Atomic Structure, Bonding and Periodicity Contents • • • • Atomic Structure Amount of Substance Bonding Periodicity Atomic Structure • • • • • Fundamental Particles Mass Number and Isotopes Mass Spectrometer Electron Arrangement Trends Ionisation Energies Fundamental Particles Atoms have a small central nucleus made up of protons and neutrons around which there are electrons. Neutron Atomic Particle Electron + + Relative Charge neutron 1 0 proton 1 +1 negligible -1 electron Proton Relative Mass Isotopes • Atoms of the same element always have the same number of protons. • Isotopes are atoms of the same element but with different number of neutrons. This gives rise to different mass numbers. • Relative abundance is the amount of each isotope as the percentage for that element occurring on the Earth Mass number: this is the number of protons and neutrons (A) 4 2 He Helium Atomic number: this is the number of protons (Z) Mass Spectrometer The sample is put in via an injector An electron gun ionises the atoms Ionisation Charged plates accelerate positive ions acceleration •The mass spectrum gives the relative masses of each isotope and the abundance of each isotope. •The relative atomic mass is the weighted average mass of an atom of an element compared with 1/12 of the mass of 12C. •We can calculate this from mass spectra A curved magnet deflects the positive ions. The lighter ions are deflected more deflection The ions are detected to produce a mass spectrum detection Example: Copper 69% of copper atoms have a relative mass of 63. 31% of copper atoms have a relative mass of 65. The weighted average is calculated as follows: (0.69 x 63) + (0.31 x 65) =63.62 Arrangement of Electrons • Energy Levels or Shells – The simplest model of electrons has them orbiting in shells around the nucleus. Each successive shell is further from the nucleus and has a greater energy. • Sub Shells and Orbitals – This model can be further refined by the concept of sub shells and orbitals. – Sub shells are known by letters s, p, d, and f. The s sub shell can contain 2 electrons, p 6, d 10 and f 14. – Electrons occupy negative charge clouds called orbitals, each orbital can hold only 2 electrons. Each type of shell has a different type of orbital. Arrangement of Electrons • How we write electron configurations – Electrons fill the lowest energy level first this means it is generally easy to predict how the electrons will fill the orbitals (it gets more complicated with the transition metals). Element Electron configuration H 1s He Li Be B C 1s2 1s22s 1s22s2 1s22s22p1 1s22s22p2 N 1s22s22p3 O 1s22s22p4 F 1s22s22p5 Element Electron configuration Ne 1s22s22p6 Na 1s22s22p63s1 Mg 1s22s22p63s2 Al 1s22s22p63s23p3 Si 1s22s22p63s23p3 P 1s22s22p63s23p3 S 1s22s22p63s23p4 Cl 1s22s22p63s23p5 Ar 1s22s22p63s23p6 Trends in Ionisation Energies 1000 First ionisation energies of period 3 elements Energy (Kj/mol) Energy (Kj/mol) The first ionisation energies of group 2 electrons 800 600 400 200 0 2000 1500 1000 500 0 Be Mg Ca Sr Ba Group 2 elements Going down a group in the periodic table there are more filled energy levels between the nucleus and the outer most electrons these shield the outer electrons from the attractive force of the positive nucleus. as the radius of the atom increases, the distance between the nucleus and the outer electron increases and therefore the force of attraction between the nucleus and outer most electrons is reduced. These factors mean that less energy is needed to remove the first electron from an atom at the bottom of the group compared to one at the top of the group. Na Mg Al Si P S Cl Ar Period 3 elements Going across a period of elements: there are more protons in each nucleus so the nuclear charge in each element increases, this increases the attractive force acting on the outer most electrons. there is no significant increase in shielding as each successive electron enters the same energy level as the one before. Overall more energy is needed to remove the first electron from its outermost shell. Amount of Substance • • • • Amount of Substance Calculations Balancing Equations Reacting Quantities Amount of Substance • Different atoms have different masses. 1g of carbon has far fewer atoms than 1g of hydrogen atoms. Chemists need a method of quantifying atoms. • We use a quantity called the amount of substance which is measured in moles. • One mole contains 6.02 x 1023 particles. • The relative atomic mass is the mass of one mole of that element’s atoms. • The relative molecular mass is the mass of one mole of molecules Calculations • Amount of substance, n – n = mass Mr • Solution calculations – The concentration of solutions are measured in mol dm3 • c= concentration in dm3 • V = volume in cm3 – n = Vc 1000 P r • The Ideal Gas equation • Providing that the pressure and temperature of gases are the same, equal volumes of two gases can be assumed to have the same number of moles. – pV = nRT • p is the pressure in Pa • V is the volume in m3 • T is the temperature in Kelvin • R= 8.31 JK-1 mol-1 Balancing Equations • Chemical reactions involve the rearrangement of atoms not the making or destroying of atoms. • It is necessary to make sure that you have the same amount of atoms on both sides of the equation. • State symbols can also be added to show the physical condition of the reactants and products • (s) – solid, (l) – liquid, (g) – gas, (aq) – aqueous Reacting Quantities • The numbers in a balanced equation give the ratio of the amount of each substance in the reaction. We can use this information to calculate quantities of reactants or products. • 50g of CaCO3 are heated how much CaO will be formed • First write a balanced equation: – CaCO3(s) CaO(s) + CO2(g) • Then calculate the Mr of the substances we are interested in: – CaCO3 40 + 12 + (3 x 16) = 100 – CaO 40 + 16 =56 • Calculate the number of moles of CaCO3 used. – n=mass/Mr 50/100 = 0.5 mol • From the equation we can see that one mole of CaCO3 produces one mole of CaO. Therefore 0.5 mol of CaCO3 produces 0.5 mol of CaO. • Finally calculate the mass of 0.5 mol