–––

Figure 5.4The metallic bonding model. A chemist’s view of how the particles are arranged in a solid sample of a metal. This diagram shows just one layer of metal ions.

Positive sodium ionsoccupy fixed positionsin the lattice.

‘sea’ ofdelocalised electrons

Na+ Na+ Na+ Na+

Na+ Na+ Na+ Na+

Na+ Na+ Na+ Na+

Figure 5.5 A sodium metal lattice. Each sodium atom provides its one valence electron to form a ‘sea’ of delocalised electrons.

!Energy must be provided to remove the valence electron from each sodium atom. However, energy is released again as the delocalised electrons are then attracted to all the positive ions in their region.

–

–

Metals have relatively high boiling temperatures.

Metals are good conductors of electrcity.

Metals are malleable and ductile.

Property Explanation

+ + + + +

+ + + + +

+ + + + +

– – –

–

–

–

–

– – – –

– –

–

–

+ + + + +

+ + + + +

+ + + + +

– – –

–

–

–

–

– – – –

– –

Strong electrostatic forces of attraction between positive metal ions and ‘sea’ of delocalised electrons holds the metallic lattice together.

Free-moving delocalised electrons will move towards a positive electrode and away from a negative electrode in an electric circuit.

When a force causes metal ions to move past each other, layers of ions are still held together by delocalised electrons between them.

+ –

+ + + ++ + + +

+ + + ++ + + +

+ + + ++ + + +

+ + + ++ + + +

Figure 5.6How the metallic bonding model can be used to explain some of the properties of metals.

Ionic Compounds• Metals donate electrons • Non-metals accept electrons Lattice is held together by electrostatic force of attraction forms a strong bond• Bond between oppositely charged cations and anions SOLID STATE: ions in fixed positionsLIQUID STATE: ions are free to move

Properties of Ionic Compounds• Solid under normal conditions

• Strong electrostatic forces • High melting and boiling points

• Electrostatic forces• Brittleness

• Forces are directional; ions of same charge are aligned = repulsion • Non-conduction of electricity in solid state

• Fixed positions • Electrical conduction when molten or dissolved in water

• Free to move as a liquid • Regularly shaped crystals

• Regular arrangement of particles

OtherMetal + Water = Metal Hydroxide + Hydrogen Gas

Electrical Conductivity

Melting Point Type of Bond

Ionic Compounds Only when molten or dissolved in water

HIGH Strong electrostatic force. Free electrons in molten or dissolved state

Metals In solid state and molten

HIGH Metallic bond and electrostatic attraction holds lattice together

Covalent Molecular Compounds

No LOW Covalent bonds and weak intermolecular forces make very weak

Covalent Lattices Only covalent layer lattices

VERY HIGH Strong covalent bonds hold togetherFree electrons to conduct

GlossaryIsomer: Compounds with the same formula but different arrangements of atoms in the molecule Isotope: Elements that have equal number of protons but different numbers of neutrons and hence different relative atomic massHDPE: Has little branching, giving it stronger intermolecular forces. Also harder and more opaqueLDPE: More branching, giving it weaker intermolecular forces. Also softer and less denseAllotropes: Different physical forms of the same element Compound: A pure substance composed of atoms of two or more elements which have chemically combined Cross-links: Covalent links between polymer chainsHydracarbon: An organic compound consisting of Hydrogen and Carbon

AOS1 & AOS2!