of CaO: – n=mass/Mr, mass=Mr x n, 0.5 x 56 = 28g • Therefore 50g of CaCO3 produces 28g CaO Bonding • • • • The nature of bonds Bond polarity and the polarisation of ions Intermolecular forces Hydrogen Bonding The Nature of Bonds Covalent Bonding H H • H C • When non-metals react together both atoms need to gain electrons to obtain a full shell of electrons they do this by forming a covalent bond. The atoms are held together by shared pairs of electrons from the highest energy level of both atoms. H Ionic bonding _ + Cl Na Metallic Bonding - ++ - - + - +- + - + - + + -+ + - -+ - + Atoms lose or gain electrons to attain a complete outer shell of electrons. An ionic bond is formed when electrons are lost and gained by two or more atoms. When atoms lose electrons they become positive ions, when they gain electrons they become negative ions. It is the electrostatic forces of attraction which hold the ions together In metals, positive metal ions are held - clouds. + + These - +together- by+electron ++through the + electrons -are+free to-move - + + + - metals+conduct + - structure, + this- is why + - + electricity. Bond Polarity and the Polarisation of Ions • In reality not all bonds are perfectly covalent or ionic. To explain why we have to define a concept called electronegativity. • Electronegativity is the ability of an atom to attract the bonding electrons. • In hydrogen fluoride the fluorine atoms are much more electronegative than the hydrogen. It pulls the electrons toward it creating what is called a polar bond. • Ionic bonds can also show polarity, this can happen if the electron cloud is distorted by strong charges on one of the ions. If a cation is highly charged it will exert a strong electrostatic attraction on the anion and distort the electron cloud. If the anion has a large electron cloud it will be easily distorted. If the electron cloud is distorted there will be electron density between the two ions giving the bond some covalent characteristics • Molecules with asymmetric charge distribution are said to be polar molecules and to possess a dipole Intermolecular Forces • Permanent dipolepermanent dipole interactions occur between polar molecules. This happens when the negative end of one molecule is attracted to the positive end of another. This force is much weaker than intramolecular bonding d+ H Cl dWeak electrostatic forces d- Cl H d+ •Temporary dipole –induced dipole interactions exist between non-polar molecules and monatomic species such as the noble gases. The distribution of the electron cloud on a molecule is not constant and at any given time it can asymmetric. this confers a temporary asymmetry on the charge distribution. The molecule is said to possess a temporary dipole, this temporary dipole can induce another temporary dipole in an adjacent molecule. There is a resulting weak electrostatic force between the two molecules Hydrogen Bonding d- O H d+ H d+ d- O H d+ H d+ d- O H d+ d- O H d+ H d+ H d+ • Hydrogen bonds are a special case of permanent dipole-permanent dipole bonding. • It exists where an electronegative element such as oxygen, chlorine fluorine or nitrogen is bonded to hydrogen. • Hydrogen bonding causes stronger intermolecular bonds than would otherwise be predicted this increases the boiling point of substances such as water. Periodicity • Chemists classify elements according to their position in the periodic table. • Periodicity is the term used to describe the repeating pattern of properties observed within the periodic table. H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cu Zn Ga Ge As Se Br Kr Rb Sr Nb Mo Tc Cs Ba La Hf Fr Ra Ac s- block Y Zr Cr Mn Fe Co Ni Ta W Ru Rh Pd Ag Cd In Re Os Ir d-block Pt Au Hg Tl Sn Sb Te I Xe Pb Bi Rn Po At p-block Trends in Group 2 Compounds • Progressing down group 2 the atomic radius increases due to the extra shell of electrons for each element. • Going down the group the first ionisation energy decreases there is more shielding between the nucleus and the outer electrons and the distance between the nucleus and the outer electron increases and therefore the force of attraction between the nucleus and outer most electrons is reduced. • Generally the melting point of the metals decreases down the group this is because as the metal ions get larger the distance between the bonding electrons and the positive nucleus gets larger and reduces the overall attraction between the two. For similar reasons the electronegativity decreases. • The reactions of the elements with water become more vigorous down the group. When they do react they produce hydroxides and hydrogen. • The solubilities of the hydroxides of the elements increase going down the group. • The solubilities of the sulphates of the elements decreases down the group. • Barium sulphate is insoluble and is used as a qualitative test to identify sulphate ions. Trends in Period Three of the Periodic Table Property Trend from left to right Explanation Atomic radius decreases because the nuclear charge increases First ionisation energy increases because the nuclear charge increases Electronegativity increases because the nuclear charge increases Electrical conductivity Increases until the non metals because the metals have an increased number of delocalised electrons Boiling point and Melting point Increases until the middle then decreases Because these properties depend on the forces between the particles. This depends on the structure of the element which varies from metallic to giant covalent to simple molecular. Summary • Atomic Structure – We consider atoms to be formed from three fundamental particles, we can determine the relative atomic mass using a mass spectrometer. • Amount of Substance – Chemists use this concept to count atoms. Using this concept we can calculate reacting quantities, in a given reaction. • Bonding – Bonding within molecules can be described as covalent, ionic or metallic. Often a bond is a hybrid between ionic and covalent. • Periodicity – Trends within the periodic table can frequently be explained by concepts of nuclear charge and shielding.